Multiple Choice
Which of the following is the correct pH of a solution formed by mixing 1.3 L of 0.45 M HCl with 5.5 L of 1.95 M HI?
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17. Acid and Base Equilibrium
pH of Weak Acids
Multiple Choice
A student titrated a 100.0 mL sample of 0.100 M acetic acid with 0.050 M NaOH. (For acetic acid, Ka = 1.8 × 10^-5 at room temperature.) Calculate the initial pH of the acetic acid solution before any NaOH is added.
A
3.75
B
5.25
C
2.87
D
4.74
Verified step by step guidance1
Start by understanding that acetic acid (CH₃COOH) is a weak acid, and its dissociation in water can be represented by the equilibrium: CH₃COOH ⇌ CH₃COO⁻ + H⁺.
Use the acid dissociation constant (Ka) to set up the equilibrium expression: Ka = [CH₃COO⁻][H⁺] / [CH₃COOH]. Given Ka = 1.8 × 10⁻⁵.
Assume that the initial concentration of acetic acid is 0.100 M and that the change in concentration due to dissociation is 'x'. Therefore, at equilibrium, [CH₃COOH] = 0.100 - x, [CH₃COO⁻] = x, and [H⁺] = x.
Substitute these expressions into the equilibrium expression: 1.8 × 10⁻⁵ = (x)(x) / (0.100 - x). Since acetic acid is a weak acid, x is small compared to 0.100, so you can approximate 0.100 - x ≈ 0.100.
Solve for 'x', which represents [H⁺], and then calculate the pH using the formula: pH = -log₁₀([H⁺]).
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