Lecture Notes on Chemical Equilibrium and Acid-Base Chemistry

Chemical Equilibrium

• Key Concepts Regarding Chemical Equilibrium:
  • At equilibrium, Q (reaction quotient) can be larger than or less than K (equilibrium constant) depending on direction of reaction.
  • Both forward and reverse reactions occur at the same rate, ensuring no net change in concentrations.
  • K is defined as the ratio of concentrations of products to reactants at equilibrium:
    K = \frac{[products]}{[reactants]}
  • K is expressed in units of molarity (M), equivalent to mol/L.
  • The rates of the forward and reverse reactions are equal at equilibrium: Rate{fwd} = Rate{rev}
  • The value of K varies greatly for different reactions.
  • Pure solids and liquids do not change the value of K.
  • Changes in concentration (adding/removing substances), temperature (adding/removing heat), or pressure/volume can disrupt equilibrium.

Reaction Stoichiometry

• Determining Reaction Progress:
  • For the reaction: Cl2(g) + H2(g) \rightleftharpoons 2 HCl(g), the change in moles of gas, ∆n, is calculated from products minus reactants:
  • ∆n = 0

Chemical Kinetics and Reactions

• Le Chatelier's Principle:
  • Changes in conditions affect the direction of equilibrium. Knowing how certain changes affect equilibrium is essential:
  1. Removing product (C): Shifts right to produce more product.
  2. Adding heat: For an exothermic reaction (ΔHrxn = -126 kJ/mol), equilibrium shifts left (endothermic direction).
  3. Increasing pressure: Shifts towards the side with fewer moles of gas.
  4. Cooling the reaction: Shifts right (favoring formation of products).
  5. Increasing volume: Shifts towards the side with more moles of gas.
  6. Decreasing pressure: Shifts towards the side with more moles of gas.

Acid-Base Equilibria

• Introduction to Chapter 18 Acid-Base Equilibria:
  • Topics: Sections 18.1 to 18.8
  • pH and Definitions: Some concepts covered in General Chemistry 1.

Acid-Base Concepts

• Acid Rain:
  • Normal rain has pH due to CO₂ (pH = 5.2)
  • Acid rain is characterized by pH less than 5, primarily due to pollutants like NOx and SOx.
  • Reactions involving acid rain formation include:
    $ H_2O (l) + CO_2 (g) ⇌ H_2CO_3 (aq)$
$ H_2O (l) + SO_3 (g) ⇌ H_2SO_3 (aq)$

Acid-Base Definitions

• Common Acid-Base Definitions:
  1. Arrhenius:
  • Acid: Produces H_3O^+ in water.
  • Base: Produces OH^- in water.
  1. Bronsted-Lowry:
  • Acid: Proton donors.
  • Base: Proton acceptors.
  1. Lewis:
  • Acid: Electron pair acceptor.
  • Base: Electron pair donor.

Reactions and Strengths of Acids/Bases

• Neutralization Reaction:

  • Reaction of a strong acid and a strong base produces water and a salt:
  • Net Reaction:
    H^+(aq) + OH^-(aq) ⇌ H_2O(l)
    \Delta H° = -55.9 \, kJ

• Strong Acids and Bases:

  • Strong acids completely dissociate in water.
  • Example: HNO3 + H2O \rightarrow H3O^+ + NO3^-
  • Weak acids partially dissociate in water.
  • Example:
    HNO2 + H2O \rightleftharpoons NO2^- + H3O^+

Acid-Dissociation Constant (KA)

• Definition: Measures acid strength.
• General equation for KA is:
  Ka = \frac{[H3O^+][A^-]}{[HA]}
• Relationship between strength of the acid and KA:
  • Stronger acids have higher [H3O^+] and larger Ka.
  • Weaker acids have lower Ka and lower % dissociation of HA.

Acid Strength Tables

• Weak Acids and their K values:
  • Chlorous Acid (HClO2): K = 1.1 x 10^{-2}
  • Hydrofluoric Acid (HF): K = 6.8 x 10^{-4}
  • Formic Acid (HCOOH): K = 1.8 x 10^{-4}
  • Acetic Acid (CH3COOH): K = 1.8 x 10^{-5}
• Strong Acids Include:
  • Nitric Acid (HNO3), Sulfuric Acid (H2SO4), Hydrochloric Acid (HCl).

pH Scale and Equilibrium

• pH Scale: Measures [H3O+].
  • pH = -\log[H_3O^+]
  • Neutral solution: pH at 7.0, with [H3O+] and [OH-] equal to 1.0 x 10^{-7} M.
• Relationship with pOH:
  • pOH = 14 - pH.

Calculating [H3O+], pH, [OH-] and pOH

• Use relations involving logarithmic functions to calculate these values from one another:
  • pHo = -log[OH^-]
  • Relationship: Kw = [H3O^+][OH^-], where K_w = 1.0 × 10^{-14} at 25 °C.
• Example Calculation:
  • pH of a solution of 0.15 M HCl:
  • Complete dissociation gives: [H_3O^+] = 0.15.
  • pH = -\log [0.15] = 0.823.

Acids and their Strengths

• Brønsted-Lowry Definitions:
  • Acid: Proton donor (H+ donor).
  • Base: Proton acceptor (H+ acceptor).
• Conjugate Pairs:
  • Acid-Base reactions produce conjugate acids/bases. An acid becomes its conjugate base after donation.

Calculating pH for Weak Acids/Bases

• Use an ICE table to analyze weak acid or base equilibria and calculate pH from H3O+ concentrations or use given Ka/Kb for other solutions.
• QC = [H3O+][A-] / [HA] to determine Ka for the case of weak acids.

Polyprotic Acids

• Concept: Acids with more than one protons to donate, with each step in dissociation being weaker.
  • H3PO4 \rightleftharpoons H2PO4^- + H_3O^+
  • Dissociation constants K{a1} > K{a2} > K_{a3} for these reactions.

Salt Solutions and Their pH

• Effect of Salts on pH: Salts may yield acidic, neutral, or basic solutions depending on components.
• Acidic solutions: contain conjugate acids.
• Basic solutions: contain conjugate bases from weak acids.
• Neutral solutions: both strong acid and base components present.

Conclusion

• Understanding equilibrium principles, acid-base definitions, pH calculations, and linked concepts are fundamental for mastery in chemical equilibria and acid-base chemistry.
• Master formulas are to apply them in problem-solving contexts and practice calculating various equilibrium states in chemical systems.