Question

(a) Is the standard free-energy change, \(\Delta G^{\circ},\) always larger than \(\Delta G ?(\mathbf{b})\) For any process that occurs at constant temperature and pressure, what is the significance of \(\Delta G=0\) ? (c) For a certain process, \(\Delta G\) is large and negative. Does this mean that the process, necessarily has a low activation barrier?

Short answer

1. There is no general rule that \(\Delta G^{\circ}\) is always larger or smaller than \(\Delta G\). The relationship between them depends on specific reaction conditions and follows the equation \(\Delta G = \Delta G^{\circ} + RT \ln Q\). 2. If \(\Delta G = 0\), the reaction is at equilibrium, meaning the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant. 3. A large negative \(\Delta G\) indicates a highly spontaneous process, but it does not provide any information about the activation barrier or the reaction rate. The relationship between \(\Delta G\) and activation energy is not directly linked, with kinetics determined by the activation energy and thermodynamics determined by \(\Delta G\).

Step-by-step solution

1. Comparing \(\Delta G^{\circ}\) and \(\Delta G\)

The standard free-energy change, \(\Delta G^{\circ},\) represents the free-energy change under standard conditions (1 atm pressure and typically, 25°C temperature). The free-energy change, \(\Delta G,\) represents the change in free energy under any given set of conditions. These values may or may not be the same, depending on the specific reaction and the conditions under which it is happening. There is no general rule that the standard free-energy change is always larger or smaller than the free-energy change. For example, consider the following reaction: \[A + B \rightarrow C + D\] The standard free-energy change, \(\Delta G^{\circ},\) represents the energy change for this reaction when all the substances are at their standard states. The free-energy change, \(\Delta G,\) takes into account the actual concentrations (or partial pressures) of the substances involved in the reaction. Thus, \(\Delta G\) varies with changes in concentration or pressure, according to the relationship: \[\Delta G = \Delta G^{\circ} + RT \ln Q\] Where R is the gas constant, T is the temperature, and Q is the reaction quotient. As long as conditions are not standard, there is no sure relationship between the two and one can be larger or smaller than the other.

2. Significance of \(\Delta G=0\)

For any process occurring at a constant temperature and pressure, the value of the free-energy change, \(\Delta G,\) determines whether a reaction is spontaneous or not. If \(\Delta G\) is negative, the reaction is spontaneous. If \(\Delta G\) is positive, the reaction is non-spontaneous, and the reverse reaction is spontaneous. If \(\Delta G = 0,\) this indicates that the reaction is at equilibrium. At equilibrium, the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant. An example of a reaction at equilibrium is the dissociation of water: \[2H_{2}O \longleftrightarrow H_{3}O^{+} + OH^{-}\] At equilibrium, the rate of dissociation of water into hydronium and hydroxide ions is equal to the rate of the reverse reaction (recombination of the ions to form water), and the system is in a state of dynamic balance.

3. Large negative \(\Delta G\) and activation barrier

A large negative value of \(\Delta G\) indicates that a process is highly spontaneous. It means that the final state is much lower in free energy than the initial state. However, this information alone doesn't provide any information about the activation barrier (or activation energy) for the reaction. The activation energy represents the minimum energy required for reactants to transform into products; it is the height of the energy barrier that must be overcome. A reaction can have a large negative \(\Delta G\) and still have a high activation barrier, meaning it may happen very slowly. Conversely, a reaction can have a small negative \(\Delta G\) and a low activation barrier, meaning it may happen quickly. For the most part, the relationship between \(\Delta G\) and activation energy is not directly linked. In general, the kinetics (rate) of a reaction is determined by the activation energy, while the thermodynamics (extent) of a reaction is determined by \(\Delta G.\)

Key concepts

Understanding Standard Free-Energy Change

The concept of standard free-energy change, denoted as \(\Delta G^\circ\), is fundamental for predicting how a chemical reaction will behave under 'standard' conditions, usually set at 1 atm pressure and 25°C temperature. This value captures the energy changes when all reactants and products are in their standard states.

However, it's essential to recognize that real-world conditions often deviate from these ideals. This leads to the non-standard free-energy change, \(\Delta G\), reflecting the situation where concentrations, pressure, and other variables differ from the standard. They are connected by the equation \(\Delta G = \Delta G^\circ + RT \ln Q\), where \(R\) is the gas constant, \(T\) is the temperature in Kelvin, and \(Q\) is the reaction quotient.

Consequently, \(\Delta G^\circ\) is not necessarily larger or smaller than \(\Delta G\); it rather serves as a comparative benchmark and helps determine the direction and extent of a reaction under standard conditions.

Deciphering Reaction Spontaneity Through Free-Energy Change

The spontaneity of a chemical reaction is not always an easy concept to grasp, but free-energy change, \(\Delta G\), makes it clearer. A negative \(\Delta G\) indicates that a reaction will occur without external energy input; it's spontaneous. Conversely, a positive \(\Delta G\) implies that energy must be put into the system to make the reaction proceed; it's non-spontaneous.

When \(\Delta G = 0\), the reaction has reached a state of chemical equilibrium, where no net change happens in the concentrations of reactants and products over time. Understanding this can empower students to predict a reaction's behavior just by considering the sign of \(\Delta G\).

Chemical Equilibrium and Free-Energy

At chemical equilibrium, a reaction has settled into a state where the rate of the forward reaction equals the rate of the reverse reaction, leading to no overall change in the concentration of reactants and products. This state corresponds to \(\Delta G = 0\).

Equilibrium does not imply that the reactions have stopped, but rather that they are proceeding at equal rates in both directions—a dynamic balance. The equilibrium constant, \(K\), derived from the concentrations of reactants and products at equilibrium, provides a quantifiable insight into the position of equilibrium.

Linking Reaction Kinetics to Free-Energy

Reaction kinetics refers to how quickly or slowly a chemical reaction proceeds towards its final products. Unlike thermodynamics, which is concerned with the direction and extent of a reaction, kinetics focuses on the rates of reactions. A large negative \(\Delta G\) indeed implies a thermodynamically favourable process, but it does not necessarily mean the reaction will occur quickly.

That's where the activation barrier, or activation energy, comes in—the energy hurdle that must be cleared for reactants to convert into products. This concept helps explain why some reactions with a highly spontaneous drive (large negative \(\Delta G\)) may still occur slowly if the activation energy is high. Therefore, a comprehensive understanding of both thermodynamics and kinetics is crucial to grasp the full picture of reaction behaviour.