Question

The following interconversion is catalyzed by the enzyme triose phosphate isomerase: Explain how you might use bond energies to estimate whether the equilibrium constant, \(K\), for this reaction would be greater, or less, than unity.

Short answer

Calculate energy change; exothermic reactions likely have \(K > 1\).

Step-by-step solution

Understand Bond Energies

First, consider the concept of bond energies. Bond energy is the energy required to break a bond between two atoms. In chemical reactions, bonds in reactants are broken and new bonds in products are formed. The change in the bond energy determines whether the reaction releases or absorbs energy.

Calculate the Net Change in Energy

For the interconversion, calculate the total bond energy of the reactants and products. Compare the energy required to break bonds in the reactants to the energy released when new bonds are formed in the products. If the reaction releases more energy than it absorbs (exothermic), the products are more stable, suggesting a tendency for the reaction to go forward.

Relate Energy Change to Equilibrium Constant

The equilibrium constant, \(K\), relates to the Gibbs free energy change \(\Delta G\) of the reaction by the equation \(\Delta G = -RT \ln K\). If the reaction is exothermic (releases energy), \(\Delta G\) is negative, making \(\ln K\) positive and suggesting \(K > 1\). Conversely, if the reaction absorbs energy (endothermic), \(\Delta G\) is positive, resulting in \(K < 1\).

Key concepts

Triose Phosphate Isomerase

Triose phosphate isomerase is an important enzyme in the glycolytic pathway that catalyzes the conversion between two molecules: dihydroxyacetone phosphate (DHAP) and glyceraldehyde-3-phosphate (G3P). This conversion is crucial because only G3P continues in the glycolysis pathway to eventually yield energy in the form of ATP.

By allowing this interconversion, triose phosphate isomerase maximizes the use of glucose, making the process more efficient. The enzyme achieves remarkable efficiency due to its ability to facilitate the reaction speed and balance the concentrations of DHAP and G3P when needed. Triose phosphate isomerase is often cited as a "perfect enzyme" because it is diffusion-limited, meaning the reaction occurs as fast as the molecules can collide.

Understanding the function of this enzyme helps highlight the dynamic balance that is maintained in metabolic pathways. It exemplifies the body's ability to manage energy resources efficiently, ensuring that the conversion allows the best outcome for energy extraction from glucose.

Equilibrium Constant

The equilibrium constant, denoted as \(K\), plays a crucial role in dictating the direction and extent of a chemical reaction. It quantifies the ratio of products to reactants at equilibrium, offering insight into the reaction’s tendency under a given set of conditions.

When \(K > 1\), it indicates that at equilibrium, products are favored over reactants, meaning the reaction proceeds significantly in the forward direction. Conversely, \(K < 1\) implies that reactants are favored, suggesting the reaction doesn't proceed much beyond the initial stages.

For reactions catalyzed by enzymes like triose phosphate isomerase, the equilibrium constant helps in understanding how enzymes can shift the position of equilibrium. Although enzymes do not change the equilibrium constant itself, they speed up the attainment of equilibrium. Thus, the equilibrium constant becomes a key to predict the outcome of reactions under enzymatic action.

Gibbs Free Energy

Gibbs free energy, symbolized as \( \Delta G \), is a fundamental concept in thermodynamics that helps determine whether a chemical process will occur spontaneously. It combines enthalpy and entropy of a system in the form of an equation \( \Delta G = \Delta H - T\Delta S \), where \( \Delta H \) is the change in enthalpy, \( T \) is the temperature, and \( \Delta S \) is the change in entropy.

A negative \( \Delta G \) indicates a spontaneous reaction, often associated with exothermic processes where energy is released. When \( \Delta G \) is positive, the reaction is non-spontaneous, suggesting energy absorption or an endothermic nature.

The relationship between Gibbs free energy and the equilibrium constant is given by \( \Delta G = -RT \ln K \), linking thermodynamics with reaction dynamics. This equation shows that a negative \( \Delta G \) correlates with \( K > 1 \), meaning products are favored at equilibrium, while a positive \( \Delta G \) leads to \( K < 1 \), favoring reactants. Understanding this connection underscores the prediction of reaction spontaneity and direction.