#AP Chemistry: Acids and Bases - The Night Before 🌙

Hey, future chemist! Let's get you prepped and confident for your AP Chemistry exam. We're diving into the world of acids and bases, focusing on those tricky weak ones. Buckle up, and let's make this stick!

#Acid-Base Basics: Strong vs. Weak

#

• Arrhenius Definition: Acids produce H⁺ ions in water, and bases produce OH⁻ ions.
• Bronsted-Lowry Definition: Acids are proton (H⁺) donors, and bases are proton acceptors.
  • For our purposes, we'll primarily use the Arrhenius definition to keep things simple.

#

• Strong Acids: Completely dissociate in water (virtually 100% ionization).
  • Memorize these 7: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, HClO₄
  • They break down into H⁺ and their conjugate base (e.g., HCl → H⁺ + Cl⁻).
• Strong Bases: Also completely dissociate in water.
  • Alkali and alkaline earth metal hydroxides: NaOH, KOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂
  • They break down into metal cation and OH⁻ (e.g., NaOH → Na⁺ + OH⁻).

Image: Visual representation of strong acids fully dissociating and weak acids partially dissociating.

#

• Weak Acids/Bases: Incompletely dissociate in water. They exist in equilibrium between the undissociated and dissociated forms.
• Weak Acid Example: Acetic acid (CH₃COOH) is a classic weak acid.
  • Only some of it breaks down into H⁺ and CH₃COO⁻.
• Weak Base Example: Ammonia (NH₃) is the most common weak base.
  • Reacts with water to form NH₄⁺ and OH⁻.
• Key takeaway: Any acid not in the list of 7 strong acids is weak.

#Equilibrium: A Quick Refresher

#

• Dynamic State: Forward and reverse reactions occur at the same rate. No net change in concentrations of reactants and products.
• Not Static: Reactions are still happening, but the overall concentrations remain constant.

Image: Graph showing reactant and product concentrations reaching a steady state at equilibrium.

#

• Ratio of Products to Reactants: K = [Products]^coefficients / [Reactants]^coefficients
• Indicates Reaction Extent:
  • Large K: Favors products
  • Small K: Favors reactants
• ICE/RICE Tables: Used to find equilibrium concentrations. Remember to use initial concentrations, changes, and equilibrium concentrations.

#Weak Acid-Base Equilibria: Putting It All Together

#

• Ka Expression: HA ⇌ H⁺ + A⁻. Ka = [H⁺][A⁻] / [HA]
• Smaller Ka = Weaker Acid: Indicates less dissociation.

Example Problem: Find the pH of a 2M solution of CH₃COOH (Ka = 1.8 × 10⁻⁵).

1. Write the Equilibrium: CH₃COOH ⇌ H⁺ + CH₃COO⁻

2. Set up ICE Table:

   	CH₃COOH	H⁺	CH₃COO⁻
   Initial (I)	2	0	0
   Change (C)	-x	+x	+x
   Equil (E)	2-x	x	x

3. Ka Expression: 1.8 × 10⁻⁵ = x² / (2-x) ≈ x² / 2 (since x is small)

4. Solve for x: x = [H⁺] ≈ 0.006

5. Calculate pH: pH = -log(0.006) = 2.22

Image: ICE table setup for a weak acid dissociation.

#

• Kb Expression: B + H₂O ⇌ BH⁺ + OH⁻. Kb = [BH⁺][OH⁻] / [B]
• Smaller Kb = Weaker Base: Indicates less dissociation.

Example Problem: Find the pH of 1M ammonia (NH₃). (Kb = 1.8 × 10⁻⁵).

1. Write the Equilibrium: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

2. Set up ICE Table:

   	NH₃	NH₄⁺	OH⁻
   Initial (I)	1	0	0
   Change (C)	-x	+x	+x
   Equil (E)	1-x	x	x

3. Kb Expression: 1.8 × 10⁻⁵ = x² / (1-x) ≈ x² / 1

4. Solve for x: x = [OH⁻] ≈ 0.0042

5. Calculate pOH: pOH = -log(0.0042) = 2.38

6. Calculate pH: pH = 14 - 2.38 = 11.62

Image: ICE table setup for a weak base dissociation.

• High-Priority Topics:
  • Distinguishing between strong and weak acids/bases
  • Setting up and solving equilibrium problems using ICE tables
  • Calculating pH for both weak acids and weak bases
  • Understanding the meaning of Ka and Kb
• Common Question Types:
  • Multiple choice questions asking for pH calculations
  • Free response questions involving equilibrium calculations and analysis of weak acid/base systems
• Time Management:
  • Practice setting up ICE tables quickly.
  • Make sure you are comfortable with log calculations.
  • When possible, use the approximation (x is small) to simplify calculations.
• Strategies:

  • Always write out the balanced equation first.

  • Double-check whether you're solving for pH or pOH.

  • Pay close attention to the given values (Ka, Kb, initial concentrations).

Practice Question

#Practice Questions

#Multiple Choice Questions

1. Which of the following is a strong acid? (A) CH₃COOH (B) HF (C) HNO₃ (D) H₂CO₃

2. A 0.1 M solution of a weak base has a pH of 11. What is the pOH of the solution? (A) 1 (B) 3 (C) 7 (D) 11

3. Which of the following statements is true regarding the dissociation of a weak acid in water? (A) The concentration of H+ ions is equal to the initial concentration of the acid. (B) The equilibrium favors the formation of the undissociated acid. (C) The equilibrium constant Ka is very large. (D) The acid completely dissociates into ions.

#Free Response Question

A 0.20 M solution of a weak acid, HA, has a pH of 3.00. (a) Write the equilibrium reaction for the dissociation of HA in water.

(b) Calculate the [H+].

(c) Calculate the Ka of the weak acid.

(d) If the concentration of the weak acid is increased, how will the pH of the solution change? Explain.

Scoring Breakdown:

(a) 1 point for correct equilibrium reaction: HA(aq) ⇌ H+(aq) + A-(aq)

(b) 1 point for correct [H+] calculation: [H+] = 10^-3 M = 0.001 M

(c) 3 points for correct Ka calculation: * 1 point for setting up the ICE table * 1 point for correct Ka expression: Ka = [H+][A-] / [HA] * 1 point for correct calculation: Ka = (0.001)(0.001) / (0.20 - 0.001) ≈ 5.0 x 10^-6

(d) 2 points for correct explanation: * 1 point for stating that the pH will decrease. * 1 point for explaining that increasing the concentration of the weak acid will increase the [H+] in the solution, thus lowering the pH.