Glossary
Arrhenius Definition
Defines acids as substances that produce H⁺ ions in water and bases as substances that produce OH⁻ ions in water.
Example:
According to the Arrhenius Definition, HCl is an acid because it releases H⁺ ions when dissolved in water.
Bronsted-Lowry Definition
Defines acids as proton (H⁺) donors and bases as proton acceptors.
Example:
In the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, water acts as a Bronsted-Lowry acid by donating a proton to ammonia.
Conjugate Base
The species formed after an acid donates its proton.
Example:
When the strong acid HCl donates its proton, the resulting Cl⁻ ion is its conjugate base, which is extremely weak.
Equilibrium
A dynamic state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in no net change in reactant and product concentrations.
Example:
In a closed bottle of soda, the dissolved CO₂ and gaseous CO₂ reach a state of equilibrium, where the rate of CO₂ dissolving equals the rate of CO₂ escaping.
Equilibrium Constant (K)
A ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients.
Example:
A large equilibrium constant (K) for a reaction indicates that products are heavily favored at equilibrium.
ICE/RICE Tables
A systematic method (Initial, Change, Equilibrium) used to calculate equilibrium concentrations of reactants and products in a reversible reaction.
Example:
To determine the final concentrations of ions in a weak acid solution, you would typically set up an ICE table.
Ka Expression
The mathematical representation of the acid dissociation constant, showing the ratio of dissociated ions to the undissociated acid at equilibrium.
Example:
For acetic acid (CH₃COOH), the Ka expression is [H⁺][CH₃COO⁻] / [CH₃COOH].
Kb Expression
The mathematical representation of the base dissociation constant, showing the ratio of protonated base and hydroxide ions to the unreacted base at equilibrium.
Example:
For ammonia (NH₃), the Kb expression is [NH₄⁺][OH⁻] / [NH₃].
Strong Acids
Acids that completely dissociate (ionize 100%) in water, releasing all their H⁺ ions.
Example:
When you add HNO₃ to water, it's a strong acid so every molecule breaks apart, making the solution very acidic.
Strong Bases
Bases that completely dissociate in water, releasing all their OH⁻ ions.
Example:
NaOH is a strong base; if you spill it, it will fully dissociate and create a highly alkaline solution.
Weak Acid Dissociation (Ka)
The equilibrium constant for the dissociation of a weak acid in water, indicating the extent to which it ionizes.
Example:
A very small Ka value for a carboxylic acid means it undergoes very little weak acid dissociation in water.
Weak Acids
Acids that only partially dissociate in water, existing in equilibrium between their undissociated and dissociated forms.
Example:
Citric acid in lemons is a weak acid; it gives lemons their sour taste but doesn't fully ionize in water.
Weak Base Dissociation (Kb)
The equilibrium constant for the reaction of a weak base with water, indicating the extent to which it produces hydroxide ions.
Example:
The Kb value helps predict how much a weak base dissociation will occur, influencing the pH of the solution.
Weak Bases
Bases that only partially dissociate in water, existing in equilibrium and accepting protons from water to a limited extent.
Example:
Ammonia (NH₃) is a common weak base used in cleaning products, as it only partially reacts with water to produce hydroxide ions.
pH
A measure of the acidity or alkalinity of a solution, defined as the negative logarithm (base 10) of the hydrogen ion concentration ([H⁺]).
Example:
A solution with a pH of 7 is neutral, while a solution with a pH of 2 is highly acidic.
pOH
A measure of the alkalinity of a solution, defined as the negative logarithm (base 10) of the hydroxide ion concentration ([OH⁻]).
Example:
If a solution has a high concentration of OH⁻ ions, its pOH will be low, indicating a basic solution.