Chemistry Study Notes on Equilibrium, Acids, Bases, and Titrations

Chemistry 6.1 Equilibrium and Le Chatelier’s Principle

• Reversible Reactions: Reactions where the products can revert back to the reactants.

  • Forward & Reverse: Molecules continuously shift between reactants and products.

• Equilibrium: The state where concentrations of reactants and products remain constant over time, with equal rates for forward and reverse reactions.

• Le Chatelier’s Principle:

  • If an external change is applied to a system at equilibrium, the system will shift in a way that counteracts the effect of the change.

Factors Affecting Equilibrium

1. Concentration

   • Adding a substance shifts the equilibrium to reduce its concentration.

   • Example: Adding more N2 to the system N2 + 3H2 ⇌ 2NH3 causes a shift to the right (producing more NH_3).

2. Pressure

   • Changes can occur by:

   1. Adding/removing a gaseous species (shifts similarly to concentration changes).

   2. Adding an inert gas (no effect on equilibrium).

   3. Changing a container’s volume:

      • Decreasing volume shifts equilibrium towards the side with fewer moles of gas.

      • Increasing volume shifts towards the side with more moles.

   • Example for volume increase: In N2 (g) + 3H2 (g) ⇌ 2NH_3 (g), it shifts left.

3. Temperature

   • Heating a system adds energy:

     • Endothermic Reaction: Energy treated as a reactant (shifts right with added heat).

     • Exothermic Reaction: Energy treated as a product (shifts left with added heat).

     • Example: Heating CaCO3 ⇌ CaO + CO2 (endothermic) shifts right.

Chemistry 6.2 Properties of Acids and Bases

• Properties:

  • Acids: Sour taste, react with metals, turn blue litmus red, neutralize bases, have H^+ ions.

  • Bases: Bitter taste, do not react with metals, turn red litmus blue, neutralize acids, have OH^- ions.

• Definitions:

  • Bronsted-Lowry Model:

  • Acids: Proton donors.

  • Bases: Proton acceptors.

  • Protons are defined as H^+ or hydronium H_3O^+.

Chemistry 6.3 Dissociation

• Dissociation: Process where an acid releases H^+ in water.

  • Not all hydrogen atoms in an acid are acidic; polyatomic ions may retain hydrogen.

  • Example: In acetic acid HC2H3O_2, only the first H is acidic.

• Acid Dissociation Reaction:

  • HA (aq) + H2O (l) ⇔ H3O^+ (aq) + A^- (aq).

• Strong vs. Weak Acids:

  • Strong acids fully dissociate in water, while weak acids partially dissociate.

Chemistry 6.4 pH Scale

• pH: Indicates acid/base strength by measuring hydronium concentration.

  • pH = - ext{log}([H^+])

  • Scale: 1-6: Acidic, 7: Neutral, 8-14: Basic.

• Calculating pH:

  • Example: For [H^+] = 0.5 M: pH = - ext{log}(0.5).

  • For given pH=2, calculate [H^+]:
    [H^+] = 10^{-2}.

• Calculating pOH:

  • pOH = - ext{log}([OH^-]),

  • Relation: 14 = pOH + pH,

  • 1 imes 10^{-14} = [H^+][OH^-].

Chemistry 6.6 Neutralization Reactions

• Neutralization Reaction: An acid reacts with a base to form water and salt:

  • ext{Acid} + ext{Base}
    ightarrow ext{H}_2O + ext{Salt}.

  • Example: HCl + NaOH
    ightarrow H_2O + NaCl.

Ionic and Net Ionic Equations

• Complete Equation: Shows all species (states included):
  HCl (aq) + NaOH (aq)
  ightarrow H_2O (l) + NaCl (aq).

• Ionic Equation: Breaks compounds into ions:
  H^+ (aq) + Cl^- (aq) + Na^+ (aq) + OH^- (aq)
  ightarrow H_2O (l) + Na^+ (aq) + Cl^- (aq).

• Net Ionic Equation: Removes spectator ions:
  H^+ (aq) + OH^- (aq)
  ightarrow H_2O (l).

Types of Reactions

1. Strong Acid + Strong Base: Always gives a neutral product (H^+ + OH^-
   ightarrow H_2O).

2. Weak Acid + Strong Base: Shifts toward basic solution.

3. Strong Acid + Weak Base: Produces acidic solutions.

4. Weak Acid + Weak Base: Behavior may vary; generally ignored.

Chemistry 6.7 Titrations

• Titration: Technique to determine the molarity of an unknown acid/base.

  • Requires one known concentration solution (standard).

  • Endpoint: Point of neutralization marked by an indicator color change.

Titration Process Steps:

1. Write a balanced equation.

2. Use known molarity to calculate moles of standard solution.

3. Use stoichiometric ratios to find moles of the unknown.

4. Calculate the molarity of the unknown solution.

• Example 1: Titrating LiOH to neutralize HBr.

• Example 2: Titration of H2SO4 with NaOH to endpoint.