Biology Olympiad

Section 1: Cell chemistry and cell biology

Module 1: Introduction
Learning objectives

At the end of this learning module, you should be able to:

 Define chemical elements, compounds, and atoms and provide examples of
the subatomic particles that make up atoms
 Explain how molecules are formed
 Describe the four emergent properties of water
 Describe how water molecules dissociates and the importance of pH, acids
and bases in this process
 Describe the general principles of molecular structure and define the most
common functional groups
 Define the four macromolecules and discuss their structure and function
 Briefly describe metabolism and explain how enzymes regulate this process
The chemical context of life

Chemical elements and compounds

The term ‘matter’ refers to anything that takes up space and has mass. It consists of
chemical elements in pure form (i.e. singular elements) and also in combinations,
called compounds. An element is a substance that cannot be broken down to other
substances by chemical reactions, eg: sodium (Na). A compound is a substance
consisting of two or more elements combined in a fixed ratio, eg: sodium chloride
(NaCl), which exhibits emergent properties not found in Na or Cl individually.
Life requires approximately 25 chemical elements. Carbon (C), hydrogen (H),
oxygen (O) and nitrogen (N) make up the primary macromolecules necessary
for life, along with smaller quantities of phosphorus (P), sulfur (S), calcium (Ca),
potassium (K), sodium (Na), chlorine (Cl), magnesium (Mg) and trace elements such
as fluorine (F) and manganese (Mn).

Atoms and molecules

The atom is the smallest unit of matter and can be thought of as the smallest particle
of a chemical element that can exist. A molecule is a group of atoms bound together
and is the smallest unit of a chemical compound.
The atomic structure of an element determines the behaviour of that element. Atoms
have the following properties:
 Every atom is composed of electrons, neutrons and protons. (Figure 1)
 Atoms have the same number of electrons as they do protons.
 All atoms of a particular element have the same number of protons and this
gives them their atomic number. For example, C has 6 protons and therefore
the atomic number for C is 6, which is denoted as 6C.
 The sum of the number of neutrons and protons of an atom gives the atomic
mass of that atom. For example, the number of protons and neutrons in C is
12, which is denoted as 12C.
12
When the atomic number and atomic mass are combined it is denoted as 6 C .
Isotopes are a version of an atom that has the same number of protons but a different
number of neutrons. Because the name of the atom is determined by the number of
protons, the name of the atom will not change, but the atomic mass will (remember
atomic mass= number of protons + number of neutrons). 99% of all carbon is the
standard 12C with 6 neutrons, however a few carbon isotopes exist: carbon 13 (13C)
with 7 neutrons, and carbon 14 (14C) with 8 neutrons.
Most isotopes are very rare but can be used for diagnostic purposes. For example,
14C (also called radiocarbon) is a radioactive isotope that forms at a constant rate. Its

emission is used in the technique of carbon dating to determine the age of once-living
matter.
Figure 1 – Basic structure of the atom.
Source: https://upload.wikimedia.org/wikipedia/commons/d/d8/Atom_diagram.png

Subatomic particles
Particles that are smaller than the atom are called subatomic particles (sub=
“below”). Protons, neutrons, and electrons are all subatomic particles and it is these
particles that make up the atom.
The atomic mass unit (amu) Standard unit of measurement for indicating mass of
subatomic particles. The dalton (Da) is another term for the amu and is currently the
more widely favoured unit of measurement. 1 amu = 1 Da.
Protons and neutrons have a mass of around 1 amu or 1 Da. For example, hydrogen
(1 proton and 0 neutrons) has a mass of 1 Da and carbon (6 protons and 6 neutrons)
has a mass of 12 Da. Electrons have a mass of ~1/2000th Da, therefore the
contribution to the overall mass of the atom is insignificant and not considered when
determining atomic mass.
Electrons are the only subatomic particles that interact between atoms. They have
the following characteristics:
 Electrons vary in the amount of energy they possess. Electrons have different
amounts of potential energy based on their position in relation to the nucleus
due to electrostatic attraction between the nucleus and electrons.
 Electrons exist at fixed energy levels, called electron shells. The first electron
shell is closest to the nucleus, second is the next closest shell, and so on.
(Figure 2)
 Electrons can move up shells by absorbing energy, and move down by
releasing energy. For example, as light energy is absorbed electrons move up
a shell and as heat is released electrons move down a shell.

Figure 2 – Electron shell configuration. Source: Unknown.

The chemical behaviour of an atom is determined by its electron configuration, that

is how the electrons are distributed in the electron shells. The first shell holds a
maximum of 2 electrons, the second holds 8, the third holds 18, etc. Electrons will fill
the lowest shell possible. Reactions occur between electrons in the outermost shell
(valence shell) and these electrons are called valence electrons.
For example: Carbon has 6 electrons in total; 2 inner shell and 4 valence electrons in
the outermost shell. Because there are 4 electrons in the valence shell we say that
carbon has a valence of 4.
Formation of molecules
Atoms typically form compounds by giving, taking, or sharing valence electrons. These
chemical bonds can be ionic (where electrons are transferred from one atom to
another) or covalent (where electrons are shared between two atoms). The polarity of
these bonds is important in biology.
Weak chemical bonds play important roles in the chemistry of life. Hydrogen bonding
is an example of a weak chemical bond; this type of bonding is an attractive
intermolecular force between molecules with a H atom bound to a small, highly
electronegative atom (eg: F, O, N) that has lone electron pairs. van der Waals forces
are also considered weak forces.

Relationship between molecular shape and function

A molecule’s biological function is related to its shape. Molecules have 3D
shapes/conformations which determines their functional capability. For example,
water has a bent molecular conformation due to the hybridisation of the electron
orbitals after covalent bonding, and methane has a tetrahedral shape also due to
orbital hybridisation. Larger, more complex molecules have complex shapes which are
crucial in biology, allowing specific interactions between particular molecules, such as
morphine and endorphins.
Water

There are several properties of water that are essential for sustaining life, these are
termed the emergent properties:
1. Water has a cohesive behaviour
2. Water is able to moderate temperature
3. Water expands upon freezing
4. Water is a versatile solvent

1. Cohesion
Water is a polar covalent molecule with a bent structure that is capable of forming
hydrogen bonds. Each water molecule can form 4 hydrogen bonds with other water
molecules. The polarity of water gives rise to other properties of cohesion, adhesion
and surface tension.
Cohesion is the ability of water molecules to bond to other water molecules.
Adhesion is the ability of water molecules to bond with other molecules. When water
is spilt over certain surfaces, such as glass, the adhesive forces are stronger than the
cohesive forces and the molecules will spread out over the surface.
The cohesion of water also contributes to its surface tension property. Surface
tension is the tension of the surface of a liquid caused by the attraction of the water
molecules on the top layer to the water molecules below. In this case cohesive forces
are greater than adhesive forces and the molecules on the surface are attracted
towards the molecules below, resulting in water trying to reduce its surface area by
bunching tightly together.
The water strider is a great example of surface tension. The water strider has flexible
legs that are able to deform at the surface of the water so they do not ‘pierce’ the water
and disrupt the surface tension (Figure 3).
Figure 3 – Water strider. Source: https://www.britannica.com/animal/water-strider

2. Moderation of temperature
Kinetic energy is the energy an object has when in motion. Heat is the total quantity
of kinetic energy in a body of matter and temperature is a measure of the average
kinetic energy of matter.
Water stabilises air temperatures by absorbing heat from the warm air during the day
and releasing it to cooler air at night. Water is an effective heat store because it can
absorb a relatively high amount of heat without its own temperature changing
significantly.
Specific heat is the amount of heat required to change the temperature of 1g of a
substance by 1°C. This is a measure of how resistant the substance is to changing its
temperature when absorbing and releasing heat. The unit of measurement for specific
heat is the calorie. The heat of vaporization is the amount of heat a liquid must
absorb for 1g of the liquid to be converted from a liquid to a gaseous state
(approximately 580 cal/g). Heat must be absorbed in order to break hydrogen bonds
and heat is released when hydrogen bonds are formed. Ice freezes and is less dense
than liquid water, see fig 3.5, floats so ‘life’ forms continue beneath the ice. Water
allows buffering of body temperatures.

3. Expansion on freezing
Water is less dense in solid form than it is in liquid form. This is important when you
consider large bodies of water. Ice that has formed in winter on the surface of a lake
will not sink as it is less dense than the liquid water beneath it. This is important to
maintain life; the ice layer on top of a body of water and glaciers that float in oceans
help to keep the water beneath insulated, ensuring a favourable environment for
animals that live in the water.
4. Versatile solvent
Water is considered a solvent because many substances will dissolve in it. A
substance that dissolves in water is termed a solute, and the liquid that is formed from
the combination of solvent and solute is a solution. Solutions in which water is the
solvent are called aqueous solutions. For example, when dissolving ionic NaCl that
is covalently bound together the polar water molecules surround the cations and
anions to form hydration shells. Any substance that has an affinity for water is called
hydrophilic (even if the molecules do not dissolve in water, eg: hair) and substances
that repel water are called hydrophobic (non-ionic and nonpolar molecules, eg: oil).
Dissociation of water molecules

Molarity
Most biologically important chemical reactions occur in water, which is why it important
that we know how to calculate the concentrations of solutes dissolved in aqueous
solutions. Scientists measure atoms using a unit called the mole (Avogadro’s
number), which is 6.02 x 1023. That is, one mole of particular solute is equal to 6.02 x
1023 atoms of that solute.
It just so happens that a mole of carbon is equal to 12g, which is the atomic mass of
carbon. A mole of hydrogen is ~1g, a mole of methane (CH4) is 16g, a mole of O2 is
32g (2 x 16) and a mole of glucose (C6H12O6) is 6 x 12 + 12 x 1 + 6 x 16= 180g.
The measurement, mole, measures the number of atoms. To calculate the number of
atoms in a solution (i.e. number of atoms per unit volume), we use a measurement
known as molarity – the concentration of a particular solute in a solution. Molarity (M)
is the number of moles of solute per 1L of the solution.

Dissociation
Another interesting property of the water molecule is its ability to dissociate. A
hydrogen can shift from one water molecule to another – it leaves behind an oxygen
covalently bonded to the other hydrogen, and also leaves behind an electron, giving
hydroxide (OH-) and a proton (H+). The proton joins with another H2O to form
hydronium (H3O+) (Figure 4). We can write this as: 2H2O  OH- + H3O+, or put more
simply, we can just write: H2O  OH- + H+.
This process is reversible, hence the double arrows. In pure water, a very low
concentration of H+ or OH- exists – that is only a small number of water molecules are
dissociated. In fact, the concentration of each ion in pure water is 10-7 M, meaning 1
in every 554 million water molecules is dissociated at any one time.

Figure 4 – Dissociation of water.

Source: https://upload.wikimedia.org/wikipedia/commons/thumb/6/6a/Autoprotolyse_eau.svg/2000px-
Autoprotolyse_eau.svg.png
pH, acids and bases
The pH scale was first developed by chemists in Carlsberg Brewery, Denmark. The
pH scale is a logarithmic scale to express the concentration of hydrogen ions (H+) in
a solution. We can change the concentration of H+ and OH- in a solution by adding
substances called acids and bases.
An acid increases the H+ concentration of a solution, and a base decreases it. Acids
increase H+ by donating protons to the solution, eg: HCl dissociates to form H+ and
Cl-, where the H+ from the HCl increases the total concentration of protons in the
solution.
Bases decrease H+ by accepting protons, often referred to as ‘mopping up’ excess
H+ in a solution. Bases such as NaOH can dissociate to give OH-, which combines
with H+ (in a neutralisation reaction) to give H2O and decreases the H+ concentration.
Bases such as NH3 (ammonia) can also reduce H+ concentration by soaking up excess
H+ to give ammonium (NH3 + H+  NH4+).
When [H+] = [OH-], a solution is said to be neutral
When [H+] > [OH-], the solution is acidic
When [H+] < [OH-], the solution is basic
In an aqueous solution, the product of the concentrations of H+ and OH- is equal to 10-
14. That is, [H+][OH-] = 10-14. In a neutral solution at 25°C, [H+] = 10-7 = [OH-]. For

example, say we add acid to a solution, bringing the [H+] to 10-5. What would be the
[OH-]? It would be 10-9.
Because the relative concentration of these two ions is critical for life, it helps to be
able to express them in an easily understandable fashion – hence the pH scale. This
typically ranges from 0 to 14, with 0 being highly acidic, 14 being highly alkaline/basic,
and 7 being neutral (Figure 5). pH is calculated according to the concentration of H+
ions in an aqueous solution and is the -log10 of the concentration of H+ ions – hence
pH is just a mathematical expression. pH = -log10[H+]. Note that pH is inversely
proportional to H+ concentration and that for a 10-fold change in H+, pH only changes
1-fold. pH can also be calculated from [OH-], so if [OH-] = 10-5 then [H+] = 10-9, which
means pH = 9.
The internal pH of most cells is around 7 and even a slight change in pH can be
harmful, affecting biochemical reactions and the shape (hence function) of bio-
molecules.
Figure 5 – pH scale.
Source: https://upload.wikimedia.org/wikipedia/commons/2/23/216_pH_Scale-01.jpg
Carbon and the molecular diversity of life

Compounds containing carbon are said to be organic – the study of such compounds
is organic chemistry. Organic compounds range from simple things like CO2 and
methane, to proteins and DNA.
The electron configuration of carbon determines the types of bonds and number of
bonds that carbon is able to form. Carbon has 4 valence electrons, which means that
it can form 4 covalent bonds with other atoms (eg: CO2, methane, ethane, ethene,
urea). Carbon atoms are the most versatile building blocks of molecules. Carbon can
combine with other elements like hydrogen, oxygen, nitrogen, sulphur and phosphate
to form a large variety of different molecules and life can be built from this limited
collection of building blocks.

Note the molecular and structural formulae above – each single-dash line represents
a bond between two atoms, which is a pair of shared electrons. Each double-dash line
represents a double bond between two atoms. Although molecules are drawn flat, in
reality are 3-dimensional.

Molecular structure
Carbon can covalently bond with itself to form carbon skeletons (chains) of
essentially limitless variety and complexity, which form the basis of most organic
molecules. These vary in length, and can have different conformations – straight,
branched, rings, etc. Carbon-carbon bonds may be single, double, or triple and the
different bond types will affect the 3D molecular shape. One example of a carbon
skeleton molecule are hydrocarbons. These are molecules with only H attached to
the carbon skeleton, and they form compounds such as fat and petrol.

Isomers are compounds that have the same molecular formula but different structures
and most likely different properties, since structure fits function. 3 different types:
 Structural isomers – differ in covalent arrangements of atoms – number of
possible isomers increases as carbon skeletons increase in size.
 Geometric isomers – differ in spatial arrangements of atoms about a double
bond, arising from the inflexibility of double bonds that do not allow rotation –
 eg: vision necessitates light-induced change from one isomer to another of
rhodopsin.
 Enantiomers – mirror images of each other – (for example) carbon may be
attached to 4 different atoms/structures in two different ways which are mirror
images of each other – these cannot be superimposed on each other. The
middle carbon is called the chiral carbon or the asymmetric carbon.
Enantiomers are biologically important molecules as well.

Functional groups
There are distinctive properties of organic molecules that depend on the shape but
also the properties of the molecular components attached to the carbon skeleton.
Functional groups are defined as a specific structure responsible for particular
chemical behaviours of a substance, and they are typically the components of organic
molecules that are most commonly involved in biochemical reactions. Each functional
group behaves similarly regardless of the molecule it is part of. The number and
arrangement of the groups helps give an organic molecule its unique properties (eg:
testosterone vs estradiol in sex development).
The major functional groups are summarised below:
 Hydroxyl group (–OH). Organic compounds containing hydroxyl groups are
called alcohols (eg: ethanol). Hydroxyl groups are polar and therefore alcohols
are easily dissolved by water.
 Carbonyl group (–COH). Compounds are called aldehydes (when attached at
the end of the molecule) or ketones (when attached in the middle of the
molecule). The oxygen is joined to carbon by a double bond and has non-
bonding electron pairs to H-bond.
 Carboxyl group (–COOH). A carboxyl group is an oxygen and hydroxyl joined
to a carbon. These compounds are called carboxylic acids because the
carboxyl group is a source of hydrogen ions. Carboxylic acids can dissociate
thanks to the two oxygens being so electronegative and pulling electrons away
from the H, which falls off as a proton. The deprotonated form is the most
common inside living cells.
 Amino group (–NH2). An amino group is two hydrogens attached to a nitrogen
and the compounds are called amines. Amino group acts as a base, because
they can pick up a proton from the surrounding solution to form NH3+. The
protonated form is the most common in living cells.
 Sulfhydryl group (–SH). Hydrogen bonded to sulfur and the compounds are
called thiols. This functional group is important in protein stabilisation.
 Phosphate group (–OPO32-). Phosphate is an anion formed by the dissociation
of hydrogens from phosphoric acid H3PO4. Compounds are called organic
phosphates and are important in energising cellular machinery.
The structure and function of macromolecules

Polymer principles
Cells are able to join small organic molecules together to form larger molecules –
called macromolecules. There are four types of macromolecules: carbohydrates,
lipids, proteins and nucleic acids. A cell generates macromolecules primarily by
sticking smaller molecules together in a process called polymerisation. Three types
of macromolecules (all except lipids) are polymers – long molecules consisting of
many similar or identical monomers (building blocks) linked covalently together. The
chemical mechanisms used to make and break polymers are similar for the three
macromolecules.
Monomers are covalently linked through the loss of a water molecule, in a
dehydration (condensation) reaction. When a bond forms between two monomers,
each monomer contributes part of the water molecule that is lost (a hydroxyl and a
hydrogen). To make a polymer, this reaction is repeated as monomers are added to
the chain one by one. Energy must be used to create polymers from monomers, in
anabolism (biosynthesis of molecules). Polymers are disassembled to component
monomers by the addition of a water molecule in a process termed hydrolysis.
Hydrolysis is essentially the reverse of dehydration synthesis, which is catalysed by
enzymes in catabolism (metabolic breakdown of large molecules).
Macromolecules – Carbohydrates

Carbohydrate refers to simple sugars (monomers) and their polymers (two or more
monomers together). These are fuels, and have storage and structural roles.
Carbohydrates are generally sweet, and smaller carbohydrates are soluble. As the
molecular size increases, solubility decreases and sweetness is lost.
The simplest carbohydrates are known as monosaccharides, or simple sugars, and
generally have the molecular formula of (CH2O)n. Glucose is a typical
monosaccharide (most names for sugars end in –ose). Simple sugars can be ketoses
or aldoses depending on the position of their carbonyl group and are classified in terms
of the number of carbons in the molecule (eg: hexoses have six carbons).
Note that structural isomerism (eg: glucose and fructose) and enantiomerism (eg:
glucose vs galactose) are important. Glucose actually has a ring structure in aqueous
solutions due to a reaction between the carbonyl and a hydroxyl group.
Note in the figure below that the –OH attached to C1 can take an up or down
orientation upon cyclisation – giving β or α glucose, respectively (Figure 6).

Figure 6 – Alpha and beta glucose.

Source:
https://upload.wikimedia.org/wikipedia/commons/a/a0/Alpha_%2B_beta_D-
Glucose_(PYRANOSE)_V.3.png
When two monosaccharides are joined in a dehydration reaction by a glycosidic
linkage, a disaccharide is formed. Sucrose is a typical disaccharide, formed from
glucose and fructose. Other disaccharides include lactose (glucose & galactose) and
maltose (glucose and glucose).
When many monosaccharides are joined, a polysaccharide is formed. These are
macromolecules, with 100s or 1000s of monosaccharides joined by glycosidic linkage.
The structure and function of polysaccharides are determined by the type of sugar
monomer and the position of the glycosidic linkages.
Polysaccharides are typically used for storage and structural integrity. Storage
polysaccharides include starch and glycogen.
 Starch is a polymer of glucose monomers, joined by 1-4 glycosidic linkages
(carbon 1 to carbon 4) – typically used by plants for storage of excess CHO,
the bond angle makes the polysaccharide helical; examples of starch: amylose
(simplest, unbranched), amylopectin (branched).

 Glycogen is a polymer of glucose but more extensively branched than even

amylopectin – used by animals for storage (‘the starch of animals’).

Structural polysaccharides include cellulose and chitin.

 Cellulose is a polymer of glucose, but the cellulosic glucose is slightly different
from starch/glycogen glucose (uses β-glucose instead of α) – used by plants
to build cell walls. In the cellulose polymer, every β-glucose is upside down in
relation to its neighbour. Cellulose is straight and never branched, and has
exposed –OH groups for H bonding with other parallel cellulose molecules –
cellulose macromolecules can group tightly & strongly together into
microfibrils. Also note that enzymes (eg: amylase) that degrade α-glycosidic
linkages cannot break β-linkages, and vice versa.

 Chitin is another structural polysaccharide – used by arthropods to build

exoskeletons and fungi to build cell walls. The chitin monomer (actually N-
acetylglucosamine) has a nitrogen-containing addition.
Macromolecules – Lipids
Lipids are diverse, hydrophobic molecules that store energy and form phospholipids
and steroids. Lipids are not polymers and have little affinity for water (hydrophobic),
which enables them to spontaneously form membranes. Lipids consist mainly of
hydrocarbons, which are highly non-polar, hence the hydrophobic nature of these
macromolecules. The three important groups of lipids are fats, phospholipids and
steroids. (Figure 7)
Fats are large molecules assembled from 1 glycerol (an alcohol) and 3 fatty acid
(hydrocarbon plus carboxyl group) molecules by dehydration reactions. The fatty acid
can be the same or different and they can vary in length as well as the number/position
of double bonds.
Phospholipids are similar to fats but only have 2 fatty acid tails, the third connection
to glycerol is used by a phosphate group which can be in turn linked to other small
molecules (usually charged/polar). Phospholipids are ambivalent towards water
because of their non-polar hydrophobic hydrocarbon tails and polar hydrophilic
phosphate/attachment head – they are said to be amphipathic. Because of this, when
added to water, PLs can self-assemble into aggregates to shield hydrophobic regions
from water – these are micelles and phospholipid bilayers. These are the major
component of cell membranes.
Steroids are lipids with a carbon skeleton containing a characteristic 4-membered
ring. Steroids differ by the functional groups attached to the rings. For example,
cholesterol is a common component of animal plasma membranes and is a hormone
precursor (sex hormones).
Figure 7 – Comparison of fats, phospholipids and steroid hormone structure. Source:
https://upload.wikimedia.org/wikipedia/commons/2/20/Common_lipids_lmaps.png
Saturated and unsaturated fats
A fatty acid with no double bonds has the maximum number of hydrogen atoms
attached to the carbon skeleton as possible and is said to be ‘saturated’. Fatty acids
with double bonds are unsaturated (i.e. all binding sites are not saturated with H), and
they have kinks where the double bond occurs. These kinks prevent unsaturated fats
from packing together closely enough at room temperature to solidify. Saturated fats
are made of saturated fatty acids (eg: animal fats) and they are solid at room
temperature (eg: butter). Plant and fish fats are typically unsaturated and are liquid at
room temperature (eg: oils). (Figure 8)
Fats are not ‘bad’ – they are a compact energy store (2x more than CHOs) for many
organisms. For example, animals need to carry energy stores around and use fat to
cushion vital organs, and plants need compact energy stores in seeds.

Figure 8 – Saturated and unsaturated fatty acids. Source:

https://upload.wikimedia.org/wikipedia/commons/c/c8/221_Fatty_Acids_Shapes-01.jpg
Macromolecules – Proteins

Proteins are composed of polypeptides in a highly specific sequence and

conformation. The function of proteins is dependent on the structure. Proteins are
absolutely essential for everything that cells do – even viruses need proteins – they
compose 50% of a cell’s dry weight. They are used for a large variety of purposes:
structural, movement, signalling, defence, catalysing biochemical reactions. Proteins
are the most structurally complex molecules known with unique 3D
shapes/conformations, but are all fundamentally polymers made of 20 basic amino
acid monomers. Polymers of amino acids are called polypeptides – a protein
consists of one or more polypeptides. They tend to be either globular or fibrous.
Amino acids (AA) are the building blocks of proteins. The general structure of proteins
is a chiral carbon (called the α-carbon), amino group, carboxyl group, hydrogen, and
an R group. Organisms typically only use one form of amino acid, the L form. In typical
physiological conditions, the amino and carboxyl group are ionised. The R group
differs with the AA and gives each AA its unique properties. The R group is typically
called the side chain and it varies in complexity. The R groups can be classified
according to polarity and charge (you will need to be able to learn these).

Polypeptides and proteins

Amino acids are covalently linked by dehydration reactions to form polypeptides. The
carboxyl group of one AA reacts with the amino group from another AA, forming a
peptide bond, and displacing water. This reaction is repeated over and over, yielding
a polypeptide. The finished polypeptide has a polarity due to an exposed amino group
(N-terminus) at one end and carboxyl group (C-terminus) at the other.
A polypeptide is a chain of covalently linked amino acids. A protein consists of more
than one polypeptide chain, often with other groups associated.

Protein conformation
The sequence of AA determines the structure of the polypeptide chain. There are four
levels of protein structure: primary, secondary, tertiary, and quaternary (Figure 9).
 Primary structure
Primary conformation of proteins is the sequence of AA in the protein. Primary
structure was first determined by hydrolysis and chromatography. The precise primary
structure of proteins is determined by inherited genetic information. A slight change in
primary structure can drastically affect protein function (eg: sickle cell disease).
 Secondary structure
The two types of secondary proteins structure and coiled and folded regions of the
polypeptide chain. The coils and folds are the result of interactions (H-bonds) between
regions of the polypeptide backbone – particularly between the C=O and N-H.
Individual H-bonds are weak but repeated bonds over a region is enough to support a
particular shape. One such shape is the α helix, a coil formed by H-bonding between
every 4th AA (only particular AAs are able). The α helix is found in globular proteins as
well as fibrous proteins (eg: keratin). Another such shape is the β pleated sheet, a
sheet formed by polypeptide chains lying side by side, H-bonds forming between the
C=O and N-H. The β pleated sheet is also found in globular proteins as well as fibrous
proteins (eg: spider silk).

 Tertiary structure
Weak interactions and disulphide bridges contribute to the tertiary structure of
proteins. These are the irregular contortions from interactions between side chains.
Many types of interactions between the side chain can cause these contortions and
bring two parts of the polypeptide into proximity. Hydrophobic interactions occur
when hydrophobic side chains run away from water and bury themselves inside the
protein, held together by dispersion forces. Hydrogen bonds form between polar side
chains. Ionic bonds form between oppositely charged side chains. Changes in pH
typically affect the charges of ionic side chains, hence tertiary structure is most
affected by pH change. Disulfide bridges form when a strong covalent –S–S– bond
forms between two cysteine side chains (which have –SH).

 Quaternary structure
Quaternary structure is the overall protein structure resulting from aggregation of
polypeptide subunits. For example, collagen and haemoglobin are examples of the
quaternary structure of proteins. While most proteins do have a quaternary structure,
not all do (eg: lysozyme).

A protein is mainly directed to fold by its primary structure, however external factors
can also influence its final conformation such as pH, temperature, salt concentration,
etc. Inside the cell, large proteinaceous structures called chaperone proteins exist
that can help proteins fold. Chaperones work by providing a newly developing protein
a secluded environment away from harmful influences so that it can fold properly. The
chaperone then releases the properly-folded protein for use.
Figure 9 – Protein structure. Source:
https://upload.wikimedia.org/wikipedia/commons/thumb/c/c9/Main_protein_structure_levels_en.svg/20
00px-Main_protein_structure_levels_en.svg.png
Denaturation
When some aspects of the protein’s environment changes, it may unfold and lose its
native 3D conformation; a process called denaturation (Figure 10). Denaturants
include heat, organic solvents (turns the protein inside out) and reducing agents, all of
which result in biologically inactive proteins.

Figure 10 – Process of denaturation by heat. Source:

https://upload.wikimedia.org/wikipedia/commons/thumb/1/1d/Process_of_Denaturation.svg/2000px-
Process_of_Denaturation.svg.png
Macromolecules – Nucleic acids

Two types of nucleic acid exist: ribonucleic acid and deoxyribonucleic acid. These
molecules allow cells to transmit genetic information to offspring. DNA directs RNA
synthesis, and RNA directs polypeptide synthesis. This is termed the ‘central dogma’
in biology.
Nucleic acids are polymers of nucleotides which consist of 3 parts: a pentose (5C
sugar), a phosphate group, and a nitrogenous base. The nitrogenous base determines
the type of nucleic acid the molecule will be and the sugar-phosphate molecule
interaction provides the ‘backbone’ of the structure (Figure 11).
The nitrogenous base of nucleic acids is either a pyrimidine or purine. Pyrimidines
include cytosine, thymine and uracil, and purines include adenine and guanine. Note
that thymine is found exclusively in DNA, and is replaced by uracil in RNA.
The pentose sugar is ribose in RNA and deoxyribose in DNA (because 2C is missing
an O). (*NB the numbering system on the pentose – take particular notice of the 3’
and 5’ carbons)
A nucleoside is a nitrogenous base and pentose, whereas a nucleotide is a
nucleoside and a phosphate (nucleoside monophosphate). Nucleotide triphosphates
(dNTPs) are polymerised by the enzyme DNA polymerase into a polynucleotide,
joining the nucleotides with phosphodiester bonds (actually made up of the 2 ester
bonds flanking the phosphate plus the phosphate itself).

Figure 11 – Structure of nucleic acids. Source:

https://upload.wikimedia.org/wikipedia/commons/thumb/e/e2/Nucleotides_1.svg/2000px-
Nucleotides_1.svg.png
Three-dimensional structure of DNA
RNA is a single polynucleotide chain and DNA is composed of 2 chains of
polynucleotides weakly bonded with H bonds. The sugar-phosphate backbone is on
the outside of the structure and the nitrogenous bases are on the inside. The bases
hydrogen bond with each other, but only certain bases are compatible. Guanine forms
3 H-bonds with C, and adenine forms 2 H-bonds with thymine or uracil.
Because guanine and cytosine and adenine and thymine go together, we can deduce
the sequence of one polynucleotide strand using the other strand. The two strands are
said to be complementary, which allows the precise copying of DNA that is so
important for heredity. Importantly, the two polynucleotide strands are antiparallel –
note that the 5’  3’ directions of the two strands run in opposite directions (Figure
12).
Closely related species and organisms share more DNA than distantly related species
or organisms and this DNA structure can be used to document the hereditary
background of an organism.

Figure 12 – 3D structure of RNA and DNA.

Source: https://upload.wikimedia.org/wikipedia/commons/thumb/3/37/Difference_DNA_RNA-
EN.svg/2000px-Difference_DNA_RNA-EN.svg.png
Introduction to metabolism

Metabolism is the sum of all an organism’s chemical reactions. This involves the
management of the material (molecular) and energy resources of the cell. Metabolic
pathways can either be catabolic (breaking down complex molecules to simpler ones
to release energy) or anabolic (consuming energy to build complex molecules from
simpler ones). These pathways intersect so that the energy provided by catabolism
can be used to drive anabolism and are therefore said to be coupled (Figure 13).

Figure 13 – Catabolic and anabolic pathways. Source:

https://upload.wikimedia.org/wikipedia/commons/8/80/Catabolism,_energy_carriers_and_anabolism.p
ng

Energy is the capacity to do work and is fundamental to all metabolic processes. Cells
transform energy from one from into another:
 Kinetic energy – energy from movement.
 Potential energy – energy that matter has due to its location or structure (eg:
water behind a dam).
 Energy transformation – use of potential energy to create kinetic energy. This
is best represented by considering the potential (chemical) energy of food being
used to power climb up to the top of a slide (kinetic), and then releasing kinetic
energy as motion as a person slides down.
Laws of thermodynamics
Energy transformations in living things are subject to the two laws of thermodynamics.

 First law of thermodynamics (principle of the conservation of energy)

Energy of the universe is constant; energy can be transferred and transformed but
cannot be created or destroyed. For example, plants convert light energy to chemical
energy.

 Second law of thermodynamics

Every energy transfer or transformation increases the entropy (a measure of disorder
or randomness) of the universe. In performing work, living cells unavoidably convert
organised forms of energy (eg: chemical energy in food) into heat (and other forms of
energy). Heat is basically the energy of random molecular motion.
Conversion to heat is the fate of all energy. Heat can only be put to work to warm
something up (eg: a body of matter like an organism). The energy difference between
food and waste molecules can be accounted for by heat released during catabolism.
Entropy is an interesting concept but it does not practically help us to determine if
certain (spontaneous) events will happen without outside help.

Gibbs free energy

The energy available to do work is referred to as Gibbs free energy. That is, the
portion of a system’s total energy that can perform work when temperature is uniform
throughout the system. The more free energy a system has, the more unstable the
system is. This means that it has a tendency to change spontaneously to a more stable
state, and therefore a greater capacity to do work. When the system does work, it uses
the free energy and becomes more stable, with less free energy. In any spontaneous
process, the free energy decreases.
Free energy can be symbolised by the equation G = H – TS, where:
 Free energy (G) is the total energy of the system
 Enthalpy (H) is the total heat of the system
 Entropy (S) is the disorder of the system
 Temperature (T) is the average heat of the system measured in Kelvin
For a process to occur spontaneously, necessarily ΔG < 0 where ΔG is the Gfinal –
Ginitial.
To determine if a metabolic reaction is spontaneous, we use ΔG. Exergonic
reactions proceed with a net release of energy, because the chemical mixture loses
energy, therefore ΔG <0. Exergonic reactions are spontaneous. Cellular respiration
releases energy from food, converting it to CO2 and H2O and transforming the released
energy into ATP. This reaction releases energy because the reaction proceeds from
a stable (more ordered) state to an unstable (more disordered) state.
Endergonic reactions absorb free energy from the surroundings, therefore ΔG >0.
Endergonic reactions are non-spontaneous. For example, photosynthesis uses light
energy and stores it in chemical structures like sugar molecules.

Energy and ATP

Living things are open systems, meaning that energy and matter continuously flow into
and out of living things. This makes sure that metabolic pathways always act either
downhill or uphill, and are never at a standstill. Energy releasing reactions can power
the energy requiring reactions. In living cells, a molecule called adenosine
triphosphate (ATP) is primarily responsible for this coupling.
ATP has the nitrogenous base adenine bound to ribose, and 3 phosphates. The bonds
between the phosphates in ATP can be broken by hydrolysis to yield energy.
Hydrolysis of the terminal phosphate yields ADP + Pi, which releases energy.
ATP has 4 negative charges in close proximity, whereas ADP only has 3 negative
charges. This means that ADP is a more stable molecule than ATP. The hydrolysis
reaction transitions a molecule of ATP (more unstable state) to ADP (more stable
state), increasing the stability of the molecule and releasing energy in the process. Pi
is also more stable on its own than when it is incorporated into ATP because of
increased electron delocalisation (increased resonance forms).

ATP couples catabolic reactions that release energy to anabolic reactions that require
energy. By giving a phosphate group to a molecule to make it more reactive the
molecule is said to be phosphorylated. Cells use energy from catabolism (eg:
respiration) to combine Pi + ADP to form ATP + H2O – this is the basis of energy
coupling.
Enzymes

Enzymes are catalytic proteins that catalyse reactions without themselves being
consumed in the reaction. All chemical reactions are based on breaking bonds
between the reactants (or substrates) and forming new bonds to make products.
We can think of exergonic and endergonic reactions in terms of the free energy graph
(Figure 14). At the start are the reactants, whose bonds need to be broken by input of
an amount of energy. This energy input could come from heat absorbed from
surroundings. Thermal energy increases the proportion of molecular collisions that
have sufficient energy to exceed the activation energy (EA) – the free energy required
to break the bonds of the reactants.
After absorbing enough energy, molecules pass through a transition state and then
settle into the new bonding arrangements. If product free energy is lower than reactant
free energy, the reaction is exergonic (unstable to stable), and vice versa.

Figure 14 – Exergonic and endergonic reactions.

Sources:
https://upload.wikimedia.org/wikipedia/commons/thumb/9/98/Exergonic_Reaction.svg/2000px-
Exergonic_Reaction.svg.png
https://upload.wikimedia.org/wikipedia/commons/thumb/4/45/Endergonic_Reaction.svg/2000px-
Endergonic_Reaction.svg.png

Enzymes and chemical reactions

Enzymes speed up reactions by lowering the EA. In many cases, EA is large enough
to prevent spontaneous reactions happening at room temperature. For example, petrol
igniting is spontaneous but luckily it does not happen at room temperature because
the EA required to initiate the spontaneous reaction is too high.
Reactions can be made to go faster by adding more heat or by lowering the EA, which
is what enzymes do. Less free energy input is required to make the reaction go ahead.
However, enzymes only change EA, not free energy (ΔG).
Enzymes have specific 3D conformations and therefore recognise specific molecules
due to their specific shape. This is why enzymes are said to be substrate-specific. The
active site is the part of the whole enzyme protein that actually binds the substrate and
it is the enzyme’s catalytic centre. Specificity at the active site is determined by the
shape of both enzyme and substrate and the  residues that project into the space.
There are two models of enzyme-substrate binding: lock and key (Figure 15) and
induced fit (Figure 16). Substrates interact with the active site by weak interactions
(eg: H-bonds, ionic bonds, dispersion, etc.).

Figure 15 – Lock and key model of enzyme action.

Source:
https://upload.wikimedia.org/wikipedia/commons/9/9f/Lock_and_key.png

Figure 16 – Induced fit model of enzyme action.

Source:
https://upload.wikimedia.org/wikipedia/commons/thumb/2/24/Induced_fit_diagram.svg/2000px-
Induced_fit_diagram.svg.png
Enzymes lower EA by:
 Bringing substrates into the correct orientation with respect to each other for a
reaction to occur
 Stressing and bending molecules, making chemical bonds more likely to break
 Providing conducive microenvironment eg: low pH with acidic side chains
 Directing participation with temporary covalent bonding between substrate and
enzyme side chains
Enzymes remain unchanged after catalysing a reaction. Small amounts of enzyme
can induce large changes. Enzymes can catalyse reactions in both directions,
however this does not often happen in most biological systems as the product is
usually rapidly removed, and so the process always goes towards the product.

Rate of catalysis is determined by:

 The amount of enzyme available. More enzyme = more catalysis.
 How fast the enzyme is. Fast enzymes = more catalysis.
 How much substrate is available. More substrate = more accessing enzymes
at any one time.
Enzymes (cont.)
Adding more substrate to a fixed concentration of enzyme will only make the reaction
go faster up to a limit. Once all of the available enzymes have their active sites filled
they are said to be ‘saturated’. The point at which enzyme saturation is reached is
also known as the maximum velocity of the enzyme, Vmax (Figure 17). The Km of an
enzyme is the substrate concentration when the rate is half the maximum, or Vmax/2.
Note that Vmax is directly proportional to [E].

Figure 17 – Enzyme saturation curve.

Source:
https://upload.wikimedia.org/wikipedia/commons/2/26/MM_curve.png

Taking a double reciprocal of the same data in the above graph, gives a linear
relationship, the Lineweaver-Burk plot. This provides a more precise way to determine
Vmax and Km.
* Vmax is determined by the point where the line crosses the 1/ Vi = 0 axis (so the [S]
is infinite).
* Note that the magnitude represented by the data points in this plot decrease from
lower left to upper right.
* Km equals Vmax times the slope of line. This is easily determined from the intercept
on the x axis.

Being proteins, enzymes are affected by their environment. Increasing temperature

increases collisions and makes chemical processes go faster up to a point. The speed
of these processes drops rapidly once an enzyme denatures. An optimal temperature
exists which allows for the greatest number of collisions without denaturing the
enzyme. pH can also denature enzymes and proteins. Enzymes have an optimal pH
(typically 6-8), which varies depending on where they are found.
Other factors that affect enzyme action
Cofactors are non-protein helpers required for catalytic activity. These can be bound
tightly to the enzyme, or associate loosely with the substrates (eg: metal ions like zinc).
Coenzymes are organic cofactors (eg: vitamins).
Inhibitors are agents that selectively inhibit the action of specific enzymes. They can
bind covalently (irreversible binding) or by weak interactions (reversible binding).
Inhibitors irreversibly bind and dissociate very slowly from enzymes. A good example
of this is penicillin, which binds covalently to modify transpeptidase thus preventing
the synthesis of bacterial cell walls. Examples of inhibitors that reversibly bind include:
 Competitive inhibitors often resemble the normal substrate and compete for
access to the active site. For example, statins reduce high cholesterol levels by
inhibiting a key enzyme in cholesterol biosynthesis. (Figure 18)
 Uncompetitive inhibitors bind to the enzyme-substrate complex so that no
product is formed. Glycophosphate is an uncompetitive inhibitor in the pathway
for aromatic amino acids.
 Non-competitive inhibitors impede enzymes by binding to another part of the
enzyme (non-active site) causing a conformation chang and resulting in a
reduction in turnover. Lead reacts with sulfhydryl groups in enzymes.
 Allosteric inhibitors do not obey the Michaelis-Menten model: binding at one
site on the enzyme can lead to conformational change and altered binding
properties on other sites of the enzyme. This type of inhibitor is critical for
metabolic control.

Figure 18 – Competitive inhibition. Source:

https://upload.wikimedia.org/wikipedia/commons/thumb/f/fb/Comp_inhib.svg/2000px-
Comp_inhib.svg.png
Control of metabolism
Cells tightly control metabolic pathways by activating and inhibiting specific enzymes
at specific times. This is done by altering the number and activity of enzymes and by
feedback inhibition.
Metabolic control often depends on allosteric regulation – regulating enzyme activity
by inducing conformation changes. Allosterically regulated enzymes are typically
composed of multiple identical subunits where the enzyme typically flip-flops between
two states: active and inactive. The allosteric site is where the regulator binds,
inducing a conformation that is either active (activator) or inactive (inhibitor). A good
example of this is the enzymes that regulate production of ATP.
Cooperativity is another form of enzymatic regulation, where the binding of a
substrate to a multi-component enzyme induces conformational changes in the other
subunits so that all subunits are more receptive to substrate. This mechanism
amplifies the response of enzymes to substrates, as seen by its sigmoidal kinetics
curve.
Revision questions
1. What is the difference between a molecule and an atom?
2. What are some of the characteristics of electrons that differentiate them from
other subatomic particles?
3. Describe how water contributes to the fitness of the environment to support
life
4. What is the difference between solute, solvent and solution?
5. Using the bicarbonate buffer system as an example, explain how buffers
work.
6. What is the difference between heat and temperature?
7. What is the pH scale? Explain how it is calculated.
8. Provide some examples of different types of carbohydrates, lipids and
proteins.
9. What is the difference between RNA and DNA?
10. What are some of the characteristics of enzymes that contribute to their ability
to regulate metabolic reactions?