Bioenergetics and Thermodynamics

• Quantitative study of energy transductions in living cells, including the chemical processes underlying these transductions, is called bioenergetics and thermodynamics.
• Energy allows work to be performed and chemical reactions can be studied from thermodynamics (focus on energy transformation by analyzing the starting and ending points of a reaction) or kinetics (examining the processes in between, previously covered in semester).

Laws of Thermodynamics

• First Law (Conservation of Energy): the total amount of energy in the universe remains constant, whereby energy can be changed, but not created or destroyed.
• Second Law: the disorder (entropy) of the universe tends to increase.

Entropy

• Entropy measures the randomness or disorder of a system and describes how molecules are arranged within a sample.
• In natural systems, entropy increases over time as molecules spread apart.
• Disorder occurs spontaneously without needing energy, while organizing a system needs energy.

Concept of Free Energy, Gibbs Equation

• Gibbs Free Energy measures the amount of usable energy stored in a molecule's bonds and it determines whether a reaction will occur spontaneously based on whether it absorbs or releases energy: ΔG = ΔH – T•ΔS
• ΔG represents the free energy in the system and is negative in spontaneous reactions.
• ΔH (Enthalpy) is the heat change.
• T is the temperature.
• ΔS (Entropy) is the randomness in the system.
• If ΔG is negative (-), the reaction releases energy and happens spontaneously (exergonic reaction).
• If ΔG is positive (+), the reaction absorbs energy, requiring energy input (endergonic reaction).
• When water exists in gas, liquid, and solid states, energy and disorder differ:
  • ΔG (Gibbs Free Energy) stays the same because it stays H2O in all states with the same bonds but can change depending on temperature and pressure.
  • ΔH (Enthalpy) varies, lowest in solid state (ice), higher in liquid, highest in gas due to energy in the molecular bonds.
  • ΔS (Entropy) is lowest in solid state (more ordered), higher in liquid, and highest in gas (more disorder).
• Each state has different ΔH and ΔS, but the same ΔG.
• By rearranging the equation to ΔH = ΔG + T•ΔS, ΔH is the total system energy.

Thermodynamics

• Thermodynamics show if the reaction is possible, and can be used to ascertain whether glucose can convert to pyruvate, depending on the energy within each molecule.

Why Chemical Reactions Occur

• Chemical reactions occur to reach stability (lower energy).
• Reactions logically go from high to low energy scales.

Exergonic Reaction

• ΔG is negative
• Indicates an exergonic reaction.
• Spontaneous and favorable reaction

Endergonic Reaction

• ΔG is positive.
• Indicates an endergonic reaction.
• Non-spontaneous and non-favorable reaction
• Energy must be added to occur, such as adding ATP.

ΔG°

• ΔG at standard conditions: 1 atm pressure, 1 M concentrations of reactants and products, 25°C temperature, and pH 7.
• Used to compare different reactions in a laboratory setting.
• ΔG will change if concentrations are changed.

ΔG is a State Function

• It depends only on the starting and ending points (concentrations), not on the reaction mechanism or intermediates.

Combustion of Glucose in Calorimeter

• Glucose + O2 → CO2 + H2O, ΔG = - 680 kcal/mol
• In the cell glycolysis and Krebs cycle are complex with many steps, but the overall ΔG is not affected.
• Glucose → CO2 + H2O, ΔG = - 680 kcal/mol
• ΔG is still the same through the pathway, so the pathway isn't important here

ΔG is Affected by Concentration

• In the reaction Glucose→Glucose 6-phosphate, the energy change (ΔG) can be calculated using ΔG = Gf - Gi, where Gf is the free energy of the final product (glucose 6-phosphate) and Gi is the free energy of the initial reactant (glucose).
• ΔG will change if glucose concentration changes due to the effect on molecules and bonds.
• Adjusting reactant or product (glucose/glucose-6 phosphate) concentrations can make a reaction spontaneous (if ΔG is negative) or non-spontaneous (if ΔG is positive).
• Changing concentrations influences whether the reaction occurs naturally or needs energy input.
• When the energy of reactants equals the energy of products it is called equilibrium.

Relation Between ΔG° and ΔG

• ΔG = ΔG° + RT ln ([Products]/[Reactants])
• ΔG = ΔG° + RT 2.3 log ([Products]/[Reactants])
• Almost the temperature and atmosphere is the same so the only thing that varies are the concentrations
• If [products] < [reactants], then ΔG < ΔG°
• If [products] > [reactants], then ΔG > ΔG°
• If [products] = [reactants], then ΔG = ΔG°
• Equilibrium is at the end.
• At equilibrium, the concentration is fixed, rates are exactly equal
• The concentrations define the equilibrium constant
• Keq = ([C]^c [D]^d) / ([A]^a [B]^b) for aA + bB ⇌ cC + dD
• ΔG can be changed by changing concentrations to make a non-spontaneous reaction spontaneous by increasing reactants or decreasing products.
• ΔG =0 (ln 1 is zero) and is at equilibrium

Reaction Favored

• If reactants are at 1M
• AG' = - RTIn = [C] [D] / [A] [B]

Multiplicative and Additive

AG is additive for multiple subsequent reactions and Keq is multiplicative such as:

• Glucose + Pi → glucose 6-phosphate + H2O then Keq1= [glucose 6-phosphate] / [glucose][P] = 3.9 x 10^-3
• ATP + H2O →ADP + Pi then Keq2 = [ADP][Pi] / [ATP] = 2.0 x 10^5 M

Effect of Changing Conditions on Equilibria

• A system at equilibrium will shift to relieve stress.
• Stress is any change that disturbs the original equilibrium.
• Changes in concentration: continuously supplying/removing reactants/products can be used by metabolic reactions.
• Changes in temperature: endothermic reactions are favored by increase in temperature, exothermic reactions are favored by decrease in temperature.
• Catalysts do not change equilibrium.