1 question, one for each idea where we can. Answer them, then see which ideas to fix.
Question 1 of 1
HA is a weak acid. Which of the numbered diagrams shown best represents a buffer solution? Water molecules, Na⁺ ions, and species present only at very low concentrations are not shown.
Answer and reasoning
ADiagram 1 A student who thinks a weak acid alone is a buffer because it is in equilibrium with its conjugate base picks this. A weak acid alone supplies almost no A⁻, so nothing is present to react with added H₃O⁺.
BDiagram 2 A student who thinks a buffer is anything that neutralizes added acid picks this. A⁻ alone can react with added H₃O⁺, but no HA is present to react with added OH⁻.
CDiagram 3Correct A buffer contains large concentrations of both a weak acid and its conjugate base. The box with four HA molecules and four A⁻ ions has both: HA can react with added OH⁻ and A⁻ with added H₃O⁺.
DDiagram 4 A student who thinks any acid with its conjugate base forms a buffer picks this. HCl is a strong acid, and Cl⁻ does not react with added H₃O⁺, so this solution of HCl and NaCl cannot stabilize pH.
Working No calculation. A buffer contains large concentrations of both members of a conjugate pair of a weak acid: HA and A⁻ together. The box with four HA and four A⁻ fits; the others show a weak acid alone, a conjugate base alone, and a strong acid with its conjugate base, Cl⁻.
In preparation: 0 of 1 sections compiled and reviewed. The rest show key terms and common mistakes from our question bank until they are.
8.8.A.1 Buffer solution Fix
Buffer solution
A solution that contains large concentrations of both members of a conjugate acid-base pair, such as a weak acid HA and its conjugate base A⁻. It stabilizes pH: adding a small amount of strong acid or strong base changes its pH much less than the same addition to water.
Conjugate acid-base pair
Two species that differ by one proton, H⁺, such as CH₃COOH and CH₃COO⁻ or NH₄⁺ and NH₃. In a buffer the acid member is a weak acid and the base member is a weak base.
Buffer reaction with added acid
The conjugate base of the buffer reacts with added strong acid: A⁻(aq) + H₃O⁺(aq) → HA(aq) + H₂O(l). The reaction proceeds essentially to completion, so the added H₃O⁺ is replaced by the weak acid HA.
Buffer reaction with added base
The conjugate acid of the buffer reacts with added strong base: HA(aq) + OH⁻(aq) → A⁻(aq) + H₂O(l). The reaction proceeds essentially to completion, so the added OH⁻ is replaced by the weak base A⁻.
Why a buffer stabilizes pH
Because added H₃O⁺ or OH⁻ is consumed by a buffer component and replaced by a weak acid or weak base, [H₃O⁺] changes much less than it would in unbuffered water; only the relative amounts of HA and A⁻ change.
Students often think A solution of a weak acid alone is a buffer, because the acid is in equilibrium with its conjugate base. In fact No. A weak acid in water is mostly un-ionized, so its conjugate base is present only at a very low concentration. A buffer needs large concentrations of both the weak acid and its conjugate base, so that there is enough of each to react with added base and added acid.
Students often think A buffer is a solution that neutralizes added acid, so a solution of a base such as A⁻ alone is a buffer. In fact No. A solution of A⁻ alone can react with added acid, but it contains almost no HA to react with added base, so added OH⁻ stays in solution and the pH rises sharply. A buffer needs large concentrations of both HA and A⁻.
4 more questions. Every wrong answer here is a real mistake students make, and you see why it is wrong as soon as you answer.
Question 1 of 4
A buffer solution contains 0.100 mol of CH₃COOH and 0.100 mol of CH₃COO⁻ (from sodium acetate) in 1.00 L. Then 0.010 mol of HCl is added, with negligible change in volume. What amounts of CH₃COOH and CH₃COO⁻ are present after the reaction is complete?
Answer and reasoning
ACH₃COOH: 0.090 mol; CH₃COO⁻: 0.110 mol A student who applies Le Châtelier's principle in the wrong direction, shifting CH₃COOH ⇌ H₃O⁺ + CH₃COO⁻ toward CH₃COO⁻ when H₃O⁺ is added, picks this. Adding H₃O⁺, a product, shifts the equilibrium toward CH₃COOH.
BCH₃COOH: 0.100 mol; CH₃COO⁻: 0.100 mol A student who thinks a buffer keeps its components unchanged when acid is added picks this. The buffer's pH changes little because the CH₃COO⁻ reacts with the added H₃O⁺, which changes the amounts of both components.
CCH₃COOH: 0.090 mol; CH₃COO⁻: 0.100 mol A student who thinks the added acid reacts with the acid component of the buffer picks this. H₃O⁺ is a proton donor and reacts with the proton acceptor, CH₃COO⁻, not with CH₃COOH.
DCH₃COOH: 0.110 mol; CH₃COO⁻: 0.090 molCorrect The added H₃O⁺, 0.010 mol, reacts with the conjugate base: CH₃COO⁻ + H₃O⁺ → CH₃COOH + H₂O. CH₃COO⁻ falls to 0.090 mol and CH₃COOH rises to 0.110 mol.
Working HCl is a strong acid, so 0.010 mol H₃O⁺ is added. CH₃COO⁻(aq) + H₃O⁺(aq) → CH₃COOH(aq) + H₂O(l) goes essentially to completion: CH₃COO⁻ 0.100 − 0.010 = 0.090 mol; CH₃COOH 0.100 + 0.010 = 0.110 mol.
Equal small amounts of HCl(aq) are added to 100 mL of pure water and to 100 mL of a buffer that is 0.10 M in CH₃COOH and 0.10 M in NaCH₃COO. The pH of the water falls by several units, but the pH of the buffer falls by much less. Which statement best explains the difference?
Answer and reasoning
AThe added H₃O⁺ reacts with CH₃COO⁻ to form CH₃COOH, so a weak acid replaces the added H₃O⁺Correct The conjugate base of the buffer reacts with added acid: CH₃COO⁻ + H₃O⁺ → CH₃COOH + H₂O. Nearly all of the added H₃O⁺ is consumed and replaced by a weak acid, so [H₃O⁺] rises only slightly.
BThe added H₃O⁺ reacts with CH₃COOH, the acid in the buffer, so that little H₃O⁺ remains free A student who thinks added acid reacts with the acid component of the buffer picks this. CH₃COOH is a proton donor and does not react with H₃O⁺; the proton acceptor CH₃COO⁻ consumes the added H₃O⁺.
CThe added H₃O⁺ is spread out among the buffer's dissolved ions, so little H₃O⁺ remains free A student who thinks a buffer works physically, by spreading out or absorbing added ions, picks this. Both liquids have the same volume; the buffer keeps [H₃O⁺] low because its CH₃COO⁻ reacts with the added H₃O⁺.
DThe added H₃O⁺ cancels out with CH₃COO⁻, forming only water, so no acid is left in the solution A student who thinks an acid and a base cancel out in neutralization, leaving only water, picks this. A proton is transferred from H₃O⁺ to CH₃COO⁻: the products are CH₃COOH and H₂O, and the weak acid CH₃COOH remains in the solution.
Working No calculation. In the buffer, CH₃COO⁻(aq) + H₃O⁺(aq) → CH₃COOH(aq) + H₂O(l) consumes nearly all of the added H₃O⁺, replacing it with the weak acid CH₃COOH, which ionizes only slightly. In water, nothing consumes the added H₃O⁺.
The diagram represents a small volume of a buffer solution of the weak acid HA and its conjugate base, A⁻, and two OH⁻ ions (from NaOH) that are added to it. Water molecules and Na⁺ ions are not shown. Which describes the acid-base particles present in this volume after the reaction is complete?
Answer and reasoning
AHA: 5; A⁻: 5; OH⁻: 2 A student who thinks a weak acid reacts with only a small fraction of added strong base picks this. HA is weak in water, but its reaction with OH⁻ goes essentially to completion, so no OH⁻ remains.
BHA: 3; A⁻: 7; OH⁻: 0Correct The weak acid of the buffer reacts with added base: HA + OH⁻ → A⁻ + H₂O, essentially to completion. The two OH⁻ ions convert two HA molecules to two A⁻ ions: 5 − 2 = 3 HA and 5 + 2 = 7 A⁻.
CHA: 3; A⁻: 5; OH⁻: 0 A student who thinks the acid and base cancel each other out in neutralization picks this. A proton moves from HA to OH⁻, forming H₂O, and each HA that reacts becomes an A⁻ ion, so A⁻ increases to 7.
DHA: 7; A⁻: 3; OH⁻: 0 A student who applies Le Châtelier's principle in the wrong direction, shifting HA ⇌ H₃O⁺ + A⁻ toward HA when base is added, picks this. OH⁻ removes H₃O⁺, so the shift is toward A⁻: HA + OH⁻ → A⁻ + H₂O.
Working Before: 5 HA, 5 A⁻; added 2 OH⁻. HA(aq) + OH⁻(aq) → A⁻(aq) + H₂O(l) goes essentially to completion: 2 HA react with the 2 OH⁻, giving 3 HA, 7 A⁻ and 0 OH⁻ (plus 2 H₂O).
Equal volumes of the two solutions in each pair are mixed at 25°C. Which mixture is a buffer solution?
Answer and reasoning
ACH₃COOH, 0.10 M, and KOH, 0.10 M A student who thinks a solution of a conjugate base alone is a buffer picks this. Equal amounts react completely, leaving CH₃COO⁻ but almost no CH₃COOH to react with added base.
BHCl, 0.20 M, and NaCl, 0.10 M A student who thinks any acid with its conjugate base makes a buffer picks this. HCl is a strong acid, and Cl⁻ does not react with added H₃O⁺, so the mixture cannot stabilize pH against added acid.
CCH₃COOH, 0.20 M, and NaOH, 0.10 MCorrect The NaOH converts half of the CH₃COOH to CH₃COO⁻ (CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O), leaving large, equal concentrations of the weak acid and its conjugate base, which is a buffer.
DHBr, 0.10 M, and KOH, 0.10 M A student who thinks buffers are neutral solutions picks this. Equal amounts of a strong acid and a strong base give a neutral KBr solution, which contains no weak acid or weak base to react with added acid or base.
Working No calculation of pH. 0.20 M CH₃COOH + 0.10 M NaOH (equal volumes): half the acid is converted to CH₃COO⁻, leaving equal large amounts of CH₃COOH and CH₃COO⁻: a buffer. Equimolar CH₃COOH + KOH: all converted to CH₃COO⁻, no CH₃COOH left. HCl + NaCl: strong acid with Cl⁻, which does not react with added acid. HBr + KOH equimolar: neutral KBr solution, no conjugate pair of a weak acid.
Compiled from the AP Chemistry Course and Exam Description (effective Fall 2024) and our question bank · Specialist review in progress. How these pages are made · Free, no account
