Acids, Bases, Kinetics, and Equilibrium: Study Guide for Introductory Chemistry

Acids and Bases

Definitions and Properties

This section covers the fundamental definitions and properties of acids and bases, including their identification and behavior in chemical reactions.

• Arrhenius Definition: An acid is a substance that increases the concentration of H+ ions in aqueous solution, while a base increases the concentration of OH- ions.

• Bronsted-Lowry Definition: An acid is a proton (H+) donor, and a base is a proton acceptor.

• Conjugate Acid-Base Pairs: When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

Example: In the reaction HCl + H2O → Cl- + H3O+, HCl is the acid and Cl- is its conjugate base.

Strength of Acids and Bases

The strength of an acid or base depends on its ability to dissociate in water.

• Strong acids/bases dissociate completely in water (e.g., HCl, NaOH).

• Weak acids/bases only partially dissociate (e.g., acetic acid, ammonia).

Example: HCl is a strong acid, while CH3COOH is a weak acid.

Identifying Acids, Bases, and Conjugate Pairs

To identify acids, bases, and their conjugate pairs in a reaction, look for the species that donate or accept protons.

• Acid-base reactions can be written as: $\mathrm{HA} + \mathrm{B} \rightarrow \mathrm{A}^- + \mathrm{HB}^+$

Predicting Acid-Base Reactions

Acid-base reactions typically proceed from stronger acids and bases to weaker ones. The direction of the reaction can be predicted by comparing the relative strengths of acids and bases involved.

pH and pOH Calculations

The pH of a solution measures its acidity, while pOH measures its basicity.

• $\mathrm{pH} = -\log_{10}[\mathrm{H}^+]$

• $\mathrm{pOH} = -\log_{10}[\mathrm{OH}^-]$

• $\mathrm{pH} + \mathrm{pOH} = 14$ \ (at 25\degree C)

Example: If [H+] = 1 × 10-3 M, then pH = 3.

Neutralization Reactions

When an acid reacts with a base, they form water and a salt in a neutralization reaction:

• $\mathrm{Acid} + \mathrm{Base} \rightarrow \mathrm{Salt} + \mathrm{Water}$

Example: HCl + NaOH → NaCl + H2O

Hydronium Ion Concentration and pH

The concentration of hydronium ions ([H3O+]) determines the pH of a solution.

• To calculate pH from [H3O+]: $\mathrm{pH} = -\log_{10}[\mathrm{H}_3\mathrm{O}^+]$

Acid-Base Titrations

Titration is a technique to determine the concentration of an acid or base by reacting it with a solution of known concentration.

• At the equivalence point, moles of acid = moles of base.

• Calculation: $M_1V_1 = M_2V_2$ (for monoprotic acids and bases)

Example: 25.0 mL of 0.10 M NaOH neutralizes 50.0 mL of HCl. The concentration of HCl is 0.05 M.

Kinetics

Factors Affecting Reaction Rate

Chemical kinetics studies the speed of chemical reactions and the factors that affect them.

• Concentration: Increasing reactant concentration generally increases reaction rate.

• Temperature: Higher temperature increases reaction rate by providing more energy for collisions.

• Catalysts: Catalysts speed up reactions without being consumed.

• Surface Area: Greater surface area increases reaction rate for solids.

Collision Theory

According to Collision Theory, molecules must collide with sufficient energy and proper orientation to react.

• Only a fraction of collisions result in a reaction (those with energy ≥ activation energy).

Equilibrium

Le Chatelier's Principle

Le Chatelier's Principle states that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance and restore equilibrium.

• Disturbances include changes in concentration, temperature, or pressure.

• The system shifts to minimize the effect of the change.

Predicting Shifts in Equilibrium

When a chemical system at equilibrium is subjected to a change, the direction of the shift can be predicted:

• Adding reactant: Shifts equilibrium toward products.

• Removing product: Shifts equilibrium toward products.

• Increasing temperature: Shifts equilibrium in the endothermic direction.

• Increasing pressure (for gases): Shifts equilibrium toward the side with fewer moles of gas.

Example: For the reaction $\mathrm{N}_2 + 3\mathrm{H}_2 \rightleftharpoons 2\mathrm{NH}_3$, increasing pressure shifts equilibrium toward NH3.