The Hydrogen Bond: A Bond for Life 9783110627947, 9783110628012, 9783110628043

Table of contents :
Preface
Contents
1. Before getting started
2. The “bond” in hydrogen bonding
3. The “water” in hydrogen bonding
4. The order in the ice
5. Dissolving like sugar in water
6. Soaps and cells
7. Recognise your vis-à-vis
8. Jerry Donohue and the DNA
9. How to make hydrogen bonds visible
10. “Now let’s step on the accelerator”
11. Substances that build themselves
12. Like Avalokiteśvara and Durga
13. Conclusion and acknowledgements
Bibliography
Index

Citation preview

Aloys Hüttermann The Hydrogen Bond

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Physics of Wetting. Phenomena and Applications of Fluids on Surfaces Bormashenko,  ISBN ----, e-ISBN ----

Hydrochemistry. Basic Concepts and Exercises Worch,  ISBN ----, e-ISBN ----

Aquatic Chemistry. for Water and Wastewater Treatment Applications Lahav, Birnhack,  ISBN ----, e-ISBN ----

Aloys Hüttermann

The Hydrogen Bond A Bond for Life

Author Dr. Aloys Hüttermann [email protected]

ISBN 978-3-11-062794-7 e-ISBN (PDF) 978-3-11-062801-2 e-ISBN (EPUB) 978-3-11-062804-3 Library of Congress Control Number: 2019932010 Bibliographic information published by the Deutsche Nationalbibliothek The Deutsche Nationalbibliothek lists this publication in the Deutsche Nationalbibliografie; detailed bibliographic data are available on the Internet at http://dnb.dnb.de. © 2019 Walter de Gruyter GmbH, Berlin/Boston Typesetting: Integra Software Services Pvt. Ltd. Printing and binding: CPI books GmbH, Leck Cover image: Dr. Malte Reimold www.degruyter.com

Preface Are you one of those people who usually skip the preface of a book? I am definitely one of them, which is why drafting this preface was actually the hardest part of writing this book. In my book, I would like to elaborate on hydrogen bonding. I have chosen this subject because I wanted to present to you a book about a field of chemistry that – first of all, is easy enough to understand without needing a university degree – second, also plays an important role in our everyday lives and – third, is also highly relevant in modern science, so as not to only present decades-old findings. Although these three attributes actually conflict with each other, I believe that my subject, hydrogen bonding, is an exception and indeed meets all these requirements. You will see that this type of bond is well worth studying. This book is roughly divided into three sections. After giving you an explanation of what a hydrogen bond is (first section), I will guide you through everyday phenomena (second section) all the way to modern science in which hydrogen bonds play an important role (third section). By “modern”, I mean individual research projects and even some entire areas of research that only came into being a few years ago. Of course, I also wrote this book because I enjoy writing. So I hope you will enjoy reading it! Düsseldorf, November 2018 Aloys Hüttermann

https://doi.org/10.1515/9783110628012-201

Contents Preface

V

1

Before getting started

1

2

The “bond” in hydrogen bonding

11

3

The “water” in hydrogen bonding

15

4

The order in the ice

5

Dissolving like sugar in water

6

Soaps and cells

7

Recognise your vis-à-vis

8

Jerry Donohue and the DNA

9

How to make hydrogen bonds visible

21 27

35 45 51

10 “Now let’s step on the accelerator” 11 Substances that build themselves 12 Like Avalokiteśvara and Durga

Index

121

117

71 89

97

13 Conclusion and acknowledgements Bibliography

61

115

1 Before getting started This book deals with what is called hydrogen bonding. The word “bond” implies or means that something “comes together”. In a broader sense, chemistry is basically all about bonding. A hydrogen bond is a bond between molecules. To understand this in greater detail, we first need to take a look at how bonds work to begin with. I am going to start with the very basics – but don’t worry, it’s not that hard to understand! So, if what I am about to tell you is something you already learnt in school, feel free to go straight to the next chapter. Before I get into bonds, I would like to talk about the “building blocks” that are capable of forming bonds in the first place. The essential building blocks in chemistry are elements. You have probably heard of that. These elements basically represent letters of a chemical alphabet, meaning all existing compounds are made up exclusively of these elements, just like all words of our (written) language are made up of letters of the alphabet. Just like all letters have a name – namely “A”, “B”, “C” and so on – every element has a name and a symbol that are used worldwide. Most of them are derived from Latin or Greek. For example, silver is called “Ag” (short for “argentum” = Latin for “silver”) and hydrogen “H” (short for “hydrogen” = Greek for “water-former” – which is a good name, because water actually contains hydrogen). While many metals have been known and named since ancient times – mostly metals such as iron, silver and gold – most of the other elements were named by those who discovered them. For a long time, names were chosen based on a property of the element, such as in the case of hydrogen. Oxygen, for example, has the symbol “O” for “oxygenium” = “acid-former”, because it is found in many acids that were known in those days. It was thought that the fact that acids are acidic was due to the presence of oxygen. Although this turned out to be incorrect, the name was left unchanged. Naming conventions changed slightly over the years and, mainly in the nineteenth century, many elements were named after cities or countries, usually the country of the person who discovered the element. This includes “Ga” (short for “gallium” = France) – right next to “Ge” (short for “germanium” = Germany). France is represented twice, by the way, because there is also “Fr” = francium. The element “Po” = polonium gained unfortunate fame due to the poisoning of Russian Putin critic Alexander Litvinenko in 2006. By contrast, no elements were named after Great Britain and the United States during that period (which I find surprising), while the small Swedish village of Ytterby inspired the names of as many as four elements because of the minerals found in the nearby Ytterby mine: yttrium “Y”, ytterbium “Yb”, terbium “Tb” and erbium “Er”. Today, more than 100 chemical elements are known, but only about 90 of them are stable enough to be used in chemistry on a larger scale. The other elements are https://doi.org/10.1515/9783110628012-001

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radioactive and decay within a short period of time. But, of course, they are still given a name and a symbol! Apart from the tradition of naming elements after cities or countries, the names of famous scientists have established themselves as a particularly popular choice. That is why, alongside elements such as americium “Am” (America did get the honour, after all), californium “Cf”, berkelium “Bk” (named after the Californian university town of Berkeley) and its German counterparts darmstadtium “Ds” as well as hassium “Hs” (for the German state of Hesse), there is also einsteinium “Es” (after Albert Einstein), curium “Cm” (after Marie Curie) and roentgenium “Rg” (after Wilhelm Röntgen), among many others. The most recently named element is oganesson “Og” (after Juri Ogenasjan, one of the co-discoverers), discovered at the Joint Institute for Nuclear Research in Dubna, Russia. All these elements are listed in what is called the “periodic table of elements”, because many properties of the elements repeat periodically. More than 90 elements used in chemistry are metals, such as iron or copper. The chemistry of these elements is (usually) called “inorganic chemistry”. Only a small number of elements, about ten, are of particular interest for this book, with the most important ones being carbon, hydrogen, oxygen and nitrogen. They have the symbols C (for carbon), H (for hydrogen), O (for oxygen) and N (for nitrogen). The chemistry of these elements, specifically that of carbon, is (usually) called “organic chemistry”. Dating back to a time when chemistry was still relatively young and many things were not yet fully understood, the terms “inorganic” and “organic” are actually outdated. But since they turned out to be rather practical, they continue to be used. It was long believed that “organic” material, namely, molecules that can be obtained from living organisms such as animals or plants could not be synthesised from “dead” or “inorganic” materials such as minerals. The formation of organic materials was thought to always require a kind of vital force called “vis vitalis”, allowing organic materials to form only from other organic materials. However, in 1828, German chemist Friedrich Wöhler was able to demonstrate that this was not the case by producing an organic material, urea, from inorganic starting materials. There is no essential difference between organic and inorganic materials; a “vis vitalis” is not required. But since minerals usually do have a different composition than compounds that are found in animals or plants, these terms continue to be used for classification. But there is no need to go into this any further. Instead, let us start with this essential question: Why do elements form bonds, anyway? That is a good question, because actually there are elements that hardly form any bonds and are “happy by themselves”. These elements are noble gases and are called helium, neon, argon, krypton and xenon. These noble gases exist in the form of atoms, which means “alone”. They hardly form any bonds at all, not even with themselves (hence the name “noble”). Instead, the individual atoms in every noble gas occur in the form of balls that are floating around at different speeds depending on the temperature. But every atom is completely independent and has nothing to

1 Before getting started

3

do with the others. Only when noble gases are cooled to a very low temperature do they turn liquid due to the slower movement and several other effects, which we need not discuss any further. But a noble gas has to get very cold for that to happen. Helium, for example, only turns liquid at –268 °C (−450 °F).1 That is only 5 degrees above absolute zero! The melting point of neon, the next larger noble gas, is at a “milder” −246 °C (−410 °F) and that of argon is at –186 °C (−303 °F). Due to these properties, it took a long time for noble gases to be discovered, despite the fact that argon makes up almost 1% of the atmosphere, making it more than a thousand times more common on the Earth than silver or gold! The discoverers of argon thought it was a bit “inert” because it does not react (Greek: argos). The other Greek-derived names of the noble gases also attest to the fact that they stand apart from the other elements: krypton is derived from kryptos = hidden, neon from neos = new, xenon from xenos = foreign. Helium is derived from helios = sun, because helium had already been confirmed to exist on the sun by 1868 and would not be discovered on the Earth until much later, in 1895. So, how come noble gases hardly ever form bonds?2 It is because of their internal structure. An element can be distinguished from another by the number of negatively charged electrons and positively charged protons. If nothing else has happened to the element, the number of electrons and protons is identical. The simplest element is hydrogen with one electron and one proton, followed by helium with two electrons and two protons, and so on. The protons are located in the atomic nucleus along with the neutrons, which are not relevant in this case, while the electrons are in the shell that surrounds the atomic nucleus. The elements are arranged in what is called the “periodic table of the elements” in order of increasing atomic number, comparable to the way letters are arranged in the alphabet. It has been discovered that certain numbers of electrons are special, namely 2, 10, 18, 36 and 54. Elements with these numbers of electrons have special chemical stability. Why is that? Using a model that is relatively simple but perfectly sufficient for this purpose, you can visualise how electrons are not simply present in the atomic shell, but distributed across several shells that resemble the layers of an onion. Whenever a shell is full, there is great stability.

1 Although Daniel Gabriel Fahrenheit (1686–1736) was a German physicist, Germany uses the common Celsius scale, named after Swedish physicist Anders Celsius (1701–1744). That is why all temperatures in the German original version were indicated in degree Celsius. However, this English version includes the Fahrenheit scale as well as mph. 2 I intentionally wrote “hardly ever” because there are, in fact, a few known noble gas compounds. However, it is pretty difficult to get noble gases to react and two of the five noble gases (namely helium and neon) still have no known chemical compounds.

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The shells can hold different numbers of electrons. The first shell closest to the atomic nucleus can hold two electrons, the next two shells can hold eight electrons and the two after those 18, with the shells being filled from the inside out. Explaining why exactly it is that way would be going too far. All in all, we can basically visualise it as follows: Atomic nucleus

2 electrons 8 electrons 8 electrons 18 electrons 18 electrons

It is not by chance that noble gases have these “magic numbers” where the electrons exactly fill the shells! Helium has two electrons (i.e. the first shell is full), neon 10 (i.e. the first two shells are filled), argon 18 (the first three), krypton 36 (the first four) and xenon 54 (the first five). The other elements that do not have these numbers of electrons are therefore often “tempted” to also reach a condition where one of these “magic numbers” is present – they form bonds. How can this happen? There are basically three different possibilities:

1.1 An atom loses or gains electrons This type of bond is found in sodium chloride, which is the salt you put in your food. Chemically, salt consists of two elements, namely sodium and chlorine. As shown in the periodic table, sodium has 11 electrons, which is one more than neon. While sodium’s two innermost shells are completely filled, there is only one single electron in the third shell. But this does not mean that sodium is unstable. On the contrary, sodium is essentially a stable metal. However, having ten electrons would make it even more stable, as we have seen earlier. That is why sodium “strives” to lose one electron to reach neon’s “magic number” of stability. Well, sodium does not strive, of course, because it has

1.1 An atom loses or gains electrons

5

no will of its own. But for the purpose of this book, I will to stick with this kind of figurative language. The exact opposite is the case with chlorine. It has 17 electrons, with the first two shells being completely filled and the third containing seven electrons. The next “magic number” would be reached by argon having 18 electrons, which means gaining one additional electron. When you combine sodium and chlorine to produce a reaction (which happens very easily – it is actually relatively dangerous), sodium donates one electron to chlorine. Both have now reached a “magic number”, with sodium having 10 and chlorine 18 electrons. But even in its new state, sodium still has eleven protons, because the atomic nucleus, which contains the protons, is never involved in a chemical reaction. But now it only has ten electrons. Since electrons are negatively charged and protons are positively charged, one negative charge is now lacking – in other words, there is now one excess positive charge. This means sodium is now electromagnetically charged, having a single positive charge. Such particles are called ions – simply meaning that an electric charge is present. Hence, sodium chloride does not contain (metallic) sodium, but sodium ions. This distinction is important, because sodium ions differ from metallic sodium in a great number of properties. Again, the exact opposite is the case with chlorine. It now has one extra negative charge. Sodium chloride thus contains “chloride ions”. To indicate that the chlorine is negatively charged, it is called “chloride”, hence the full name “sodium chloride”. At this point, I would now like to give you a brief overview of how chemical reactions are visually represented. Let us start with chlorine. Prior to gaining an electron, there were seven electrons in chlorine’s outermost shell – the inner ones are irrelevant, because they are completely filled – sometimes written as dots surrounding the chlorine. This is done using chlorine’s chemical symbol “Cl”. When adding one more electron (abbreviation: e–), the shell is full. But chlorine is now electrically charged, which is indicated by a superscript minus sign. Altogether, it looks something like this:

The above figure is called a chemical equation. The starting substances of the reaction, in this case the chlorineatom and the electron, are called reactants and are placed to the left of the arrow. The products, which are the substances formed by the reaction, are placed to the right of the arrow. In this case, the product is a chloride ion. Usually, chemical equations only include those substances that are involved in the reaction.

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An equivalent equation can be written for sodium. Previously, sodium’s outermost shell contained only one electron. As with chlorine, the inner shells are not relevant. When sodium loses its electron, a positively charged sodium ion is formed, expressed by the superscript plus sign:

Combining these two chemical equations produces the following equation:

As you can see, there is an electron on both sides of the equation. This can simply be crossed out, shortening the equation as follows:

This equation now looks better to a chemist. Chemical equations should ideally contain only elements and no longer include any electrons; otherwise, it is very likely that something went wrong. Since this is not the case here, this equation has turned out alright. It is important that there is no chemical “cohesion” between the sodium ions and the chloride ions in the salt. Once the electron is exchanged, ions act similar to noble gases and are “happy by themselves”. They are present in sodium chloride in the form of balls – much like the noble gases discussed earlier. However, the ions are charged. Unlike with noble gases, which have no charge, this causes an attractive force between them. Sodium ions and chloride ions attract each other like magnets. Negative attracts positive and vice versa, just like the positive and negative poles of a magnet. But the difference between the two ions and a magnet is that every magnet has a positive pole and a negative pole. That is not how it works in sodium chloride. The sodium by itself is the positive pole and the chlorine (or chloride) by itself is the negative pole. But the actual attraction is very similar. This attractive force causes sodium chloride to be a solid instead of a gas. This is called an ionic bond, although the word “bond” is actually not quite correct. It would be more appropriate to call it “ionic attraction”, but the term “ionic bond” has still become established in chemical terminology. I will address this in more detail later. This bond is very strong, by the way. Pure sodium chloride is extremely stable and only melts at 808 °C (1,487 °F). By comparison, zinc melts at 419 °C (786 °F) and silver at 960 °C (1,760 °F).

1.2 Two atoms share electrons

7

1.2 Two atoms share electrons In addition to the possibility of one atom losing an electron and another gaining it, it is also possible for two atoms to “share” an electron. This works as follows: Let us take hydrogen, which has one electron. To reach a “magic number”, hydrogen could lose this electron. Then it would have none (zero is also a “magic number”), with only the proton being left. In reality, this does happen very often. We will get to this later. The other possibility would be to gain one electron. Hydrogen would then have two electrons, reaching helium’s “magic number”. But it is relatively difficult for hydrogen to gain one electron, because hydrogen is very small. After all, that second electron needs to go somewhere! While there are some compounds in which hydrogen is negatively charged (known as hydrides), they are not that easy to produce. Hydrides were not known by chemists until after World War II. The third possibility is that two hydrogen atoms form a bond, producing a molecule. A molecule is understood to be any chemical compound that consists of more than one atom. In this bond, the two electrons, one of each hydrogen atom, are basically “available” to both hydrogen atoms. They are positioned in the (imaginary) middle of the bond. There are always two electrons in a bond and never just one or three, for example. By forming a molecule, each hydrogen atom has “access” to the other’s electron. This means that both hydrogen atoms now “possess” two electrons, so both hydrogen atoms have reached a “magic number”! This can also be visually represented using dots, with “H” being the chemical symbol for hydrogen:

The two dots positioned between the two atoms forming the bond are usually written as a dash to show that they form a bond. A dash then represents two electrons. I am sure you have seen this before. The above representation would then look as follows:

The big difference between this bond and sodium chloride is that hydrogen contains a true bond, which means that the two hydrogen atoms are firmly attached to each other. This is also called an atomic bond, to emphasise that in this case the atoms are really “stuck” together. This is also emphasised by using the term “hydrogen molecule” or H2, for short. Unlike the spherical shape of sodium or chloride

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in sodium chloride, a hydrogen molecule looks more like a dumbbell. The energy that is needed to break a hydrogen molecule up again is exceptionally great, requiring temperatures of several thousand degrees or voltages of 3,000–4,000 V! The chemically correct term for this type of bond is “covalent bond”. This name means (when translated) that two atoms “cooperate” by sharing bonding electrons, which were previously also called “valences”. I will therefore always use the term covalent in the following chapters. Atoms can also form more than one covalent bond with each other, like nitrogen, for example. Nitrogen has seven electrons, which means it is three shy of neon’s “magic number” of ten. The first shell is completely filled with two electrons while the second shell contains five electrons. The number it needs to reach is eight. Two electrons can be added per bond, with one coming from each partner. Nitrogen therefore needs three bonds to reach the “magic number”. That is exactly what happens in nature. A nitrogen atom forms three covalent bonds with another nitrogen atom, basically gaining three additional electrons by producing a nitrogen molecule (N2). The magic number has been reached!

The figure also shows the more common representation using dashes on the right. Again, one dash represents two electrons. It is also important to note about this representation that the two electrons that fully remain with a nitrogen atom are also represented as a dash. They are also called “electron pairs“. The representation is a bit figurative, but not completely wrong. As a matter of fact, these electron pairs are very important, as we will see later.

1.3 Many atoms form what is called a “metallic bond” A metallic bond – as its name implies – is found in almost all metals, such as iron, copper, gold or silver. But their bonding relationships are comparatively complicated and not easy to explain. Luckily, metallic bonds are not required for the purposes of this book. That is why I will skip metallic bonds and move on without explaining them any further. How can we predict what bond is present in a chemical substance? And how can we determine which element forms which bond? It is not that easy and depends on the situation. I will give you an example: In sodium chloride, chlorine is negatively charged, as we have discussed. In chlorine gas, chlorine forms a covalent bond with itself. It makes sense: Chlorine needs an additional electron to reach a “magic number”, as we have seen. If one chlorine atom were to simply “take away” an electron

1.3 Many atoms form what is called a “metallic bond”

9

from another, it would reach a “magic number” – but the other chlorine atom would be even further away from it than before! So, if chlorine forms a covalent bond instead, both chlorine atoms can reach the “magic number” of 18 electrons, with eight being in the third shell. The result is a chlorine molecule (Cl2). Cl2 is a green gas, hence chlorine’s name, which is derived from chloros = Greek for green.

However, chlorine still has a strong “attractive force” for electrons, which means it has a tendency to form a chloride ion with a single negative charge. That is why chlorine gas reacts extremely easily with other substances and is also very toxic. But once chlorine has received its electron, it is comparatively “friendly”. Chloride ions are relatively non-toxic. After all, salt can be eaten. Oxygen is another example. For instance, in a reaction between carbon and oxygen, when done correctly, a compound called carbon dioxide or CO2 is formed. As the name suggests, this compound contains two oxygen atoms and one carbon atom. Carbon has six electrons, with four being in its second shell, and therefore needs another four electrons to reach a “magic number”. Oxygen has eight electrons, with two being in the first shell and six in the second, and therefore needs another two electrons. In CO2, carbon (chemical symbol: C) forms two covalent bonds with each of the oxygen atoms (chemical symbol: O). In other words, carbon forms a double bond with each of the oxygen atoms, enabling all of the atoms to reach a “magic number”:

The situation is different with calcium oxide. Calcium oxide is “burnt lime”, a compound most readers are probably familiar with. It consists of calcium (chemical symbol: Ca) and oxygen (chemical symbol: O), abbreviated as CaO. Calcium has 20 electrons, which is two more than argon. Its first three shells are filled, with two electrons remaining in the fourth. The same thing that happens

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with sodium chloride happens with calcium oxide. Calcium donates its two electrons to oxygen, allowing both to reach a “magic number”.

But covalent bonds and ionic bonds may also occur next to each other. Since this will be important in a few sections later in my book, I will briefly address it below. A good example of such a mixed bonding situation is potassium cyanide, which, at least according to Peter Gabriel, is one of the ingredients in the magic wands used by cheerleaders. Potassium cyanide consists of three atoms, namely potassium, nitrogen and carbon. Since you are already familiar with nitrogen and carbon, let me now introduce you to potassium: Potassium, “K”, can be thought of as something like sodium’s “big brother”, situated in the periodic table between calcium and argon. This means it has 19 electrons, with its first three shells being filled and the fourth holding a single electron. When examining the bonding situation in potassium cyanide, it becomes apparent that carbon and nitrogen are connected by a covalent triple bond, comparable to that of nitrogen in the atmosphere. However, carbon has one fewer electron than nitrogen, so only nitrogen is happy while carbon still needs one electron.

No problem! Potassium donates its solitary electron, thereby reaching argon’s electron status of 18 electrons. Now all involved atoms have reached a “magic number”.3

This produces a total of two ions, one positively charged potassium ion and one negatively charged ion consisting of nitrogen and carbon. This is also called cyanide or cyanide ion. As described earlier, carbon and nitrogen form as many as three covalent bonds in a cyanide ion. The name is derived from Greek (kyanos = blue) and stems from the intense colour of ferric ferrocyanide, which is also known as Prussian blue. This also inspired the German name for hydrogen cyanide, “Blausäure” (“blue acid”), even though it is actually perfectly clear and colourless.

3 I would like to briefly point out that potassium cyanide cannot be produced the way the two equations suggest! They only serve to explain the bonding situation and are therefore rather “formalistic”.

2 The “bond” in hydrogen bonding The hydrogen bond was not included among the bonds discussed in the previous chapter. This was done for a reason, because this bond differs significantly from other bonds. But before we actually get to hydrogen bonding, we first need to take another closer look at the first two types of bonds from the previous chapter. In reality, bonding is still a bit more complicated. In addition to the possibilities 1 (ionic bonding, in which electrons travel from one atom to another) and 2 (covalent bonding, in which the electrons are positioned in the middle between two atoms), there are some in-between possibilities. This is best visualised by imagining that the two electrons in a bond are not positioned in the middle between the two atoms, but closer to one of the atoms. As long as the electrons are not too far away from the middle, both atoms can still “access” both electrons. If the electrons travel further to one side, a situation occurs where the bond is not yet “broken”, but basically only one atom gets to use the so-called bonding electrons. If the electrons are pulled to one side even more strongly, you can imagine the bond being broken, resulting in an ionic bond like the one we know from sodium chloride. Of course, the bond does not actually “break” in reality – a covalent bond does not form in the first place. Hence, what determines the position of the bonding electrons in the bond? It depends primarily on the two atoms involved. Logically, you can imagine that in bonds formed between two identical atoms – such as the hydrogen molecule H2 or the chlorine molecule Cl2 – the bonding electrons are positioned exactly in the middle. In a bond between carbon and hydrogen, the electrons are also pretty much exactly in the middle. By contrast, in an oxygen–hydrogen bond, the bonding electrons are distributed very unequally and positioned very close to the oxygen. Still the bond, however, exists, thus allowing hydrogen to reach the “magic number”. This is also true of chlorine-hydrogen and nitrogen-hydrogen bonds. In an effort to calculate the specific differences, a measurement called electronegativity was introduced and calculated for almost every element. So, if two atoms have identical or similar electronegativity, the electrons tend to be positioned in the middle. If the difference in electronegativity is slightly bigger – but not too big – a bond is formed where the electrons are closer to one side, namely that of the more electronegative element. If the differences in electronegativity are too big, an ionic bond is formed instead of a covalent bond. I will not go into the individual numerical values – especially because there are various calculation methods and some of the values may differ depending on the method used. For the purposes of this book, it is sufficient to distinguish between three groups: https://doi.org/10.1515/9783110628012-002

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2 The “bond” in hydrogen bonding

Chlorine, oxygen and nitrogen (group 1) have high electronegativity while hydrogen and carbon (group 2) are moderately electronegative and the vast majority of metals such as sodium and calcium (group 3) have low electronegativity. Whereas bonds between atoms of one group tend to have “equal rights”, bonds between atoms of groups 1 and 2 are “unequal”, with the electrons being closer to the element from group 1. Atoms of groups 1 and 3 form no covalent bonds at all and instead form ionic bonds like in sodium chloride. Bonds between atoms of groups 2 and 3 are a bit more complicated, but since they are not required for this book, I will leave those out. Let us take a look at the bond between oxygen and hydrogen. Oxygen is in group 1 and hydrogen in group 2, producing a covalent bond in which the electrons are closer to oxygen. So, what happens as a result of the unequally distributed electrons? What happens is something very important and crucial: The bond itself is no longer equal. This type of bond is also called polar bond. This implies that this bond basically contains a negative and a positive pole, with the negative pole being located at the more electronegative atom and the positive pole at the other atom. By contrast, bonds in which the electrons are positioned closer to the middle are known as non-polar bonds. Let us examine this using a substance you know well: water. Water consists of three atoms, namely one oxygen atom and two hydrogen atoms. As we have learned, oxygen has six electrons in its second shell. By bonding with two hydrogen atoms, which can each form one bond, water is produced. It has the following structure (shown using dots on the left and dashes on the right):

The figure on the right shows the two electron pairs on the oxygen atom that are made up of the four electrons which have remained with the oxygen atom. However, as I have mentioned, in a bond between oxygen and hydrogen, the electrons are positioned closer to the oxygen rather than being distributed equally. As a result, oxygen has a slightly negative charge and hydrogen has a slightly positive charge. However, the electrons are still in a covalent bond, without forming any ions where the electrons would be with one element only. This is indicated by adding the Greek symbol “δ” instead of just writing “–” or “+”. Hydrogen is given a δ– and oxygen a δ+:

2 The “bond” in hydrogen bonding

13

This means that the electrons in this bond are closer to oxygen than to hydrogen, producing a kind of semi-positive charge on the hydrogen and a semi-negative charge on the oxygen. However, these semi-positive and semi-negative charges cause δ+ and δ– to attract each other – similar to what happens in sodium chloride! In addition to the normal covalent bonds in water, further bonds form between the water molecules. They can basically be represented as follows:

The oxygen atom in the water molecule has another four electrons, or two electron pairs, represented by two dashes, that are not involved in the bonds with the hydrogen atoms. They are also known as “free electron pairs”. An attractive force now forms between hydrogen and each of oxygen’s free electron pairs, which is usually also called a bond in chemical terminology. More specifically, it is called a hydrogen bond – which brings us to the actual subject of our book. As an example, two of these hydrogen bonds in water are shown as dashed lines in the above figure. Why is a hydrogen bond called hydrogen bond, anyway? Quite simply because a hydrogen atom is often involved as δ+ in these bonds and because the hydrogen bond was first discovered in molecules that contain hydrogen. But a hydrogen bond is not limited to hydrogen. On the contrary, there are many so-called hydrogen bonds that do not contain any hydrogen at all. For the purpose of this book, a hydrogen bond is defined as follows: A hydrogen bond is the attractive force between a positive charge, either a true positive charge, such as the one in Na+, or a δ+ charge, such as the one in water, and a free electron pair of another atom.

But I would like to emphasise straight away that there are major differences between a hydrogen bond and a normal covalent bond. The two atoms in a covalent bond are tightly attached to each other. For example, neither of the two hydrogen atoms of a hydrogen molecule can simply detach from the bond and move away from the other. A lot of energy is required to accomplish this – and the hydrogen molecule is destroyed in the process! This is different in a hydrogen bond, where the individual bonding partners can easily move away from each other – and do so all the time. Hydrogen bonds are also more flexible in terms of length and position than covalent bonds, which are very rigid formations.

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2 The “bond” in hydrogen bonding

The best way to imagine the difference between a hydrogen bond and a covalent bond is to think of the Vienna Opera Ball, with the dancers being the molecules. When you are dancing at the Vienna Opera Ball, there is basically a true molecular bond between you and your hand. Although it is theoretically possible to separate you from your hand, it would be very cruel and difficult and you would then not really be whole anymore! But in this large ballroom, there are many couples dancing a waltz. When the waltz is finished, the couples usually separate and change dance partners. But while they are dancing, each couple also forms a bond, even though this bond may sometimes be a bit shorter or longer and does not necessarily always move in the same direction. So, when the male dancer makes a turn and knows how to dance a waltz, the female dancer will follow accordingly. Otherwise, he will step on her feet. . . but in any case, any move he makes produces a response from her. This bond between the dancers is comparable to a hydrogen bond. If we installed a camera on the ceiling and filmed the entire ballroom for an extended period of time, we would see that, on average, most of the dancers are attached to another person. The dance partners change from time to time – but almost all involved in the dance are, on average, not alone. Something similar happens in water, but now you need to imagine the dance partners switching constantly instead of splitting up all at once after each piece of music, like when dancing a waltz. In addition, each water molecule can form a bond with not just one, but several water molecules! This creates a kind of “threedimensional network of dancers”. I have to admit, this is now getting a bit difficult to imagine. But it will at least give you a rough idea of what happens in water, where every water molecule is bound to several other water molecules by hydrogen bonds. So, despite not being a “true” covalent bond, a hydrogen bond is still very much existent and important.

3 The “water” in hydrogen bonding As already mentioned in Chapter 1, the word “hydrogen” in hydrogen bond is derived from Ancient Greek, hýdor = water. More specifically, hydrogen means “water-former”. So, what significance does this hydrogen bond have for water? A crucial one! The hydrogen bonds in water are often largely responsible for the way water looks and acts in nature. Water is a truly remarkable compound. Most people are not amazed by the surprising properties of water, but that is only because water is found everywhere on our planet. If water had not been known before and were to be newly synthesised by chemists at a university’s chemistry department or a Max Planck Institute, people would be stunned by the amazing properties of water. The chemists involved would have a good chance of getting a delightful call from Stockholm, inviting them to accept a Nobel Prize. For example, one of the remarkable properties of water is the fact that it is liquid under normal conditions, such as 25 °C (77 °F) and normal pressure. The boiling point, that is, the temperature at which water becomes a gas, is remarkably high, namely, 100 °C (212 °F). The Celsius scale is even based on water’s boiling and melting point, as most of you probably know. So, the number 100 was not picked at random. But why is this property of water so remarkable? It becomes apparent when examining the behaviour of compounds that are structurally very similar to water. For example, a compound called methane or CH4 is formed when the oxygen atom in the water molecule is replaced by a carbon atom and two more hydrogen atoms are added, because carbon needs four bonds to be happy, as we have seen. It consists of one carbon atom at the centre, with four hydrogen atoms attached to it. In the next chapter, we will discuss this compound in more detail. It can be visually represented as follows:

You are also familiar with methane, although you may not necessarily know it by this name. Methane is the primary component of natural gas. As the word natural gas implies, methane is gaseous under normal conditions. It only becomes liquid at very low temperatures, namely, at −161 °C (−258 °F). That is really very cold! Liquid methane is therefore usually not found on the Earth, but it does occur in space, such as on Saturn’s moon Titan, where there are entire lakes and rivers of liquid methane. https://doi.org/10.1515/9783110628012-003

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3 The “water” in hydrogen bonding

There is also a comparable compound of nitrogen and hydrogen. Nitrogen forms three bonds, with three hydrogen atoms being attached to one nitrogen atom. The resulting compound is called ammonia or NH3 and can be visually represented as follows:

You may have heard the name ammonia before. About 100–120 years ago, ammonia was used as a refrigerant in refrigerators. Because ammonia is highly toxic, it was replaced by chlorofluorocarbons in the 1920s, which was thought to be a huge step forward at the time, until the hole in the ozone layer was discovered. But that is another story. In any case, ammonia is also gaseous under normal conditions. The temperatures required to turn it into a liquid are not quite that low, namely, only −33 °C (−28 °F). But there is still a difference of 133 °C (272 °F) to the boiling point of water. This is almost exactly the same difference as that between ammonia (boiling point: −33 °C/−28 °F) and methane (boiling point: −161 °C/−258 °F). It is starting to become apparent that the 100 °C boiling point of water is indeed pretty high. The situation is similar when replacing oxygen by yet other elements, such as sulphur to produce hydrogen sulphide (a gas that strongly smells of rotten eggs – hopefully it does, because if you cannot smell it anymore, its concentration is so high that you would be almost dead because of how toxic it is), phosphor to produce phosphine (which is also toxic and stinks pretty badly) or chlorine to produce hydrogen chloride (which is not exactly healthy, either) – all of which are gaseous compounds that look as follows:

Hydrogen sulphide

Phosphine

Hydrogen chloride

What happens when we replace hydrogen instead of oxygen? It is not much of an improvement. When replacing the two hydrogen atoms by a second oxygen atom, the result is O2, the oxygen found in the atmosphere. Oxygen becomes liquid at even lower temperatures than methane, with its melting point being −183 °C (−297 °F)! When replacing hydrogen by chlorine, a compound called dichlorine monoxide is formed. Although it is not particularly stable or easy to produce, it is also gaseous!

3 The “water” in hydrogen bonding

17

Dichlorine monoxide

Other elements are not doing any better either. As you can see, water is pretty unique in its property of being liquid under normal conditions. But this is essential for your daily life! If water were usually a gas, it would be pretty hard to imagine that you and I could even exist. So, that is just one reason to treat water with respect. Why is that? To better understand this property, let me first give you a brief overview of the terms solid, liquid and gaseous. Known as the three states of matter, they refer to the form in which matter exists at different temperatures. From a scientific point of view, the temperature of matter is simply a measurement for the movement of the atoms or molecules it consists of. The higher the temperature, the faster the atoms move. The gaseous state of matter may be the easiest to imagine. A gaseous compound is largely characterised by freedom. The individual molecules are separate from each other and are floating through space. Of course, that is an idealisation, because if the individual molecules in a gas were truly free and independent from each other, there would never be something like wind. But on the whole, imagining a gas as an accumulation of individually separate molecules (or atoms, in the case of noble gases) is not too far from reality. If the specific gas and the temperature are known, it is even possible to calculate the average speed of the individual gas particles. For example, nitrogen, the main component of air, has a speed of more than 1,800 km/h (1,120 mph) at room temperature. How come not all compounds are gaseous, anyway? That is because in many compounds, there are opposing or attractive forces that hold together the individual particles the compound consists of. I will get to that later. As a general rule, it can be assumed that the vast majority of compounds would be gaseous at room temperature if only the kinetic energy of their particles played a role. A compound is therefore naturally gaseous at higher temperatures, because that is when the individual particles move the most, enabling the resulting energy to overcome any opposing forces. The specific temperature varies from compound to compound. For example, temperatures of −195 °C (−319 °F) and higher are enough for nitrogen, while sodium chloride needs more than 1,465 °C (2,670 °F)! A solid compound has a rigid structure. The individual molecules are arranged next to each other like bricks in a wall. Ideally, there is perfect order. Of course, that is not always the case, but overall, a solid compound can be thought of that way. The structure of sodium chloride is a good example. Sodium chloride can be visualised as follows:

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3 The “water” in hydrogen bonding

The positive sodium ions (the smaller balls) are each surrounded only by negative chloride ions (the larger ones) and vice versa. This is due to the fact that positive and negative charges attract each other – but negative and negative charges repel each other (positive and positive ones do the same, of course). Because of this structure, the like-charged ions are kept apart from each other and are surrounded only by oppositely charged ions. Calcium oxide, which I introduced to you in Chapter 1, has exactly the same structure, by the way, but is made up of calcium and oxygen ions instead of sodium and chloride ions. It makes sense that solids occur mainly at low temperatures, because that is when the individual particles move the least. Solids that remain solid even at higher temperatures are characterised by strong attractive forces that prevent the individual particles from moving too far from their position and thereby breaking up the structure into which they are arranged in the solid state. In sodium chloride, this is caused by the electromagnetic attraction between the positive sodium ions and the negative chloride ions. As explained above, this force is very strong in sodium chloride, which only turns liquid at 808 °C (1,487 °F). In a liquid, the situation is different, and it is also not that easy to imagine a liquid. Many properties of liquids, including those everybody knows, are still being researched and even experts are still unsure about how a liquid is structured in all its details. One way to attempt to understand this is to imagine that even though the molecules in a liquid are constantly moving, closely adjacent molecules or atoms still form a defined structure. However, this structure does not extend throughout the entire liquid, as would be the case in a solid, but instead “breaks off” after a short length. But even in a liquid, there definitely have to be attractive forces that bond at least closely adjacent molecules to each other. Otherwise, the compound would be a gas instead of a liquid. In water, these attractive forces are almost exclusively the result of hydrogen bonds, as has been measured. Because each water molecule is bonded to several other water molecules, these bonds are in the aggregate extremely strong, causing the high boiling point of 100 °C (212 °F). According to estimates, water without any

3 The “water” in hydrogen bonding

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hydrogen bonds would have a melting point of −100 °C (−148 °F) and a boiling point of −75 °C (−103 °F), which would be 100 °C (148 °F) and 175 °C (251 °F) less, respectively, than seen in nature! Only when the temperature exceeds 100 °C (212 °F) does the resulting movement of the water particles cause the structure of the molecules to be broken up due to the increased movement. Water then becomes gaseous. At this point, it is helpful to once again compare methane, ammonia and water in greater detail. The molecules are shown again below (from left to right).

Methane

Ammonia

Water

You can see that no hydrogen bonds can exist in methane, because this would require carbon to have free electron pairs, which are not present. Moreover, the bond between carbon and hydrogen is almost non-polar, so a δ+ charge on the hydrogen atom does not exist either. You may remember from the previous chapter that hydrogen and carbon are classified into the same group (group 2) of electronegativity. This means that the electrons in a bond between these two elements would be positioned roughly in the middle. Since methane also contains no ionic bonds, there are neither hydrogen bonds nor electromagnetic forces to counteract the movement of the particles. That is why it is not surprising that methane is gaseous at room temperature. Why does methane turn liquid at some point, anyway? After all, there is a difference of more than 80 °C or 150 °F between the boiling point of methane (−161 °C/−258 °F) and the boiling point of neon (−246 °C/−410 °F). This question is actually not that easy to answer. Methane, among many other similarly structured compounds, has been found to contain the so-called van der Waals forces, which cause methane to become liquid at a sufficiently low temperature – but at a higher temperature than neon, which has no such forces. But we need not go into this any further. In any case, methane does not contain any hydrogen bonds that would attract two methane molecules to each other. van der Waals and other forces are much, much weaker, as can be measured. Even a comparatively low energy such as the kinetic energy between the particles suffices to separate the methane molecules from each other, thereby turning methane into a gas. The situation is different in ammonia. This is partly because it has polar bonds, with nitrogen being δ− and hydrogen δ+. If you remember the previous chapter, nitrogen is in group 1 and hydrogen in group 2, so the electrons are closer to nitrogen.

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3 The “water” in hydrogen bonding

Nitrogen also has another free electron pair, as we can see, proving that hydrogen bonds do exist in ammonia. But when comparing water and ammonia, you will see that the relationship between the hydrogen atoms and the free electron pairs in ammonia is a bit awkward. There are too many hydrogen atoms – namely, three – compared to the single electron pair. When applied to the Opera Ball analogy, this would mean that there are three times as many gentlemen as there are ladies. So all the ladies are dancing but more than half of the gentlemen are standing around without a partner. It is therefore not surprising that the boiling point of ammonia ranks almost exactly in the middle between those of methane and water. Ammonia does have hydrogen bonds that cause it to become gaseous at much higher temperatures than methane – but it does not have as many as water because of the awkward ratio between the bonding partners. Hence, methane has a boiling point of−161 °C (−258 °F) while ammonia boils at−33 °C (−28 °F) and water at 100 °C (212 °F).

4 The order in the ice Water has another absolutely remarkable property that is a true miracle when you look at it more closely, even though everybody may be aware of it. It is the fact that water is less dense as a solid than as a liquid. More specifically, water is most dense at 4 °C/39 °F (even more specifically, at 3.983 °C/37.4 °F). At higher or lower temperatures, the density of water starts to increase again. The difference in density between solid and liquid water is considerable, with water expanding by about 9% when freezing. That is why the ice cubes in a nicely chilled cocktail float at the top instead of the bottom. What is so miraculous about this? Just think about it: This property causes water on the Earth to be liquid even at the bottom of the deepest ocean. If water were most dense in its solid state as ice, the pressure at the bottom of the ocean would instantly cause it to turn to ice. Moreover, water freezes from the top down instead of from the bottom up. On a lake or other body of water, a layer of ice initially forms at the surface and then expands downwards. But in 99.9% of all known compounds, it is exactly the other way round. They are most dense in the solid state, meaning the most molecules are present per litre or cm3 (or mm3, depending on the volume you are looking at). This is because they have the greatest order when they are in the solid state. When they are liquid, the individual molecules need more space, because they are constantly moving. Fewer molecules fit into a given space than in the solid state, meaning the density is lower. If water were a completely normal chemical compound, it would be a disaster for life on the Earth! Even at the equator, the ocean would be solid ice below a depth that is no longer reached by enough sunlight (maybe about 50–100 m below the surface, I don’t know the exact depth) all the way down to the bottom. Considering that oceans are, on average, several kilometres deep, more than 99% of the world’s oceans would consist of ice. At the poles, the Earth would even be frozen solid! Because water freezes from the top down, it also forms a kind of natural protective barrier against freezing solid. That is not the case with all other substances, which freeze from the bottom up. Needless to say, life as we know it would be almost impossible to imagine under these conditions. Why is that? Again, this can be explained by hydrogen bonding. But to do so, I first need to explain the spatial arrangement of atoms in water in a bit more detail, again using the figures of methane, ammonia and water we have already seen in the previous chapter: https://doi.org/10.1515/9783110628012-004

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4 The order in the ice

Methane

Ammonia

Water

Let us start with methane. In the above figure, a methane molecule looks as if it were flat “like a pancake”. But this is not the correct spatial structure that methane has in nature. Its structure can be examined using appropriate methods and we know pretty exactly what it looks like. You can also picture methane’s structure by imagining that the four bonds originating from the carbon atom try to move away from one another as far as possible. This is partly due to the fact that the hydrogen atoms need space – but mainly because the bonds are formed by electron pairs. Electrons are negatively charged, so the further the electron pairs are away from each other, the better. Take a moment and try to imagine it. It’s not that easy! When expressed in mathematical terms, the object we are imagining would be called a tetrahedron. A tetrahedron looks like a pyramid, but unlike the pyramids of Giza, it has a triangle as the base instead of a square. The carbon atom sits at the centre, and the hydrogen atoms in the corners. (Any old-school fantasy role-playing gamers among you will of course be familiar with this object: a D4). Altogether, it looks like this:

It is actually amazing that the real structure of a methane molecule can be derived from this simple consideration, because it is almost a perfect match. According to the rules of mathematics, the angle which the two hydrogen atoms and the carbon atom would need to form a tetrahedron is (H–C–H) 109. 5°. This is actually the case in reality!

4 The order in the ice

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The structure of the ammonia molecule can be visualised by imagining that one of methane’s hydrogen atoms is replaced by nitrogen’s electron pair. This would look as follows:

In the above figure, the electron pair is shown looking a bit like a distorted “American football”, which – for reasons that would take too long to explain – is a fairly decent approximation. An electron pair takes up more space than a covalent bond. In somewhat simplified terms, this is because a covalent bond ends with a hydrogen atom that slightly pulls the electrons in its direction. This does not happen with an electron pair. Since the electron pairs repel each other, the structure is slightly “distorted” or “modified” to give unbound electron pairs a bit more space. This results in an H–N–H angle which is slightly smaller than that of methane, but not much. It is 106. 8°. Water has two electron pairs, so its structure looks something like this:

Accordingly, the resulting angle is still slightly smaller than that of methane, namely 104. 5°. The representation from the previous chapter, in which water was drawn in a linear manner, therefore does not conform to reality. Water has an

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4 The order in the ice

angle. But this does not fundamentally change the network of hydrogen bonds formed in water. When water turns into solid ice, something very odd happens: Ice accumulates in such a way that tetrahedrons are formed, just like in methane – the only difference being that oxygen instead of carbon is now at the centre. However, oxygen only forms two covalent bonds with hydrogen, with the two remaining spots being “occupied” by the electron pairs. But each oxygen atom in ice is still surrounded by four hydrogen atoms. How does that work? It is relatively simple: The other two bonds in the tetrahedrons formed in ice are now hydrogen bonds instead of covalent bonds. This would not be a big deal; after all, these bonds also exist in liquid water. But what is remarkable is that hydrogen bonds in ice have the same length and the same direction as covalent bonds. Covalent bonds and hydrogen bonds thus take on the same structure. When applied to the Opera Ball analogy, this would be like tying a pair of dancers together so that the dance partners can no longer move away from each other, let alone increase or reduce the distance between them. What exactly does ice look like? The exact structure is a bit complicated. For a better understanding, let us first take a look at an “ice tetrahedron”.

It is important to note that only two of the hydrogen atoms have a covalent bond with the oxygen atom at the centre. The other two are bonded to another oxygen atom, which is not shown here. These hydrogen atoms are represented by the dashed lines. To these hydrogen atoms, the oxygen atom is only bonded by the hydrogen bonds – but this is not relevant here, as we have seen. The tetrahedrons now stack up like in the figure below, which shows a fragment of the structure:

4 The order in the ice

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You have to imagine that every tetrahedron contains an oxygen atom at the centre and a hydrogen atom in each corner. When the ice melts, the individual water molecules regain some freedom to move. The water molecules now “remember” that hydrogen bonds and covalent bonds are not the same thing. This has major consequences: In a covalent bond, the distance between the bonding partners is fixed and does not change unless great energy is used. For example, it is similar to how the distance between your hand and your elbow is fixed. Trying to change it would not be that easy! But when you are dancing with someone, it is very much possible to move a little bit closer to each other or further apart during the dance, even if it is only a few centimetres. You can still keep dancing without any problem. The situation is similar with hydrogen bonds that have a higher degree of freedom than covalent bonds. But in liquid water, the individual water molecules are no longer bound by the rigid structure of the ice. This allows them to move closer to each other – if only a little bit – because the hydrogen bonds and the covalent bonds now no longer have the same structure. In keeping with the analogy, it would be as if the restraints between the dancers had been loosened and they could now move a bit closer to each other again. When looking at the structure of ice as shown above, you will notice that there is still some space between the tetrahedrons – the structure is not completely “dense”. However, water cannot make use of this space while being solid, because it is bound by the structure. This is different when water is liquid. In the liquid state, the individual molecules are far less squeezed into a rigid structure. In liquid water, everything “wiggles into place” in such a way that some of this space can still be used. As a result, more water molecules can now fit into a given

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4 The order in the ice

space – not that many, but enough to produce a difference in density that causes water to be most dense at 4 °C (39 °F). Conversely, when water freezes, the dance partners are getting tied to each other again, at a precisely defined distance and in the rigid structure shown above. This procedure of “being locked in position” takes up a bit more space than when the water molecules can position themselves as close to each other as they “like”. This causes water to expand. The resulting force is so great that it can even break rock! When heating water to temperatures above 4 °C (39 °F), the individual water molecules move increasingly faster as the temperature rises and therefore need more space. This means the density decreases as water expands. But this is completely normal and happens in all compounds. The miraculous thing about water is its state at the temperature range between 0 °C (32 °F) and 4 °C (39 °F), a phase during which its density increases rather than decreases despite the rising temperature. In somewhat simplified terms, this surprising behaviour is due to the fact that covalent bonds and hydrogen bonds in ice are treated as if they were the same, even though they are not. Only when water turns liquid does it “remember” that there is a difference. Therefore, the next time you see an ice cube in your drink, remember what a miracle it is that it floats at the surface instead of sitting at the bottom. You owe your life to this fact!

5 Dissolving like sugar in water You may have heard the expression “We are not made out of sugar”, which I always loved to hear as a kid when I was being reprimanded for not wanting to go outside in the rain. But why does sugar dissolve in water? In other words, why is water a good, and sometimes even great, solvent for some substances, but not for others? It is actually very simple: Water best dissolves those substances that are chemically similar to it and with which it can easily form hydrogen bonds. Substances with a different chemical structure dissolve less well in water. How can we find out which substance dissolves well in water which does not? “Very easy”, you might say, “I’ll just take some water and try it out.” That is true, of course, and, when in doubt, it would also be the method of choice. So, let me rephrase this question: How can we assess which substance dissolves well in water and which does not, based on its chemical structure alone? That is not that difficult either. Let us start with a simple substance: alcohol. In chemical terminology, however, the alcohol that is in your beer or wine is called ethanol instead of alcohol. That is because the term alcohol is used in chemistry to refer to a whole class of substances – the alcohol we know is just one of them. The chemical structure of ethanol looks like this:

You can see here that two carbon atoms are bonded to each other. That is nothing unusual. In organic chemistry, the branch of chemistry that deals primarily with carbon compounds, having two carbon atoms in one molecule is not that much. Carbon is one of very few elements that can form long chains or rings. In organic chemistry, this is so ubiquitous that carbon is often seen as more of a basic framework of a compound – similar to bricks in a wall. In ethanol, it can also be observed that five of the six available carbon binding sites (because carbon always forms four bonds) are taken by bonds with hydrogen. In organic chemistry, this is also a common occurrence found in the vast majority of organic compounds. The fact that carbon forms chains or rings while the remaining bonds are formed with hydrogen led chemists to radically simplify the above-shown molecular notation. I will give you a brief explanation of how this is done, because it will https://doi.org/10.1515/9783110628012-005

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5 Dissolving like sugar in water

also make things easier for me in the following chapters. When using this notation, ethanol looks like this:

Now that is really easy to write! As you can see, each bend – and the end – of this zigzag line represents a carbon atom. Unless indicated otherwise, all bonds are filled with hydrogen atoms. Sometimes things are not made quite as simple and at least the end is indicated, which then looks like this:

This notation significantly shortens ethanol at first glance. But this is permissible, because the chemical properties of organic molecules are usually not attributed to the carbon chain, but to the other atoms that are attached to the carbon chain. Hence, it was discovered that many of ethanol’s properties are still present when the carbon chain is lengthened by one atom. The resulting compound is called propanol and looks like this:

In the shorthand or bond line notation:

This is why almost all molecules that contain an OH group are collectively called alcohols. To a chemist, alcohols are therefore an incredible multitude of compounds. I will take this opportunity to insert some additional information about carbon. If you get your chemistry right, carbon forms very, very long chains. When reaching a certain length, these are usually called polyethylene in chemical terminology. You know it as the material used for plastic bags. The name is derived from the production method and need not be explained in more detail. But in this case,

5 Dissolving like sugar in water

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there are tens of thousands, sometimes even hundreds of thousands, of carbon atoms linked together! In addition, carbon has the ability to form more than a single bond with itself. This means there are also double bonds and even triple bonds. For example, adding a double or triple bond to propanol also results in the formation of stable molecules that are called allyl alcohol and propargyl alcohol and look like this:

Allyl alcohol

Propargyl alcohol

In the shorthand notation:

Allyl alcohol

Propargyl alcohol

In addition, carbon also forms rings, with pentagonal and hexagonal rings being particularly stable. Hexagonal rings with alternating single and double bonds are even more stable. These rings – which are actually “flat as a pancake” – are also called benzene rings, because the simplest of these molecules is called benzene and looks like this:

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5 Dissolving like sugar in water

You may be familiar with the name of a slightly more complex molecule in which two such rings are connected: naphthalene, which used to be the primary component of mothballs. It is also flat and has the following structure:

For all these reasons, there are far more known chemical compounds in organic chemistry than in inorganic chemistry, about 30–40 times as many. But let us get back to ethanol. The very first visual graphical depiction of ethanol is – again – not correct in terms of the real structure of ethanol. Ethanol actually looks more like it is shown in the shorthand notation, which is slightly “jagged”:

In this representation, the carbon atoms are not marked with a C, but are instead represented as “crossings”. The structure can be best visualised by imagining two methane tetrahedrons linked together, with one being upside down:

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Of course, one tetrahedron has an oxygen atom attached to it, which is not shown here. The zigzag or shorthand notation is therefore pretty close to chemical reality, which is one reason why it is so commonly used. Ethanol is divided into two parts, if you will. The left part consists of carbon– carbon and carbon–hydrogen bonds. These bonds each have the bonding electrons in the middle – the bonds are non-polar, meaning they have no positive or negative pole. The right, slightly shorter part consists of oxygen and hydrogen attached to a carbon chain. While the bonding electrons in the bond between oxygen and carbon are slightly closer to the oxygen atom (but are not really positioned on one side), the oxygen–hydrogen bond is clearly polar, with oxygen having a δ– and hydrogen a δ+. Oxygen also has two additional free electron pairs that are shown in the figure below:

Non-polar part

Polar part

Water can therefore form hydrogen bonds with the right, polar part in a similar way as it does with itself. It cannot do the same with the non-polar part of the molecule, but that is not a problem with ethanol. The hydrogen bonds with ethanol’s polar

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part are so strong that they cause ethanol to mix completely with water, as you know from experience. You can basically imagine it as follows:

The situation is different when the non-polar part becomes bigger, because then things obviously get increasingly “difficult” for water. The next longer alcohol, propanol, the structure of which is also shown above, also dissolves completely in water. The next longer alcohol after that with four carbon atoms is butanol. It no longer dissolves completely in water, having a solubility of 80 g in 1 L. With the next higher alcohols, it continues to get even more difficult until reaching those that do not dissolve at all. For example, one extreme case is the following molecule:

It still contains a somewhat lonely OH group in the bottom left corner, but apart from that, it is a true carbon–hydrogen monster. This molecule’s solubility in water is virtually zero. The name will sound familiar to you. It is called cholesterol, indicating that it contains an alcohol group. The situation is different with sugars. In chemical terminology, the name sugar means something entirely different than in everyday language, with the term sugar also standing for a whole group of substances. You may be aware of this because there are other sugars apart from common household sugar, such as glucose, maltose, fructose or lactose. These names stand for various compounds that are all classified as sugar from a chemical point of view.

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What all sugars have in common is the high number of OH groups in the molecule. (Does this mean sugars can chemically also be called alcohols? Yes, they can . . . although this may be a little confusing. Strictly speaking, sugars are a subgroup of alcohols. But pointing out this fact will not help you when the police pull you over. . .). The structure of household sugar, chemically also called sucrose , is shown below:

You will notice that this compound is relatively complex, even though it is sold by the kilo in supermarkets and consumed in even larger quantities if you have young children. That is true! I will not go into any more detail about its specifics. I only want to show that household sugar has eight OH groups and therefore dissolves in water . . . like sugar in water! The solubility of sugar in water has been measured. At 20 °C (68 °F), 1.97 kg of sugar dissolves in 1 L of water. Such a solution contains almost twice as much sugar as water! I will stay in the kitchen for a bit longer and move on from sugar to salt. Salt dissolves in water in a considerable amount, specifically 359 g in 1 L at 20 °C (68 °F), although it does not contain any OH groups at all! How is this possible? Again, it is actually quite simple if you remember the definition of a hydrogen bond from Chapter 3. Let me repeat: A hydrogen bond is the attractive force between a positive charge (either a true positive charge, such as the one in Na+, or a δ+ charge, such as the one in water) and a free electron pair of another atom.

This is why the present case offers even two opportunities for hydrogen bonds. In addition to forming between Na+ and the free electron pairs of the oxygen in water, they also form the other way around between hydrogen’s δ+ charges in water and the chloride ion’s free electron pairs. Remember, chloride has as many as four of them and is also negatively charged. A close examination of the dissolution of sodium chloride in water found that the sodium ion forms hydrogen bonds with the surrounding water molecules – and not just one, but several. Altogether, six water molecules become attached to sodium. This can be visualised as follows:

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5 Dissolving like sugar in water

The same happens in chloride. Again, each chloride ion in water is bonded to six water molecules. But the chloride ion is negatively charged and has a total of four free electron pairs, resulting in hydrogen bonds between the hydrogen atom in the water molecule and the chloride.

These hydrogen bonds cause salt to dissolve in water. There is a final topic to mention: Chloride and sodium are separated in the process. In water, sodium and chloride are almost completely separate from each other. That is not the case with sugar. Although sugar dissolves in water, each individual sugar molecule remains the whole. By the way, this difference used to puzzle chemists for a long time in the nineteenth century, and Svante Arrhenius’ discovery of the fact that sodium chloride “splits up” (or “disassociates” in scientific terminology) while sugar does not earned him a Nobel Prize in 1903.4 Arrhenius was actually a very prolific scientist who even predicted the greenhouse effect in 1895! He thought it would actually not be such a bad thing because then it would stay warm all year round . . . did I mention that he was from Sweden? 4 Bear in mind that sodium chloride is otherwise an incredibly stable substance, considering its high melting and boiling points (808 and 1,465°C, respectively), as we have discussed. But in water, sodium chloride instantly loses its stability.

6 Soaps and cells In the previous chapter, I gave you several examples of substances that dissolve well in water. So which substances dissolve poorly in water? They are mainly substances that cannot form hydrogen bonds at all or only with difficulty – namely all substances in which most of the bonds are non-polar. It is just like in the kitchen with your salad dressing: The sugar (if you use any), the salt and the vinegar dissolve well in water, but the oil does not. It floats on top. That is because salad oil is non-polar. Chemical laboratories perform most experiments in rather non-polar mixtures, by the way. That is because many compounds of chemical interest do not dissolve in water. Instead, other substances are used, such as chloroform or petroleum ether, which is pretty similar to the petrol you put in your car. Is it possible to tell whether a substance is polar or non-polar by looking at it? Yes, it is. Basically, it is really simple: The more carbon and hydrogen atoms and fewer other atoms a substance contains, the more non-polar it is and the more poorly it dissolves in water. Fats are a good example. Everybody knows that fats, just like oils, do not dissolve in water. Although there are many different fats, the vast majority of them have a very similar structure, which is as follows:

The circled portion on the left is always the same. Chemically speaking, it is called a triglyceride, but what that means exactly need not be explained here in any more detail. Only the right portion, namely the length and structure of the chains, differs from fat to fat. Oils like your salad oil, for example, basically also look like that, by the way. As you can see, fats do have oxygen atoms, but that is not enough to form a sufficient number of hydrogen bonds. The carbon chains are simply too long. But when a substance like soda lye is added to a fat, something very interesting happens – a process known as saponification. In this reaction, the fat shown above is split into four parts: https://doi.org/10.1515/9783110628012-006

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6 Soaps and cells

The part highlighted in dark grey, resembling an “m”, forms a molecule called glycerol. But this is not important in this case. The other parts, highlighted in slightly lighter grey, form what are called fatty acids. They are the ones that matter in the following section. These fatty acids differ significantly from the fat from which they form is an important aspect. I will explain this in greater detail, using one of these fatty acids, a fatty acid sodium salt, as an example to be studied more closely. It looks like this:

The big difference to the chemical structure of fat is the left part, which, even more specifically, looks like this:

The carbon chain is abbreviated here as “R” (for “residue”). That is very common in chemistry – chemists like to be lazy about that, or focus on what is essential, depending on which way you look at it. This compound is known as a carboxylic acid or, even more specifically, as a carboxylic acid salt. These compounds will be discussed in more detail later. But what is already apparent and sufficient for now is that this carboxylic acid salt has as many as five free electron pairs on the two oxygen atoms. At the same time, there is also a negative charge on one of the oxygen atoms, similar to the cyanide ion of potassium cyanide from the very first chapter. This part of the molecule is therefore very much capable of forming hydrogen bonds. The long carbon chain on the other side does not “get along” well with water.

6 Soaps and cells

37

Thus, when adding this compound to water, something very interesting happens. The molecules arrange themselves into a sphere, with the carbon chains being on the inside and the carboxylic acid groups at the surface:

The carboxylic acid salts are shown in this image as described below:

In chemical terminology, these spheres are called micelles. What is the purpose of these micelles? Well – the purpose is implied by the name of the original reaction: saponification. These fatty acid salts and the micelles that form out of them are exactly what we know as soaps! As a matter of fact, soap was made from fat using exactly the same process for centuries. There even used to be a profession called “soaper” that was devoted entirely to carrying out this reaction and making soap out of fat. Nowadays, there are hardly any soapers anymore, apart from a few exceptions, such as Tyler Durden in the movie “Fight Club”. Soaps are produced industrially and are essentially made from petroleum. But I am just mentioning this in passing. As you all know, soap can be used to clean laundry or to wash our hands, of course. Most importantly, soap removes fat from laundry. This would be difficult to

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6 Soaps and cells

do with water alone, mainly because of the poor solubility of fat in water. But soap makes it much easier, and here is why: Soaps are present in the form of micelles, which are non-polar on the inside, because that is where the carbon chains are located. This allows fats to be present inside the micelles, where they are basically dissolved. So, adding an aqueous soap solution to laundry – combined with a little bit of scrubbing – causes the fat to be stored inside the micelles:

But the micelle as a whole is soluble in water, simply because the carboxylic acid groups are arranged on the outside and are thereby able to form lots of hydrogen bonds with water. When removing the laundry from the soap solution, the micelle and the fat it contains remain in the water. This is how soap makes it possible to get the fat into water and thereby remove it from laundry! I am going to tell you about another great way to clean laundry later on in Chapter 10. But let us now move on from household chores to modern research, because researchers have been asking themselves the exact opposite question: How do we get sodium chloride out of water – and into “oil” (i.e. a non-polar solvent)? This is not that simple, because both the sodium and the chloride ions in sodium chloride are charged! How is that supposed to work? When you add salt to your salad oil, no matter how long you shake and try to dissolve it, it is not going to happen! “Well”, you might say, “it doesn’t matter, because we’ve got water to dissolve sodium chloride. We don’t need oil.” In cooking, it really may not make any difference, but, as is often the case in scientific research, there is a lot more to it.

6 Soaps and cells

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But, on that note, I first need to show you something else, because the arrangement as a micelle is not the only one that fatty acids are known to have. There is another one that is absolutely essential for life on the Earth, including your life and mine. Under certain circumstances, fatty acids arrange in such a way that the carboxylic acid groups are facing not only the outside, but also the inside, forming what is called a liposome:

As you can see, a liposome resembles a hollow sphere. It has an interior space that is filled with water and is surrounded by a kind of border. This border consists of two layers of fatty acids arranged in such a way that the carbon chains are facing in and the carboxylic acid groups are facing out:

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6 Soaps and cells

This is also known as a lipid bilayer. The word “lipid” is derived from the Greek word “lipos”, meaning “fat”. It could also be called a “fatty bilayer”, but lipid bilayer has become the established term. But why is this arrangement so important? Because an identical structure also occurs in cells, including those that are found in your body and mine. The lipid bilayer has the convenient property of preventing the water-soluble substances in its interior from reaching the outside. In order to accomplish this, they would first have to pass through the bilayer – which they cannot do because they are not soluble in the fatty layer. As a result, they are trapped inside the cell, which enables a cell to exist in the first place. However, the lipid bilayers in the cells of your body have a significantly more complex structure. For example, almost all of them contain a substance I introduced to you earlier: cholesterol. Cholesterol performs important functions, such as ensuring increased stability of the lipid bilayer. An average human body contains about 140 g of it! Moreover understanding the lipid bilayer gets more complicated when taking an even closer look at a cell in your body. All cells need minerals like sodium ions or potassium ions to be able to exist. But how do they get into the cell? Sodium and potassium are charged ions, which means their solubility in water is very high and thus their chance of crossing a lipid bilayer is virtually zero. This brings us right back to the problem we started with! Things get even more complicated when considering that sodium ions and potassium ions have completely different roles in your body, which is also responsible for closely monitoring the concentration of sodium and potassium in different parts of your body. When you are given an IV at the hospital, it should be based on 0.9% sodium chloride solution (known as an isotonic saline). If you were given a potassium chloride solution instead, you would, unfortunately, after the IV most likely be dead. As you can see, the question of how to get “salt into oil” is rather interesting. In the 1970s, the discovery that compounds like the one shown below are basically capable of locking up sodium ions caused quite a stir in the scientific community:

Compounds like these have been dubbed crown ethers. Ethers are specific chemical compounds, which I will not go into right now. But they are called “crown” ethers because, when seen from the side, their structure looks a bit like a crown:

6 Soaps and cells

41

As I have mentioned, these compounds bind sodium, with the sodium being located inside the crown. Since the oxygen atoms in crown ether have free electron pairs, several hydrogen bonds can form, as shown in the following figure:

Now what can be done with these compounds and the sodium bound in them? Very easy: They can be dissolved in “oil”! Crown ethers dissolve very well in non-polar compounds. The figure shown above and the previous image make clear how this works: The crown ether is polar on the inside, making the sodium feel “right at home”. But, towards the outside, only the carbon–hydrogen bonds are visible. They harmonise particularly well with non-polar liquids and ensure solubility. Hence, the situation is the opposite of that seen in a micelle. The sodium can “enter” the crown ether, which then dissolves in the non-polar solvent. The crown ether then serves as a kind of molecular lift for the sodium. Interestingly, (only) crown ether with ten carbon atoms, as shown above, takes sodium along. If the crown ether is slightly larger, as shown below, it no longer binds sodium:

Why? This is because sodium is too small for hydrogen bonds to form. Instead, this compound binds sodium’s “big brother”, potassium. This means that crown ethers

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6 Soaps and cells

are capable of distinguishing between sodium and potassium! These and other spectacular findings prompted the Nobel committee to award the 1987 Nobel Prize in chemistry to three scientists who, along with their doctoral candidates and teams, had been instrumental in studying crown ethers and similar compounds. But how does your body transport sodium and potassium into a cell and distinguish between them? This question is again not that easy to answer. It has been discovered that the body does this by using the so-called ion channels in which sodium and potassium are also distinguished by their size in a way that is very similar to how this happens in crown ethers. Potassium channels transport only potassium, and sodium channels only sodium. The ion channels are more or less integrated into the lipid bilayer, acting as a kind of tunnel between the interior of the cell and its surroundings. It has been discovered that the structure of certain potassium channels and the structure of crown ethers do share some similarities, although the former are obviously far more complex. In 2003, another Nobel Prize was awarded for insights into the structure of certain ion channels. To conclude this chapter, I would like to introduce you to valinomycin, a molecule found in bacteria:

As you can see, valinomycin bears a certain resemblance to crown ethers because it is also ring-shaped and has lots of inward-facing oxygen atoms. Just like the previously shown crown ether, it binds potassium very well, as opposed to sodium, which is hardly bound at all. What is valinomycin used for? It is used for the

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purpose of transporting potassium into cells, specifically hostile cells. Valinomycin is known as an antibiotic and is used in nature by certain bacteria to defend against other bacteria. Valinomycin has the ability to “store” potassium while also being able to penetrate the lipid bilayer of bacterial cells. By freely introducing potassium into hostile bacteria, valinomycin disrupts the processes in these bacteria so much that they are killed! The word antibiotic means exactly the same thing: against life. That is also the reason why crown ethers are quite toxic.

7 Recognise your vis-à-vis I would like to take a step back and briefly discuss the structure of carboxylic acids, which were introduced in the previous chapter. I will be using acetic acid as an example. You are all familiar with acetic acid. Acetic acid is found in vinegar and is responsible for giving vinegar its sour taste. Chemically, acetic acid looks like this:

It is a fairly simple compound. Unlike fatty acids, it also dissolves in water because of its short carbon chain. The exact three-dimensional structure of acetic acid is as follows:

Acetic acid is structurally divided into two parts. The left part, consisting of a carbon atom attached to three hydrogen atoms, has the tetrahedral structure we discussed earlier. The right part, by contrast, has a different structure. There is a double bond between oxygen and carbon. This is basically nothing unusual; we have already seen this in CO2 in Chapter 1. However, as a result of this double bond, the other bonds now have a bit more space, slightly expanding and flattening the entire structure. This is mainly due to the fact that with one centre and three bonds, this is the arrangement where all bonds have the greatest distance to each other. The double bond and the two other bonds on the carbon atom lie in the same plane, producing an angle of (approximately) 120°, as would be expected: https://doi.org/10.1515/9783110628012-007

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7 Recognise your vis-à-vis

120°

Tetrahedral

Flat

When examining acetic acid, it becomes apparent that acetic acid forms hydrogen bonds with itself, which looks like this:

The two flat ends fit together very well, as you can see. But what is important about this structure is that the pairing of the two acetic acid molecules is not random but strategic. The hydrogen bond is used to transmit directional information from one acetic acid molecule to the other. This directional information specifically refers to the arrangement in a line – with each second carbon atom facing outwards. That does not seem like that much (yet)? It really isn’t. Determining the direction in acetic acid any more specifically is difficult because the molecule is relatively small. In addition, it needs to be considered that a single bond in carbon can be freely rotated like an axle. That is why the following structure

cannot really be distinguished from the one below:

7 Recognise your vis-à-vis

47

Using the bond between the carbon atoms, the molecule was simply “turned around” like a wheel on an axle. The following, considerably more complex molecule makes things better:

This molecule is slightly larger, but not that difficult to understand for our purposes. It consists of four parts:

Part A is simply a carbon atom with three hydrogen atoms, like in acetic acid. This part is not relevant. Part B is similar to something I showed you in Chapter 5. It resembles a benzene ring, with one of the carbon atoms having been replaced by a nitrogen atom. This structure is called pyridine and is also very stable. It is important to note that, first, nitrogen has a free electron pair that can form hydrogen bonds and, second, that this pyridine ring is flat just like the benzene ring. Part C is a nitrogen atom that is attached to part D. Part D is again a carboxylic acid group, similar to those in fats, whose structure I showed you in the previous chapter. In fat, however, this part was attached to an oxygen atom, as indicated by the arrows in the figure below:

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7 Recognise your vis-à-vis

This is now a nitrogen atom. The combination of part C + D is also called an amide. But the only thing that matters for our purposes is that the nitrogen atom still has a bond with a hydrogen atom, so part D itself is not particularly relevant right now. What does this molecule do? It also forms two hydrogen bonds with itself:

You can tell right away that the transmitted directional information is now considerably larger. The second molecule can only attach to the first when it turns around and the Parts “D” are not on the same side, but part from each other. Hence, the first molecule determines the direction in which the second can attach to it. Systems like these have spawned an entire field of research called molecular recognition. This term stems from the fact that the individual molecules recognise each other – and, when done correctly, attach to each other the way you want them to. This can then be taken advantage of to the point of creating molecular architecture. In the next chapter, I will talk in more detail about some other benefits and uses of molecular recognition and why even our existence depends largely on the fact that molecular recognition exists. But first, I would like to show you another slightly more complex molecule, which also involves molecular recognition. It was introduced by Prof. E. W. “Bert” Meijer from the Eindhoven University of Technology in 1998 and looks like this:

Again, “R” stands for a carbon chain, which, however, is not important in this case. This molecule now forms no less than four hydrogen bonds with itself in the following way:

7 Recognise your vis-à-vis

49

This number of hydrogen bonds in a compound was a world record for a “synthetically” produced molecule at the time, so this molecule caused quite a stir. But there is a small trick to it, which needs to be explained. Just like in acetic acid, this molecule can be rotated at many positions – including the one shown in the following figure:

Rotating the molecule at this position produces the following structure:

It is instantly apparent that things are now getting a bit difficult for hydrogen bonds. So why does the molecule not rotate around this axis? Because in addition to the four hydrogen bonds, there is a fifth running through the molecule itself:

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7 Recognise your vis-à-vis

This additional hydrogen bond holds the molecule in place!

8 Jerry Donohue and the DNA At this point, I would like to bring up a molecule that shows how essential hydrogen bonds are in nature. It is DNA, or deoxyribonucleic acid. Do not expect me to explain what DNA means. The name is not essential to understanding DNA, so I won’t go into it any further. In biology, DNA is the molecule in which genetic information is stored, making it the most essential molecule without which “nothing works”. DNA has the following structure:

But this is only a small section. DNA molecules that can be isolated in nature are much, much bigger. After all, DNA can store the entire genetic material of a human being, for example. Now if you think that DNA looks pretty complicated, I agree https://doi.org/10.1515/9783110628012-008

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with you – but I will try to simplify things a little bit. DNA is usually divided into what are called the backbone and the bases. The backbone provides the framework along which the bases are arranged, and is not that important for this chapter. The bases are the essential part of DNA. Similar to how information on a computer is stored as a series of zeros and ones, DNA stores information as a series of bases. But instead of using just one and zero, there are four different bases, namely adenine (A), thymine (T), cytosine (C) and guanine (G). The backbone and the bases are structured as follows:

Base #3: Adenine (A)

Base #2: Cytosine (C)

one Backb

Base #1: Cytosine (C)

The backbone, which always repeats itself anyway, is therefore simply “ignored” in practice and only the bases are listed (e.g. CCA – which would be the molecule shown above). Readers of the German newspaper Frankfurter Allgemeine Zeitung

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may remember that on 27 June 2000, an entire page was filled with such a DNA sequence. When abbreviating the backbone as “R” (for residue), DNA is shortened as follows, depending on the base:

Adenine

Cytosine

Guanine

Thymine

This makes things a bit easier to read, so I will stick with it. But I would like to go into some more detail about the chemical structure of these four bases. These compounds are called bases because, when isolated in water, they react as bases, similar to the way soda lye does. But that does not really matter for the role of the bases in DNA. By itself, DNA (as its name implies) is actually even acidic, which is the exact opposite. But this term has still become established in scientific language, so I will stick with it. When looking at adenine, for example, you will see that it also contains alternating single and double bonds similar to those in benzene (Chapter 5) and pyridine (Chapter 7). In the figure below, the double bonds are highlighted:

Just like in benzene and pyridine, they add great stability to this molecule. Moreover, it basically becomes flat. The fact that there are not just bonds between

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carbon atoms but also quite a few between carbon and nitrogen atoms does not make much of a difference. Something similar is true of the other three bases. I would now like to take you back to the year 1953 to the laboratory in Cambridge, where James Watson and Francis Crick were attempting to understand the structure of DNA, called the “double helix” – of course, the two did not know at the time that the double helix actually was the correct structure. Apart from the fact that DNA is so incredibly important for biology, I believe one of the reasons why the story of how the structure of DNA was discovered has become so well known is that it is a bit like a great movie.5 According to popular opinion, it features all the important characters of a classic movie: – The two youthful heroes facing challenge with little money but great enthusiasm are James Watson and Francis Crick. When they published the correct structure, James Watson was just 25 years old! Watson at least presented himself and Crick that way in his book The Double Helix, published in 1968, sparking a not insignificant amount of criticism. – The seemingly unbeatable and superior opponent is Linus Pauling, then recognised as the greatest chemist of his time, who operated a giant lab in the United States. – The tragic hero Rosalind Franklin provided the decisive X-ray images of DNA that made the discovery possible in the first place, but died before she could receive the Nobel Prize. – The “supporter” of the two heroes is Maurice Wilkins, who, presumably by chance, made said X-ray images available to Watson and Crick, but not to Pauling. As in a classic thriller, the two heroes already seem to have lost the race when Linus Pauling publishes a proposal for the structure of DNA in early 1953. But after reading this publication, they realise that the proposed structure cannot be correct. This is generally explained by the fact that Linus Pauling did not see Rosalind Franklin’s DNA X-ray images. But it still seems like there is a jinx on it. They feel like they are on the right track, but somehow it all just does not fit together. Feel free to try what Watson and Crick did. Your task is to figure out the arrangement of the bases in DNA. The following information will help you: – The structure of DNA is always made up of at least two DNA molecules, also called strands. The structure proposed by Linus Pauling was based on three strands. However, Watson and Crick believed that there were only two. – Research by Erwin Chargaff had shown that the number of adenine units is equal to the number of thymine units and the number of cytosine units is equal to the number of guanine units, leading Watson and Crick to deduce that adenine and thyme as well as cytosine and guanine must somehow be connected.

5 By that I mean movies in general – not the existing TV movie made in 1986 on the same subject.

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– Based on Rosalind Franklin’s X-ray images and their own experiments,6 Watson and Crick assumed that the bases of DNA are positioned “on the inside”, that is, opposite each other. From this, they deduced that there must be hydrogen bonds between adenine and thymine as well as between cytosine and guanine. With adenine coming from one DNA strand and thymine from the other, the same must be true of cytosine and guanine. So what do these hydrogen bonds look like? Now that is your task. Adenine and thymine as well as guanine and cytosine should naturally be arranged in such a way that the backbone, “R”, is always in the same position, meaning the “Rs” of the adenine–thymine pair are pointing in the same direction as those of the guanine– cytosine pair. Moreover, the recognition pattern of adenine/thymine should be a different one than that of guanine/cytosine. The molecules are shown again below, with all usable H atoms being highlighted and all usable free electron pairs being marked by arrows. A hydrogen bond always forms between an electron pair and an H atom. Of course, the individual molecules can be turned around and also be mirrored.

Adenine

Cytosine

Thymine

Guanine

Did you try it? Seriously? And you were able to produce a structure? If that is true: Congratulations! This makes you an even greater scientific genius than Watson and Crick put together!

6 They largely consisted of constructions made from cardboard and wire. A model that two of them built is on display in London’s Science Museum.

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After all, the honest answer is: It does not fit. There is no solution. And that was Watson and Crick’s problem for a long time. But things changed when Jerry Donohue entered the story. Like in every movie, there is also the role of the “best supporting actor”, for which there is even an Oscar given out every year. And “supporting” was precisely what Jerry Donohue did. We also know the exact date: 27 February 1953. On that day, there was a conversation between Watson and Donohue.7 Like Watson, Donohue had come to Cambridge as a researcher visiting from America and was working in the same laboratory. Watson was doing experiments, unsuccessfully trying to figure out the structure of DNA. On said 27 February 1953, Donohue gave him the crucial tip. Watson had assumed the chemical structure of the bases to be like the structures published in a book by author James N. Davidson in 1950, which is exactly how the bases in this chapter have been depicted so far. Donohue informed him that those structures were not necessarily correct; on the contrary, they were most likely even wrong. More recent research had shown that the structure of two of the four bases, namely guanine and thymine, should differ from Davidson’s as follows:

Guanine according to Davidson

Thymine according to Davidson

Alternative structure

Alternative structure

7 The sources I have read about this are somewhat conflicting. Sometimes it is Watson who talked with Donohue, sometimes Crick, and sometimes even both of them. The exact content of the conversation also varies slightly from source to source. But it is undisputed that, without Donohue’s help, Pauling would probably have discovered the structure first.

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Now give it another try! It only took Watson and Crick one more day to complete a first experimental model of their double helix structure on 28 February 1953. The correct structures as well as the usable hydrogen atoms and electron pairs are shown again below:

Adenine

Cytosine

Thymine

Guanine

Well? Did you find an arrangement this time? If you did, it should look something like this, because this is how the hydrogen bonds in DNA are arranged: Thymine

Adenine

Cytosine

Guanine

By the way, it is not that easy to explain why the structure of guanine and thymine is different from the one originally proposed. In any case, it made a lot of sense to assume

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that guanine and thymine looked the way James Davidson described them. That is why without Donohue’s tip, the realisation that these structures could be wrong would probably have come too late for Watson and Crick or may never have come at all. About 1 month later, on 2 April 1953, the two submitted their now-famous manuscript to Nature, a scientific magazine. This manuscript was published on 25 April 1953 and is considered one of the most important scientific publications of the last century. In this publication, they expressed their gratitude to Jerry Donohue for “constant advice and criticism”, but that is not the most famous excerpt from this article, which goes as follows: “It has not escaped our notice that the specific pairing we have postulated immediately suggests a possible copying mechanism for the genetic material.” And that is correct, as we know today. The hydrogen bonds and the resulting molecular recognition between the bases ensure that whenever there is, for example, adenine in a DNA strand, thymine will be inserted at the same position into a second, newly forming DNA strand during replication – and vice versa; wherever there is cytosine, guanine will be inserted – and vice versa. DNA replication involves the formation of what is called a complimentary strand. This works exactly like a negative in “analogue” photography, where a fixer and a developer are used to produce photographs – some of the older readers among you may remember how this used to be done. In a negative, all the dark areas are actually light and vice versa, just like in a complementary strand, where thymine originally was adenine, cytosine was guanine and so on. When this new DNA strand is copied again, the original one is recreated, just like in “analogue” photography, where the negative produces a picture on photo paper through exposure. All of this is only possible because of hydrogen bonding – truly a “bond for life”! Considering how important this is, I will explain it in a bit more detail. This is what the hydrogen bonds in DNA look like: Thymine

Adenine

Cytosine

Guanine

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Everything fits perfectly. It is not only important that adenine/thymine and guanine/cytosine fit, but also that combinations like adenine/cytosine and guanine/ thymine do not. Because that would result in the following structure: Cytosine

Adenine

Thymine

Guanine

It is instantly apparent that hydrogen bonding is not possible between adenine and cytosine or between thymine and guanine, because in both possible pairs, hydrogen atoms would be directly opposite to each other. An adenine/guanine pair is also not possible – due to the fact alone that the structure does not offer enough space for it. The opposite is the case with cytosine/thymine, because the distances are too large. But the hydrogen bonds would not fit either. Feel free to give it a try. Hence, only adenine matches with thymine and only guanine with cytosine, making the copying mechanism unambiguous. Well, maybe not that unambiguous, because it has been known to happen in the body that the wrong base is inserted by mistake. This is estimated to happen in about 0.1% of all cases, but it depends on various factors. However, the DNA copying mechanism in your body has to be far more precise, because otherwise there is always the risk of unexpected and undesired consequences – such as developing cancer. That is why there is a veritable “proofreading machinery” to ensure that the error rate is reduced even further. But the accuracy must not be too high either, because otherwise there would be no evolution and we would all still be bacteria floating around on the Earth. Watson, Crick and Wilkins received the Nobel Prize in medicine in 1962. By the way, Linus Paul also won a Nobel Prize the same year, but he was awarded the Nobel Peace Prize in recognition of his efforts to end nuclear testing. Interestingly, he received the prize in 1963 retroactively for 1962. He had already received a scientific Nobel Prize, the one in chemistry, in 1954. Jerry Donohue returned to America

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where he went on to become a highly respected professor, without ever working with DNA again. I would like to mention one more thing. It has been discovered that the hydrogen bonds in DNA described by Watson and Crick are not the only ones found in nature. For example, there is another arrangement known as Hoogsteen base pairing, named after Karst Hoogsteen, a Dutch biochemist. However, this only works with one of these four bases, guanine. Under favourable conditions, it forms a structure that looks as follows:

Although this structure almost looks like a piece of art, it has turned out to be rather common8 and the presence of such structures, called G-quadruplexes or G-tetrads, has important control functions in specific genes. For this reason, they are the subject of intense research.

8 There is almost always another potassium atom at the centre, like in some of the crown ethers from Chapter 6. This structure has been found to add further stability. For the sake of simplicity, potassium is not shown in the figure.

9 How to make hydrogen bonds visible All the hydrogen bonds I have introduced to you so far seemed to come “out of nowhere”. I simply told you that they exist and where they are located. But that kind of approach is, of course, not an option in science, where you cannot just claim that there are hydrogen bonds between molecules of A and B. You also have to prove it. So I will now explain to you how to do this. There are several ways to prove the presence of hydrogen bonds. You are already familiar with one of them: X-ray analysis, as was used by Rosalind Franklin to analyse DNA. However, this method is comparatively complex. An easier method is proton nuclear magnetic resonance spectroscopy, abbreviated as 1H-NMR. Explaining how this works in detail is rather complicated, so I am not even going to attempt it. But this measuring technique is one of the most commonly used methods in organic chemistry for the structural analysis of substances. It delivers (at least) one signal for each hydrogen atom in a compound, depending on the atom’s “chemical environment”. The signal also depends on the number of hydrogen atoms, with two hydrogen atoms producing a signal twice as strong as that of one. The number of hydrogen atoms present can therefore be directly derived from the signal. When the measurement is performed properly and the signals are then read correctly, 1H-NMR can determine the structure of a molecule. All of this also applies to the following molecule, which I will call “A”:

The important hydrogen atom, which we will be focusing on below, is highlighted. This molecule A forms two hydrogen bonds with a second molecule B, which looks as follows9:

9 The two molecules A and B were not chosen at random, but I will get to this later in Chapter 11. https://doi.org/10.1515/9783110628012-009

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Producing this result:

This will hopefully make it easier for you to understand that the hydrogen bond changes the “chemical environment” of the highlighted hydrogen atom. The bound hydrogen atom now produces a different signal when performing a measurement. When measuring only molecule A shown at the top, a different signal would therefore be expected than when analysing a mixture of the two molecules A and B. A measurement is performed in the same way. First, the molecule A shown at the top is measured alone, usually in an amount of a few milligrams. Several mixtures of A and B are then analysed while gradually increasing the amount of B. What would now be expected?

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When only the molecule A shown at the top was measured, a signal would be expected for the highlighted hydrogen atom, looking something like this:

Of course, the other hydrogen atoms also produce signals, but let us ignore those for now. Assuming a mixture of A and B in a ratio of 5:1 is analysed, the highlighted hydrogen atom would be expected to produce two signals: one signal for the hydrogen atoms that are now bound by the hydrogen bond between A and B, and a second – slightly stronger – signal from the highlighted hydrogen atoms that are still free. This signal is stronger because there is five times as much A as there is B, which is why not all molecules of A can be bound. Altogether, the following signal is expected:

This actually happens, but only when the sample is cooled to a very low temperature. At normal temperatures, such as 20 °C (68 °F), however, the hydrogen bonds are constantly moving back and forth and the signals become “blurred”. Two signals are only visible when preventing this movement by strong cooling. But this is not done in practice, because it is very complicated and also not necessary. What is really seen is still only one signal, but slightly shifted (a bit to the left, in this example) – the “average”, so to speak.

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When analysing a mixture of A and B containing a higher percentage of B, the visible signal is even further shifted to the left. The spectrum below shows such a series of measurements using various mixtures of A and B, in which the amount of A was kept constant while varying the amount of B:

The use of the ppm measuring scale and the odd fact that this scale increases to the left, which is the exact opposite of how it normally works, is due to the measuring technique. Please do not let this bother you. The spectrum shows very clearly how the signal of the highlighted hydrogen atom travels further to the left, the higher the percentage of molecule B. The percentage of B is shown on the far right. Well, anyone who is used to reading these kinds of spectra will see right away that it is very clear. When looking at such a series of measurements for the first time, it may not be instantly noticeable. I will therefore explain the above figure in a bit more detail:

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Shown at the very bottom are the signals that are measured when analysing the first molecule A alone, indicated by an arrow in the figure below:

Shown above that are the signals that are produced when adding 20% (“0.2 eq.” = 0.2 equivalents, i.e. 20%) of B. On the far right, highlighted in grey in the figure below, a few weak signals are starting to appear; these are the signals of the new molecule B – as could be expected:

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In addition, a new signal has appeared on the far left. It is the shifted signal of the highlighted hydrogen atom shown in the figure at the very beginning of this chapter:

By the way, if you look very carefully, it can also be spotted in the first spectrum on the far right in the first “forest” of signals at 8.0 ppm:

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So, what happened? By adding 20% of molecule B, the signal of the highlighted hydrogen atom has shifted from the original 8 to about 8.7. This is a strong indication that the presence of molecule B does indeed change the “chemical environment” of this hydrogen atom, meaning a hydrogen bond involving this hydrogen atom has formed. As you can see, the other signals do not change. This is therefore a strong indication that these hydrogen atoms are not involved in the hydrogen bond that has formed. This information is of course also important to determine the exact location of the hydrogen bond. The signals above that are produced when adding 40% of the second molecule B, indicated by an arrow in the figure below. The signal on the far left has been shifted even further to the left.

Above that is the signal at 60% B – with the signal being even further to the left and so on. When adding a surplus of the second molecule B, indicated by “2 eq.” for twofold surplus up to “10 eq.” for tenfold surplus, nearly every single molecule A is now bound by hydrogen bonds. However, the other signals do not change. This is not surprising, because these hydrogen atoms do not form hydrogen bonds. But as the concentration increases, the signals of the second molecule B are gaining the “upper hand”. That is also why the signal of the highlighted hydrogen atom is dropping slightly, as can be seen clearly in the next figure:

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Hence, the shifting of the signal can be used to determine the presence of a hydrogen bond and to identify which hydrogen atoms in the individual molecules are involved in it. Another advantage of this method is that it can determine not only the presence of a bond but also its strength by evaluating the individual measurements more closely. In addition, such a measurement does not take very long, only about one day, and requires only a small amount of substance. So it is no wonder that virtually all chemists worldwide use this method to examine hydrogen bonds. You may be wondering: Why is a tenfold surplus of the second molecule B necessary to bind all of the first (A)? Shouldn’t it be enough to just add the same amount? It depends on the strength of the bond and the fact that what we are dealing with here is called equilibrium. I will go into this very briefly – and then I will also tell you how to take advantage of such equilibria in chemistry. When analysing the concentration of the free first molecules of A, the concentration of the free second molecules of B and the concentration of the bound molecules – also called a duplex – in our system, it becomes apparent that, regardless of the conditions, the following equation applies: A*B = K *AB In this equation, A stands for the concentration of the free molecules of A, B for the concentration of the free molecules of B and AB for the concentration of the duplex. K is a constant. This constant K indicates the strength of the bond between the molecules. It is extremely dependent on the specific conditions, such as temperature and solvent. When trying to bind as many molecules of A as possible which in the equations means to make A as small as possible, it is helpful to rearrange the equation a little bit: A=

K * AB B

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You can see that A automatically becomes smaller as the concentration of B increases. When trying to bind as many molecules of A as possible, it is therefore necessary to add as much of the bonding partner B as is reasonably possible. That is exactly what was done in the above measurement. You can see that it is very reasonable to add a tenfold surplus of B to really bind almost all molecules of A. Thinking in terms of equilibria takes some getting used to, because it somewhat contradicts everyday experience – but a lot of things can be done with it. To illustrate this point, I would like to take a look at another very simple (but completely hypothetical) equilibrium. Let us assume there are 10,000 people living in a small city, 9,900 of whom are Schalke 04 supporters and the remaining 100 are Borussia Dortmund supporters.10 The relationship between them would be a chemical equilibrium, producing the following equation: Schalke = 99 × Dortmund This means there is a ratio of one Dortmund supporter to 99 Schalke supporters, two Dortmund supporters to 198 Schalke supporters and so on. Let us suppose that despite this difficult starting situation, Borussia Dortmund was interested in converting all residents of that town into Dortmund supporters. How could Borussia go about it? In real life, it would be difficult, of course, because a Schalke 04 supporter would never betray her or his club and defect to the other side, no matter how well Borussia may be playing. But in chemistry, there is a surprisingly simply solution: Borussia Dortmund would convince the Dortmund supports to move away from that city! So, what would happen then? Since the equilibrium always (!) applies, a few Schalke supporters would now automatically become Dortmund supporters. When all Dortmund supporters have moved away, there would only be 9,900 people left in that city, but since the equilibrium continues to apply, 9,801 people would now support Schalke and 99 Dortmund, meaning 99 of the original Schalke supporters have now changed their minds. The above-described relationship between Dortmund and Schalke supporters also applies the other way round: There automatically must be a ratio of one Dortmund supporter to 99 Schalke supporters, so in a population of 9,900 people, there must be 99 Dortmund supporters, regardless of whether they were previously Schalke supporters or not. Now, if Borussia Dortmund were to carry out another “moving-away” campaign, there would be 9,702 people left in the city. According to the equilibrium, 9,605 of them would be Schalke supporters and 97 Dortmund supporters.

10 For those readers who are not familiar with Schalke 04 and Borussia Dortmund: They are German football clubs that share a healthy rivalry. Alternatively, you could think of Manchester United and Manchester City.

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As you can see, a respectable 100 + 99 + 97 = 296 of all (original) residents of the city have converted to being Dortmund supporters, meaning the percentage of people who go into every season dreaming of celebrating a championship in Borsigplatz, the birthplace of Borussia Dortmund, has now almost tripled. If Borussia Dortmund were to keep going this way, the club’s management could succeed in turning all residents of the city into Dortmund supporters, even though 99% of the mood is always set against them. However, there would ultimately be no residents left in the city. Well, you can’t have everything. What does that mean for chemistry? Let us suppose the equilibrium between Schalke and Dortmund would be like that between the starting product and the end product in a chemical reaction, where we would, of course, only be interested in the end product (no offense, Schalke supporters). Even with a very unfavourable starting situation where only 1% of our desired product exists, we can still carry out 100% of the reaction – if we can manage to successfully remove the end product from the reaction (and ensure that the starting product does not react otherwise, which is sometimes a problem). Once this has happened, the starting product reacts with an automatic response, forming more of the end product. Many chemical reactions can only be carried out for this single reason and would otherwise not work at all.

10 “Now let’s step on the accelerator” . . . says the boy called “the professor” in the classic “Emil and the Detectives” by Erich Kästner to the other boys when they talk about recovering Emil’s money and supporting him in fighting evil Mr Grundeis. In other words: Let’s get going – quickly! In chemistry, there are comparably helpful accelerators, which are useful when it comes to setting things in motion at an accelerated rate. However, these are usually called catalysts. Those catalysts that are likely to be known to (almost) all readers are catalytic converters, that is, porous metallic structures used in cars. I do not mean those; the function of these catalytic converters is really difficult to explain and is not dealt with in this book, since they do not form hydrogen bonds. The catalyst I would like to address in this chapter looks like this:

It has been found that it can accomplish truly amazing things. However, we should start with some background information. In Chapters 4–7, I showed you that methane and other carbon compounds have the so-called tetrahedral structure. This has unexpected consequences. Look at the two-molecule models below:

What exactly A, B, C and D stand for does not play a role here; it only matters that they are all different. Each of the two molecules has one A, one B and so on. Only A and B have been switched. A is highlighted to make this more visible. But are the two molecules identical? You can try to rotate them until they are congruent. I tell you right away: It doesn’t work. These two molecules are almost identical, but not quite. They are like an image and its reflexion or like your left and right hand. They are also almost identical, but just almost. https://doi.org/10.1515/9783110628012-010

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This fact is called chirality and the corresponding molecules are referred to as chiral, from Greek chiros = hand. It simply results from the mathematical properties of a tetrahedron and was recognised (independently of each other) by Jacobus van’t Hoff and Joseph le Bel as early as 1874. Since the two aforementioned molecules differ only in the spatial arrangement of A, B, C and D, they have the same name. To yet distinguish between the two chiral forms, a system was developed in which one form is designated as the R form and the other as the S form. There are clear rules to determine which form is R and which is S. For historical reasons, there is unfortunately another system in which the two forms are designated with D and L, with different rules for distinguishing between them. Chirality would not be that important if it did not have unexpected impacts in real life. An example of that is carvone, which looks like this:

Carvone

(D) or (S) carvone

(L) or (R) carvone

On the left, carvone is shown in the way molecules have so far been depicted; next to it, you can see the two chiral forms of carvone in a sort of three-dimensional (3D) view. The D or S carvone is shown in the middle, and the L or R carvone on the right. The two chiral forms differ only in the arrangement of the group below the ring, which points to the left in the first case (in D and S carvone) and to the right in the second case (in L and R carvone). D carvone smells of caraway and is also the substance that accounts for the primary smell of caraway on your spice rack. By contrast, L carvone smells of mint. This minor difference is sufficient for the olfactory sensors in your nose to associate the two molecules with different smells, although they otherwise have a completely identical structure. “Well, it’s no big deal when substances smell different”, you might say. But the whole thing can have even more far-reaching consequences. I would like to demonstrate this by comparing two types of sugar (please do not bother too much about the structural complexity, only the difference matters!):

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Just like carvone above, the two sugars are shown in a 3D view. The sugar on the left is grape sugar (also called glucose), and the one on the right is galactose, a subunit of lactose, or milk sugar, which is contained in milk, as the name implies. The two sugars differ only in the arrangement of a single alcohol group, which is highlighted. The alcohol group points “downwards” in glucose and “upwards” in galactose.11 What consequences does this difference have, which appears to be minor at first glance? Very significant ones – perhaps even for you personally. Glucose can be eaten by all people – however, there are many people who do not tolerate lactose. This is also referred to as lactose intolerance. About 75% of all people worldwide are affected by it, among Chinese even 95% and among Thais 98%! Only in Central and Northern Europe, the United States and Russia can the majority of people eat lactose, that is, tolerate larger amounts of milk. Lactose intolerance is even associated with the settlement history of Africa and other regions. The reason for this is that in certain regions there are people who do tolerate lactose well or at all, although other people in surrounding areas commonly do not. It is assumed that the ancestors of these first people had an advantage due to their genetic make-up, as they were able to drink the milk of their riding animals, such as camels, while travelling through deserts or other inhospitable areas and thus populated regions that were inaccessible to other people who did not have this option. This is already a huge effect of chirality – given that glucose and galactose differ only in the arrangement of a single group. In many medicines, too, only one form is effective (i.e., R or S, D or L); the other is inactive or sometimes even toxic. You see, chirality is indeed important. The question of what exact structure a molecule has, that is, what the exact chiral form looks like and, of course, how it can be produced in a targeted manner has kept chemistry busy since the discovery by van’t Hoff and Le Bel, and many Nobel Prizes, among them the very first in chemistry to van’t Hoff in 1901, have been awarded for discoveries and synthesis methods in this field. To make life easier, chemists have also introduced a shorthand notation for chiral molecules. A bold line indicates that this residue projects forwards, a dashed line

11 “Upwards” was deliberately put in quotation marks, because the real spatial arrangement is yet a bit different. This description was chosen for the sake of clarity.

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that this residue projects backwards and a solid line that this residue projects neither forwards nor backwards, but lies in the plane of the paper. This looks like this:

Chiral molecule

The same molecule in the exact shorthand notation

The catalyst I described at the very beginning of this chapter is called proline. It has the advantage that it is commonly found in nature and is inexpensive, which is an undeniable benefit in chemistry. More specifically, this catalyst is also chiral – and the form that is mostly used in practice is the so-called S proline or L proline, which has the following structure12:

The highlighted part on the right, which is, by the way, a carboxylic acid – you might remember it – projects forwards. At this position, a hydrogen atom projects backwards as usual, which was left out in the above figure for the sake of simplicity. When adding it, the catalyst looks like this:

This catalyst accelerates a whole range of reactions, including the following reaction:

12 Of course, R or D proline could theoretically be used as well, but it is not that common in nature and is therefore more expensive.

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This reaction or this type of reaction in general is called “aldol reaction“ and was described as early as 1872 by Russian chemist Alexander Borodin (who was also a famous classical composer of what is referred to as the Russian School). You may have never heard of this reaction, but every chemistry student who has ever visited a lecture in organic chemistry is familiar with it. After all, this reaction is one of the most important reactions in this field of chemistry. To give you an example, the synthesis of the medicine atorvastatin, which is marketed by the company Pfizer under the names “Sortis” and “Lipitor” and used to be the best-selling medicine in the world for some time, involves two aldol reactions. Atorvastatin has the following structure:

Atorvastatin

The two alcohol functional groups highlighted in grey are produced in two successive aldol reactions. What makes the aldol reaction interesting is that the two reactants A and B are not chiral, but the product C very well is, namely, at the highlighted position:

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This fact and the significance of the aldol reaction as such have challenged many chemists to find a method that not only delivers the product but also delivers only one chiral form of it straightaway. This does not happen automatically. When the reaction is carried out “just like that”, it produces equal amounts of the product in S and R form. But if the exact chiral form is important, half of the product is basically wasted. If you want one of the chiral forms to be produced in excess, the reaction needs to be adapted, for example, by using an additional aid. Around 2000, Benjamin List (back then at the Skaggs Institute in San Diego, today at the Max Planck Institute for Carbon Research in Mülheim an der Ruhr) found out that L proline was such an aid. When adding L proline to the above reaction, more than 90% of the reaction product will be produced in one form, namely, the following, in which the resulting alcohol projects forwards:

(But only in one form!)

This really was a great success, as most of the previously available methods required reagents that were much larger, more complex and, most importantly, expensive. By contrast, one kilogram of L proline costs well below 100 euros, even in high-purity form, which is a true discount price. How exactly does this reaction work? This has been extensively investigated and the following mechanism is assumed to be most likely: Proline first reacts with molecule A above to form the following molecule and water:

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As you can see, the nitrogen atom here forms four bonds – the scientific correct term would be that the nitrogen is “tetravalent” – and is additionally positively charged. This is also why this molecule is not very stable and immediately reacts again in such a way that the H atom highlighted in the figure below is donated as a so-called proton, forming a double bond. As a result, the nitrogen atom again forms three bonds (or to say it correctly, it becomes “trivalent”) and everything is alright again. As usual, however, the nitrogen atom still has an electron pair, which will be important later on.

This resulting molecule, which I will call “D”, has three important properties. The first one is that it reacts with molecule B from the original equation, which is this one:

faster than molecule A. This property is the most important one, because otherwise A and B would also react “just like that” – and nothing would be gained. The second property, resulting from proline now being “integrated”, is that it is chiral, but in such a way, and this is the third property, that exactly this chiral site, as it is a carboxylic acid, can form a hydrogen bond with molecule B. When taking a closer look at the structure in 3D view, you will notice that the carboxylic acid group in molecule D on one side of the double bond is quite precisely located above it.

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This entails that in the following reaction molecule B approaches “from below” – first, because “from above” the carboxylic acid group is simply in the way and secondly, because a hydrogen bond forms between the oxygen atom of B and the hydrogen atom in the carboxylic acid group of D, which “holds” molecule B “in place”. This looks something like this:

The two molecules then react with each other to form a new molecule in a spatially precisely defined way, since the arrangement of the carboxylic acid group and the hydrogen bond spatially controls the reaction. The next reaction has three steps. I will explain them one after the other, but in reality, and this is important, they all take place simultaneously. In the first step, one bond of the carbon–oxygen double bond in molecule B “folds over” towards the hydrogen atom of molecule D. The carboxylic acid group, in turn, detaches from the hydrogen atom and thereby becomes negatively charged:

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In addition, one double bond of molecule D “folds over” towards molecule B13:

This happens at the same time as the bond “folds over” towards the hydrogen atom; consequently, the resulting alcohol projects forwards and the hydrogen atom, which is located on the same carbon atom, backwards. For the sake of simplicity, this is no longer shown in the figures below. The free electron pair of the nitrogen atom also “folds over” downwards:

13 The proline group in the following figures is shown slightly tilted to make the resulting alcohol more visible. In a realistic arrangement, it would be covered by the carboxylic acid group.

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The result is the following molecule shown again below, in which – this is important and therefore I would like to repeat it – the newly created alcohol group then projects forwards:

The nitrogen atom in this molecule is now again “tetravalent” and positively charged. At the same time, the carboxylic acid group is negatively charged, because the hydrogen atom was donated to the resulting alcohol group. As a result, it instantly reacts with the water created at the beginning to form the following two molecules:

These are exactly the desired reaction product C and proline! This means that the proline is fully recovered during the reaction and can “step on the accelerator” in a second reaction, or, in chemical terms, act as a catalyst. It does not even have to be available in the same amount as the two other molecules. This is not the case in practice either. It is usually used in a percentage of about 10–20%. Of course, it is essential for the reaction that proline with as high a chiral purity as possible is used, since proline virtually transfers its chirality to the product in the course of the reaction. However, this is not a problem, since high-purity proline is cheaply available. So, the “trick” is to use the cheap availability of chiral proline for other reactions as well to easily synthesise new chiral compounds that cannot be simply obtained from nature or other sources in chiral form. Unfortunately, as it is often the case in life, such tricks work extremely rarely, and this is why the above reaction attracted so much attention.

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As already described in the first chapter, proline is also referred to as an organic molecule. Findings like this have led to the development of an entire field of chemistry called “organocatalysis”, that is, the use of organic molecules as catalysts. These organocatalysts have by now produced spectacular results, which have until recently been thought of as hardly possible. One aspect of this success story is particularly interesting: Benjamin List was by far not the first to recognise the potential of proline as a catalyst. As early as 1971, two working groups separately published results regarding the use of proline as a catalyst for chiral syntheses in a very similar type of reaction. Hence, it would have been possible to recognise much earlier that proline was suitable for other reactions as well. There are now more than 10 different types of reactions that are catalysed by proline. Why did it take nearly 30 years – which is ages in modern chemistry – for the potential of proline and organocatalysis to be discovered in the first place? We can only speculate on that. The most obvious question is, of course: Why was this field of research not established by the original discoverers? There is an answer to this question: Both original publications came from industrial chemists, one from researchers of Schering and the other from HoffmannLa Roche. They were actually aiming for something entirely different, namely, the synthesis of pharmaceutically active ingredients. When they recognised that an important intermediate for these active ingredients was much easier to produce using proline, they published this finding to subsequently continue working on other aspects of active ingredient synthesis, which was exactly what their employers paid them for. So the original discoverers were not necessarily expected to further investigate this field. But other chemists could have picked up the ball and continued playing – why did they not? Malicious gossip has it that there may be two major reasons for that: First, publications by industrial chemists are not always taken quite seriously by academic scientists. On top of that, one of these two publications is not even an academic paper, but a patent application – academics tend to look down on those. Secondly, both publications were published in German. Unfortunately, however, many academics, especially in the English-speaking world, think that foreign language skills are unnecessary – “learning a different language is as flamboyant as wearing a crown in a bus” is regrettably the attitude of many English and Americans. Benjamin List is now a serious contender for the Nobel Prize. If he were to get it, he might drink a toast to all the chemists before him who were too arrogant to take German patent applications seriously and therefore did not beat him to it. But like I said, it is all just speculation. Other discoveries that were later on recognised as revolutionary also had precursors that were not followed up on. Research does not always move in a straight line. There are often discoveries where people ask themselves later on why they were not discovered much earlier by

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someone else. In any event, proline has by now become one of the most commonly used catalysts in modern chemistry. Now I would like to introduce you to another catalyst system that can be regarded as more of a model system – but demonstrates very clearly how a catalyst works. It involves the following reaction, in which a molecule, which I will again call A, practically reacts with itself as follows:

Product

You will probably think that the reaction looks somewhat chaotic and complex, so I will explain it in a bit more detail. This is a so-called Diels–Alder reaction, named after the discoverers Otto Diels and Kurt Alder, who were awarded a Nobel Prize for the discovery of this type of reaction in 1950. In this type of reaction, two parts of A, namely, the two highlighted areas, react with each other to form a ring:

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You can imagine that each of the three double bonds “folds over”, creating as many as two new bonds, which are circled, and one double bond moves “one position forward”:

By this reaction, a much more complex structure is obtained and that is why Diels– Alder reactions became so popular in chemistry. Andrew Hamilton (back then at the University of Pittsburgh, and now at Oxford University) and his working group then investigated what happens when carrying out this reaction, but adding either the molecule shown below, which is marked with B, or the molecule C.

As you can see, the only difference between B and C is the number of carbon atoms located on the central ring between the two “arms”, namely, two in B and only one in C – or, in other words, the angle is 180° in B and 120° in C.

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Now the following effect becomes apparent: With molecule C, the reaction takes place 3.5 times faster than “just like that” – with molecule B, however, 10 times slower. This means that this minor difference in the molecular structure has a tremendous effect on the reaction! How can this be explained? Actually, it is quite simple. Molecule A above forms two times two hydrogen bonds with both molecule B and C – similar to those shown in Chapter 9. With molecule B, this looks like this:

For the reaction to take place, however, the following highlighted areas of molecule A have to be located in spatial proximity to each other, as already mentioned above:

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Owing to the fact that molecule A forms hydrogen bonds with molecule B via its two carboxylic acid groups, it becomes slightly stretched and is kept in this position. However, this makes a reaction practically impossible! The two highlighted areas are simply too far apart from each other. For this reason, only the molecules that are currently unbound will enter into a reaction. But as these molecules are in the minority, it takes much more time than in the absence of molecule B until they all react. With molecule C, it is exactly the other way round. When A forms hydrogen bonds with this molecule, it looks like this:

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In this case, the hydrogen bonds push together the reactive areas of molecule A. It is thus logical that the reaction is accelerated in the presence of molecule C, making C a true catalyst! I would like to briefly get back to proline. Proline is a so-called amino acid. This means that proline has an amine group, basically meaning that a nitrogen atom is present, as well as a (carboxylic) acid group. When many of these amino acids are linked together via the so-called amide groups (which you already know from Chapter 7), compounds with the following structure are created:

The figure above shows six amino acids, differing in the residues R. In this case, proline was not among them, but it is commonly found in nature. When linking together 100 or more of these amino acids – it can be up to 30,000 (!) – giant molecules referred to as “proteins” are created. You may have heard of them, for example, in connection with meat that contains particularly high amounts of protein compared to pasta or rice. However, proteins are much more important, as the name implies. This term was used as early as 1839 by Dutch biochemist Gerardus Mulder, after having taken up a suggestion from Swede Jöns Jakob Berzelius, one of the greatest chemists of all times. He, in turn, derived the term protein from Greek – from “protos” meaning the first/most important and “proteios” meaning fundamental. And this is indeed true of proteins! Proteins have essential functions in the human body. For instance, they are catalysts, the best ones known in chemistry and biology. Proteins with a catalytic effect are also referred to as enzymes. There are also other proteins, which I do not want to address in greater detail at this point. Enzymes make sure that molecules in the human body are converted into other molecules. Enzymes are so effective that they manage to accelerate reactions not only a hundred-fold, which would already be pretty good, but often even a million-fold. Without enzymes, we would not survive a single day. How do enzymes do that? Basically, they often act similar to proline. They temporarily bind to a type of molecules and make sure by forming hydrogen bonds that other molecules react with them, the reactions being precisely controlled in terms of space and time. Or they act like molecule C in Andrew Hamilton’s system and make sure by forming hydrogen bonds that reactive areas of molecules get closer to each other and thus react faster. Since proteins are much larger than proline, they can virtually tailor themselves to specific reactions and form not only one but

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several hydrogen bonds. Since proteins are so much larger than proline, I cannot show that to you in a figure, because either I would need as much space as half a dining table or the figure would be illegibly tiny. So you will have to try imagining it or just believe me. For example, a protein called streptavidin forms as many as 15 hydrogen bonds with its target molecule called biotin. No wonder biotin and this protein are “as thick as thieves”.14 By the way, if you would like to see enzymes, not just described in a book, but as a substance in real life, you do not need to visit a chemical laboratory. Just look at your detergent! The detergent industry is one of the major “users” of enzymes, which are contained in every modern detergent in significant amounts. But to tell it right away, enzymes are basically a white powder, whose excellent properties are not apparent at the first glance. The reason why enzymes are contained in detergents is easier to understand when you briefly go back to Chapter 6 and once again have a look at the structure of a fat. As discussed fats can be split up into glycerol and fatty acids using substances like soda lye, which is referred to as saponification. Glycerol and the individual fatty acids can be removed more easily from laundry, like a pair of jeans, than the fat itself, because they are much smaller and are also better soluble in water. You could now be tempted to wash a stained pair of jeans with soda lye to saponify the fat. But this would be a bad idea. You may have no fat anymore – but no jeans either, because even highly diluted soda lye causes quite large holes in jeans (as many chemistry students – including me – had to learn after a day of practical training in the laboratory). No wonder soda lye is contained in many drain cleaners. When adding an appropriate enzyme to the detergent, however, it will also “split” the fat without destroying the jeans. For other types of soiling, such as albumin or red wine stains, there are also special enzymes that decompose the corresponding substances. The detergent industry has taken advantage of these properties of enzymes, ensuring that detergents have indeed improved significantly. This progress has led to an unintentionally funny situation: Enzymes have a maximum temperature at which they work best. Above this temperature, their structure changes and they no longer work properly.15 The optimum temperature usually ranges between 30 and 40 °C, which is the body temperature of most mammals. However, many users of washing machines are used to boil-wash laundry to get it really clean, which is quite reasonable when using detergents that do not contain enzymes. Since enzymes are produced industrially using genetic engineering

14 Experts do not quite agree as to the exact number, because it can be disputed which hydrogen bonds point to biotin and which do not. A distinction is made between something like a “first circle” and “second circle”. In any case, it is at least eight, which is still quite a few. 15 You have certainly experienced this phenomenon in your laundry room, but with wool. Wool also consists almost exclusively of proteins – and if woollen jumpers are washed too hot, the structure of these proteins also changes irrevocably: the jumper “shrinks”.

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methods, which did not become commonly available until the mid-1980s, this is not that long ago. When using detergents that contain enzymes, which nowadays even the cheapest “no-name products” do, this has an adverse effect. At 95 °C (i.e., boil wash), normal enzymes are destroyed within a few minutes – so the effect of the detergent does not improve, but deteriorates significantly! So, the next time you use your washing machine, you had better set the temperature to 40 °C (or what is suggested on the package of the detergent). Your laundry will thank you.

11 Substances that build themselves In the previous chapter, I discussed about catalysts, that is, substances that accelerate reactions like good Samaritans. Of course, it would now be interesting to look into reactions where a catalyst helps itself, meaning that it catalyses a reaction in which it forms as a product. There is a whole range of such reactions. They are referred to as “autocatalytic” (from Greek: auto = self) reactions. These reactions automatically become faster over time, as more and more catalyst is produced. It would be even more interesting if this reaction were not only autocatalytic, but were to also involve “transmission of information”, meaning that the reaction product is more complex than the starting substances. Reactions of this type are not called autocatalytic, but self-replicating. What information can be transmitted? In the simplest case, the directional information in terms of molecular recognition, as already described in Chapter 7. The reason why such reactions are extraordinarily interesting is that they – at least in the opinion of all experts in this field – must have existed when life began on the Earth. It is largely unclear how life originated on the Earth. Something like the following is widely agreed: 4.2 billion years ago, some kind of a primeval ocean, also referred to as “primeval soup”, existed on the Earth. At the beginning, the primeval ocean contained a whole smorgasbord of chemicals dissolved in it. But, basically, that was it. There was no life. Around 400 million years later, that is, 3.8 billion years ago, the first cells and bacteria emerged, something you can certainly call life. However, even this point in time is disputed; some researchers assume that life already existed as early as 4 billion years ago. What happened in between? That’s the big question! Somehow, because otherwise you and I would not even exist, a process or a whole range of processes must have taken place within the primeval soup on the Earth’s surface, which caused more complex molecules and then regular structures to gradually come into being, culminating in something like cells and bacteria. Nobody knows exactly what processes were involved. But, of course, we want to know it, because the evolution of life out of simple chemicals, virtually out of “nothing”, is something incredibly fascinating. The only thing known with certainty in this entire field of research is that it has to work somehow, as evidenced by your and my existence. Apart from that, practically everything is subject to controversy, as you may have already noticed.

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When taking a look at a cell, you will find that it is already a highly complex system. Something like that does not just “fall from the sky”. The probability of a cell developing spontaneously out of a primeval ocean is immeasurably small. One thing life is characterised by is the ability to reproduce, that is, to copy or to replicate, even though not always completely identically. After all, your children do not look the same as you either – but this is another story. It is assumed that before there were cells capable of replicating or copying, there must have been much simpler reaction systems in which molecules copied and replicated – and that before there was an evolution among living organisms, something like a “chemical evolution” had taken place. These reaction systems are called selfreplicating. Of course, it is very difficult for scientists investigating this subject to fully reconstruct the processes that led to the evolution of life on the Earth. One reason for this is that the Earth is already populated and organisms such as bacteria thus intervene in possible processes all the time. We cannot just turn back time and watch what happens (unfortunately). On the other hand, many of these processes are most likely to have lasted very long, perhaps centuries, if not millennia. We simply do not have so much time. Therefore, scientists use simulations and model systems, for example, model systems of self-replicating reactions. A schematic model of such a self-replicating reaction system would look something like this:

The two molecules A and B react with each other to form the product C. However, A and B each also have recognition sites, enabling them to mutually form hydrogen bonds, similar to adenine and thymine in DNA or the molecules described in Chapters 7 and 9. Of course, the product C also has both recognition sites, allowing A and B to attach to product C. As a result, a complex of A, B and C forms, which is marked with ABC on the top left of the above figure. A and B then react to form C. As a result, a complex of two C molecules forms, which is marked with C2 on the

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top right of the above figure. This reaction is accelerated by the formation of the complex ABC, simply because as a result of this complex formation, A and B are located in close proximity to each other, similar to the catalyst of Andrew Hamilton in the previous chapter. C is thus also a catalyst – catalysing its own formation! The complex C2 now breaks up, allowing molecule C to attach to A and B again. However, there are now two molecules C, which is why two molecules C can each attach to a molecule A and a molecule B two times. As a result, four molecules C form, then eight, then 16 and so on, resulting in what is called exponential growth. Things get interesting if there are two replicators, that is, one from A, B and C – and another from, let’s say, A, B’ and C’. A and B react to form C, but likewise A and B’ react to form C’. Hence, the two compete for molecule A. Let us assume we fill B and C as well as B’ und C’ into a vessel, but only add as much A that the two replicators have to compete for A. When done correctly, one of these replicators should prevail (at least this is how it has been demonstrated in theory), because it would be faster than the other and would take away molecule A from the other. This would then represent some kind of a chemical evolution. At least since the 1980s, researchers have extensively looked for self-replicating reactions, and in 1986 Günter von Kiedrowski from the University of Göttingen presented the first reaction system of this kind. It is actually chemically very similar to DNA. A chemical “sister” of DNA, RNA, which has a very similar chemical structure, is capable of not only self-replicating but also catalysing many reactions. Sidney Altman was awarded the Nobel Prize in 1989 for this finding. In addition, it was found that the primeval soup most likely contained molecules that could serve as starting substances for RNA synthesis. Numerous models of the evolution of life on the Earth now assume some kind of an “RNA world”, that is, a primeval ocean dominated by RNA. Following synthesis of the first DNA replicator, it was attempted to build not only further replicators on the basis of DNA or RNA (although these may indeed have once existed in nature), but also systems that are based on other reactions and chemical structures. The reason for this is, first, that there is an interest in understanding the fundamental mode of action of replicators and, secondly, that “synthetic” systems can be, as is hoped, better “tuned” to achieve desired effects. By the way, the first synthetic replicator of this kind, which has nothing to do with DNA anymore, was also presented by Günter von Kiedrowski, together with his degree candidate Andreas Terfort in 1992. I would like to introduce you to a more recent synthetic replicator, which was synthesised in 2000 by Maik Kindermann in the laboratory of Günter von Kiedrowski, now professor at the Ruhr University of Bochum. I chose this replicator for two reasons: First, because it is not very large and therefore easier to explain and, secondly, because self-replication is not that simple.

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When taking another look at the above scheme, you will see where a major snag is. If the two molecules C fit together so well, why would they separate? But if the complex C2 does not break up, then C cannot act as a catalyst – it is virtually rendered inactive. When Günter von Kiedrowski tested his first replicator, he found that only part of the products C is actually active, another part is not. As he had demonstrated in theoretical experiments, however, a real chemical evolution requires replicators that work better. As big a sensation the first replicator might have been, it is unfortunately not suitable to bring about a chemical evolution. The replicator developed by Maik Kindermann comes very close to such an ideal replicator and this is what makes it so interesting. It is shown in the figure below:

As in the previous chapter, the actual reaction here is again a Diels–Alder reaction. For a reaction to take place, the following highlighted areas have to be located in close spatial proximity to each other:

Just like in the previous chapter, three double bonds “fold over” in this reaction:

When transferring this reaction to the general scheme shown above, the picture looks like this:

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As you can see, it is just two hydrogen bonds each that bind molecule A and B to the product C – but that is enough! When taking a look at the structure of the complex ABC on the top left of the above figure, you will see the essential thing: Similar to how it works in Andrew Hamilton’s system from the previous chapter, product C again makes sure that the reactive areas of A and B are located in close proximity to each other so that the reaction between them can take place at an accelerated rate. This reaction has also been examined using the 1H-NMR method, and the course of the reaction can be clearly seen in the spectrum below:

This image shows a fragment of the entire spectrum, because there are naturally many more hydrogen atoms in the respective molecules, but it clearly shows the course of the reaction. The signals for product C are shown on the left and the signals of molecule A on the right. At the beginning, only very little C forms – but the reaction then really gets going! The fact that the system is indeed self-replicating can be recognised by the rate at which product C forms. But it is also possible – as has been tested – to simply modify molecule B, for example, to the effect that it does not form any hydrogen bonds with A and the new product C, like this:

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This reaction proceeds at a much slower rate than that of the replicator, because the molecules do not form any bonds. The situation is similar when leaving molecule B as it is, but modifying molecule A. Hence, it is the hydrogen bonds that matter. Why is this replicator so close to the ideal growth? Of course, this has been tested as well, and one of the reasons is that complex C2 is a bit “crooked”, meaning that although two molecules C find each other and form hydrogen bonds, they do not fit together perfectly:

As you can see in the above figure, the two recognition sites, marked with the grey bars, do not fit together accurately, but form an angle. Therefore, the complex C2 breaks up relatively easily. Molecules A and B are much smaller and more mobile and can thus attach to molecule C more easily – and this is exactly what is intended. When taking another look at the molecules from Chapter 9, where I described how hydrogen bonds can be measured, you will notice that these are quite similar to the two molecules A and B. However, they do not react with each other and can therefore be used as model substances to obtain certain data needed to properly examine the replication system. Is a molecular evolution known today? Unfortunately not. The big difficulty is that replication systems are very hard to find and further conditions have to be met for molecular evolution to take place. So far, nobody has figured out how to fit everything together. Even if the replicator described above is modified only slightly, self-replication does not take place anymore, sometimes not even a reaction. So, there is still enough to do for researchers who want to know how life originated on the Earth.

12 Like Avalokiteśvara and Durga In this chapter, I address an aspect in greater detail that I have previously “withheld”, namely, the fact that most of the reactions described in earlier chapters do not take place in a vacuum, but in a solution. This is only logical, given that the molecules need to meet somehow, and this works best when they are dissolved in a liquid, however, not in water. The reactions, for example those described in Chapter 11, do not take place in water and the hydrogen bonds described in Chapters 7 and 9 cannot be measured either. Why is that? First, this is because some of the described molecules are not readily soluble in water, but this is not the main reason. The main reason is that water forms hydrogen bonds itself, preventing the molecules to find each other. And since water is the solvent, there are naturally much more water molecules than other molecules present. When applied to the Opera Ball analogy, this would mean that a male dancer is trying to find a dance partner, but the men outnumber the women hundred to one. Obviously, this makes dancing quite difficult. This is why solvents that do not form hydrogen bonds such as chloroform are mostly used to examine molecular recognition. But how can molecules find each other in water as well (or in other solvents forming hydrogen bonds)? What can our dancer do? If he is a normal gentleman, there is relatively little he can do. The situation is different if he happens to be like the bodhisattva Avalokiteśvara, a kind of “spiritual being” in Buddhism. Avalokiteśvara has 11 heads and, most importantly, a thousand arms! This makes it a lot easier to catch a lady. If that lady then happens to be like the Hindu Goddess Durga, who can have up to 20 arms, depending on the situation, things get even easier. Exactly this phenomenon is found in DNA, for example. As explained earlier, DNA recognises itself in water. However, even DNA requires a length of at least 10–20 bases and thus more than 20 hydrogen bonds – the exact number depends on the concentration and other conditions – to actually form a double helix structure. Otherwise, the DNA helix would be broken up by the water molecules or does not form in the first place. Giving their molecules many “arms” is exactly what several researchers and, most notably, their working groups have attempted to do. I am going to address the results of their work in this chapter. However, first I would like to answer the question why it is interesting to examine molecular recognition in water in the first place. The answer is simply because nearly all reactions in the human body also take place in water. When intending to use molecular recognition systems, for example, to develop new medicines or diagnostic agents, these have to be effective in water as well. Therefore, many researchers have attempted to develop corresponding systems that also work in water or similar solvents. https://doi.org/10.1515/9783110628012-012

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One of these researchers is Carsten Schmuck from the University of DuisburgEssen. For instance, he examined the following molecule,16 which he presented in 1999:

When adding this molecule to water, it actually finds itself and forms hydrogen bonds with itself. Of course, the main reason for this is that there are as many as six of it at a very short distance like this:

Another, much larger molecule with the same “basic architecture” was presented by Prof. Schmuck in 2005 and referred to as a “molecular flytrap”. It looks like this:

16 This molecule is both negatively charged on one of the oxygen atoms and positively charged on one of the nitrogen atoms, which is additionally “tetravalent”. Yet – for reasons I would like to just skip for now – it is, unlike the molecules formed in the proline-catalysed reaction in Chapter 10, very stable.

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As you can see, it has a “central unit” and three “arms” – and binds the subsequent molecule so strongly that it will (almost) never let go of it, even in water. You will at least know it by name, because it is citric acid:

Citric acid

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How does that work? The figure below shows the bond between one of the three carboxylic acid groups of citric acid, which is present as the so-called citrate, meaning that it is negatively charged, and the “flytrap”:

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However, not only one of the carboxylic acid groups is bound to the “flytrap”, but also all three, with the three arms “folding over” towards the citric acid. Schematically, you can imagine this as follows:

This threefold “clutch” ensures the incredibly strong attractive force, with the citric acid being practically “trapped”. You may already recognise what such molecules can be used for, namely, for detecting substances or “fishing” them out of solutions. Some time earlier, in the 1990s, another researcher, George Whitesides from Harvard University, published a strategy of arranging molecules so as to create veritable molecular architectures with networks of hydrogen bonds. He took the following two molecules as a basis:

When combining these two molecules, a kind of a two-dimensional network is created with a pattern that is reminiscent of an Oriental carpet:

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This network makes sure that these two molecules together form a solid – and do not dissolve in virtually any liquid, rendering chemical reactions impossible. Prof. Whitesides then slightly modified the two molecules by inserting some carbon chains at relevant positions, simply marked with “R” as usual, and replacing a nitrogen atom by a carbon atom as follows:

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The result is a network that still looks impressive, but is slightly smaller and thus also soluble, making it suitable for chemical reactions:

Prof. Whitesides refers to this structure as a “rosette”, and presented a whole range of such rosettes. Jean-Marie Lehn17 from the College de France in Paris developed a very similar system in 1996. However, he used only one instead of two different molecules, namely, the following one:

17 He is one of the Nobel Prize winners who were distinguished for their research into crown ethers, see Chapter 6.

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It also forms a kind of hexagon with the following structure:

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Steven C. Zimmerman from the University of Illinois chose yet another approach. The starting molecules of his system, which he also published in 1996 together with his team, are these two:

The architecture of the molecules, the left one of which is straight and the right one slightly bent, is not random. When combining these two molecules, they recognise each other by the following hydrogen bonds, as shown on an example complex of three molecules:

Of course, other molecules can attach on the left and right, creating a kind of “molecular screw”, which looks something like this:

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To conclude this chapter, I once again discuss two topics that were already addressed in the previous chapters. The first topic is DNA. As you remember, the information in DNA is transmitted via the structure of the hydrogen bonds, where only adenine matches with thymine and guanine with cytosine. Unfortunately, DNA in the body is sometimes attacked, resulting in a change in the molecular structure. This can have disastrous consequences for the copying mechanism, which is, in turn, one of the main causes of cancer. For example, if two thymines succeed each other in DNA, they may react with each other in the presence of UV light as follows:

This often entails that these thymines are no longer recognised during a later copying process, entailing that a “bad copy” is produced at this position. Such bad copies are one of the main causes of skin cancer. Therefore, a UV protection cream should be used when the skin is exposed to sunlight for extended periods of time, as sunlight contains UV radiation. Such a protection cream absorbs most of this UV component, thereby largely preventing this reaction.

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Another important change is the reaction of guanine with environmental toxins, or the like to form a compound named 8-oxoguanine, which looks like this:

Guanine

8-Oxoguanine

The area where the guanine has changed is highlighted. 8-Oxoguanine has the unpleasant property that it forms hydrogen bonds not only with cytosine but also with adenine, given as follows:

As a disastrous consequence, it is no longer ensured that a cytosine is inserted during a DNA copying process. Likewise, the new DNA may contain an adenine at this position, which is a mistake. Unfortunately, 8-oxoguanine and guanine differ only insignificantly, which makes it relatively difficult to detect 8-oxoguanine in DNA. In 2011, Prof. Dr Shigeki Sasaki from Kyushu University in Fukuoka and his team developed a “detector” for 8-oxoguanine, which I would like to introduce to you. It looks like this:

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As you can see, the molecule consists of two parts, namely, an adenine (bottom, the molecule highlighted in light grey), to which a second molecule (for those interested: diazaphenoxazine, the molecule highlighted in dark grey) is attached. This detector can be inserted into a DNA strand to detect 8-oxoguanine. To this end, a “detector DNA” is added to a test DNA that is suspected to contain 8-oxoguanine. This detector DNA is synthesised in such a way that, if the assumption

8-Oxoguanine

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is correct, the 8-oxoguanine is located opposite this detector molecule in the DNA double strand. The detector molecule now forms a very stable complex with the 8-oxoguanine, in which the detector molecule “wraps around” the 8-oxoguanine, forming as many as five hydrogen bonds: The presence of 8-oxoguanine can thus be verified on the basis of the hydrogen bonds. This is already a good thing, but this detector can do even more. The detector molecule has the pleasant property that it emits an intense blue-green light when irradiated with the appropriate light. But this only applies as long as the molecule is “free”. When the above complex of 8-oxoguanine and the detector molecule forms, the electronic states of the detector molecule are slightly changed by the hydrogen bonds, which entails that the molecule no longer emits any light. Prof. Sasaki and his team were able to demonstrate that this “blackout” actually only takes place if the detector molecule (in the “detector DNA”) is located opposite an 8-oxoguanine (in the “test DNA”) in a DNA double strand. With any other base opposite to it, the detector molecule continues to emit blue-green light. This makes it possible to optically verify that hydrogen bonds have actually been formed between 8-oxoguanine and the detector. Why is this so important? Simply because such optical methods allow for detecting molecules in tiny concentrations. Other detection methods, such as 1H-NMR described in Chapter 9, are more difficult and also require larger amounts of substance. Finally, I would get back to a molecule that I described earlier in Chapter 7, namely, the system developed by E.W. “Bert” Meijer. You might remember what it looked like:

It already has quite a few “arms”. Prof. Meijer and his working group have somewhat modified, or rather extended this molecule, meaning that they slightly altered the residue “R”, which simply consisted of a carbon chain in the molecule from Chapter 7:

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As you can see, the old molecule has more or less been doubled. However, a closer look reveals that this new molecule cannot bond to itself congruently, because the recognition sites do not quite fit together:

First, the molecule would have to rotate around its own axis at the centre; second, the two hydrogen atoms are always located exactly at the centre and the two other binding sites on the outside. Of course, it is intended to be this way! Now how does this molecule behave? It nevertheless recognises itself, because the bonding with as many as four hydrogen bonds is extremely strong. However, this results in an arrangement where the second molecule is slightly displaced:

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This means that only four of the eight possible binding sites (four on the left and four on the right) are used, which is why binding sites are left over to which other molecules can attach. In the end, a long chain is created:

Measurements have shown that this chain can consist of more than one hundred molecules. The chain becomes so long that a structure with properties that are macroscopically similar to plastics is created. The kind of plastics that plastic bags are made of also consists of molecules forming very long chains, as already explained in Chapter 5. However, there is a significant difference between the chain of Prof. Meijer and the plastics used in plastic bags, which is called polyethylene in chemistry. In polyethylene, all bonds in the chain are covalent and are therefore difficult to break. In Prof. Meijer’s chain, only some of the bonds are covalent, namely, those within the chain molecule. The bonds between the molecules are hydrogen bonds. Under the appropriate conditions, these are almost as strong as a proper covalent bond, but when the conditions are altered, the hydrogen bonds break apart and the chain collapses. This means that you can determine whether or not a plastic structure is created merely by selecting the external conditions. This plastic material is highly interesting. Another group of scientists, namely, those headed by Christoph Weder from the University of Fribourg, went a step further and extended this system by synthesising a molecule that has not only two but three recognition sites. It looks like this:

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It is again the highlighted recognition sites that are important – the rest of the molecule is less interesting and was only selected because it is easy to synthesise. This molecule now forms a network with itself, which looks something like this:

The interesting thing about it is that the network structure of this molecule causes a sort of glass to form, that is, a solid you can look through. It is also relatively stiff and can be picked up, for example, with pincers, without bending. This fact alone would not be that spectacular, given that glass already exists. Moreover, this molecule is not quite as stable as solid glass; for example, it can be cut with a razor blade. But it can do something normal glass cannot. When cutting real glass, for example, with a glass cutter, the resulting cuts cannot be simply smoothened out again. However, this can easily be done with Christoph Weder’s molecule. For instance, when you cut through the “Weder glass” with a razor blade, the cut can be removed again within a few hours by heating up the molecule or exposing it to UV light (which is also just heating, but more targeted) and letting it cool down again. The “Weder glass” then becomes just as smooth as before! Why is that? The reason is that the hydrogen bonds break when exposed to heat. When the material is cooled down, they are recreated, ensuring that the structure is “as good as new” again:

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Hence, the “Weder glass” is self-healing: a truly amazing property owed to hydrogen bonds.

13 Conclusion and acknowledgements At this point, I thank you for having made it this far – after all, I have guided you, as promised, all the way from very simple molecules through DNA and enzymes to current research. This book is devoted only to one subject, namely, hydrogen bonding, but I hope that you still did not get bored and that you have learnt what an important role hydrogen bonds play in your life, even if you had not been quite aware of that before. In addition, I thank people who supported me in writing this book. First of all, I thank Dr Maik Kindermann for allowing me to use parts of his doctoral thesis is this book, including some figures. He also did the proofreading of the original German version of this book. Dr Malte Reimold, who possesses a perfect combination of scientific understanding and graphic design skills, made sure that my raw templates became images and graphs that are nice to look at. Max Düren, Dr Stefan Höppner, Tianqiao Pan and, most importantly, Ute Hüttermann provided me with valuable suggestions on how to make this book easier to understand, especially for “non-scientists”, after having reviewed and corrected the first German version of the this book. I thank Mrs Berber-Nerlinger from the Oldenbourg Publishing House for the copy-editing and her support in creating the first German version book. I thank everyone at DeGruyter, namely, Ria Fritz, Maureen Pagel and Lena Stoll first for suggesting the English version of this book and the support with the realization. Prof. Dr Shigeki Sasaki and Prof. Dr Carsten Schmuck provided me with research data and results as well as manuscripts of yet unpublished works. Finally, I thank my wife, E. Hyun, for indulging me when I spent quite a few weekends in an effort to further improve this book – as well as for reviewing this book and giving me suggestions for improvement.

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Bibliography Below I have compiled a bibliography in case you would like to learn more about certain subjects addressed in this book. The majority of the basic information provided in Chapters 1 to 5 is taken from the “HollemannWiberg”: Hollemann, Wiberg, A Textbook of Inorganic Chemistry, 102nd edition, Berlin 2007. For Chapters 6 to 12, I used my own textbook as well as the “Voet–Voet”: Bräse, Bülle, Hüttermann, Organische und Bioorganische Chemie [Organic and Bioorganic Chemistry], 2nd edition, Weinheim 2008. Voet, Voet, Pratt, Beck-Sickinger, Hahn, Lehrbuch der Biochemie [A Textbook of Biochemistry], 2nd edition, Weinheim 2010.

Chapter 1 If you are interested to learn more about the 112 chemical elements, I recommend that you read: Ulf von Rauchhaupt, Die Ordnung der Stoffe: Ein Streifzug durch die Welt der chemischen Elemente [The Order of Substances: A Ramble through the World of Chemical Elements], 2nd ed., Frankfurt, 2009 – it is however, so far only published in German. “A cheerleader waves a cyanide wand, there’s a smell of peach blossom and bitter almond” in Genesis, “Fly on a windshield/Broadway Melody of 1974”, from: The lamb lies down on Broadway 1974. If the song is to be believed, Howard Hughes (indirectly) “produces” potassium cyanide himself, because that substance is contained in cigarette smoke.

Chapter 6 Fight Club, directed by: David Fincher, USA 1999, based on the novel by Chuck Palahniuk published in 1996 For crown ethers, see the Nobel lecture of Charles Pedersen Charles Pedersen, The Discovery of Crown Ethers (Nobel Lecture), Angew. Chem. 1988, 100, 1053–1059. For ion channels, see the Nobel lecture of Roderick MacKinnon Roderick MacKinnon, Potassium Channels and the Atomic Basis of Selective Ion Conduction (Nobel Lecture), Angew. Chem. 2004, 116, 4363–4376.

Chapter 7 The molecule of “Bert” Meijer is published in F. H. Beijer, R. P. Sijbesma, H. Kooijman, A. L. Spek and E. W. Meijer, Strong Dimerization of Ureidopyrimidones via Quadruple Hydrogen Bonding, J. Am. Chem. Soc. 1998, 120, 6761–6769.

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Bibliography

Chapter 8 The story of how DNA was discovered can be read in J. D. Watson, The Double Helix. A Personal Account of the Discovery of the Structure of DNA, Hamburg, 1997 (reprint). J. D. Watson, A. Berry, DNA. The Secret of Life, New York 2003. Jerry Donohue’s biography can be read on Wikipedia at http://en.wikipedia.org/wiki/Jerry_Donohue The original publication by Watson & Crick: J. D. Watson, F. H. C. Crick, A Structure for Deoxyribose Nucleic Acid, Nature 1953, 171, 737–738. See also http://www.sns.ias.edu/~tlusty/courses/landmark/WatsonCrick1953.pdf The (incorrect) structure according to Linus Pauling is published in: L. Pauling, R. B. Corey, A Proposed Structure for the Nucleic Acids, PNAS 1953, 39 (2), 84–97. Concerning the accuracy of the DNA copying mechanism, I recommend (if you understand German) that you read the following article by Prof. Dr. Gottfried Schatz, published in the Neue Zürcher Zeitung on 17 February 2011: http://www.nzz.ch/nachrichten/kultur/aktuell/schoepfer_zufall_1. 9562044.html The G-quadruplex is described, among other publications, in M. Gellert, M. N. Lipsett, D. R. Davies, Helix Formation by Guanylic Acid, PNAS, 1962, 48, 2013–2018.

Chapter 9 The described measurement is taken from the doctoral thesis of Dr. Maik Kindermann: Maik Kindermann, Kleine organische Replikationssysteme und kristalline Filme durch AmidiniumCarboxylat-Wechselwirkungen an der Luft-Wasser-Grenzschicht [Small Organic Replication Systems and Crystalline Films through Amidinium-Carboxylate Interactions at the Air-Water Interface], Doctoral Thesis, Bochum, 2001.

Chapter 10 Just for the sake of completeness: Erich Kästner, Emil and the Detectives, Overlook Press, 2014 (first published in 1929). If you are interested to learn how automotive catalytic converters and similar catalysts of this type work, I recommend that you read the Nobel lecture of Gerhard Ertl: Gerhard Ertl, Reactions at Surfaces: From Atoms to Complexity (Nobel Lecture), Angew. Chem. 2008, 120, 3478–3590. Wikipedia has some articles worth reading, such as Lactose intolerance: https://en.wikipedia.org/wiki/Lactose_intolerance Aldol reaction: https://en.wikipedia.org/wiki/Aldol_reaction Organocatalysis: https://en.wikipedia.org/wiki/Organocatalysis Proline:https://en.wikipedia.org/wiki/Proline For carvone, see G. F. Russel, J. I. Hills, Odor Differences between Enantiomeric Isomers, Science, 1971, 172, 1043–1044.

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The proline-catalysed aldol reaction was first described in B. List, R. A. Lerner, C. F. Barbas, III., Proline-Catalyzed Direct Asymmetric Aldol Reactions, J. Am. Chem. Soc. 2000, 122, 2395. The mechanism is published in Linh Hoang, K. N. Houk, S. Bahmanyar, B. List, Kinetic and Stereochemical Evidence for the Involvement of Only One Proline Molecule in the Transition States of Proline-Catalyzed Intraand Intermolecular Aldol Reactions, J. Am. Chem. Soc. 2003, 125, 16–17. The “precursor” reactions from 1971 are published in U. Eder, G. Sauer, R. Wiechert, New Type of Asymmetric Cyclization to Optically Active Steroid CD Partial Structures, Angew. Chem. 1971, 10, 492–493. Z. G. Hajos, D. R. Parrish: Asymmetric Synthesis of Optically Active Polycyclic Organic Compounds, German patent DE 2102623 (patent filed on 20 January 1971). For the Diels-Alder reaction used by A. Hamilton, see S. C. Hirst, A. Hamilton, Complexation Control of Pericyclic Reactions: Supramolecular Effects on the Intramolecular Diels-Alder Reaction, J. Am. Chem. Soc. 1991, 113, 382–383. For streptavidin-biotin, see A. W. Hendrickson, A. Pähler, J. Smith, Y. Satow, E. A. Merritt, R. P. Phizackerley, Crystal Structure of Core Streptavidin Determined from Multiwavelength Anomalous Diffraction of Synchrotron Radiation, PNAS, 1989, 86, 2190–2194.

Chapter 11 A good introduction to the “origins of life” can be found, for example, in the doctoral thesis of Dr. Kindermann (see above) E. Szathmáry and J. Maynard Smith, The Major Transitions in Evolution. Oxford University Press, New York 1995. The first DNA-based replicator is published in G. von Kiedrowski, A Self-Replicating Hexadeoxynucleotide, Angew. Chem. 1986, 86, 932–934. The first fully synthetic replicator is published in A. Terfort, G. von Kiedrowski, Self-Replication by Condensation of 3-Aminobenzamidines and 2-Formylphenoxyacetic Acids, Angew. Chem. 1992, 104, 626–628. The replicator theory is published in G. von Kiedrowski, Minimal Replicator Theory I: Parabolic Versus Exponential Growth, Bioorganic Chemistry Frontiers 1993, 3, 113–146. For RNA as a catalyst, see the Nobel lecture of Sidney Altman: Sidney Altman, Enzymatic Cleavage of RNA by RNA (Nobel Lecture), Angew. Chem. 1990, 102, 735–744. The “Kindermann replicator” is published in the doctoral thesis of Dr. Kindermann (see above) as well as in M. Kindermann, I. Stahl, M. Reimold, W.M. Pankau, G. von Kiedrowski, Systems Chemistry: Kinetic and Computational Analysis of a Nearly Exponential Organic Replicator, Angew. Chem. 2005, 117, 6908–6913.

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Bibliography

Chapter 12 To learn more about Carsten Schmuck’s work, see C. Schmuck, Highly Stable Self-Association of 5-(Guanidiniocarbonyl)-1H-Pyrrole-2-Carboxylate in DMSO – The Importance of Electrostatic Interactions, Eur. J. Org. Chem. 1999, 2397–2403. C. Schmuck, M. Schwegmann, A Molecular Flytrap for the Selective Binding of Citrate and Other Tris-carboxylates in Water, J. Am. Chem. Soc. 2005, 127, 3373–3379. To learn more about George Whitesides’ work, see G. M. Whitesides, E. E. Simanek, J. P. Mathias, C. T. Seto, D. N. Chin, M. Mammen, D. M. Gordon, Noncovalent Synthesis: Using Physical-Organic Chemistry to Make Aggregates. Acc. Chem. Res. 1995, 28, 37. To learn more about Jean-Marie Lehn’s work, see A. Marsh, M. Silvestri, J.-M. Lehn, Self-Complementary Hydrogen Bonding Heterocycles Designed for the Enforced Self-Assembly into Supramolecular Macrocycles, J. Chem Soc. Chem. Commun. 1996, 13, 1527–1528 To learn more about Steven Zimmerman’s work, see P. M. Petersen, W. Wu, E. E. Fenlon, S. Kim, S. C. Zimmerman, Synthesis of Heterocycles Containing Two Cytosine or Two Guanine Base-Pairing Sites. Novel Tectons for Self-Assembly, Bioorg. Med. Chem. 1996, 4, 1107. To learn more about Shigeki Sasaki’s work, see Y. Taniguchi, R. Kawaguchi, and S. Sasaki, Adenosine-1,3-Diazaphenoxazine Derivative for Selective Base Pair Formation with 8-Oxo-2′-Deoxyguanosine in DNA, J. Am. Chem. Soc. 2011, 133, 7272–7275. To learn more about “Bert” Meijer’s work, see R. P. Sijbesma, F. H. Beijer, L. Brunsveld, B. J. B. Folmer, K. J. H. K. Hirschberg, R. F. M. Lange, J. K. L. Lowe and E. W. Meijer, Reversible Polymers Formed from Self-Complementary Monomers Using Quadruple Hydrogen Bonding, Science 1997, 278, 1601–1604. To learn more about Christoph Weder’s work, see Diederik W. R. Balkenende Christophe A., Monnier Gina L., Fiore, and Christoph Weder, Optically responsive supramolecular polymer glasses; Nature Communications 2016, 7, 10995.

Index 1

H-NMR 61 – in the Kindermann replicator 94 8-oxoguanine 107 Acetic acid 45 – Hydrogen bonds 46 – Molecular recognition 46 – Three-dimensional structure 45 Adenine 52 – Hydrogen bonding with thymine 59 Alcohol 27 Alder, Kurt 82 Aldol reaction 75 – Chirality 75 Allyl alcohol 29 Americium 2 Amino acids 86 Ammonia 16 – Boiling point 20 – Hydrogen bonds 20 – Three-dimensional structure 23 Antibiotics 43 Argon 2 Arrhenius, Svante 34 Atorvastatin 75 – Synthesis using aldol reactions 75 Autocatalysis 89 – Self-replication 89 Benzene 29 Benzene rings 29 Berkelium 2 Berzelius, Jöns Jakob 86 Biotin 87 Borodin, Alexander 75 Borussia Dortmund 69 Butanol 32 Calcium 9 Calcium oxide 9 – Structure 18 Californium 2 Carbon 2 – as the basic element of organic chemistry 27 – in carbon dioxide 9 Carbon dioxide 9

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Carvone 72 – Chirality 72 Catalysts 71 – Autocatalysis 89 – in aldol reactions 76 – in Diels-Alder reactions 82 – Proline 74 – Proteins 86 – Self-replication 89 Cells 40 – Lipid bilayer 40 Chargaff, Erwin 54 Chemical equilibria 68 Chemical evolution 91 Chirality 72 – Lactose intolerance 73 – Nomenclature 72 Chlorine 4 – Chloride 5 – Chlorine gas 8 – in sodium chloride 4 – Structure of the chlorine molecule 8 Chlorine gas 8 Cholesterol 32 – in the lipid bilayer 40 Citric acid 99 – Molecular recognition 99 Covalent bond 7 Crick, Francis 54 Crown ethers 40 – Comparison with ion channels 42 – Comparison with valinomycin 42 – Sodium 40 – Toxicity 43 Curium 2 Cyanide ion 10 Cytosine 52 – Hydrogen bonding with guanine 59 Darmstadtium 2 Detergents 87 – Enzymes 87 Dichlorine monoxide 16 Diels, Otto 82 Diels-Alder reaction 82 – in the Kindermann replicator 92

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Index

DNA 51 – 8-oxoguanine 107 – Arrangement of hydrogen bonds 57 – Backbone 52 – Bases 52 – Copying mechanism 58 – G-tetrad 60 – Hoogsteen base pairing 60 – Molecular recognition 57 – Reaction with UV light 106 – Self-recognition in water 97 – Story of how the structure of DNA was discovered 54 – Structure 51 Donohue, Jerry 56 Einsteinium 2 Electron pairs 8 Electronegativity 11 Elements 1 Enzymes 87 – Proteins 87 – Streptavidin 87 Equilibria 68, 69 – in chemistry 68 Erbium 1 Ethanol 27 – Three-dimensional structure 30 Fats 35 – Saponification 35 – Structure 35 Fatty acids 36 – Liposome 39 – Micelles 37 – Soaps 37 – Structure 36 Francium 1 Franklin, Rosalind 54 Gabriel, Peter 10 Galactose 73 Gallium 1 Germanium 1 Glucose 73 Glycerol 36 G-quadruplex See G-tetrad Grape sugar See Glucose G-tetrad 60

Guanine 52 – Correct structure 56 – G-tetrad 60 – Hydrogen bonding with cytosine 59 – Structure according to Davidson 56 Hamilton, Andrew 83 Hassium 2 Helium 2 Hoogsteen base pairing 60 Hoogsteen, Karst 60 Hydrides 7 Hydrogen 1, 2 – Origin of name 1 – Structure of the hydrogen molecule 7 Hydrogen bonding 13 – Comparison with the Vienna Opera Ball 14 – Definition 13 – Detection by 1H-NMR 61 – in ammonia 20 – in proline-catalysed aldol reactions 78 – in self-replication 94 – in the Kindermann replicator 94 – Role in the boiling point of water 18 – Role in the density anomaly of water 25 Hydrogen chloride 16 Hydrogen cyanide 10 Hydrogen sulphide 16 Ice 24 – Structure 24 Inorganic chemistry 2 Ion channels 42 Ionic bond 4 Isotonic saline 40 Kiedrowski, Günter von 91 Kindermann, Maik 91 Kindermann replicator 91 Krypton 2 Lactose intolerance 73 Le Bel, Joseph 72 Lehn, Jean-Marie 103 Lipid bilayer 40 Lipitor See Atorvastatin Liposomes 39 – Structure 39 List, Benjamin 76

Index

Meijer, E.W. “Bert” 48 – “Switchable chain” 109 – Ureidopyrimidones 48 Metallic bond 8 Methane 15 – Boiling point 19 – Natural gas 15 – Tetrahedral structure 22 Micelles 37 – as soaps 37 Molecular flytrap 98 Molecular recognition 48 – “Molecular flytrap” 99 – in DNA 57 – in self-replication 94 – in the Kindermann replicator 94 – Ureidopyrimidones 48 Mulder, Gerardus 86 Naphthalene 30 Neon 2 Nitrogen 2 – Structure of the nitrogen molecule 8 Noble gases 2 Oganesson 2 Organic chemistry 2 – Organocatalysis 81 Organocatalysis 81 Oxygen 1 – Boiling point 16 – in calcium oxide 9 – in carbon dioxide 9 – Origin of name 1 Pauling, Linus 54 Periodic table of elements 3 Phosphine 16 Polonium 1 Potassium 10 – Crown ethers 41 – in potassium cyanide 10 – Ion channels 42 Potassium cyanide 10 – Cyanide ion 10 – Hydrogen cyanide 10 Primeval ocean 89 Primeval soup 89 Proline 74

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– as amino acid 86 – Catalyst in aldol reactions 76 Propanol 28 Propargyl alcohol 29 Proteins 86 – Amino acids 86 – as catalysts 86 – Enzymes 87 – Streptavidin 87 Proton nuclear magnetic resonance spectroscopy See 1H-NMR Roentgenium 2 Saponification 35 Sasaki,Shigeki 107 Schalke 04 69 Schmuck, Carsten 98 Self-replication 89 – Chemical evolution 91 – Origins of life on Earth 89 Silver 1 Soap 37 Sodium 4 – Crown ethers 40 – in sodium chloride 4 – Ion channels 42 – Sodium ions 6 Sodium chloride 4 – Dissociation 34 – Ionic bond 4 – Solubility in water 33 – Structure 17 – Structure in aqueous solution 33 Sortis See Atorvastatin States of matter 17 – Gases 17 – Liquids 18 – Solids 17 Streptavidin 87 Sucrose 33 Sugar 32 – Chirality 73 – Solubility in water 33 Terbium 1 Terfort, Andreas 91 Tetrahedron 22 – Chirality 71

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Index

– in ethanol 30 Thymine 52 – Correct structure 56 – Hydrogen bonding with adenine 59 – Structure according to Davidson 56 Urea 2 Ureidopyrimidones 48 – Glass 111 – Polymer 109 Valinomycin 42 van-der-Waals forces 19 Van’t Hoff, Jacobus 72 vis vitalis 2 Water 12 – as solvent 27 – Boiling point 15 – Density anomaly 21 – Ethanol 31 – Melting point 15

– Three-diemensional structure 23 Watson, James 54 Weder, Christoph 111 Whitesides, George 101 Wilkins, Maurice 54 Wöhler, Friederich 2 Xenon 2 Ytterbium 1 Yttrium 1 Zimmerman, Steven C. 105 δ− 12 – in ammonia 19 – in water 13 – Role in hydrogen bonding 13 δ+ 12 – in ammonia 19 – in water 13 – Role in hydrogen bonding 13