Chapter 1 - Biochemistry in Space and Time

Biochemistry

The study of the chemistry of life processes

A possible timeline for Biochemical Evolution

• Key metabolic processes are common to many organisms

• Different organisms have macromolecules of similar structure and common biochemical processes.

• Suggests all organisms evolved from a common ancestor

Eukaryotes

Multicellular organisms

• ex: animals, plants, human beings, and many microscopic unicellular organisms such as yeast

Prokaryotes

Unicellular organisms

• lack a nucleus

All organisms can be placed in one of three domains

Eukarya, Bacteria, or Archaea

• based on biochemical characteristics

DNA

• is constructed from 4 building blocks

• Is a linear polymer composed of monomers consisting of a sugar (deoxyribose), a phosphate, and one of the 4 nitrogenous bases (bases)

• Has polarity which is important for many processes

Covalent Structure of DNA

• the backbone of DNA contains linked sugars and phosphates

• Variable bases extend from the backbone

The Double Helix

• Base-paired DNA forms a double helix structure

• Bases form specific base pairs held together by hydrogen bonds

• The two strands are antiparallel

Watson-Crick Base Pairs

• A-T pairs form two hydrogen bonds

• G-C pairs form three hydrogen bonds

DNA Structure

Explains heredity and the storage of information

• each strand serves as a template for a new partner

• Due to specific base pairing

• Allows for the generation of two identical daughter double helices from one parent strand

Two complementary DNA strands

Spontaneously assemble to form a double helix

The biochemical timescale

The timescale for biological interactions and processes

• on the order of picoseconds to microseconds

Covalent bonds

• formed by electron sharing between two adjacent atoms the strongest bonds

• The strongest bonds

• typical C-C covalent bonds have a distance of 1.54 Angstrom and a bond energy of 355 kJ mol-1 (85 kcal mol-1)

• Angstrom is 0.1 nm, or 10-10 m.

Resonance

Some molecules, such as adenine, exhibit multiple covalent structures called resonance structures

Ionic Interactions

Noncovalent interactions that occur between fully charged atoms or molecules

Coulomb energy

The energy (E) of electrostatic attraction between opposite charges or repulsion between like charges is given by ___

E = kq1q2/Dr

Where:

• k is a proportionality constant

• q1 and q2 are the charges on the two atoms

• D is the dielectric constant of the solvent

• k is a proportionality constant r is the distance between atoms

Electrostatic Interactions in Water have

• a bond distance of ~ 3 Angstrom

• a bond energy, given by Coulomb energy, of 5.86 kJ mol-1 (1.4 kcal mol-1)

Electric Dipoles

• molecules with no overall charge can have regions where electron distribution is uneven

• Leads to ____

• Dipoles can interact with ions or with other dipoles

Hydrogen Bonds

• occur between an electronegative atom and a hydrogen covalently bonded to another electronegative atom.

• vary in bond distance from 1.5 Å to 2.6 Å.

• have bond energies from 4 to 20  kJ mol−1 (1–5 kcal mol−1).

Hydrogen-bond donor

the group that includes both the atom to which the hydrogen atom is covalently bonded and the hydrogen atom itself

hydrogen-bond acceptor

The lone pair of electrons that is on the atom less tightly linked to the hydrogen atom

Van der Waals Interactions

• occur when two atoms are sufficiently close.

• occur when transient asymmetry in electron distribution in one atom induces complementary asymmetry in a neighboring atom.

• involve neighboring atoms attracting each other.

• are weak.

• have bond energies from 2 to 4  kJ mol−1 ( 0.5–1.0  kcal mol−1).

Van der Waals Contact Distance

• Attraction increases as two atoms come closer to each other, until they are separated by the van der Waals distance.

• At distances shorter than the van der Waals contact distance, strong repulsive forces become dominant.

Properties of Water

• water is a polar molecule with a partial positive and partial negative end

• Water is highly cohesive

• A large number of hydrogen bonds are formed in liquid water, and the maximum number of hydrogen bonds are formed in crystalline ice

The Hydrophobic Effect

• Nonpolar molecules in water can be driven together by the ____

  –powered by the increase in entropy of water

  –associated interactions are called hydrophobic interactions

The Double Helix is an Expression of the Rules of Chemistry

When a double helix forms, charge repulsion occurs between the negatively charged phosphates of the backbone.

• These repulsive forces are reduced by the high dielectric constant of water and interaction of positively charged ions with the phosphate groups.

Hydrogen Bonds Between Complementary Bases

Explain the Specificity of Sequence Pairing

• each individual nitrogenous base hydrogen bonds equally well with water as with its complementary base

• Hydrogen bonding explains the specificity of sequence pairing

DNA Base Pairs Are the Optimal van der Waals Distance Apart

• In the interior of the helix, bases are stacked and interact through van der Waals interactions.

• The hydrophobic effect also contributes to the favorability of base stacking.

• Surface complementarity occurs when hydrogen-bond donors align with hydrogen-bond acceptors and nonpolar surfaces

First Law of Thermodynamics

the total energy of a system and its surroundings is constant

Second Law of Thermodynamics

the total entropy of a system plus that of its surroundings always increases

states that, for a process to take place, the entropy of the universe must increase, which is possible only if

ΔSsystem > ΔHsystem /T    

or  TΔSsystem > ΔHsystem

•In other words, entropy will increase if and only if

ΔG = ΔHsystem − TΔSsystem < 0    

•Biochemical reactions will occur only if the ΔG is negative, which is only when the entropy of the universe increases.

Entropy

can decrease locally in the system if there is a corresponding increase in entropy in the surroundings.

Gibbs Free Energy

• −TΔStotal is the free energy or Gibbs free energy.

• The change in Gibbs free energy is used to describe the energetics of biochemical reactions

ΔG = ΔHsystem − TΔSsystem

The Formation of the Double Helix

• Heat is Released

The spontaneous formation of a double helix reduces the entropy of the DNA molecules.

• appears to violate the Second Law of Thermodynamics

Heat released by helix formation increases the entropy of the surroundings.

Acid-Base Reactions

Are Central in Many Biochemical Processes

• involve the addition or removal of a hydrogen, H+, ion.

• pH is a measure of the H+ concentration and is defined by

• pH = −log [H+]

• H+ and OH− ions are formed upon the dissociation of H2O

• H2O ↔ H+ + OH−

The Equilibrium Constant of Water

• The equilibrium constant (K) for the dissociation of water is defined as

K = [H+][OH−]/[H2O]

• KW, the ion constant of water, is defined as

KW = K[H2O]

• This can be simplified to

KW = [H+][OH−]

Proton or Hydroxide Ion Concentration Can Be Calculated if the Other Is Known

• KW has a known value

       KW = [H+][OH−] = 10−14

• From this, we can calculate

       [H+] = 10−14/[OH−] and [OH−] = 10−14/[H+]

• At exactly pH 7.0,  [H+] = [OH−] = 10−7 M.

Acid-Base Reactions Can Disrupt the Double Helix

• As base is added to a solution of double helical DNA, the helix is disrupted or denatured.

High pH Causes Loss of Hydrogen-Bond Donors in DNA

• The chemical basis of the denaturation is the disruption of base-pairing.

• example: the loss of a proton by the base guanine prevents base-pairing with cytosine.

The pKa Value Describes the Susceptibility of Proton Removal

• Proton dissociation for a substance HA has an equilibrium constant defined by the expression

Ka = [H+][A−]/[HA]

• The pKa value indicates the susceptibility of proton to  removal by reaction with a base:

pKa = −log(Ka)

• When pH is equal to the pKa,

−log[H+] = −log([H+][A−]/[HA])

and

[H+] = [H+][A−]/[HA]

When the pH is at the pKa, the Group is 50% Likely to be Deprotonated

• Dividing [H+] = [H+][A−]/[HA] by [H+] reveals that

1 = [A−]/[HA] or [A−] = [HA]

• When the pH is equal to the pKa, the concentration of the protonated form of HA is equal to the deprotonated form A−.

The N-1 Proton of Guanine

• The N-1 proton of guanine has a pKa of 9.7.

• When pH is near to or exceeds 9.7:

  • the proton is increasingly likely to be lost

  • Base-pairing is disrupted

  • The helix becomes denatured

Buffers Regulate pH in Organisms and in the Lab

• Buffers resist changes in the pH of a solution.

• Buffers are most effective at a pH near its pKa.

When a Buffer is Present

• pH Change is Gradual

Titration

• gradually adding known amounts of reagent to a solution with which the reagent reacts while monitoring the results

The ionization reaction of a weak acid is given by ___ & the equilibrium constant for this reaction is ___

• Taking the logarithms of both sides yields ___

HA ↔ H+ + A−
Ka = [H+][A−]/[HA]

• log(Ka) = log([H+]) + log([A−]/[HA])

The Henderson-Hasselbach Equation

• Recalling the definitions of pKa and pH and rearranging yields the Henderson–Hasselbalch equation

pH = pKa + log ([A−]/[HA])

• Weak acids are most effective as buffers at pH near the pKa value of its acid component.

A Buffer Functions Best

Close to the pKa Value of Its Acid Component

Phosphoric Acid

Is an Important Buffer in Biological Systems

• Physiological pH is typically near 7.4.

• Given the pKa values shown, inorganic phosphate exists as a nearly equal mixture of H2PO4− and HPO42−  in physiological systems

Genomic Sequences

Encode Proteins

• The most fundamental role of DNA is to encode the sequences of proteins.

• Proteins are built from 20 building blocks, called amino acids, rather than 4, as in DNA.

• Proteins spontaneously fold into elaborate three-dimensional structures, determined almost exclusively by their amino acid sequences.