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BIOMOLECULES

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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

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• 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 of building blocks

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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 maintained by non-covalent forces, not by covalent bonds

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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 charged (ionic bonds)

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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

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 blocksThe 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 U

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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• 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 + wqP= 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 workApplication 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 re-established

• 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’0

eq

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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