Glossary
Concentrations of reactants and products (constant at equilibrium)
At equilibrium, the amounts of starting materials and resulting substances no longer change, even though reactions continue to occur.
Example:
Once a weak acid solution reaches equilibrium, the concentrations of reactants and products like the undissociated acid and its conjugate base will remain stable, not increasing or decreasing.
Equilibrium (Kinetic Definition)
The state where the forward and reverse reaction rates are equal, leading to constant concentrations of reactants and products over time.
Example:
In a closed bottle of soda, CO2 gas dissolving into the liquid and CO2 gas escaping from the liquid reach a dynamic equilibrium, where the amount of dissolved CO2 remains constant.
Equilibrium (Thermodynamic Definition)
The state of a system where the free energy (G) is at its minimum, and there is no net tendency for the reaction to proceed in either direction.
Example:
A chemical reaction will proceed until it reaches equilibrium, the lowest energy state, much like a ball rolling to the bottom of a valley.
Equilibrium constant (K)
A value that expresses the ratio of product concentrations to reactant concentrations at equilibrium, indicating the extent to which a reaction proceeds.
Example:
A very large equilibrium constant for a reaction like the combustion of methane indicates that at equilibrium, there will be overwhelmingly more products than reactants.
Forward and reverse reaction rates
The speed at which reactants are converted to products (forward) and products are converted back to reactants (reverse) in a reversible reaction.
Example:
When hydrogen and iodine react to form HI, the forward reaction rate is how fast H2 and I2 combine, while the reverse reaction rate is how fast HI decomposes back into H2 and I2.
Minimum free energy (G)
The lowest possible Gibbs free energy value a system can achieve, indicating the most stable state for a reaction at a given temperature and pressure.
Example:
For any spontaneous reaction, the system will naturally evolve towards the point of minimum free energy, where the reaction is most stable and no longer has a driving force.
Nonspontaneous reaction (ΔG > 0)
A reaction that requires continuous external energy input to proceed under a given set of conditions, indicated by a positive change in Gibbs free energy.
Example:
Converting water into hydrogen and oxygen gas through electrolysis is a nonspontaneous reaction because it requires a constant supply of electrical energy to occur.
Product-favored
Describes a reaction where, at equilibrium, the concentration of products is significantly greater than the concentration of reactants, typically associated with K > 1.
Example:
The formation of rust is a product-favored reaction, meaning that at equilibrium, there's a lot of rust and very little unreacted iron and oxygen left.
Reactant-favored
Describes a reaction where, at equilibrium, the concentration of reactants is significantly greater than the concentration of products, typically associated with K < 1.
Example:
The decomposition of water into hydrogen and oxygen is a reactant-favored process at room temperature, meaning that at equilibrium, most of the substance remains as water.
Reaction quotient (Q)
A measure of the relative amounts of products and reactants present in a reaction at any given time, used to determine the direction a reaction will shift to reach equilibrium.
Example:
If you mix reactants and products in a beaker, calculating the reaction quotient tells you if the reaction will proceed forward or backward to reach equilibrium.
Spontaneous reaction (ΔG < 0)
A reaction that proceeds without continuous external energy input under a given set of conditions, indicated by a negative change in Gibbs free energy.
Example:
The rusting of iron is a spontaneous reaction because it occurs naturally over time without needing an external push, as its ΔG is negative.
Standard free energy change (ΔG°)
The change in Gibbs free energy for a reaction when all reactants and products are in their standard states (1 atm for gases, 1 M for solutions, 298 K).
Example:
When comparing the inherent favorability of different reactions, chemists often look at the standard free energy change to see which one is more likely to proceed under ideal conditions.
ΔG (Free energy change under nonstandard conditions)
The change in Gibbs free energy for a reaction under any specific set of conditions, which determines the spontaneity of the reaction at that moment.
Example:
Even if a reaction is spontaneous under standard conditions (ΔG° < 0), its ΔG might become positive if product concentrations are very high, making it nonspontaneous at that specific point.