BIOMOLECULES
&
BIOENERGETICS
In this house, we obey the laws of thermodynamics!
Homer Simpson
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THE CHEMICAL COMPONENTS OF A CELL
There are 89 naturally occurring elements
However, only 15 are present in all living things, and a further
8–10 are only found in particular organisms
More than 99% of the atoms in animals’ bodies are
accounted for by just four elements— hydrogen (H), oxygen
(O), carbon (C) and nitrogen (N)
Hydrogen and oxygen are the constituents of water, which
alone makes up 60–70% of cell mass
Together with carbon and nitrogen, hydrogen and oxygen
are also the major constituents of the organic compounds on
which most living processes depend
Carbon is the biggest contributor to the dry weight of the
body
The prevalence of C is due to its unparalleled versatility in
forming stable covalent bonds by electron-pair sharing 2
• Carbon can form as many as four such bonds by sharing each
of the four electrons in its outer shell with electrons
contributed by other atoms
• Carbon has the potential to form an enormous variety of
linear, branched and cyclic compounds
• In addition to C, H, O and N, many biomolecules also contain
sulfur (S) or phosphorus (P)
• The above six macroelements are essential for all organisms
• A second biologically important group of elements, which
together represent only about 0.5% of the body mass, are
present almost exclusively in the form of inorganic ions
• This group includes the alkali metals sodium (Na) and
potassium (K), and the alkaline earth metals
magnesium(Mg) and calcium(Ca)
• The halogen chlorine (Cl) is also always ionized in the cell
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Composition of the human body
after the removal
of water
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• All other elements important for life are present in such small
quantities that they are referred to as trace elements
• These include transition metals such as iron (Fe), zinc (Zn),
copper (Cu), cobalt (Co) and manganese (Mn)
• A few nonmetals, such as iodine (I) and selenium (Se), can also
be classified as essential trace elements
THE BIOMOLECULAR HIERARCHY
• The major precursors for the formation of biomolecules are
water, carbon dioxide, and three inorganic nitrogen
compounds-ammonium (NH4+),nitrate (NO3-) and dinitrogen
(N2)
• Metabolic processes assimilate and transform these inorganic
precursors through more complex levels of biomolecular order
• In the first step, precursors are converted to metabolites,
simple organic compounds that are intermediates in cellular
energy transformation and in the biosynthesis of various sets
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• The major building blocks are amino acids, sugars,
nucleotides, fatty acids and glycerol
• By covalent linkage of these building blocks, the
macromolecules are constructed: proteins, polysaccharides,
polynucleotides (DNA and RNA) and lipids
• Strictly speaking, lipids contain relatively few building blocks
and are therefore not really polymeric like other
macromolecules; however, they are important contributors to
higher levels of complexity
• Interactions among macromolecules lead to the next level of
structural organization, supramolecular complexes
• Supramolecular assemblies include multifunctional enzyme
complexes, ribosomes, chromosomes, cytoskeletal elements
and viruses
• Supramolecular assemblies are an interesting contrast to their
components because their structural integrity is maintained7by
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NON-COVALENT INTERACTIONS IN BIOMOLECULES
• In aqueous solutions, covalent bonds are 10-100 times
stronger than the other attractive forces between atoms,
allowing their connections to define the boundaries of one
molecule from another
• But much of biology depends on the specific binding of
different molecules to each other; this binding is mediated
by a group of non-covalent attractions that are individually
quite weak, but whose energies can sum to create an
effective force between two separate molecules
• These non-covalent interactions are mainly of four types:
1. Electrostatic attractions: These result from the attractive
forces between oppositely charged atoms; they are quite
strong in the absence of water
• They readily form between permanent dipoles, but are
greatest when the two atoms involved are fully charged12
2. Hydrogen bonds: represent a special form of polar
interaction in which an electropositive hydrogen atom is
partially shared by two electronegative atoms
• Water weakens these bonds by forming competing
hydrogen-bond interactions with the involved molecule
3. van der Waals attractions:
• The electron cloud around any non-polar atom will
fluctuate, producing a flickering dipole
• Such dipoles will transiently induce an oppositely
polarized flickering dipole in a nearby atom
• This interaction generates a very weak attraction
between atoms. But since many atoms can be
simultaneously in contact when two surfaces fit closely,
the net result is often significant
• Water does not weaken van der Waals attractions
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4. Hydrophobic force: is not, strictly speaking, a bond at all
• However, a very important force is caused by the pushing
of non-polar surfaces out of the hydrogen-bonded water
network, where they would otherwise physically interfere
with the highly favorable interactions between water
molecules
• Bringing any two non-polar surfaces together reduces
their contact with water; in this sense, the force is
nonspecific
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THE USE OF ENERGY BY CELLS
Cells create and maintain order, in a universe that is
tending always to greater disorder
To create this order, the cells in a living organism must
perform a never-ending stream of chemical reactions
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Two opposing streams of chemical reactions occur in cells:
 The catabolic pathways break down foodstuffs into smaller
molecules, thereby generating both a useful form of energy
for the cell and some of the small molecules that the cell
needs as building blocks
The anabolic, or biosynthetic pathways use the energy
harnessed by catabolism to drive the synthesis of the many
other molecules that form the cell
• Together these two sets of reactions constitute the
metabolism of the cell
• For a cell to grow or to make a new cell in its own image, it
must take in free energy from the environment, as well as
raw materials, to drive the necessary synthetic reactions
• This consumption of free energy is fundamental to life; when
it stops, a cell decays towards chemical equilibrium and soon
dies
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Bioenergetics is a discipline that quantifies the energy
transfer that occurs in living cells and describes the nature of
the chemical reactions that bring about this transfer
• Biological energy transductions are governed by the same
rules that govern other natural processes –The Laws of
Thermodynamics
• Thermodynamics is a study of energy transformations
• It aims to describe and relate – in relatively simple
mathematical terms – the physical properties of systems of
energy and matter
• In studying thermodynamics, there are certain terms that
one has to be familiar with:
 A system is defined as that part of the universe chosen for
study. The surroundings are simply the entire universe
excluding the system. The system and surroundings are
separated from each other by a boundary
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 A system is at any time in a certain thermodynamic state or
condition of existence (which types of molecule are present
and in what amount, the temperature, the pressure, etc.)
 A system is said to be closed if it can exchange heat with the
surroundings but not matter. If matter can be exchanged
between the system and the surrounding, the system is said
to be open. A living system is an open system
 An isolated system is one in which the boundary permits
neither matter nor energy to pass through
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A bomb calorimeter is used to measure the heat given off in
the oxidation of a combustible substance like food
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• What is being measured in the bomb calorimeter is heat
 Heat, or thermal energy, q, is a form of kinetic energy;
that is, energy arising from motion
• It is the change in energy of a system that results from a
temperature difference between it and the surroundings
• Heat is said to flow from a region of higher temperature,
where the average speed of molecular motion is greater,
to one of lower temperature
• Thermodynamics :thermo – heat , dynamics –movement
• The Zeroth Law of Thermodynamics: if A is in thermal
equilibrium with B, and B is in equilibrium with object C,
then C is also in thermal equilibrium with A
• This law justifies the concept of temperature and the use
of thermometers
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The First Law of Thermodynamics is a conservation law:
energy can be changed from one form to another, but in all
its transformations energy is neither created nor destroyed
The energy of a system plus its surroundings is constant in
time
To see more clearly how the First Law operates, internal
energy and work have to be defined
As with heat, both internal energy and work are measured
in units of joules (or calories)
The internal energy, U, is the energy within a system
It represents only those kinds of energy that can be
modified by a chemical process – translational, vibrational,
rotational, bonding, and non-bonding energies
The internal energy of a system cannot be measured
directly; it is calculated from other measured properties
Moreover, it is not U that is measured but a change in U22
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Both heat and work are forms of energy that can be
transferred across the boundary of a system
The difference between the two is that work, w, is the
equivalent of a force (e.g. gravity) acting through the
displacement of an object, while heat is the transfer of
energy owing to a temperature difference
Work involves energy transfer through the non-random
movement of particles, and heat, through the random
movement of particles
ΔU=q + w
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The above equation is an expression of the first law: loss or
gain of internal energy through heat transfer and work
The internal energy of a system will increase either by
transferring heat to it or by doing work on it
The internal energy of a system will decrease either by
transferring heat from it or by doing work on the surrounding
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The work achieved by movement through some
distance caused by the application of a force is called
mechanical work
In biochemical systems, work is often concerned with the
pressure and the volume of the system under study
And since biochemical reactions are assumed to take
place at constant pressure, the work done by the system
is an expression of the change in volume of the system
 the usual expression, w=-FΔS is replaced by w=-PΔV
 the negative signs because work is done against an
opposing force or pressure
Enthalpy, H, is the internal energy of a system plus the
product of its volume and the external pressure exerted
on the system
H=U+pV
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qP represents heat transferred at constant pressure
ΔU=qP + w
qP= ΔU-w
When the pressure is constant and the system expands from
state 1 to state 2, the system does work on the surroundings.
If the only type of work is pV-work :
qP= U2-U1+p(v2-v1)=ΔU+p Δv
Differences in enthalpy can be calculated as:
ΔH= Δ(U+pV)
ΔH = ΔU+ pΔV+VΔp
Assuming the external pressure is constant:
ΔH = ΔU+ pΔV
Substituting for ΔU:
ΔH=(qP – pΔv)+ pΔV
ΔH=qP
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• Heat absorbed by a process at constant volume measures ΔU,
and the heat absorbed by a process at constant pressure
measures ΔH
• These quantities of heat will in general differ, because in a
change at constant pressure some energy exchange will be
involved in the work done in the change of volume of the
system
• The thermochemistry of biological systems is almost always
concerned with ΔH, since most natural biochemical processes
occur under conditions more nearly approaching constant
pressure than constant volume
• However, since most such processes occur in liquids or solids
rather than in gases, the volume changes are small. To a good
approximation, the difference between ΔH and ΔU is
neglected in biochemistry and we simply talk about the
energy change accompanying a given reaction
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• A process for which the change in enthalpy is negative is
called exothermic, as heat is let out of the system into the
surroundings
• A process for which the change in enthalpy is positive is called
endothermic, as heat is let into the system from the
surroundings
• Changes in enthalpy (and other functions) are generally given
for processes occurring under a standard set of conditions
• The standard state is usually defined as one mole of a pure
substance at 298.15 K (25 0C) and 1 atm
• Simply put, The First Law states that, if a system does work,
w makes a negative contribution to ΔH; the system loses
energy. This implies that not even the most sophisticated
known “machine” – the human body, as far as known – can do
work without an energy source (that is food)
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The Second Law of Thermodynamics
The First Law of Thermodynamics deals with the energy
changes that accompany chemical and physical processes; but
it does not indicate the direction in which changes will occur
Chemical and physical processes occur in such a way that,
matter and energy - given enough time - tend to achieve a
state of equilibrium ( a state of no change)
The equilibrium state is also the state of higher probability; this
is the reason why it is favored and processes tend to move
toward it
The Second Law of Thermodynamics states that systems will
change spontaneously from states of lower probability to states
of higher probability
Both matter and energy have the tendency to go from being
concentrated ( ordered) to being distributed (disordered) –
spontaneously (irreversibly)
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• The Second Law can be restated: the universe constantly
changes so as to become more disordered
 But what is it that determines the position of equilibrium and
the spontaneous (irreversible) process that leads to it?
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• Entropy, S, is the measure of the probability or equivalently,
the degree of disorder of a state
• When heat is transferred to/from a system, part of it is used
to do work and the rest is released as waste heat
• The waste heat increases the number of different
arrangements that the molecules in the system can have; it
increases their entropy
• A fixed quantity of heat energy has a greater disordering
effect at low temperature than at high temperature; and
entropy and heat are related as:
S= ΔH/T; where T is temperature in degree Kelvin
• Change in the entropy of the universe (Δsuniv) is the measure
of the spontaneity of a reaction
• Δsuniv has two components: change in the entropy of the
system ( Δssys) and the surrounding (Δssurr)
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• Instead of measuring two values of ΔS and adding them up
to know whether a reaction is spontaneous or not, a
function has been devised that allows one to deduce Δsuniv
without having to calculate Δssurr
Δsuniv = Δssurr+Δssys
• Δssurr is the change in entropy in the surrounding caused by
the heat released from the system at a certain temperature
Δssurr = -ΔHsys/ T (the negative sign because heat is
released)
Δsuniv = -ΔHsys / T + Δssys
• Multiplying both sides by -T:
-T Δsuniv = ΔHsys -T Δssys
• -T Δsuniv is referred to as the Gibbs free energy change (ΔG)
ΔG = ΔHsys -T Δssys
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• Free energy is energy that is available in a form that can be
used to do work at constant temperature and pressure
• A spontaneous reaction loses some free energy in doing
work and the ΔG would be negative
• Since T is always positive, Δsuniv has to be negative in order
to make ΔG negative
 This is in accordance with the fundamental measure of
spontaneity, that is, a positive Δsuniv
• Simply put, the Second law states that living systems can
not change a given amount of heat (ΔHsys ) into an
equivalent amount of work. They use part of it (ΔG) to do
work and the rest is spent on increasing ΔSuniv
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• Living cells-by surviving, growing, and forming complex
organisms-are generating order and thus might appear to
defy the second law of thermodynamics; how is this
possible?
• The answer is that a cell is not an isolated system: it takes in
energy from its environment in the form of food, or as
photons from the sun (or even, as in some chemosynthetic
bacteria, from inorganic molecules alone), and it then uses
this energy to generate order within itself
• In the course of the chemical reactions that generate order,
the cell converts part of the energy it uses into waste heat
• This heat is discharged into the cell's environment and
disorders it, so that the total entropy-that of the cell plus its
surroundings-increases, as demanded by the laws of
thermodynamics
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• The Third Law of Thermodynamics states: any system not at
absolute 0 (-273 K) has some amount of T Δssys which is the
property of the system at that temperature
• Simply put, the Third Law states that unless a system is at
absolute 0, some of its energy would be unavailable to do work
Application of ΔG to biochemical reactions
• If ΔG for a reaction is negative, the reaction can proceed in that
direction
• If positive, the reverse reaction has a negative ΔG so that the
reaction will occur in the opposite direction
• If ΔG is zero, the reaction proceeds in neither direction and is
said to be in a state of equilibrium
• Consider the reaction in which substance A is converted to
substance B and the reaction is at equilibrium:
A B
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• If the concentration of either A or B changes, the equilibrium
would be displaced and the reaction would proceed in the
specific direction to maintain the equilibrium
• For example, increasing the concentration of A will cause the
reaction to move towards the right, producing substance B and
lowering the concentration of A, until the equilibrium is reestablished
• Since ΔG determines the direction in which a reaction
proceeds, it follows that the value ΔG must depend on reactant
concentrations
• Mass action ratio is the ratio between the concentration of
product (s) and reactant (s)
• The mass action ratio at equilibrium is known as the
equilibrium constant, Keq and it indicates the position of
equilibrium
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• The relation between change in free energy and concentration
of reactants and products is given as:
ΔG= ΔG0+ RT ln [B]/[A]
• Every reaction has its own standard free energy change value
(ΔG0 ) defined at 250c (298K) and 1 atm with all solutes at 1
molar concentration
• At equilibrium ΔG = 0 and [B]/[A] is Keq
ΔG0= -RT ln Keq
• Standard free energy change is the energy that drives the
reaction to equilibrium under the stated conditions
• Summarizing, the sign and magnitude of ΔG (and hence the
direction of a reaction) depends on two factors: the actual
concentrations of the substrates and products and the value of
the constant ΔG°
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• The value of ΔG° alone is, therefore, not sufficient to determine
the direction in which a reaction will proceed in a living cell
• Even if a reaction is endergonic by virtue of its positive ΔG° , it
could be made exergonic by decreasing the mass action ratio,
making RT ln [B]/[A] more negative
• This way, ΔG (the actual free energy change), would be made
negative; it is ΔG that determines the spontaneity of a reaction
 For biochemical reactions , there are additional components of
the standard state:
 The concentration of H+ is 10-7 molar (pH=7), the
concentration of water is 55.5 molar and the concentration of
Mg+2 is 1 mM (Mg+2 is essential for stabilizing ATP
molecules; and ATP is a major player in biochemical reactions)
 Transformed standard free energy change (ΔG’0 ) is the free
energy change under the new set of conditions = -RT ln K’370eq
Gibbs free energy and the coupling of biochemical reactions
• Many of the reactions necessary to keep cells and organisms
alive must run against their thermodynamic potential
• A thermodynamically unfavorable (endergonic) reaction can be
driven in the forward direction by coupling it to a highly
exergonic reaction
• There are two ways of coupling reactions:
1. Coupling-in-series: when the product of one reaction is the
substrate for the next reaction and so on
• As previously stated, a positive ΔG’0 can be made to give a
negative ΔG by adjusting the concentrations of substrates
and products
• The concentrations of substrates and products can be
adjusted by making them parts of a series of reactions
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• For example, in the reactions S→ A → B → P, A → B can have
a positive ΔG’0
• A → B can be made exergonic, if the concentration of A is
kept high enough by making the Keq of S→ A very large
(concentration of A would be high) and the the Keq of B→ P
very high (concentration of B would be low)
• The principle of coupling-in-series underlies all biochemical
pathways, such as glycolysis, …
2. Coupling in parallel
• Adjusting concentrations may not always be an option: very
high/low concentrations may not be compatible with
physiology
• For example, glucose + phosphate → glucose-6-phosphate
has a positive ΔG’0
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• Large increases in the glucose concentration would lead to
unwanted side reactions a dangerous osmotic effect
• Large decreases in the concentration of glucose-6-phosphate,
on the other hand, would affect the many metabolic processes
that require glucose-6-phosphate
• The problem is solved by coupling the reaction with the
hydrolysis of ATP (which is highly exergonic)
• Since free energy changes are additive, the formation of
glucose-6-phosphate from glucose, using the phosphate
derived from ATP is made possible
• The reactions in this case do not occur at tandem; they occur at
the same time
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Glycolysis
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