INTRODUCTION: PowerPoint

Living organisms are composed of "inert" molecules whose interactions give the organism special properties:

1) They exchange energy and matter with their surroundings; this allows them to go against the tendency of systems to become more disordered (2nd Law of Thermodynamics: organism ---> more ordered but overall the universe ---> more disordered). Ultimate energy source is sunlight.

2) They transmit and use energy in chemical form since they must function at constant T, P, and V and unlike steam or internal combustion engines can't do PV work. They store energy in the form of chemical bonds, e.g. ATP. Energy is released by breaking high energy bonds and forming lower energy bonds.

3) Living organisms replicate themselves accurately.

In Biochemistry we study the chemical reactions of these processes

ELEMENTS: Relatively few of the 92 naturally occuring elements are used extensively in living organisms C, O, H, N, P and S are used almost exclusively.

Why are these elements chosen?? -- Let's look at their chemistry.

C -- needs 4 e's to complete its outer shell ---> forms 4 covalent bonds by sharing e's with the other common elements ---> relatively equal sharing of electrons ---> strong covalent bonds. The 4 atoms bonded to carbon lie at the corners of a tetrahedron.

Thus, carbon has a great deal of flexibility in possible bonding geometry ---> enormous variety of possible molecules (organic chemistry).

H -- needs 1 e' ---> one bond, e.g. H--H , --C--H etc.

O -- needs 2 e's ---> 2 bonds e.g. --C--O--H, H--O--H, >C==O

4 sp3 hybrid orbitals -- 2 have lone electron pairs and 2 form covalent bonds

P and S have more complex electronic configurations ===> they form a variety of compounds

These elements form:

PRECURSOR MOLECULES: CH4, CO2, H2O, NH3, N2, PO4, SO3, H2S (MW's 18 -50) these were probably present in the prebiotic earth and form...

BUILDING BLOCK MOLECULES: amino acids, nucleotides, monosaccharides (sugars), fatty acids, glycerol etc. (MW's 100 - 400)

these can be produced from precursors in simulations of prebiotic earth -- precursor molecules + energy source such as electric discharge (lightening), radiation (X-ray, UV), heat

which form
MACROMOLECULES:

which form

SUPRAMOLECULAR ASSEMBLIES:

which form...

ORGANELLES:

Which form...

CELLS ====> TISSUES


Princples of Thermodynamics

To analyze thermodynamic processes we divide the universe into two components

  1. System: the part we're interested in
  2. Surroundings: all of the rest

The System may be open and exchange matter and energy with the surroundings or it may be close and cannot exchange matter and energy with the surroundings. Living organisms are open systems.


First Law of Thermodynamics: Energy is Conserved

DU = Ufinal - Uinital = q - w

q = heat absorbed by the system from the surroundings

w = work done by the system on the surroundings

heat is random molecular motion while work is force times distanced moved under its influence

Exothermic Processes release heat and have q<0

Endothermic Processes absorb heat and have q>0

Energy: The SI unit is joule (J) although we will frequently use calorie ; 1 cal = 4.2 J


State Functions: are independent of the path or the way the system or universe gets to them. They depend on the state of the system and not on its history. Energy, U, is a state function, therefore a cyclic process (which returns to its original state) has DU = 0

Heat and work are not state functions because they depend on the path.

Enthalpy, H = U + PV is a useful function for living systems which operate at constant P (mostly).

from DH = DU + PDV and substituting for DU = q - w the two PDV terms cancel and we get DH = qp

So what? Since DH is a State Function, we can measure it for any path and it will be valid for any other path

==> the DH for glucose + O2 --> CO2 + H2O can be measured in a calorimter and the value will be valid for metabolism in a cell.


Second Law of Thermodynamics: The Universe tends toward Maximum Entropy

In any spontaneous process the disorder (Entropy) of the universe increases (Universe = System + Surroundings). Disorder is defined as the number of equivalent ways, W, of arranging the components. We define a measure of this called Entropy, S, which is more manageable in expressing the many ways of arranging the components:
Entropy = S = k
b ln W (kb = Boltzman's constant) and S is a State Function.
Since the energy of the universe is constant, the 2nd law of Thermodynamics is:

For any spontaneous process -- DSsystem + DSsurroundings = DSuniverse > 0

(Note: Creationists sometimes claim that life violates the 2nd Law of Thermodynamics, and therefore a higher power must be involved in creating life. But while life requires creation of order (DSlife < 0) this is accompanied by an increase in entropy elsewhere; order is created by living organisms by the input of energy from the surroundings.

See Table 1-3

Free Energy: We can't measure DSuniverse easily and some processes occur when DSsystem < 0. We need another indicator of spontaneity for the system.

at constant pressure, P, and temperature T: DS qP / T = DH / T

==> for a spontaneous process: DH - TDS < 0 first recognized by J. Willard Gibbs

so we call the Gibbs Free Energy, DG = DH - TDS and if DG < 0 a process will be spontaneous
A process which releases enough heat, DH < 0, can overcome a decrease in entropy of the system.

DG is the maximum work we can obtain from a process. Furthermore, we must emphasize that a DG < 0 does not mean that a process occurs rapidly. Proteins, Nucleic Acids, and polysaccharides have DG's << 0 with respect to hydrolysis yet are perfectly stable in the absence of enzymes (catalysts) to increase the speed of the process.


Chemical Equilibria: G depends upon concentration (S changes with concentration)

GA = GA° + RT ln[A]; GA° is the standard free energy of A, R = 2.0 cal/deg-mole = 8.3 J / deg-mole.

For the reaction aA + bB == cC = dD ==>DG = DG° + RT ln ([C]c [D]d) / [A]a [B]b)

DG° is the standard free energy change for this reaction when reactants and products are in their standard states.

This principal is very important, because it shows that a reaction can be caused to go in either direction by changing the concentrations of products and reactants; this often occurs in living organisms.

At equilibrium, DG = 0 ==> D = - RT ln ([C]c [D]d) / [A]a [B]b) = -RT ln Keq

Note: ln Keq = -DH°/R (1/T) + DS°/R ==> a van't Hoff plot of LnKeq vs. 1/T allows one to calculate DH° and DS°, and, of course D

Note that a 10x change in Keq corresponds to a 5.7 kJ / mole (1.3 kcal/mole) change in DG° which is less than half the energy of a hydrogen bond.

Standard States: By definition the free energy of pure elements in their standard states (25°C, 1 atmosphere pressure, in their most stable form) is 0. The free energy of formation of a compound, DGf, is the change in G for formation of one mole from its component elements, all in their standard states.
Then for a chemical reaction
D = DGf (products) - DGf (reactants).

To the physical chemist, the standard state of compounds is a concentration of unit activity (activity is the concentration corrected for non ideal behavior), 25°C, and 1 atmosphere pressure. In Biochemistry we can usually assume that concentrations are low enough (≤ mM) that activity = concentration (M)

Biochemists modify this definition to account for the fact that nearly all biochemical processes take place in water near neutral pH, so the concentration of pure water, 55.5 M, and [H+] = 1 x 10-7 M at pH = 7 are incorporated into DG° ==> DG°'


Coupled Reactions: An endergonic reaction can be driven by an exergonic reaction; for example

A + B == C + D
DG1
D + E == F + G
DG2
A + B + E == F + G
DG3

DG3 = DG1 + DG2, so if |DG2| > |DG1| their sum, DG3 will be < 0 and the overall process is spontaneous. Reaction 2 keeps [D] low so reaction 1 continues to make more. These reactions are said to be couple by a common intermediate, D.


Biomolecules

What DETERMINES the PHYSICAL and CHEMICAL PROPERTIES of BIOMOLECULES??

A. Their FUNCTIONAL GROUPS:

Two functional groups can react together to form a new type of functional group; these reactions are very important in the formation of biological molecules

a) esters -- an acidic group (carboxylic acid or organic phosphate) can react with an alcohol (or thiol) to form an ester:

b) amides -- an acidic group reacting with an amino group


B. Their ENVIRONMENT: Water -- most of life processes occur in solutions of other molecules in H2O

Fig. 2-1

DIPOLAR: O contains 2 pairs of non-bonding electrons and tends to pull electrons away from H's leading to partial + charges (d+'s) on H's and partial - charges (2d-) on O

This makes H2O a good solvent for charged molecules (salts etc.) and other dipolar molecules (alcohols, amines, ketones, acids etc.) and a poor solvent for apolar molecules

The partial charges lead to electrostatic attractions between H2O molecules
===> Hydrogen Bonds

  • Fig. 2-2
  • The boiling point of water is 264°C higher than that of CH4 which has a similar molecular weight.

    On O, 2 orbitals contain lone electron pairs forming H-bonds to 2 other H2O molecules; they point toward 2 corners of a tetrahedron. The other 2 orbitals form covalent bonds to H-atoms and point to the other 2 corners of a tetrahedron.

    Fig. 2-3: The structure of ice makes it less dense that liquid water ==> ice floats which means that oceans and lakes remain liquid even when the temperature is below freezing.

    H-Bonds are weak chemical bonds 3-5 kcal/mole (20 kJ/mole) vs. 110 kcal/mole for --O-H covalent bond; but they are very important in biology because:

    a) H2O forms many H-bonds with itself (3.4 H-bonds/H2O ---> great cohesiveness of molecules leading to high boiling point, heat of vaporization, specific heat etc.

    b) Many (most) functional groups of biomolecules can form H-bonds with H2O and with other functional groups. This makes H20 an excellentg solvent for polar molecules (Fig. 2-7) and ions (Fig. 2-6); anions are attracted to the positively charged H’s and cations to the negatively charged O’s. Also H2O reduces the force attracting the ions to one another. Coulonb’s Law says this force = F = k q1q2/(D r2) where q’s are the charges on the two ions, r is the distance between them and D is the dielectric constant of the media; water has a very high dielectric constant ==> the force between the charges is weaker.

    --O-H----O=C< | --O-H----N< | --O-H----O< | >N-H----O==C< | >N-H----N<

    General FOrm D-H---A where D-H is the H-bond donor and A is the acceptor.

    AMPHIPHILES: Many biological molecules have both polar and non polar functional groups ==> they are both hydrophobic and hydrophilic and are said to be Amphiphilic or Amphipathic. These molecules often form specific structure to accommodate their hydrophilic and hydrophobic components; see Fig. 2-10 for the structures of micelles and bilayers

    HYDROPHOBIC EFFECT:

    This is spontaneous, therefore DG < 0; however in all cases DH > 0 so DS > 0 (i.e. -TDS < 0) Table 2-2

    The Hydrophobic Effect or Hydrophobic Bond is Entropy Driven: Entropy decreases when a hydrophobic molecule eneters water; i.e. the water becomes more ordered, because H2O can't H-bond with hydrophobic molecules and forms an ordered cage around them Fig. 2-8


    DISSOCIATION: Bond between --O-H sometimes breaks down forming H+ and OH-

    H-O-H <--> H+ + OH- | OH- gets an electron from H. (hydrogen atom)

    H+ actually exists as a hydrated proton (H3O+)

    This is a reversible chemical reaction whose equilibrium lies far to the left

    Keq = = 1.8 x 10-16M at 25 °C

    Since this is so small, we consider [H2O] to be constant = 55.5 M and incorporate it into Keq to give a new constant, Kw = [H+][OH-] = (55.5M) (1.8 x 10-16 M) = 1 x 10 -14 M2 @25°C

    Note: If no other source of H+ or OH- is added, then ---> [H+] = [OH-] and [H+]2 = 10-14 M2

    -----> [H+] = 10-7 M = [OH-] at neutrality

    Thus, the concentrations of H+ and OH- we're concerned with in biological systems are very small; but they can change by orders of magnitude (powers of 10). Therefore, we want to express them on a LOGARITHMIC SCALE e.g.

    pH = log10 {1 / [H+]} = - log10[H+] ====> at neutrality, pH = -log(10-7) = -(-7) = 7

    the -sign makes physiological pH's positive

    [OH-] = pOH = -log = -(log 10-14 - log[H+]) = 14 - pH

    or pH + pOH = 14

    Many functional groups of biomolecules can also dissociate; They are weak acid/base pairs.

    Bronsted -Lowry definitions: Acid ---> Proton Donor and a Base ---> Proton Acceptor
    An example is Acetic Acid/Acetate anion: CH
    3COOH ´ CH3COO- + H+ ; Acetic acid is obviously a proton donor; Acetate is its conjugate base (proton acceptor) because it can bind protons: CH3COO-+ H2O ´ CH3COOH + OH-

    For Ammonia: NH3 + H2O --> NH4+ + OH- and NH4+ --> NH3 + H+

    Strong Acid: HCl ---> H+ + Cl- Strong Base: NaOH ---> Na+ + OH-

    Both completely dissociate producing either H+ (strong acids) or OH- (strong base) which are the strongest acids and bases which can exist in aqueous solutions.

    Weak Acid: HA --> H+ + A- ===> Keq = Ka =

    Typically Ka << 1 (e.g. 1.7 x 10-4 for acetic acid to 5.6 x 10-10 for NH4+)

    take -log of both sides and rearrange
    ---> -logKa = pKa = -log [H+][A-] / [HA] = -log[H+] - log [A-] / [H+ ]

    pKa = pH - log [A-] / [HA] or pH = pK3+ log [A-] / [HA]

    This is the Henderson-Hasselbach Equation; describes the titration curve of a weak acid/base conjugate pair.

    Fig. 2-15 see also Table 2-5 for pKa’s of common buffers.

    Weak Acid/Base Pairs are pH buffers; in the pH region around the pKa value (4.6 for acetic acid), the pH of a solution of a buffer changes slowly when acid or base is added. This is because there is a reservoir of A- to react with XS H+ added and of HA to react with XS OH- added.

    When pH = pKa of a weak acid, the buffering capacity is a maximum because [HA] = [A-]. This can be seen in the titration curve; the slope of the curve, pH vs. vol. base (or vol. of acid if the titration curve is for titration of a weak base) is near zero meaning the pH changes very slowly when acid or base is added.

    The titration curve provides a method for measuring the concentration of acids or bases. At the equivalence point, the amount of OH- which has been added equals the amount of acid, HA, originally present. The equivalence point is detected when the pH changes rapidly as all HA has been converted to A- and the solution is no longer buffered.


    Physiological Buffers: weak acid conjugate base

    Cells (cytoplasm) are buffered primarily by phosphate H2PO4- <--> H+ + HPO4-2

    Ka = 1.38 x 10-7 ===> pKa = 6.86, this is a good buffer for an intracellular pH = 7.

    (Note: the first pKa of phosphoric acid is 2.15; when this proton dissociates it leaves behind a negative charge on H2PO4- making it more difficult to remove the next proton, thus it has a mujch higher pKa = 6.86. HPO4-2 has two negative charges making the final pKa even higher at 12.38)

    Blood is buffered by bicarbonate H2CO3 <--> H+ + HCO3- with a Ka = 1.7 x 10-4 = K1

    carbonic acid bicarbonate pKa = 3.77

    How can this buffer at pH = 7.4 !!!

    H2CO3 is in equilibrium with dissolved CO2(d) <---> CO2(d) + H2O <--> H2CO3 ==> K2

    The actual equilibrium system is the sum of these three reactions:

    CO2(g) + H2O <--> H+ + CO3-==> Ka' = K1 . K2 . K3

    This results in a buffer system which holds pH = 7.4. The pH of the blood can be altered by changing [CO2] measured as the partial pressure of CO2, pCO2, in the lungs; pCO2 can be regulated by the rate of breathing (note: CO2 is a product of metabolism and is removed by breathing).

    If blood pH decreases, we breathe more rapidly, pCO2 decreases and blood pH increases.

    If blood pH increases, we breathe more slowly, pCO2 increases and blood pH decreases.