LAURA BATMANIAN

JUSTIN RIDGE SIMON WORRALL

BIOCHEMISTRYfor health professionals

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PART A BIOLOGICAL CHEMISTRY

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Atomic number and atomic massThe atoms of different elements have different numbers of subatomic particles but all the atoms of a single element have the same number of protons in their nuclei. The number of protons is unique for each element and is referred to as its atomic number. The atomic number is written as a subscript on the left of the symbol for the element. Thus, 6C tells us that an atom of carbon has six protons in its nucleus. Since atoms do not carry a net charge the carbon atom must also have six electrons.

The mass number of an element allows us to determine the number of neutrons in the nucleus. The mass number is written as a superscript to the left of the element’s symbol. So, using carbon as an example 12

6 C tells us that the nucleus of a carbon atom contains six protons (from its atomic number) and six neutrons (mass number – atomic number). The mass number also gives us a close approximation of the atomic mass (in daltons).

IsotopesAll the atoms of an element have the same number of protons, but sometimes the number of neutrons varies. These different forms of a single element are referred to as isotopes. For example, carbon naturally occurs as a mixture of three isotopes with atomic masses of 12, 13 and 14. The most common form is 12

6 C which accounts for 99% of naturally occurring carbon. The remaining 1% consists mainly of 13

6 C (6 protons and 7 neutrons) with a small amount of 14

6 C (6 protons and 8 neutrons).

Both carbon-12 and carbon-13 are stable isotopes whose nuclei do not lose particles. However, carbon-14 is unstable and is radioactive. Radioactive isotopes have nuclei which spontaneously lose particles and give off energy. This process is often referred to as radioactive decay, and can result in a change in the atomic number such that a different element is formed. For example, 14C decays to produce stable nitrogen.

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7

In this decay reaction a neutron becomes a proton, which remains in the nucleus, an electron, and excess energy, which is released.

Radioactive isotopes can be useful in clinical diagnosis and in therapy. Isotopes that are inten- tionally introduced into the body are called radiopharmaceuticals. Depending on the type, the isotope will collect in one or more areas of the body. Since the isotope emits radiation, it is easily tracked and can be followed through the body and used to check if organs are healthy. Radioactive isotopes are also given to cancer patients in an attempt to damage cancerous tissue. However, radiation from decaying isotopes can also damage healthy tissues leading to cellular injury, often resulting in cellular death.

How are electrons organised in atoms?Simple models of an atom overemphasise the size of the nucleus relative to the whole atom. For a small atom such as helium, if the nucleus was the size of a small marble then the atom would have a radius of 50–60 m. At this scale the electrons would only be a few millimetres in diameter. From this you can see that the majority of the volume occupied by an atom does not contain anything. This means that when two atoms come together to undergo a chemical reaction, their nuclei are widely separated, and that only the electrons are involved.

The electrons associated with an atom have differing amounts of energy. Electrons close to the nucleus have the lowest amount of energy and are strongly attracted by the positively charged nucleus. Electrons further away from the nucleus are said to have higher energy because energy has to be expended to push them against the attraction of the nucleus. The energy levels of the electrons are not continuously distributed, instead occurring in discrete steps. If there was a continuous distribution of energy levels then the electrons would act like a ball rolling down a slope. However, because of the discontinuous distribution of the energy levels, electrons act more like a ball on a staircase. When a ball rolls down stairs it can spend time on each step but must drop quickly from step to step. Similarly, electrons do not spend appreciable time between energy levels. Thus, electrons are found in electron shells whose energy is relative to their distance from the nucleus (Fig 1-2). Electrons can move from one energy level to a higher one by absorbing

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1ELEMENTS AND COMPOUNDS, CHEMISTRY AND LIFE

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Orbital theoryInitially electrons were thought to orbit the nucleus in the same way that planets orbit a sun. However, this planetary model does not give a real picture of an atom. Electrons do not circle the atom in fixed, circular orbits. To get a better picture of atomic structure, chemists describe orbitals—regions around the nucleus where an electron is likely to be found most of the time. This orbital model is represented as an electron cloud surrounding the nucleus of the atom that represents the probable region of greatest electron density.

Each electron shell can now be thought of as an electron cloud containing electrons with a specific energy level that are distributed in a specific number of orbitals arranged in three-dimensional space (Fig 1-4). The first electron shell (1s) is spherical, the second has four orbitals of which one (2s) is spherical and three are dumbbell-shaped (2p orbitals). The next shell also has one ‘s’ and three ‘p’ orbitals, as well as others with more complex shapes. The shapes of these orbitals are important because they determine the shape of molecules when they are used to form chemical bonds (see below).

Each orbital is occupied by a maximum of two electrons. The first electron shell can hold two electrons in its s orbital whereas the next shell can hold a maximum of eight electrons in its four orbitals. Each of these electrons basically has the same energy but occupies a different volume of space. Chemical reactivity arises from the presence of unpaired electrons in one or more of their outermost shells.

Chemical bonds and compounds

Atoms with incomplete valence (outermost) shells can share or transfer valence electrons to or from another atom such that both atoms complete their valence shells. This normally results in the atoms staying close to each other (Fig 1-5). This interaction is termed a chemical bond, of which covalent and ionic are the strongest (Table 1-2).

FIGURE 1-2 Electrons exist at different energy levels in atoms.

Electrons closest to the nucleus have the lowest energy whereas

those furthest away have the highest. The energy levels are not

continuously distributed but exist in discrete steps. An electron may

absorb energy from the environment and jump one or more levels,

a process called excitation. Later it can return to its initial state by

giving up the energy it previously absorbed.

energy (e.g. light). This process is called excitation. Later, when the electron returns to its original energy level, the excess energy it possessed is released to the environment (e.g. as heat).

How electrons are distributed into their shells determines the chemical reactivity of the atom. The difference between one element and the next in the periodic table of the elements (a table made by arranging the elements according to their atomic number, part of which is shown in Fig 1-3) is the addition of a proton, an electron and one or more neutrons. The first electron shell can hold one pair of electrons whereas the next two can hold four pairs of electrons. This means that the first three levels hold 18 electrons.

The chemical properties of an element largely depend on the number of electrons in the outermost shell. These are often referred to as valence electrons. Atoms with the same number of valence electrons have similar properties. Atoms with full outermost shells are generally unreactive, being unable to easily react with other atoms. These atoms are also said to be inert. Atoms with incomplete outer shells are reactive.

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FIGURE 1-3 The initial elements in part of the periodic table of the elements. This figure shows how each element relates to the next. Each

element differs from the next by the addition of a proton and a variable number of neutrons. Elements with similar electron distributions such

as hydrogen, lithium and sodium, or helium, neon and argon, have similar chemical reactivity. A full periodic table is shown on page XXX.

[Based on Campbell & Reece, 2005, Biology, 7th Edition, Pearson Benjamin Cummings]

FIGURE 1-4 Electrons really occupy defined volumes of space. To better define the behaviour of electrons, the concept of orbitals—

volumes in which electrons spend 90% of the time—was developed. Electrons are distributed into shells of differing energies, with electrons

in each shell occupying defined orbitals.

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1ELEMENTS AND COMPOUNDS, CHEMISTRY AND LIFE

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Covalent bondsA covalent bond is formed when two atoms share a pair of valence electrons. The simplest example of this is to look at the formation of a molecule of hydrogen (H2) from two hydrogen atoms. Hydrogen atoms have a single valence electron in a shell (1s) that can hold two electrons. When two hydrogen atoms approach each other they reach a point where their electron orbitals overlap. At this point they can share electrons such

TABLE 1-3 Representative values for covalent and noncovalent bonds

Class of bond Type of bond Bond length (nm)

Strength (kJ/mole)

In vacuum In water

Covalent Covalent 0.15 380 380

Noncovalent Ionic 0.25 335 13

Hydrogen 0.30 17 4

van der Waals interaction (per atom) 0.35 0.1 0.1

Note: In water, covalent bonds are much stronger than the other attractive forces between atoms. Thus they define the boundaries of one molecule from another. However, many of the important biological interactions between molecules are mediated by noncovalent interactions that are individually quite weak, but together can create effective interactions between two molecules. These noncovalent forces are ionic bonds, hydrogen bonds and van der Waals interactions. The strengths of all noncovalent bonds are less than that of covalent bonds, in both the presence and the absence of water. The strength of a bond can be measured as the energy required (kilojoules; kJ) to break all the bonds in one mole of a molecule that contains only one bond, that bond being of one type only. The values in water are more representative of their relative importance in biological systems, whereas in vacuum values are really the maximum value for each bond type.

that each atom now has two associated electrons, with complete valence shells. Two or more atoms interacting by covalent bonds constitute a molecule (Table 1-3).

A similar story can be told for the formation of a molecule of oxygen (O2) from two oxygen atoms. However, since an oxygen atom has six valence electrons in a shell that requires eight to be complete, the two oxygen atoms share two pairs of electrons to complete their valence shells. The sharing of a single pair of electrons is referred to as a single bond, and the sharing of two pairs is termed a double bond.

The number of pairs of electrons that an atom needs to share to fill its valence shell, that is, the number of covalent bonds it generally needs to form to do this, is termed its binding capacity or valence. This can be used to explain the valences of elements such as hydrogen, oxygen and nitrogen, but not for some other elements. In naturally occurring compounds phosphorus often has a valence of five, not three as would be predicted using the rule outlined above. This is because a phosphorus atom, which has five electrons in its valence shell, can use its three unpaired electrons to make single bonds but can also use its outermost pair of electrons to make a double bond.

So far the examples of bonding that have been examined are between two atoms of the same element. However, atoms of different elements can also interact to form molecules. One of the simplest examples of two different elements combining to form a molecule is water (H2O). In this molecule, oxygen completes its valence shell by forming single bonds with two

TABLE 1-2 How atoms interact to make simple

molecules

CompoundDistribution of electrons

Structural representations

Covalent bond type

Hydrogen(H2)

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Methane(CH4)

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Note: This table demonstrates how three elements can be used to make molecules through chemical bond formation. The first two examples show molecules made from two atoms of the same element whereas the last two show molecules made from two different elements

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PART A BIOLOGICAL CHEMISTRY

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Molecular shapesAny molecule has a distinctive size and shape that is dependent on the atoms used to make it, and on the pattern in which they are bonded to each other. As we will see in later chapters, the functionality of many biological molecules is often determined by their three-dimensional shape.

Molecules made from two atoms such as H2 or O2 are always linear. However, molecules comprising three or more atoms have much more complex shapes. Their shapes are derived from the orbitals used to form the bonds between the atoms. When an atom forms covalent bonds, the orbitals in the valence shell rearrange. An atom with valence electrons in both s and p orbitals hybridise to form four new hybrid orbitals that are teardrop-shaped and extend from the region of the nucleus. These orbitals delineate a volume of space called a tetrahedron, a shape similar to a pyramid. An example of a tetrahedral molecule is methane (Fig 1-9A). The carbon nucleus sits at the centre of the tetrahedron and the four hydrogen atoms bonded to the carbon sit at its four corners. Water is also a tetrahedral molecule though it is less easy to see why (Fig 1-9B). The shape of water is derived from the formation of two single bonds between the central oxygen atom and two

FIGURE 1-9 Examples of simple molecules. A shows the

structure of a molecule of methane. When an atom with valence

electrons in both the s and p orbitals forms a covalent bond, the

orbital hybridises to give four teardrop-shaped hybrid orbitals that

delineate a tetrahedron. In the case of methane the carbon sits at the

centre of the tetrahedron and the four covalent bonds it makes with

single hydrogen atoms are the four corners. B shows the structure of

water. Oxygen also makes single bonds with hydrogen atoms (two)

and these sit at opposing corners of the tetrahedron. The other two

corners are occupied by pairs of electrons that are not used to make

bonds. Thus, water is also a tetrahedral molecule.

hydrogen atoms. Each of these sit at opposite corners of the tetrahedron while the other corners are occupied by non-bonding orbitals containing pairs of electrons which are totally derived from the oxygen atom.

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1ELEMENTS AND COMPOUNDS, CHEMISTRY AND LIFE

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CHAPTER SUMMARY

Content relating to this chapter is available online at:

http://evolve.elsevier.com/AU/Batmanian/biochemistry/

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