Biochemical concepts
3.2 Chemical bonding
First electron shell
Hydrogen (H) Atomic number = 1 Mass number = 1 or 2 Atomic mass = 1.01
Sodium (Na) Atomic number = 11 Mass number = 23 Atomic mass = 22.99
Atomic number = number of protons in an atom
Mass number = number of protons and neutrons in an atom (boldface indicates most common isotope) Atomic mass = average mass of all stable atoms of a given elements in daltons
Chlorine (Cl) Atomic number = 17 Mass number = 35 or 37 Atomic mass = 35.45
Potassium (K) Atomic number = 19 Mass number = 39, 40, or 41 Atomic mass = 39.10
Iodine (I) Atomic number = 53 Mass number = 127 Atomic mass = 126.90 Carbon (C)
Atomic number = 6 Mass number = 12 or 13 Atomic mass = 12.01
Nitrogen (N) Atomic number = 7 Mass number = 14 or 15 Atomic mass = 14.01
Oxygen (O) Atomic number = 8 Mass number = 16, 17, or 18 Atomic mass = 16.00
Third electron shell
Fourth electron shell
Fifth electron shell Second
electron shell
1p+ 6p+
6n0
7p+ 7n0
8p+ 8n0
53p+ 74n0 19p+
20n0 17p+
18n0 11p+
12n0
Figure 3.2 Basic diagrammatic representation of the atomic structure of several stable atoms. Note the difference in atomic structures between atoms. You should also note that because these atoms can exist as stable isotopes, the mass number can differ because they can differ in the number of neutrons present in the nucleus. (adapted from Tortora and Derrickson,Principles of Anatomy and Physiology, Twelfth Edition, 2009, reproduced by permission of John Wiley & Sons Inc.)
together they form an oxygen molecule which is symbolized as O2.
A compound is a substance which has molecules comprised of atoms of two or more different elements. For example, water is a compound because it consists of two hydrogen atoms joined together with an oxygen atom, and it is thus symbolized as H2O. It is important to note that while all compounds are molecules, not all molecules are compounds.
Most of the chemical elements in our bodies exist in the form of compounds which can be further classified asinorganic compounds (lack- ing carbon) or organic compounds (containing carbon). Some of the most important organic compounds that are relevant tosport and exercise metabolism include carbohydrates, fats and pro- teins. Collectively, these compounds are called the macronutrients, and it is through the action of specific biochemical processes (i.e. chemical
reactions) that our bodies utilize these food sources to provide our muscles with the energy to exercise. Figure 3.3 shows the general flow of chemical organization from atom to macronutrient.
Neutrons Protons
form
ATOMS
gain/lose protons gain/lose electrons gain/lose neutrons
DIFFERENT ELEMENTS such as
such as are the main elements in
BIOMOLECULES
HYDROGEN OXYGEN NITROGEN
CARBON
IONS ISOTOPES
Electrons
CARBOHYDRATES FATS PROTEINS
Figure 3.3 The basic flow of matter. Atoms of elements such as oxygen, carbon, hydrogen and nitrogen ultimately combine to make biomolecules, examples of which include the foodstuffs and fluids that we eat and drink in order to fuel our muscles during exercise and, more importantly, to sustain daily life
3.2.1 Ionic bonds
We have seen above how atoms can becomeions by losing or gaining an electron, which ultimately results in a positively or negatively charged ion, respectively. If you have two atoms that can achieve a stable valence shell by either donating or gaining an electron, the result is a force of attrac- tion which can bond the oppositely charged atoms together via an ionic bond. This is most simply illustrated through the bonding of sodium (Na) and chlorine (Cl) atoms to make the compound sodium chloride (NaCl), which is the chemical name for common table salt (see Figure 3.4). It is important to remember that whenever atoms
bond via ionic bonds, the net charge of the newly formed compound is always zero.
3.2.2 Covalent bonds
In contrast to ionic bonds, covalent bonds (the strongest of all chemical bonds) work on the principle of atoms sharing electrons. Atoms can form covalent bonds by sharing one, two or three pairs of their valence electrons to make a single, double or triple covalent bond, respectively (see Figure 3.5). The larger the number of electron pairs shared, the stronger the covalent bond.
Covalent bonds can be further classified as nonpolar or polar. In nonpolar covalent bonds,
Na Na
Cl Cl
Electron donated
Electron accepted
Atom Ion
−
+
Atom Ion
Sodium: 1 valence electron
Chlorine: 7 valence electrons (a)
(b)
Na Cl
+ −
Ionic bond in sodium chloride (NaCl) Packing of ions in a crystal of sodium chloride
Na+ Cl−
(c) (d)
Figure 3.4 Ionic bond formation. In this example, sodium forms an ionic bond with chlorine by donating an electron to the chlorine atom. The sodium atom now becomes positively charged and the chlorine atom now becomes negatively charged. (adapted from Tortora and Derrickson,Principles of Anatomy and Physiology, Twelfth Edition, 2009, reproduced by permission of John Wiley & Sons Inc.)
H
+ δ−
δ+
δ+ H
H H
H2O
H H
O
Oxygen atom Hydrogen atoms Water molecule
O O
Figure 3.5 Covalent bond formation. In this example, electrons are shared equally between oxygen and hydrogen atoms. This bond is referred to as a polar covalent bond because the oxygen atoms attract the electrons more strongly. As such, the oxygen end of the water molecule has a partial negative charge (δ−)and the hydrogen ends have a partial positive charge (δ+). This polar covalent bond makes water an excellentsolvent, as discussed later in this chapter. (adapted from Tortora and Derrickson,Principles of Anatomy and Physiology, Twelfth Edition, 2009, reproduced by permission of John Wiley & Sons Inc.)
the atoms share the electrons equally, meaning that one atom does not attract the shared electrons more than the other. When two or more atoms of the same element form a covalent bond, it is always nonpolar.
In contrast, a polar covalent bond is one where the sharing of electrons between atoms is unequal, so that one atom attracts the shared electrons more than the other. An important example of polar covalent bond is a water molecule (H2O); here, it is the oxygen atom which has the greater power to attract electrons to itself (greaterelectronegativity) from the two hydrogen atoms.
3.2.3 Molecular formulae and structures In briefly recapping what we have covered so far, you should now appreciate that atoms from elements can combine together, largely through the action of ionic or covalent bonds, to make molecules and compounds. In biochemistry, we can depict the atoms which form the molecule or compound through writing itsmolecular formula, which involves writing the chemical symbols of the atoms involved and moreover, the number of atoms of each element involved.
For example, one molecule of glucose (a vital energy source for exercise) contains six carbon atoms, 12 hydrogen atoms and six oxygen atoms
and can therefore be written as C6H12O6. From the molecular formula, we can also calculate the molecular weight of the molecule, i.e. the sum of the atomic masses of the elements which make up the molecule. Note that we begin this process by multiplying the atomic mass of the element by the number of atoms present in the compound. For example, since glucose contains six carbon atoms, 12 hydrogen atoms and six oxygen atoms and the atomic masses of each element are 12, 1 and 16 respectively, the molecular weight can be calcu- lated as follows:
C6H12O6=(12×6)+(1×12)+(16×6)=180 In addition to molecular formula, we are also interested in knowing the molecular structure (i.e.
constitutional formula) of the compound – the structural arrangement by which the atoms have bonded to form the compound.
Figure 3.6a shows the molecular structure of a glucose molecule. A single line symbolizes a single covalent bond between atoms and a double line symbolizes a double covalent bond. Similarly, if there were a triple covalent bond between atoms present in this molecule, it would be symbolized by a triple line. In this example, the structure of glucose is shown in an open chain format which appears to exist as a two-dimensional structure. However, it is important to note
H C C C C C C H
O OH H OH OH OH H
H H HO H
(a)
C C
C C
H O
HO
H
OH H
OH H
OH H H
OH H
(b)
O
HO OH
OH OH
CH2OH
(c)
C C
Figure 3.6 (a) Molecular structure of a glucose molecule. Note that the single and double lines rep- resent a single or double covalent bond, respectively.
(b) Molecular structure of glucose as shown as a ring like structure. (c) Molecular structure of glucose as written in shorthand method. In this example, carbon atoms are understood to be at locations where two bond lines interact and single hydrogen atoms are not shown that in reality, molecules and compounds exist as three-dimensional structures. Furthermore, many compounds are also formed in special structural shapes. For example, the majority of glucose molecules in our bodies are stored in a ring-like structure comprising a ring-shaped
‘carbon skeleton’ with hydroxyl groups attached (hydroxyl groups consist of an oxygen and carbon atom joined by a single covalent bond) (see Figure 3.6b). Given that organic molecules and compounds are often relatively big, we sometimes draw their constitutional formula using shorthand methods (see Figure 3.6c).
In understanding molecular formulae and struc- tures, it is important to note that some compounds can have the same molecular formula (i.e. the same number and type of atoms) but can differ in constitutional formula because the atoms have
bonded in a different structural arrangement. Such molecules are called isomers, and one example is that of glucose and fructose, which both have a formula of C6H12O6. We will re-visit these compounds in Chapter 5.
3.2.4 Functional groups
Since we have now covered the basic processes of chemical bonding, it is also important to note that many atoms can bond repeatedly within compounds in certain combinations to yield specific functional groups. Indeed, during chemical reactions (see next section) the atoms contained within these functional groups tend to move between compounds as a unit rather than as individual atoms. Much of these functional groups can bond to a carbon atom (or to a chain of carbon atoms known as a carbon skeleton) via a single covalent bond to form ring-like structures or straight or branched chains. Some of the most important functional groups are shown in Table 3.2, where we also show the shorthand notation and bond structure.
Table 3.2 Common functional groups found within biological compounds. R denotes carbon skeleton to which the functional group is attached to
Name Shorthand notation Bond structure
Carboxyl –COOH O
C
R OH
Hydroxyl –OH R O H
Amino –NH2
R N H H
Phosphate –PO4
P O
O− O− O R
Carbonyl –CO O
C
R R
Sulphydryl –SH R S H