Physical Biochemistry

What are the energy converting processes in biological systems?

• Organisms use energy obtained from the environment to maintain homeostatic conditions that are far from equilibrium; reaching equilibrium with the environment is equivalent to death.

• Solar energy is converted to chemical energy through the processes of photosynthesis and carbon fixation; chemical energy is converted by cells into useful work (osmotic work, chemical work, and mechanical work).

• The primary energy conversion processes in cells are oxidation–reduction (redox) reactions that involve the transfer of electrons between molecules.

• All biological processes follow the laws of thermodynamics. The first law of thermodynamics states that energy cannot be created or destroyed, only converted from one form to another. The second law of thermodynamics states that entropy (S), or the dispersion of energy, in the universe is always increasing.

• Enthalpy (H) is the heat content of a molecule and is reflected in the number and type of chemical bonds. Exothermic reactions give off heat, and the change in enthalpy (ΔH) is negative; endothermic reactions absorb heat, and ΔH is positive.

• Change in Gibbs energy (ΔG) describes the spontaneity of a reaction in terms of absolute temperature (T) and changes in enthalpy (ΔH) and entropy (ΔS) by the relationship ΔG = ΔH − TΔS. Exergonic reactions (ΔG < 0) are favorable for the forward reaction, whereas endergonic reactions (ΔG > 0) are unfavorable for the forward reaction.

• ΔG°′ denotes the change in standard Gibbs energy under biochemical standard conditions—namely, when the temperature (298 K) and pressure (1 atm) are held constant and the aqueous solution is pH 7 (that is, [H⁺] = 10⁻⁷ M) with the concentration of H₂O held constant at 55.5 M.

• The equilibrium constant (Keq) can be used to determine the ΔG°′ value of a reaction by the equation ΔG°′ = −RT ln Keq, in which Keq is the ratio of the equilibrium concentrations of products over reactants, R is the gas constant, and T is the temperature in kelvins.

• The ΔG°′ value for a reaction is a constant and corresponds to the amount of energy needed to go from the biochemical standard condition, where all reactants and products are present initially at 1 M concentrations, to the condition at which all reactants and products have reached equilibrium.

How do coupled reactions use the chemical energy of ATP to drive otherwise unfavorable metabolic reactions?

• The ΔG value of coupled reactions is equal to the sum of the ΔG values for each individual reaction. Cleavage of a phosphoanhydride bond in ATP is often coupled to an otherwise unfavorable reaction to drive the reaction forward (overall ΔG < 0).

• The ΔG value of metabolic reactions cannot always be determined because the steady-state concentrations of all reactants and products may not be known. Therefore, it is common to use the biochemical standard Gibbs energy change, ΔG°′, which is a constant, to determine the overall favorability of coupled reactions to a first approximation.

• The ΔG°′ for hydrolysis of the phosphoanhydride bond between the β and γ phosphates of ATP is −30.5 kJ/mol and that of the phosphoanhydride bond between the α and β phosphates is −32.3 kJ/mol. The actual Gibbs energy change, ΔG, for the hydrolysis of the β and γ phosphoanhydride bonds, which is called the phosphorylation potential, ΔGp, is −50 kJ/mol.

• The formation of new bonds as a result of phosphoryl group transfer in reactions involving ATP is an exergonic process and provides the energy for ATP-coupled reactions; hydrolysis of the phosphoanhydride bonds, which is an endergonic process, does not provide this energy.

• Phosphoanhydride bond energy in ATP is also used for cellular work, most often in the form of initiating protein conformational changes. Two examples of ATP-dependent conformational changes are the muscle myosin protein and the Na⁺–K⁺ ATPase transporter protein.

• Three chemical properties of ATP account for the large −ΔG that occurs upon phosphoryl transfer: (1) electrostatic repulsion destabilizes ATP; (2) free phosphate ion following ATP hydrolysis has more resonance forms than phosphate covalently attached to ADP, which is entropically favored; and (3) the phosphate ion and ADP have a greater degree of solvation than ATP and are therefore more stable.

• Breaking the phosphoanhydride bond in ATP is itself endergonic because bond breakage requires energy, whereas bond formation is exergonic and releases energy. However, the overall reaction of ATP hydrolysis is exergonic because the energy released by the formation of hydrogen bonds between Pi and H₂O is greater than the amount of energy required to break the phosphoanhydride bond. Increases in entropy also contribute to the overall favorability of the ATP hydrolysis reaction.

• Researchers have proposed that at an early stage of biochemical evolution, the compound acetyl phosphate (AcP) may have been the chemical precursor to ATP. Because AcP can phosphorylate ADP in the presence of Fe³⁺, but cannot phosphorylate other nucleoside diphosphates, this may have led to ATP becoming the superior energy currency of the cell.

• The energy charge (EC) of the cell reflects the relative concentrations of ATP, ADP, and AMP. When the EC is high, anabolic pathways (biosynthesis) are favored because ATP is plentiful; when the EC is low, catabolic pathways (degradation) are favored.

What are the properties of H₂O that make it so critical for life?

• Water has four properties that make it essential for life on Earth: (1) the solid form of water is less dense than the liquid form, which is why ice floats; (2) water is liquid over a wide range of temperatures; (3) water is an excellent solvent because of its hydrogen-bonding abilities and polar properties; and (4) water has a high specific heat capacity and functions as a temperature buffer.

• The molecular structure of H₂O gives rise to a separation of charge, with two partial negative charges on the oxygen atom (2δ⁻) and one partial positive charge on each of the hydrogens (δ⁺ and δ⁺).

• Extensive hydrogen bonding between H₂O molecules gives water its unusually high viscosity, boiling point, and melting point compared with those of other molecules of a similar molecular mass.

• The transfer of H⁺ between H₂O molecules is called proton hopping and gives rise to an electric current through a “water wire.”

• Biochemical processes depend on weak noncovalent interactions, which permit structures to exist for short periods of time. The three basic types of weak noncovalent interactions in nature are hydrogen bonds, ionic interactions, and van der Waals interactions. In addition, hydrophobic effects contribute to the stability of macromolecules in aqueous solutions.

• The addition of nonpolar compounds to water breaks hydrogen bonds between H₂O molecules without replacing them and leads to the formation of ordered cage-like H₂O structures, which is energetically unfavorable.

• Hydrophobic effects result from nonpolar compounds associating with each other to minimize the amount of H₂O that must be ordered at the interface between the hydrophobic region and H₂O.

• The colligative properties of aqueous solutions (freezing point depression, boiling point elevation, vapor pressure lowering, and osmotic pressure) are affected by the concentration of solute molecules, not by their chemical properties. Osmotic pressure is especially important in biological systems, where solute concentrations modulate cell size as a result of water diffusion across a semipermeable plasma membrane.

What are definitions of pH and pKa with regard to the ionization of H₂O?

• The ionization of water gives rise to hydrogen ions, H⁺, which are protons, and hydroxide ions, OH⁻. Protons do not exist free in solution, but instead combine with H₂O to generate hydronium ions, H₃O⁺.

• The water ionization constant Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ M² at 25 °C. The concentration values, [H⁺] and [OH⁻], are reciprocally related, so [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M in pure water.

• The pH scale runs from 0 to 14 and is a convenient method to describe H⁺ concentration in aqueous solutions by using the expression pH = −log[H⁺]. Solutions with pH [OH⁻]), solutions with pH > 7.5 are basic ([H⁺] < [OH⁻]), and neutral solutions have a pH in the range 6.5–7.5 ([H⁺] ≈ [OH⁻]).

• Acids are proton donors, whereas bases are proton acceptors. The ionization reaction of an acid–base conjugate pair can be written as HA ⇌ H⁺ + A⁻.

• The acid dissociation constant, Ka, is derived from the ionization reaction of an acid and can be defined as pKa = −log Ka. Acid–base conjugate pairs with low pKa values are able to dissociate protons at low pH, whereas acid–base conjugate pairs with high pKa values dissociate protons at high pH.

• The Henderson–Hasselbalch equation relates pH and pKa and can be used to determine the ratio of A⁻ (proton acceptor) to HA (proton donor) at a given pH if the pKa is known: pH = pKa + log([A⁻]/[HA]).

• In aqueous solutions with pH values below the pKa, the acid–base conjugate pair is mostly in the protonated form ([HA] > [A⁻]), whereas in aqueous solutions with pH values above the pKa, the acid–base conjugate pair is mostly in the deprotonated form ([A⁻] > [HA]).

• Buffers are aqueous solutions that resist changes in pH because of the protonation or deprotonation of an acid–base conjugate pair present at a high enough concentration to absorb small changes in H⁺ or OH⁻ concentration.