Question

The relationship of free-energy change (ΔG) to the concentrations of reactants and products is important because it predicts the direction of spontaneous chemical reactions. Consider, for example, the hydrolysis of ATP to ADP and inorganic phosphate (Pi). The standard free-energy change (ΔG°) for this reaction is -7.3 kcal/mole. The free-energy change depends on concentrations according to the following equation:

ΔG = ΔG° + 1.42 log10 ([ADP] [Pi]/[ATP])
In a resting muscle, the concentrations of ATP, ADP, and Pi are approximately 0.005 M, 0.001 M, and 0.010 M, respectively. At [Pi] = 0.010 M, what will be the ratio of [ATP] to [ADP] at equilibrium?
(a) 1.38 × 106
(b) 1
(c) 7.2 × 10-8
(d) 5.14

Answer 1

Free-energy change (ΔG) is fundamental in determining the spontaneity of reactions, which can be exergonic (energy-releasing) or endergonic (energy-absorbing). Calculations adjust for real-world conditions using the reaction quotient, Q. The Gibbs free energy equation (ΔG = ΔH - TΔS) also factors into understanding ΔG.

Step-by-step explanation:

The relationship between the free-energy change (ΔG) and the concentrations of reactants and products is crucial for predicting the spontaneity of chemical reactions. When the standard free-energy change (ΔG°) is negative, as in the hydrolysis of ATP, the reaction is considered exergonic, meaning it releases free energy. Conversely, if ΔG° is positive, the reaction is endergonic and absorbs free energy.

To calculate the actual free-energy change (ΔG') in real-world conditions where concentrations are not standard, one can use the equation ΔG' = ΔG° + RT ln(Q), where R is the gas constant, T is the absolute temperature, and Q is the reaction quotient. The equilibrium constant (Keq) is used when the reaction is at equilibrium. This equation is derived from the Gibbs free energy equation, which integrates changes in enthalpy (ΔH) and entropy (ΔS).

Understanding these principles allows us to determine the standard free energy change and predict the behavior of a chemical reaction under various conditions, which is essential for fields ranging from biochemistry to industrial chemistry.