12-Week 2025 Review

Vocab

100

This is an acid that does not fully dissociate.

What is a weak acid?

200

This is what a buffer is made of.

What is a weak acid/base and its conjugate?

300

Solutions that resist changes in pH when a strong acid or base is added

Buffer solution

400

This term that is used to describe a substance that is able to donate a hydrogen cation.

What is a Bronsted-Lowry acid?

500

What is the point in a titration at which the reaction between titrant and sample has just been completed called?

Equivalence Point

**CONCEPTS**

100

You would add this to a solution of ammonia (NH₃) to make it a buffer.

What is ammonium (NH₄⁺)?

200

2 NO₂ (g) <--->  N₂O₄ (g)     ΔHº = -58.0 kJ

Name four ways to favor the forward reaction (right).

add NO₂; remove N₂O₄

increase pressure; decrease volume

decrease temperature

300

What is the dominant species present at the equivalence point when titrating HNO₂ with NaOH?

NO₂^-

400

What type of titration results in a pH < 7 at equivalence point?

Titrating weak base with strong acid

500

Which direction will the reaction shift when the pressure is decreased on the following system?
3 H₂(g) + N₂(g) <--> 2 NH₃(g)

Left (reactants)

Acid / Base Math

100

pH of a 0.01M NaOH

[OH-] = 0.01M
pOH = -log (.01) = 2
pH = 14-2 = 12

200

pH of a 50 mL sample of 0.005 M HF at the beginning of a titration before NaOH is added. K_a = 1.43x10^-5

ICE Table: HA <-> H₃O⁺ + A^- 
K = x²/(.005-x)
x = [H₃O⁺]
-log (x) = 3.57

300

pH of 0.005M NH₃ K_b = 1.43x10^-5

B <-> OH^- + BH⁺ ICE Table
K = x²/(.005-x)
x = [OH-]
pOH = -log (x); 14-pOH = pH
pH = 10.43

400

12.5 mL of 3.0 M HCl is required to neutralize 25.0 mL of an unknown sample of NaOH. What is [NaOH]?

1. mol HCl = 3 mol/L * .0125L = 0.0375 mol
2. * 1mol OH-/mol HCl
3. / 0.025L = 1.5M

500

6.8 mL of 0.005 M Hg(NO₃)₂ is needed to reach the eq. point for a 25. mL sample. What is the original [Cl-] in ppm?
Hg²⁺ + 2Cl^- <-> HgCl₂ (aq)

0.005 mol Hg/L * 0.0068L * 2 mol Cl/mol Hg / .025L = 0.0027 M Cl

0.0027mol/L * 35.45 g/mol * 1000 mg/g = 96.4 ppm

Buffer Math

100

This is the pH of 50. mL buffer solution containing 2.0 M CH₃COOH and 2.0 M NaCH₃COO. K_a =1.8x10^-5.

4.74
pH = pKa + log(base/acid)
pH = -log(1.85E-5) + 0
pH = 4.74

200

What is the pH of a buffer solution containing 0.10 M propionic acid (C₂H₅COOH) (Ka = 1.35x10^-5) and 0.13 M of sodium propionate (NaC₂H₅OO)?

pH = pKa + log (base/acid)

pH = 4.87 + log(.13/.1)
pH = 4.98

300

0.200 M of acetic acid (Ka = 1.8x10^-5) is titrated with 0.200 M NaOH. What is the pH at half-eq. point?

pH = pKa = 4.74

400

To prepare a buffer at pH = 4.83, how many mol NaF must be added to 0.1 L of 0.123 M HF. (K_a of HF = 6.8 x 10^-4)

4.83 = 3.17 + log (mol F-/.0123 mol HF)
0.57 mol NaF

500

What is the pH for a buffer containing 1.00 mol NH₄⁺ and 1.00 mol NH₃ when 0.10 mol HCl is added? K_a = 5.9x10^-10

pH = pKa + log (base/acid)
pH = 9.22+log(0.9/1.1)
pH = 9.22 -0.0869
pH = 9.13

Salts and Ksp

100

NaNO₃ is a ______ salt.
KF is a _______ salt.
NH₄Br is a ______ salt.

neutral, basic, acidic

200

The K_sp expression for a saturated solution of Ca₃(PO₄)₂ is:

K_sp= [Ca²⁺]³[PO₄^3-]²

300

K_sp of MgF₂ = 6.4x10^-9. Determine [Mg²⁺] and [F^-].

MgF₂ <-> Mg²⁺ + 2F^-
K_sp = x(2x)² = 4x³
[Mg²⁺] = x = 1.2x10^-3 M
[F^-] = 2x = 2.4x10^-3 M

400

Determine K_sp of NiCO₃ if [Ni²⁺] = 3.5x10^-4 M at equilibrium.

NiCO₃(s) <-> Ni²⁺ + CO₃^2-
K_sp = x² =(3.5E-4)²
K_sp = 1.2x10^-7

500

Addition of acid will _____ the solubility of BaCO₃.

increase