Essay On Chemistry And Biochemistry

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The elements of biochemistry

Biochemical systems carry out an enormous variety of chemical reactions with great efficiency. These reactions can be catabolic, breaking down larger molecules to release energy and to generate precursors for further reactions, and/or anabolic, combining molecules together to generate biologically useful molecules. When a chemist carries out a reaction in the laboratory, they have numerous different techniques that can be applied to increase the yield or rate of a reaction; they can alter the temperature or pressure or perhaps add a catalyst. In contrast, biological systems have to carry out the reaction at the temperature maintained by the organism, and, with the exception of organisms living deep in the ocean, at atmospheric pressure. In biological systems, enzymes are employed to increase the rate of reaction; enzymes are proteins whose substrate-binding site acts to lower the energy of high energy species along the reaction pathway from starting material to product. They can achieve enormous enhancement in the rate of reaction compared with the uncatalysed reaction. The classic example is the enzyme triose phosphate isomerase which interconverts dihydroxyacetone phosphate and d-glyceraldehyde-3-phosphate during the breakdown of glucose. This reaction occurs 109 times faster in the presence of the enzyme than when uncatalysed [1].

These rate enhancements are particularly remarkable when we consider that less than a third of the naturally occurring elements are used by biological systems. In order to be exploited in a biological system, elements must be sufficiently abundant in a form that can be taken up by living things. Thus, many catalytic species that are in common use in the laboratory are simply not available for biochemical reactions, for example palladium, used in the addition of hydrogen across a double bond. Of the elements in the periodic table, 28 are essential for animal life (Figure 1); the most recent element found to be essential is bromine, which was found to be required for the proper formation of networks by the protein collagen IV in 2014 [2]. Of the 28 essential elements, 11 make up 99.9% of the atoms in the human body. The other 17 are known as trace elements and are present in very small amounts, ranging from a milligram to gram quantities in an adult human.

Figure 1Periodic table illustrating the elements essential for life

The 28 elements essential for animal life are indicated by coloured squares; trace elements are shown in yellow and those present in larger quantities are shown in green. The six most abundant elements in the human body are carbon, hydrogen, oxygen, nitrogen, phosphorus and calcium, accounting for almost 99% of the mass of an adult human. Carbon, hydrogen, oxygen and nitrogen are the building blocks of organic biomolecules, calcium is present in large amounts in bones and teeth (in addition to being vital for cell signalling in smaller amounts), phosphorus is likewise found in bones and teeth (smaller quantities are a vital part of DNA, adenoside triphosphate – the energy currency of the cell – and play an important role in cell signalling).

Atomic structure and bonding

Electrons in atoms are organized into a series of shells with successively higher energies (and greater distance from the nucleus). The shells are identified by the principal quantum number which takes integer values from one to seven for the elements of the current periodic table. The shell with principle quantum number 1 has the lowest energy with 2 being the next highest in energy, and so on. Within a shell there are subshells designated by the letters s, p, d and f, and within each subshell electrons occupy orbitals: regions of space that may be occupied by up to two electrons, and whose energy and shape can be described mathematically by an equation known as the wave function. It is important not to confuse orbitals with an orbit; electrons do not move around the nucleus along a fixed path as they would in an orbit. Instead the wave function allows us to calculate the probability of finding an electron at a particular position around the nucleus. Different subshells have different numbers of orbitals; s subshells have just one orbital and can therefore accommodate only two electrons, p subshells have three orbitals which can be filled by six electrons in total, and d subshells have five orbitals accommodating up to ten electrons. Shells are filled according to the Aufbau principle, i.e. filling the shell with the lowest principal quantum number first.

Arrangements of electrons or electron configurations, in which the outer shell (the occupied shell with highest principle quantum number) is full are more favourable than those in which it is partially filled. Of the elements in the periodic table, only the noble gases in group 18 have full outer shells. Bond formation, the movement of electrons between atoms, allows other elements to achieve this configuration. Chemical bonds can be covalent, where electrons are shared between atoms or ionic, where electrons are transferred from one atom to another resulting in one positively and one negatively charged species. Looking at the number of electrons in the outer or valence shell enables us to work out how many bonds an atom would need to form in order to fill its outer shell. It is important to note, however, that a bond will only actually form if the energy of the electrons in the bond is lower than the energy of those electrons in the isolated atoms.

As the majority of biological chemistry relates to covalently bonded molecules composed primarily of the elements carbon, hydrogen, nitrogen and oxygen, it is particularly important to know how many bonds each of these elements form. Hydrogen, with its single valence electron requires one additional electron to achieve the noble gas configuration and therefore makes only one bond. H+ ions (often just called protons), in which the hydrogen atom has lost an electron, also play many roles in biological systems – for example the formation of ATP is driven by a concentration gradient of H+ ions across a membrane. Carbon, with four valence electrons, achieves a full outer shell by forming four covalent bonds, for example by sharing an electron with each of four other atoms. Nitrogen has five valence electrons and forms three covalent bonds leaving one pair of non-bonded electrons; a lone pair. The lone pair is important for the reactivity of nitrogen as it can be used to make a new bond with electron-poor species in chemical reactions. The lone pair can also be shared with an H+ ion leading to the formation of ammonium, i.e. the lone pair allows nitrogen to act as a base. Finally, oxygen with six electrons makes two covalent bonds and has two lone pairs of electrons. As with nitrogen, the presence of lone pairs on oxygen makes this atom reactive towards electron-poor species, including H+. Although we often refer to H+ ions as though they exist in that form, in water free H+ forms the hydronium ion H3O+ very rapidly.

When an ionic bond is formed, electrons are transferred completely from one atom to the other. The interaction in an ionic bond is entirely Coulombic in nature (i.e. only due to the force of attraction between positive and negative particles) and there is no directionality to the interaction. Such a bond occurs when the elements differ widely in their ability to attract electrons or more formally, when there is a very large energy difference between the valence orbitals (the outermost orbitals containing the electrons available to participate in bonding) in the two atoms. Covalent bonding involves sharing of electrons between atoms and occurs when the two atoms are more similar in their ability to attract electrons; i.e. when the valence orbitals have similar energy. Sharing of electrons in a covalent bond requires atomic orbitals on each of the atoms to interact with each other. One of the consequences of this is that, in contrast with ionic bonds, covalent bonds are directional: a new bond cannot form in a region of space where the orbitals that interact to form the new bond have no electron density. Two classes of covalent bonds occur commonly in biological chemistry, σ bonds and π bonds (Figure 2A,B). Although a pair of electrons are shared when a covalent bond is formed, however, this sharing is not necessarily equal. Atoms with a strong ability to attract electrons, i.e. low energy valence orbitals, are referred to as electronegative. If the two atoms forming a covalent bond differ in electronegativity then there will be greater electron density closer to the more electronegative atom. This results in a permanent dipole, with one atom partially negatively charged and the other partially positively charged (Figure 2C).

Figure 2σ and π bonds

In these images, atomic orbitals and bonds are depicted as line drawings indicating shape and as isosurfaces, regions of space enclosing a defined fraction of the electron density. (A) When two s-orbitals overlap, a σ-bond is formed. In σ bonds, the atomic orbitals overlap ‘head on’ and there is electron density between the two nuclei in the bond. As a result, these bonds are typically strong compared with π bonds between the same elements. (B) In a π bond, two p-orbitals overlap laterally, and the nuclei lie in a plane in which there is no electron density. There is less overlapping between orbitals in a π bond than in a σ bond and so these bonds are typically weaker. It should be noted that a double bond, consisting of a σ and π bond is stronger than a single bond between the same two elements as the strength of both the σ and π components are included in the bond strength. (C) When a bond is formed between carbon and oxygen, there will be more electron density located near the oxygen atom, as illustrated here in the isosurface for a carbon–oxygen π bond.

Delocalization of electrons

The description of covalent bonding has so far assumed that a covalent bond involves the sharing of one pair of electrons between two atoms, i.e. that the bond is localized. For the majority of biological molecules this description is adequate, however in some cases this description of bonding does not explain the observed properties of a molecule.

A well-known example is benzene. It was originally thought that benzene contained three alternating single and double bonds, however measurements showed that the bonds were all of equal length. We now consider the bonding in benzene not as three pairs of p-orbitals each interacting to make one double bond, but six p-orbitals each interacting with its neighbours to create a ring of electron density above and below the plane of the carbon atoms (Figure 3A). Compounds with delocalized rings of electrons are of major importance in biological systems. For example, the bases in DNA all contain delocalized rings as a part of their structures and favourable interactions between the electrons in these rings (referred to as π–π stacking) contribute to the stability of the DNA double helix (Figure 3B).

Figure 3Delocalized structures

(A) Each of the six carbon atoms in benzene contribute a p-orbital to the delocalized system. A circle in the centre of the ring can be used to highlight the fact that the system is delocalized, however many biochemists prefer to use the alternating bond representation. (B) The four bases found in DNA, all have delocalized rings of p-orbitals; in this diagram atoms shown in red each contribute a p-orbital to the delocalized system. It is not immediately obvious that the atoms indicated with an arrow can contribute a p-orbital, however using a more sophisticated approach to bonding we can show that this is the case. Adenine and guanine each have ten electrons in a delocalized ring, while cytosine and thymine have six. (C) Representation of the delocalized carboxylate anion. In this system, four electrons two from the double bond and two from the negatively charged oxygen are delocalized over three atoms. (D) Retinal has a linear delocalized system including 12 p-orbitals. Each of the atoms that contributes a p-orbital to the delocalized system is shown in red.

Delocalization of electrons does not only occur in rings; another type of system where delocalization occurs is where three (or more) parallel p-orbitals are adjacent. Consider the carboxylate anion (discussed in The carbonyl functional group section) where a carbon atom makes a double bond with one oxygen atom and a single bond with a negatively charged oxygen atom. In this structure, we can visualize the negative charge on the oxygen being used to make a new double bond with the carbon atom and the existing double bond breaking to leave a negative charge on the oxygen atom (Figure 3C). Although we can visualize single and double bonds exchanging position within the molecule, this is not an accurate description of bonding in the molecule. In reality, the electron density is spread over three p-orbitals, and a higher electron density exists on the two oxygen atoms than on the central carbon atom. Delocalization of electron density across three p-orbitals is also important in explaining why the bond formed between two amino acids in a protein chain is planar (see Functional groups found in amino acids section). Molecules with electrons are delocalized over a larger number of adjacent parallel p-orbitals are also common in biology. These molecules are usually referred to as conjugated and can be identified by their alternating chain of single and double bonds. For example, retinal, the light-absorbing molecule that is bound to the protein opsin in the photoreceptor cells responsible for vision in mammals, has electron density delocalized across 12 p-orbitals (Figure 3D). The long delocalized system is essential for the absorption of light in the visible region of the electromagnetic spectrum.

Non-covalent interactions

Non-covalent interactions, such as the π–π stacking mentioned in the above section, arise due to electrostatic interactions between two different molecules or within the same molecule between atoms that are not bonded together, without the sharing of electrons via a covalent bond. These interactions are much weaker than the covalent bond but they occur very frequently and, as a result, can have a huge influence on the properties of a molecule. Many biomolecules are macromolecules with thousands of atoms and therefore make many hundreds of thousands of non-covalent interactions. Non-covalent interactions are particularly important in proteins. Proteins are polymers of amino acids, synthesized in a linear chain on the ribosome. Each protein chain folds into a specific 3D structure that is essential for its function; non-covalent interactions between the constituent amino acids determine the 3D structure. Non-covalent interactions are also important in DNA where they help to ensure that the sequence of DNA is preserved upon replication; in the lipid bilayer where non-covalent interactions between lipids create a barrier around the cell; and in molecular recognition (discussed in more detail in [3]). There are several classes of non-covalent interactions; here we discuss van der Waals interactions, hydrogen bonds and briefly, ionic interactions. We also discuss the ‘hydrophobic effect’ which is commonly invoked to explain why non-polar molecules do not disperse in water and why proteins fold.

van der Waals interactions

A dipole is an uneven distribution of electron density within a molecule such that one region of the molecule has a higher electron density than the other and the two regions are equally but oppositely charged. van der Waals interactions occur when a dipole on one molecule interacts with the dipole on another molecule. These dipoles can be permanent or instantaneous. Permanent dipoles occur due to the uneven charge distribution in a covalent bond between two elements that differ greatly in electronegativity. The partially positively charged (δ+) atom on one molecule can interact favourably with the partially negatively charged (δ) atom on another. Interactions between instantaneous dipoles are called London dispersion forces. They are the weakest among the non-covalent interactions, but also the most prevalent. London dispersion forces occur because the electron density in an atom or molecule does not have an even distribution; at any one time the electron density may be higher in one region than the other. The electron density is redistributed with time, thus the regions of high electron density are different from one moment to the next. The uneven charge distribution is called instantaneous dipole. The distribution of the electron density in neighbouring molecules is influenced by the dipole of the first molecule; areas of relative high electron density on one molecule induce an area of low electron density on the neighbouring molecule and vice versa; thus neighbouring molecules form instantaneous dipoles that attract each other. Typically, each of these interactions has a strength of only ∼2 kJ mol−1 (compared with covalent bonds which typically have a strength of hundreds of kJ mol−1) and the magnitude of the interaction is strongly distance dependent, approaching zero at a separation of ∼8–10 Å. When molecules have large surface areas that can come into close contact, for example in biological macromolecules, these interactions can make a huge contribution to the total free energy.

Hydrogen bonding

Hydrogen bonds are a special case of dipole–dipole interaction, but are considered separately here as they are vital for the function of many biochemical systems. A covalent bond between hydrogen and an electronegative atom, such as oxygen, nitrogen or fluorine, is polar, with electron density in the vicinity of the hydrogen much lower than that around its bonding partner. The favourable interaction between the δ+ hydrogen of one X–H bond (where X is an electronegative element) and the δ X atom of another is called as a hydrogen bond. Hydrogen bonds are stronger than van der Waals interactions, but weaker than covalent bonds, with a typical strength between 5 and 40 kJ mol−1. It is important to note that the strength of a hydrogen bond depends heavily on the geometry of the atoms involved; the bonds are strongest when the three atoms involved in the bond are collinear. Hydrogen bonds are responsible for the specificity of base pairing, A to T and C to G, in DNA strands. They also play a key role in formation of structural elements during protein folding.

Ionic interactions

Ionic interactions occur between species that have full, permanent charges, i.e. ions. These interactions are much stronger than hydrogen bonds and London dispersion forces, but are much less common in biological systems. Ionic interactions between oppositely charged amino acids play an important role in stabilizing protein structure, as demonstrated by changes in protein shape with pH. Proteins have evolved to form the correct 3D structure at the pH of the environment where they function. At pH values far above and below the physiological range the charges of some of the amino acids forming the protein are changed (see section on pH and pKa for details). This changes the ionic interactions within the protein, in many cases causing the protein to unfold. Ionic interactions are particularly important in stabilizing proteins found in thermophilic organisms – those that thrive at temperatures above 40°C. Ionic interactions also play a key role in the binding of charged molecules such as ATP to their macromolecular partners.

The ‘hydrophobic effect’

The ‘hydrophobic effect’ is not a separate class of non-covalent interaction, but it is discussed here as the factors that give rise to this effect are very important for structure formation by biological macromolecules and especially proteins. Many molecules or parts of molecules, are hydrophobic or ‘water-hating’ and tend to cluster together when placed in water; this behaviour is often referred to as the hydrophobic effect. The hydrophobic effect is thought to be primarily entropic (see Why do chemical reactions happen? section), arising due to changes in hydrogen bonding within water in the presence of a non-polar molecule. Water forms an extensive network of transient hydrogen bonds which are broken and formed as the water molecules move. If a non-polar species is placed in water, it disrupts the hydrogen bonding networks. To minimize the number of hydrogen bonds lost due to this disruption water forms an ordered shell around the non-polar species. To make a large number of these ordered shells is entropically very unfavourable; clustering non-polar molecules together minimizes the non-polar surface area exposed to water and hence also minimizes the number of ordered water molecules (Figure 4). This is believed to be the primary driving force for burying non-polar groups on the inside of globular proteins and arranging polar groups on the outside. It is also the driving force for the self-assembly of lipid bilayers (discussed in detail in [4]).

Figure 4The hydrophobic effect

Water molecules form an ordered structure around a hydrophobic molecule (shown in grey). When two hydrophobic molecules aggregate the surface area exposed to water is reduced; this reduces the number of ordered water molecules in the hydration shell. Having more water molecules disordered in solution is entropically favourable.

Functional groups in biochemistry

The chemical reactivity of an organic compound depends upon the way in which its atoms are bonded together. Certain collections of atoms having the same connectivity occur frequently in organic compounds and these collections are called functional groups. Functional groups have a characteristic chemical behaviour, and it is possible to predict some of the properties of a molecule based on which functional groups are present. When discussing organic compounds we will use the skeletal representation; those unfamiliar with this representation should refer to Figure 5. Functional groups that are commonly found in biological molecules are shown in Figure 6.

Figure 5Displaying molecules using a skeletal representation

The skeletal representation makes use of the fact that carbon and hydrogen are the most common elements in organic compounds, and also that carbon almost always forms four covalent bonds. (A) The process of converting a displayed formula into a skeletal formula has three stages: (i) the molecule is drawn with the carbon chain ‘staggered’ in a zigzag; this is important to allow correct identification of the number of carbon atoms in the final molecule and is a better representation of the actual bond angles. (ii) The ‘C’ label for carbon atoms is omitted. (iii) The ‘H’ label for hydrogen atoms is omitted only for hydrogen atoms bonded to carbon. When interpreting a skeletal formula, we assume that carbon will make four bonds; if a carbon appears to make fewer than four bonds in the skeletal representation, then the ‘missing’ bonds are assumed to be to hydrogen atoms. Using the skeletal representation highlights the heteroatoms present in the molecule and becomes particularly useful once we consider how to display different isomers. (B) The guidelines for drawing skeletal structures are relatively flexible and in some cases, for example where we wish to draw attention to a particular group of atoms, the ‘C’ symbol for some of the carbon atoms may be shown. It is important not to omit any hydrogen atoms bonded to carbon atoms that are explicitly displayed in this way. (C) Wedge and hashed bonds can be used to illustrate the direction of bond vectors relative to the plane of the page.

Figure 6Functional groups

(A) This figure illustrates the some of the functional groups that occur frequently in biological systems. In these structures, the wavy bonds represent connections to the rest of the molecule. *In both amine and amide, one or both of the hydrogen atoms bonded to the nitrogen may be replaced by a carbon atom. (B) Many biological molecules include a large number of different functional groups, as illustrated by the structure of coenzyme A with the functional groups labelled in blue and red.

Hydrocarbon functional groups

Functional groups containing only carbon and hydrogen may be saturated (containing only carbon–carbon single bonds) or unsaturated (containing one or more carbon–carbon double or triple bonds). Alkanes are saturated molecules and are chemically unreactive compared with other functional groups. Alkanes are non-polar and hydrophobic, interacting less favourably with water than with other non-polar molecules. Alkenes are unsaturated hydrocarbons and they are also hydrophobic. They are, however, more reactive than alkanes, in particular towards the addition of atoms across the double bond. Alkanes and alkenes are abundant in cell membranes, where their hydrophobic nature is important in creating a barrier around the cell that is impermeable to molecules including water, ions and other polar molecules.

Functional groups in which carbon forms a single bond with an electronegative atom

The alcohol, ether, amine and thiol (–SH) functional groups all contain a carbon atom forming a single bond with an electronegative heteroatom (any atom that is not carbon or hydrogen). The difference in electronegativity between the carbon and the heteroatom makes these bonds polar; there is a higher probability of finding the electrons in the bond close to the electronegative atom than to the carbon atom. To indicate this polarization, we often show carbon with a partial positive charge and the electronegative atom with a partial negative charge. This uneven charge distribution greatly influences the reactivity of these molecules, with the carbon atom potentially being reactive towards electron-rich species known as nucleophiles. The electronegative atoms themselves may also be reactive. Oxygen, nitrogen and sulphur all have lone pairs of electrons that can form new covalent bonds in a chemical reaction; amines in particular are reactive towards protons, forming –NH3+ at physiological pH (see pH and pKa section). The uneven charge distribution also influences the way these molecules interact with water; these functional groups form favourable van der Waals dipole–dipole interactions with water molecules, as described earlier, and their presence will increase the solubility of a compound.

The carbonyl functional group

The carbonyl group, a carbon atom forming a double bond with an oxygen atom, is a part of many different functional groups. Carbonyl groups are found in a vast range of biological molecules, including proteins, DNA and sugars and they take part in many biological reactions. The reactivity of the carbonyl group varies depending upon which other atoms are bonded to the carbonyl carbon. The simplest carbonyl compounds are aldehydes, which have at least one hydrogen atom bonded to the carbonyl carbon, and ketones in which the carbonyl carbon is bonded to two other carbon atoms. As is the case with alcohols, the carbon–oxygen double bond in carbonyl groups is polarized, with the carbon atom partially positively charged and the oxygen atom partially negatively charged. The carbonyl carbon is particularly susceptible to attack by nucleophiles and this makes the carbonyl group useful for the formation of new carbon–carbon bonds. One important example where carbonyl chemistry plays a key role is the formation of the ketone bodies (acetoacetate, β-hydroxybutyrate and the breakdown product acetone). Ketone bodies are water-soluble species that can act as an alternative energy source to glucose in some tissues when glucose is scarce, for example during starvation or intense exercise. The ketone bodies are formed in the liver and then released into the bloodstream for use as a fuel in the heart, brain and muscle.

Carboxylic acids, esters, amides and acyl phosphates are known as carbonyl derivatives. As the name suggests, carboxylic acid functional groups readily lose a proton (the proton bonded to the oxygen atom); the resulting species is referred to as a carboxylate group or carboxylate anion. In these functional groups, the carbonyl carbon forms an additional single bond with an electronegative element. Although at first glance, it might seem as though this should make these carbonyl compounds more susceptible to nucleophilic attack than aldehydes or ketones (because the carbonyl carbon is bonded to more electronegative atoms) this is not actually the case due to delocalization of electrons within the functional group, as described earlier (see section: Delocalization of electrons). In biological chemistry, reactions involving the interconversion of carbonyl derivatives are very common, for example in protein synthesis where carboxylic acid groups are converted into amides.

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