CHM1046 Exam 3 Review Notes

CHM1046 Exam 3 Review Notes

Lecture 16: Basics of Acid-Base Chemistry

• Definitions:

  • Arrhenius Theory:
  • Acid: Donates H+ ions
  • Base: Donates OH- ions
  • Brønsted-Lowry Theory:
  • Acid: Donates H+ ions
  • Base: Accepts H+ ions
  • Lewis Theory:
  • Acid: Accepts electron pairs
  • Base: Donates electron pairs

• Ionization Constant (Kw):

  • Core Formula: Kw = [H+][OH-] = 1.0 x 10^-14 (at 25 °C)
  • Important Note: 1 unit change in pH = 10x change in acidity level (e.g., pH 1 is 10x more acidic than pH 2).

• Equations:

  • pH = -log[H+]
  • pOH = -log[OH-]
  • [H+] = 10^(-pH)
  • [OH-] = 10^(-pOH)
  • Shortcut: -log(10.67) = 1.02816 (significant figures apply).

• Strong Acids/Bases:

  • Strong acids and bases completely dissociate in water.

Lecture 16 Questions

1. • Given [H+] = 4.0 x 10^-7 M, calculate pOH.
2. • Find pH for a 0.001 M solution of H2SO4.
3. • Identify Brønsted-Lowry bases in the reaction: HCO₃⁻(aq) + CN⁻(aq) ⇌ CO₃²⁻(aq) + HCN(aq).

Lecture 17: Neutralization and Expected pH

• Neutralization:

  • Reaction of acids and bases to form salt and water.
  • Steps:
  • Identify acid (H) and base, look for products (H2O + salt).
  • Convert pH to H+ concentration.
  • Account for volume (mL to L) and using molarity to find moles.
  • Balance equations using molar ratios.

• Calculating pH:

  • Use acid dissociation constant (Ka) or base dissociation constant (Kb).
  • Note: Use Kw for conversions between acid and base concentrations.
  • Determine if a species is basic or acidic based on its dissociation in water.

Lecture 17 Questions

4. Determine if LiCNO is neutral, acidic, or basic, and write equilibrium equations as necessary.
5. Calculate pH of a 0.320 M solution of Ca(NO₂)₂ (Ka of HNO₂ = 4.5 × 10⁻⁴).
6. Calculate volume of 7.6 M HI needed for a solution with pH of 3.20.

Lecture 18: Weak Acids and Bases

• Weak Acids/Bases:

  • Dissociation equations:
  • Weak Acid: HA + H2O ⇌ A- + H3O+
  • Weak Base: B- + H2O ⇌ HB + OH-
  • Strength Indicators:
  • Larger Ka or Kb indicates stronger acids or bases.
  • Smaller pKa or pKb indicates stronger acids or bases.

• Polyprotic Acids:

  • Have multiple dissociation steps, each with different Ka values.
  • Ka1 > Ka2 > Ka3.
  • Amphoteric substances can act as acid or base, e.g., H2O.

Lecture 18 Questions

7. Find percent ionization in 0.300 M solution of formic acid (Ka = 1.78 × 10⁻⁴).
8. Calculate pH of 0.47 M carbonic acid (H₂CO₃) using given Ka values.

Lecture 19: Buffers

• Buffers:

  • Mixtures of a weak acid and its conjugate base or vice versa, resist pH changes during acid/base additions.
  • Buffer capacity refers to the amount of acid/base a buffer can resist before a drastic pH change occurs.

• Finding Buffer Capacity:

  • Eliminate non-buffer pairs.
  • Evaluate moles in conjugate pairs.
  • Compare the smallest numbers to find the largest buffer capacity.

• Henderson-Hasselbalch Equation:

  • Requires knowledge of pKa and concentration ratios of acid/base.

• Adding strong acids/bases to buffers alters concentrations of the conjugate pairs but does not drastically shift pH.

Lecture 19 Questions

9. Identify combinations for buffers.
10. Calculate pH of given buffer systems with added NaOH.

Lecture 20: Titration and Equivalence Points

• Titration concepts: pH = pKa at the half equivalence point.
• Know how to read titration curves and identify equivalence points in reactions.

Lecture 20 Questions

11. Calculate pH for mixing HBr and CH₃NH₂.
12. Determine pH during titration of acetic acid with NaOH.
13. Find the pH at the equivalence point in the titration of HCOONa with HCl.

Lecture 21: Solubility and Ksp

• Solubility Product Constant (Ksp):

  • Indicates how much solute can dissolve in a solution.
  • Higher Ksp means more soluble substances.
  • Molar solubility refers to the molarity of solute that can dissolve.

• Ksp Calculations:

  • Set up concentration expressions based on balanced dissolution reactions.
  • Use ICE tables where needed to relate concentrations.

Lecture 21 Questions

15. Write Ksp expressions for silver cyanide and SrF2.
16. Calculate Ksp given molar solubility for Ni3(PO4)2.

Lecture 22: Le Chatelier's Principle

• Applying Le Chatelier's Principle to solubility, consider how changes in conditions (such as pH) can impact solubility of compounds.
• Example: Adding NaF decreases the solubility of PbF2.

Lecture 22 Questions

17. Analyze the effects of various changes on solubility of ammonium nitrate.

Lecture 23: Multiple Equilibria

• Understand how to manipulate reactions to find new equilibrium constants, including reversing reactions, adjusting coefficients, and combining reactions.

Lecture 23 Questions

18. Find the new equilibrium constant for reactions with modified coefficients.

Lecture 24: Spontaneity & Entropy

• Spontaneity: Refers to whether a process can occur without external energy.
• Entropy: Measures randomness, generally increasing from solids to gases; calculated as products minus reactants.

Lecture 24 Questions

19. Calculate entropy for vaporizing CCl₂F₂ at 25°C.
20. Find ∆S° for the reaction of NH₃ and HCl.

Lecture 25: Gibbs Free Energy

• Understanding Gibbs free energy (∆G) is essential to predict spontaneity and reaction direction:

  • ∆G < 0: spontaneous
  • ∆G > 0: non-spontaneous

• Delta G formula manipulations for different conditions, including equilibrium and concentration dependencies.

Lecture 25 Questions

21. Estimate equilibrium constant based on Gibbs free energy.
22. Calculate ∆G for the reaction at specified conditions.