AP Chemistry 8.3: Acid-Base Reactions and Neutralization

Acid-base reactions involve proton transfer, but calculating the final pH may require both stoichiometry and equilibrium.

• The most important rule for these problems is:
• Do the reaction stoichiometry first. Then do equilibrium calculations if necessary.
• This prevents many common mistakes.

Consider:

• HCl + NaOH → NaCl + H₂O
• The net ionic equation is:
• H₃O⁺ + OH⁻ → 2H₂O
• or in simplified form:
• H⁺ + OH⁻ → H₂O
• Because both acid and base are strong, they react essentially completely.
• Suppose:
• 50.0 mL of 0.100 M HCl
• is mixed with:
• 30.0 mL of 0.100 M NaOH.
• Calculate moles.
• HCl:
• 0.0500 L × 0.100 mol/L = 0.00500 mol H⁺
• NaOH:
• 0.0300 L × 0.100 mol/L = 0.00300 mol OH⁻
• They react 1:1.
• Remaining H⁺:
• 0.00500 − 0.00300 = 0.00200 mol
• Total volume:
• 50.0 + 30.0 = 80.0 mL = 0.0800 L
• Therefore:
• [H⁺] = 0.00200/0.0800
• [H⁺] = 0.0250 M
• Then:
• pH = −log(0.0250)
• pH ≈ 1.60
• The major lesson is that you cannot simply use the original acid concentration after mixing.
• You must account for:
• reaction + leftover moles + total volume

Consider:

• HA + OH⁻ → A⁻ + H₂O
• This reaction proceeds strongly toward products because OH⁻ removes a proton from the weak acid.

Before the equivalence point, some HA remains while A⁻ has been produced.

• That means the solution may become a:
• buffer
• containing a weak acid and its conjugate base.

At equivalence, essentially all HA has been converted into A⁻.

• The conjugate base A⁻ can react with water:
• A⁻ + H₂O ⇌ HA + OH⁻
• Therefore the equivalence-point solution is generally:
• basic
• for a weak-acid/strong-base titration.

Consider:

• B + H₃O⁺ → BH⁺ + H₂O

Before equivalence, the mixture can contain:

• B and BH⁺
• forming a buffer.

At equivalence, the weak base has been converted into its conjugate acid BH⁺.

• BH⁺ can react:
• BH⁺ + H₂O ⇌ B + H₃O⁺
• Therefore the equivalence point is generally:
• acidic
• for a weak-base/strong-acid titration.

When a weak acid reacts with a weak base, the direction and final pH depend on the relative strengths of the acid and base and the acid-base properties of the products.

• These situations require comparison of equilibrium constants rather than simply assuming the final solution is neutral.

Acid-base reactions generally favor formation of the:

• weaker acid and weaker base
• If one side contains a much stronger acid than the other, proton transfer tends to proceed away from that stronger acid.

Comparing Ka or pKa values can help predict the favored direction.

Never immediately construct a Ka or Kb ICE table when a strong acid/base is reacting stoichiometrically with another species.

First determine how many moles react.

Then identify what remains.

Then determine what controls pH.

For acid-base mixture problems:

1. Convert to moles

2. Perform reaction stoichiometry

3. Determine what remains

4. Divide by total volume if needed

5. Calculate pH using the chemistry of the remaining species

Whenever strong acid or strong base is added, you often need to perform a nearly complete reaction before doing equilibrium calculations.

• This gives one of the most important Unit 8 rules:
• Stoichiometry first. Equilibrium second.
• For example:
• HA + OH⁻ → A⁻ + H₂O
• First calculate how many moles of HA and OH⁻ react.
• Then determine how much HA and A⁻ remain.
• Only then should you use a buffer equation or equilibrium calculation.