Glossary
Common Ion Effect
The decrease in the solubility of a sparingly soluble ionic compound when a soluble salt containing a common ion (an ion already present in the solution) is added.
Example:
Adding NaCl to a saturated solution of AgCl will decrease the solubility of AgCl due to the common ion effect from the chloride ion.
Dynamic Equilibrium
A specific type of equilibrium where the forward and reverse reactions are still occurring at equal rates, meaning the system is active and not static, even though macroscopic properties remain constant.
Example:
In a saturated sugar solution, sugar molecules are constantly dissolving and crystallizing, but the net amount of dissolved sugar remains constant, illustrating dynamic equilibrium.
Equilibrium
A state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, leading to constant concentrations of reactants and products.
Example:
When a bottle of soda is opened, CO2 gas escapes until the system reaches a new equilibrium where the rate of CO2 dissolving equals the rate of CO2 escaping.
Equilibrium Constant (K)
A numerical value that expresses the ratio of product concentrations to reactant concentrations at equilibrium, with each concentration raised to the power of its stoichiometric coefficient. It is temperature-dependent.
Example:
For the reaction 2NO₂(g) ⇌ N₂O₄(g), the equilibrium constant K would be expressed as [N₂O₄]/[NO₂]².
Gibbs Free Energy (ΔG)
A thermodynamic potential that measures the 'useful' or process-initiating work obtainable from an isothermal, isobaric thermodynamic system. A negative ΔG indicates a spontaneous process.
Example:
For a reaction at equilibrium, the Gibbs Free Energy (ΔG) is zero, indicating no net driving force for either the forward or reverse reaction.
ICE Table
An acronym for Initial, Change, Equilibrium, a systematic method used to organize and calculate equilibrium concentrations of reactants and products in a reversible reaction.
Example:
To find the equilibrium concentrations of H₂ and I₂ after HI decomposes, you would set up an ICE table to track how their amounts change from initial conditions.
Le Châtelier's Principle
States that if a system at equilibrium is subjected to a stress (change in concentration, temperature, or pressure), the system will shift in a direction that relieves that stress.
Example:
If you add more reactant to a system at equilibrium, according to Le Châtelier's Principle, the reaction will shift to the product side to consume the added reactant.
Molar Solubility
The concentration of the metal cation (or anion, depending on stoichiometry) in a saturated solution of a sparingly soluble ionic compound, expressed in moles per liter (mol/L).
Example:
If the molar solubility of PbCl₂ is 0.01 M, it means 0.01 moles of PbCl₂ dissolve per liter of solution to form Pb²⁺ ions.
Product-favored reaction
A reaction at equilibrium where the concentration of products is significantly greater than the concentration of reactants, indicated by an equilibrium constant (K) value greater than 1.
Example:
A reaction with a K value of 1000 is a product-favored reaction, meaning at equilibrium, there will be a much higher amount of products than reactants.
Reactant-favored reaction
A reaction at equilibrium where the concentration of reactants is significantly greater than the concentration of products, indicated by an equilibrium constant (K) value less than 1.
Example:
If K for a reaction is 0.001, it's a reactant-favored reaction, implying that very little product forms at equilibrium.
Reaction Quotient (Q)
A value calculated using the concentrations of reactants and products at any given point in time, not necessarily at equilibrium, used to predict the direction a reaction will shift to reach equilibrium.
Example:
If you mix reactants and products and calculate a reaction quotient (Q) that is less than K, the reaction will proceed towards the products to reach equilibrium.
Reversible Reaction
A chemical reaction that can proceed in both the forward direction (reactants to products) and the reverse direction (products to reactants) simultaneously.
Example:
The synthesis of ammonia, N₂(g) + 3H₂(g) ⇌ 2NH₃(g), is a classic reversible reaction used in industrial processes.
Solubility Product (Ksp)
The equilibrium constant for the dissolution of a sparingly soluble ionic compound in water, representing the extent to which the solid dissolves to form ions.
Example:
The solubility product (Ksp) for AgCl is very small, indicating that only a tiny amount of AgCl dissolves in water.