Glossary
5% Approximation
A simplifying assumption used in equilibrium calculations, valid when the equilibrium constant (K) is very small, allowing 'x' to be neglected when added to or subtracted from a relatively large initial concentration.
Example:
When solving for 'x' in (1.0 - x) if K is tiny, the 5% approximation allows you to treat (1.0 - x) as approximately 1.0, thereby avoiding the quadratic formula.
Change in Concentration
The amount by which the concentration of a reactant or product increases or decreases as a reaction proceeds from its initial state to equilibrium, typically represented by 'x' in ICE tables.
Example:
In the reaction A → B, if A decreases by 0.1 M, then the change in concentration for A is -0.1 M, and for B it is +0.1 M.
Equilibrium
The state in a reversible chemical reaction where the rates of the forward and reverse reactions are equal, leading to constant concentrations of reactants and products.
Example:
When a bottle of soda is sealed, CO₂ gas dissolving into the liquid and CO₂ liquid escaping into the gas phase reach equilibrium, so the amount of dissolved CO₂ remains constant.
Equilibrium Concentrations
The final, stable concentrations of all reactants and products in a reversible reaction once the system has reached a state of chemical equilibrium.
Example:
After a reaction has settled, measuring the amounts of each substance gives you their equilibrium concentrations, which will remain constant over time.
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 large equilibrium constant (K > 1) for a reaction means that products are favored at equilibrium, while a small K (K < 1) means reactants are favored.
Equilibrium Expression
A mathematical equation that relates the equilibrium concentrations of products to reactants, with each concentration raised to the power of its stoichiometric coefficient.
Example:
For the reaction 2NO₂(g) ⇌ N₂O₄(g), the equilibrium expression is K = [N₂O₄] / [NO₂]².
ICE Tables
A systematic method (Initial, Change, Equilibrium) used to organize concentration data and solve for unknown equilibrium concentrations in chemical reactions.
Example:
To determine the final concentrations of N₂ and H₂ in the Haber process, you would set up an ICE table to track their changes from initial amounts to equilibrium.
Initial Concentrations
The concentrations of reactants and products present in a reaction mixture at the very beginning, before any significant reaction has occurred.
Example:
If you begin a reaction with 0.5 M of reactant A and 0 M of product B, these are your initial concentrations for A and B, respectively.
Products
The substances formed as a result of a chemical reaction.
Example:
When hydrogen and oxygen combine to form water (2H₂ + O₂ → 2H₂O), water is the product.
Quadratic Formula
A mathematical formula used to solve quadratic equations (ax² + bx + c = 0), which may be necessary in equilibrium problems when the 5% approximation is not valid.
Example:
If the 5% approximation fails, you'll need to rearrange your equilibrium expression into a quadratic equation and use the quadratic formula to solve for x.
Reactants
The starting substances in a chemical reaction that are consumed to form products.
Example:
In the combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O), methane and oxygen are the reactants.
Stoichiometric Coefficients
The numerical coefficients in a balanced chemical equation that indicate the relative number of moles of each reactant and product involved in the reaction.
Example:
In 2H₂ + O₂ → 2H₂O, the stoichiometric coefficient for H₂ is 2, meaning two moles of H₂ react for every one mole of O₂.
Weak Acid/Base Problems
Equilibrium problems involving acids or bases that only partially ionize in solution, requiring the use of ICE tables to determine equilibrium concentrations and pH.
Example:
Calculating the pH of a 0.1 M acetic acid solution is a classic weak acid problem that relies on an ICE table and the acid dissociation constant (Ka).