General Chemistry Study Guide: Chemical Equilibrium, Acids and Bases, and Thermodynamics

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:

$K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Kc: Uses concentrations (mol/L).

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

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

$K = \frac{[SO_3]^2}{[SO_2]^2[O_2]}$

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 $K_1$ and reaction 2 has $K_2$, the overall reaction's K is $K_1 \times K_2$.

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 $A \rightleftharpoons B$, if initial [A] = 1.0 M, [B] = 0, and $K = 0.5$, 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; $pH = -\log[H^+]$

• pOH: Measures basicity; $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$ (at 25°C)

• pKa and pKb: $pK_a = -\log K_a$, $pK_b = -\log K_b$

• Example: If $[H^+] = 1 \times 10^{-3}$ M, $pH = 3$.

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.

$\Delta G = \Delta H - T\Delta S$

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

• Example: Combustion of glucose is spontaneous.

Summary Table: Acid and Base Strengths

Type	Examples	Ionization/Dissociation
Strong Acid	HCl, HNO3, H2SO4	Complete
Weak Acid	CH3COOH, HF	Partial
Strong Base	NaOH, KOH	Complete
Weak Base	NH3, C5H5N	Partial

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.