Chemical Equilibrium

Definition of Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products over time.

• Dynamic process: Both forward and reverse reactions continue to occur.

• No net change: Concentrations of all species remain unchanged.

• Example: In the reaction N2(g) + 3H2(g) → 2NH3(g), equilibrium is reached when the rate of ammonia formation equals its decomposition.

Equilibrium Constant Expression

The equilibrium constant (K) quantifies the ratio of product and reactant concentrations at equilibrium for a given reaction.

• General form: For a reaction aA + bB → cC + dD:

• Kc: Uses concentrations (mol/L).

• Kp: Uses partial pressures (atm) for gaseous reactions.

• Example: For 2SO2(g) + O2(g) → 2SO3(g):

Manipulating Equilibrium Constants

• Multiplying a reaction by a number: Raise K to that power.

• Reversing a reaction: Take the reciprocal of K.

• Adding reactions: Multiply their K values.

• Example: If reaction 1 has and reaction 2 has , the overall reaction's K is .

Le Châtelier's Principle

Le Châtelier's Principle predicts how a system at equilibrium responds to disturbances.

• Adding/removing reactants or products: Shifts equilibrium to counteract the change.

• Changing temperature: Shifts equilibrium depending on endothermic/exothermic nature.

• Changing pressure (gases): Shifts toward side with fewer moles of gas.

• Example: Increasing pressure in N2 + 3H2 → 2NH3 shifts equilibrium toward ammonia.

Calculating Equilibrium Concentrations

• ICE Tables: Used to organize Initial, Change, and Equilibrium concentrations.

• Steps:

  1. Write balanced equation and K expression.

  2. Set up ICE table.

  3. Solve for unknowns using K value.

• Example: For , if initial [A] = 1.0 M, [B] = 0, and , use ICE table to find equilibrium concentrations.

Acids and Bases

Definitions of Acids and Bases

• Arrhenius: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry: Acids donate protons (H+); bases accept protons.

• Lewis: Acids accept electron pairs; bases donate electron pairs.

• Example: NH3 is a Brønsted-Lowry base because it accepts H+.

pH, pOH, and pK Scales

• pH: Measures acidity;

• pOH: Measures basicity;

• Relationship: (at 25°C)

• pKa and pKb: ,

• Example: If M, .

Strengths of Acids and Bases

• Strong acids: Completely ionize in water (e.g., HCl, HNO3).

• Weak acids: Partially ionize (e.g., CH3COOH).

• Strong bases: Completely dissociate (e.g., NaOH).

• Weak bases: Partially dissociate (e.g., NH3).

• Example: HCl is a strong acid; acetic acid is a weak acid.

Polyprotic Acids

Polyprotic acids can donate more than one proton per molecule.

• Examples: H2SO4 (diprotic), H3PO4 (triprotic).

• Each ionization step has its own Ka value.

• Example: H2SO4 ionizes in two steps: first is strong, second is weak.

Determining Acidic or Basic Nature

• Salts: Some salts form acidic or basic solutions depending on their ions.

• Example: NaCl forms a neutral solution; NH4Cl forms an acidic solution.

• Rule: Salts from strong acid and strong base are neutral; salts from weak acid or base are not.

Chemical Thermodynamics (Brief Overview)

Thermodynamics in Chemical Reactions

Thermodynamics studies energy changes in chemical reactions, focusing on enthalpy, entropy, and free energy.

• Enthalpy (ΔH): Heat change at constant pressure.

• Entropy (ΔS): Measure of disorder.

• Gibbs Free Energy (ΔG): Determines spontaneity.

• Spontaneous reactions: Occur without external energy input (ΔG < 0).

• Example: Combustion of glucose is spontaneous.

Summary Table: Acid and Base Strengths

Additional info: Thermodynamics is only briefly mentioned in the study guide, so coverage here is limited to basic definitions and the Gibbs free energy equation.