AP Chemistry Quiz: Introduction To Titration

What this quiz covers

This quiz focuses on Introduction To Titration, giving you a quick way to practice the rules, question types, and explanations that matter most for AP Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

Question 1

1. $\mathrm{B}$
2. $\mathrm{BH^+}$ (correct answer)
3. $\mathrm{OH^-}$
4. $\mathrm{H_3O^+}$
5. Equal amounts of $\mathrm{HCl}$ and $\mathrm{B}$ remain unreacted

Explanation: This question tests understanding of introduction to titration, specifically identifying the predominant species at the equivalence point of a weak base-strong acid titration. When titrating a weak base B with HCl, the reaction is B + H₃O⁺ → BH⁺ + H₂O. At the equivalence point, all of the original weak base B has been converted to its conjugate acid BH⁺, so BH⁺ is the predominant species (ignoring spectator ions like Cl⁻). The solution will be acidic because BH⁺ is a weak acid that can donate protons to water: BH⁺ + H₂O ⇌ B + H₃O⁺. A common misconception is that B remains at equivalence (choice A), but by definition, the equivalence point is where all B has been neutralized. To determine species at equivalence, consider what products form when the limiting reactant is completely consumed—for weak base titrations, this produces the conjugate acid.

Question 2

1. The pH decreases slightly because added acid converts some $\mathrm{B}$ to $\mathrm{BH^+}$. (correct answer)
2. The pH increases slightly because added $\mathrm{Cl^-}$ makes the solution more basic.
3. The pH stays exactly constant because buffers prevent any pH change.
4. The pH increases sharply because the solution is near the initial region.
5. The pH decreases sharply because excess $\mathrm{H_3O^+}$ must already be present before equivalence.

Explanation: This question tests understanding of introduction to titration, specifically predicting pH changes in the buffering region of a weak base-strong acid titration. Before equivalence, the solution contains both B (weak base) and BH⁺ (conjugate acid), forming a buffer. When a small amount of HCl is added, the reaction B + H₃O⁺ → BH⁺ + H₂O converts some B to BH⁺, decreasing the [B]/[BH⁺] ratio. According to the Henderson-Hasselbalch equation, this decreases the pH slightly, but the change is small because the buffer resists large pH changes. A common misconception is that buffers prevent any pH change (choice C), but buffers only minimize changes, not eliminate them entirely. When analyzing buffer behavior during titration, remember that added acid or base shifts the ratio of buffer components, causing small but measurable pH changes.

Question 3

1. HCl and Cl−
2. B and BH+ (correct answer)
3. H3O+ and OH−
4. Cl− and H2O
5. BH+ and OH−

Explanation: This question tests understanding of introduction to titration, specifically identifying the buffer components in a weak base-strong acid titration. When HCl is added to the weak base B, the reaction B + H⁺ → BH⁺ occurs, converting some base to its conjugate acid. Before the equivalence point, both the weak base B and its conjugate acid BH⁺ are present in solution, and this B/BH⁺ conjugate pair is responsible for the buffering behavior. This buffer system can neutralize small amounts of added H⁺ (through B + H⁺ → BH⁺) or OH⁻ (through BH⁺ + OH⁻ → B + H₂O), thereby moderating pH changes. A common misconception (option A) is that HCl and Cl⁻ form a buffer, but this is incorrect because HCl is a strong acid that completely dissociates, leaving no HCl molecules to participate in buffering. The key strategy is to identify the weak acid-base conjugate pair that exists when partial neutralization has occurred.

Question 4

1. Excess $\text{OH}^-$ is present because $\text{NaOH}$ is a strong base
2. Excess $\text{H}_3\text{O}^+$ is present because the acid has not yet been completely neutralized (correct answer)
3. Only $\text{Na}^+$ and $\text{Cl}^-$ are present because neutralization is complete
4. A buffer of $\text{HCl}$ and $\text{Cl}^-$ controls the pH
5. Only water determines the pH because strong acids and bases cancel each other out at all points

Explanation: This problem involves introduction to titration for a strong acid-strong base system before equivalence. Before the equivalence point, HCl is in excess because insufficient NaOH has been added to neutralize all the acid. The solution contains H₃O⁺ from the unreacted HCl, making it acidic. Option D (buffer of HCl/Cl⁻) is incorrect because HCl is a strong acid that completely dissociates, and Cl⁻ has no basic properties to form a buffer. Remember to track which reagent is in excess at different stages of the titration to determine what controls the pH.

Question 5

1. The pH is less than 7 because the solution contains the conjugate acid $\text{NH}_4^+$ formed during the titration (correct answer)
2. The pH is equal to 7 because equivalence always corresponds to neutral pH
3. The pH cannot be predicted qualitatively without calculating $K_a$
4. The pH is greater than 7 because $\text{Cl}^-$ ions hydrolyze to produce $\text{OH}^-$
5. The pH is greater than 7 because the solution contains excess $\text{NH}_3$ at equivalence

Explanation: This problem focuses on introduction to titration at the equivalence point of a weak base-strong acid system. At equivalence, all NH₃ has been converted to NH₄⁺ through the reaction NH₃ + H⁺ → NH₄⁺. Since NH₄⁺ is the conjugate acid of the weak base NH₃, it acts as a weak acid, producing H₃O⁺ and making the solution acidic (pH < 7). Option B (pH = 7 at equivalence) is incorrect because this only occurs for strong acid-strong base titrations; weak base-strong acid titrations have acidic equivalence points. Remember that the pH at equivalence depends on the acid-base properties of the products formed.

Question 6

1. The pH decreases, but only slightly, because both $\mathrm{NH_3}$ and $\mathrm{NH_4^+}$ are present. (correct answer)
2. The pH decreases abruptly because the solution cannot resist pH change until equivalence.
3. The pH increases because added acid converts $\mathrm{NH_4^+}$ into $\mathrm{NH_3}$.
4. The pH remains exactly constant because buffers prevent any pH change.
5. The pH becomes neutral because the buffering region corresponds to pH 7.

Explanation: This question assesses the introduction to titration. Titration curves for weak base-strong acid display an initial high pH, a buffering region with slow pH decrease as conjugate acid forms, a sharp drop at equivalence, and low pH with excess acid. In the buffering region, the mixture of NH3 and NH4+ resists pH changes, so adding small acid amounts causes only slight pH decreases by shifting the equilibrium. This qualitative behavior arises from the buffer's capacity to absorb added H3O+ without drastic composition changes. A common misconception is that buffers keep pH exactly constant (choice D), but they only minimize changes, not prevent them. When evaluating buffer response, consider how added species react with the conjugate pair.

Question 7

1. Point I
2. Point II
3. Point III (correct answer)
4. Point IV
5. Points II and IV equally

Explanation: This question assesses the introduction to titration. Titration curves map pH versus titrant volume, highlighting regions where composition changes dictate behavior: initial acid, slow-rising buffer, rapid equivalence shift, and excess base plateau. The equivalence region features a rapid pH change over small volume additions because the buffer is depleted, and the solution transitions sharply from acidic to basic dominance. Qualitatively, this steep segment contrasts with the gradual changes elsewhere, reflecting high sensitivity when neither species is in excess. A common misconception is that the slow pH increase (point II) is the equivalence region (implying choice B), but that's actually the buffer. To identify regions, look for where pH sensitivity is highest to small titrant additions.

Question 8

1. The pH is greater than 7 because the conjugate base formed at equivalence makes the solution basic.
2. The pH is approximately 7 because all titrations have neutral equivalence points.
3. The pH is less than 7 because the conjugate acid of the weak base is present at equivalence. (correct answer)
4. The pH is less than 7 because excess strong acid must be present at equivalence.
5. The pH is greater than 7 because the weak base remains unreacted at equivalence.

Explanation: This question assesses the introduction to titration. Titration curves for weak base-strong acid systems start with a high pH from the weak base, show a buffering region where pH decreases gradually due to the mixture of base and conjugate acid, then a sharp drop at equivalence. The composition at equivalence is the conjugate acid of the weak base, which hydrolyzes to produce H3O+, making the solution acidic with pH less than 7. Post-equivalence, excess acid further lowers the pH slowly, as the curve reflects the dominance of strong acid. A common misconception is that all titrations have pH 7 at equivalence (choice B), but this only holds for strong-strong systems. To determine pH at equivalence, identify if the conjugate species is weak and thus affects neutrality.

Question 9

1. Mostly $\text{HA}$ remains because weak acids are not fully neutralized by strong bases
2. Mostly $\text{H}_3\text{O}^+$ remains because the acid is still in excess
3. Mostly $\text{A}^-$ is present because all $\text{HA}$ has been converted to its conjugate base (correct answer)
4. Mostly excess $\text{OH}^-$ is present because the equivalence point occurs after the base is in excess
5. Equal amounts of $\text{H}_3\text{O}^+$ and $\text{OH}^-$ are present because equivalence always means neutral pH

Explanation: This question tests introduction to titration concepts at the equivalence point of a weak acid-strong base titration. At equivalence, all the weak acid HA has been converted to its conjugate base A⁻ through the reaction HA + OH⁻ → A⁻ + H₂O. The solution contains primarily A⁻ (and Na⁺ as a spectator ion), making it basic since A⁻ is the conjugate base of a weak acid. Option E (equal H₃O⁺ and OH⁻ means neutral) is incorrect because equivalence doesn't mean pH = 7 for weak acid-strong base titrations; the pH depends on the basicity of A⁻. Always consider what species remain after the neutralization reaction to determine the pH at equivalence.

Question 10

1. The pH remains nearly constant because the solution is buffered by $\text{Cl}^-$ ions
2. A small addition of titrant causes a large change in pH because the acid and base are nearly completely consumed in stoichiometric proportions (correct answer)
3. The pH decreases because adding base increases the concentration of $\text{H}_3\text{O}^+$ through hydrolysis
4. The solution contains significant amounts of both a weak acid and its conjugate base, so it resists pH change
5. The pH changes linearly with added $\text{NaOH}$ because strong acids and bases dissociate completely

Explanation: This question requires understanding of introduction to titration for a strong acid-strong base system. In the equivalence region, the moles of H⁺ from HCl are nearly equal to the moles of OH⁻ from NaOH, meaning both reactants are almost completely consumed. A tiny additional amount of NaOH causes a dramatic pH change because there's no buffer to resist the change—the solution transitions from slightly acidic to slightly basic. Option D (buffered by Cl⁻) is incorrect because Cl⁻ is the conjugate base of a strong acid and has no buffering capacity. Remember that sharp pH changes occur when there's no buffer system present to resist the addition of acid or base.

Question 11

1. The pH increases, but only gradually compared with the sharp change near the equivalence region (correct answer)
2. The pH decreases because additional $\text{NaOH}$ shifts water autoionization to produce more $\text{H}_3\text{O}^+$
3. The pH remains constant because nitrate ions form a buffer with $\text{HNO}_3$
4. The pH increases sharply because the solution is in the equivalence region
5. The pH decreases sharply because the solution contains excess strong acid

Explanation: This question involves introduction to titration for a strong acid-strong base system past the equivalence point. After equivalence, all HNO₃ has been neutralized, and the solution contains excess OH⁻ from the added NaOH. Adding more NaOH increases [OH⁻] and thus increases pH, but the change is gradual because we're simply diluting a basic solution with more base. Option D (sharp increase due to equivalence region) is incorrect because we're already past equivalence where sharp changes occur. The key insight is that pH changes are gradual when adding more of the excess reagent, unlike the sharp change at equivalence.

Question 12

1. The pH at equivalence must be 7 because equal moles of acid and base have reacted.
2. The pH at equivalence is typically greater than 7 because the conjugate base $\mathrm{A^-}$ affects the solution. (correct answer)
3. The pH at equivalence is typically less than 7 because the strong base always dominates at equivalence.
4. The pH at equivalence is undefined because all acid and base are gone.
5. The pH at equivalence is always the same as the initial pH of the weak acid.

Explanation: This question tests understanding of introduction to titration, specifically the pH at equivalence for different types of titrations. In a weak acid-strong base titration, the equivalence point occurs when all HA has been converted to A⁻. Since A⁻ is the conjugate base of a weak acid, it has basic properties and reacts with water: A⁻ + H₂O ⇌ HA + OH⁻. This hydrolysis reaction produces OH⁻, making the solution basic with pH > 7 at equivalence. This contrasts with strong acid-strong base titrations where pH = 7 at equivalence because neither product ion affects pH. A common misconception is that equivalence always means pH = 7 (choice A), but this is only true when neither product has acid-base properties. To predict pH at equivalence, identify the species present and determine if they can act as acids or bases in water.

Question 13

1. The solution contains excess $\mathrm{H_3O^+}$ because the weak acid is not fully neutralized.
2. The solution contains a significant mixture of $\mathrm{HC_2H_3O_2}$ and $\mathrm{C_2H_3O_2^-}$, so it is a buffer.
3. The solution contains primarily $\mathrm{C_2H_3O_2^-}$ (the conjugate base) and spectator ions, with no excess strong base. (correct answer)
4. The solution contains excess $\mathrm{OH^-}$ because the equivalence point occurs only after base is in excess.
5. The solution contains equal amounts of strong acid and strong base, producing a neutral buffer.

Explanation: This question assesses the introduction to titration. Titration curves reflect composition shifts as base is added: from weak acid to buffer mixture, then to conjugate base dominance at equivalence, and excess base afterward. At the equivalence point, all weak acid is converted to its conjugate base (C2H3O2-), resulting in a basic solution due to hydrolysis, with no excess strong base yet. This qualitative state produces pH greater than 7, distinguishing it from strong-strong neutral equivalence. A common misconception is that the equivalence point is still a buffer with acid and conjugate (choice B), but all acid is neutralized. To characterize equivalence, verify if only the conjugate species remains without excess titrant.

Question 14

1. Initial solution region, because the strong acid and its conjugate base are both present.
2. Buffering region before equivalence, because strong acid–strong base titrations form a buffer.
3. Equivalence region, because pH changes very slowly there.
4. Post-equivalence region, because excess $\mathrm{OH^-}$ neutralizes added base.
5. None of the regions; a strong acid–strong base titration does not have a true buffering region. (correct answer)

Explanation: This question assesses the introduction to titration. Titration curves for strong acid-strong base reveal composition shifts from excess acid (low pH) to neutral salt at equivalence (sharp pH jump) to excess base (high pH), without a flat buffering segment. The lack of a buffering region occurs because strong acids and bases do not coexist with their conjugates in a way that resists pH changes effectively. Qualitatively, pH changes gradually pre- and post-equivalence but abruptly at equivalence due to the sudden shift from acid to base dominance. A common misconception is that strong-strong titrations have a buffering region before equivalence (choice B), but no weak species are involved to create a buffer. To spot buffering, check for the presence of a weak acid/base and its conjugate in comparable amounts.

Question 15

1. The pH decreases slightly because added base neutralizes the conjugate base.
2. The pH increases because additional $\mathrm{OH^-}$ remains in excess. (correct answer)
3. The pH remains nearly constant because the solution is a buffer.
4. The pH changes abruptly by several units because the system is still at the equivalence region.
5. The pH becomes exactly 7 because post-equivalence solutions are neutral.

Explanation: This question assesses the introduction to titration. Titration curves illustrate how adding titrant alters solution composition, from weak acid dominance to a buffer mixture, then to conjugate base at equivalence, and finally excess base post-equivalence. In the post-equivalence region, excess $\mathrm{OH^{-}}$ from the strong base controls the pH, causing it to increase with additional titrant as the logarithmic relationship amplifies small concentration changes. The qualitative behavior shows a flatter curve here compared to the steep equivalence region, reflecting reduced sensitivity to added volume. A common misconception is that pH remains constant post-equivalence like a buffer (choice C), but without both conjugate forms, it's not buffered. Always identify if excess titrant is present to predict pH response to further additions.

Question 16

1. Excess $\mathrm{OH^-}$ from the titrant is present because the base has been added for a long time.
2. Only water and spectator ions are present because all acid has been neutralized exactly.
3. Only $\mathrm{HA}$ remains because the titrant has not yet reacted.
4. A mixture of $\mathrm{HA}$ and $\mathrm{A^-}$ is present, with no excess strong base. (correct answer)
5. Only $\mathrm{A^-}$ is present because equivalence has effectively been reached.

Explanation: This question assesses the introduction to titration. Titration curves track pH changes as titrant addition modifies the solution from weak acid to a mixture with conjugate base, showing a slow pH rise in the buffering region before a steep increase near equivalence. Just before the equivalence region, the composition is still a mixture of HA and A-, with HA not yet fully neutralized, leading to buffered, gradual pH behavior. This pre-equivalence point lacks excess base, distinguishing it from post-equivalence where OH- dominates and pH is higher. A common misconception is that only A- is present just before equivalence (choice D), but some HA remains until exact neutralization. Focus on the ratio of acid to conjugate base to pinpoint the titration stage.

Question 17

1. The solution contains excess $\mathrm{H_3O^+}$ because the acid is still in excess.
2. The solution contains excess $\mathrm{OH^-}$ because the base has just begun to dominate.
3. The solution contains primarily spectator ions and water, with neither $\mathrm{H_3O^+}$ nor $\mathrm{OH^-}$ in excess. (correct answer)
4. The solution is a buffer containing significant amounts of both $\mathrm{HCl}$ and $\mathrm{Cl^-}$.
5. The solution contains primarily the conjugate acid of the base, making the solution acidic.

Explanation: This question assesses the introduction to titration. Titration curves for strong acid-strong base systems show an initial low pH due to excess acid, a gradual pH increase as base neutralizes acid, a sharp rise at equivalence, and a high pH plateau with excess base. The composition changes from excess H3O+ to a neutral salt solution at equivalence, where neither H3O+ nor OH- dominates, leading to pH around 7. Post-equivalence, excess OH- causes the pH to rise slowly, reflecting the logarithmic dependence on excess titrant concentration. A common misconception is that the solution at equivalence is a buffer with HCl and Cl- (choice D), but strong acids do not form buffers with their conjugates. When analyzing equivalence points, first confirm if the resulting salt is neutral or hydrolyzes to affect pH.

Question 18

1. The pH decreases because adding base increases the concentration of $\mathrm{H_3O^+}$.
2. The pH increases because added $\mathrm{OH^-}$ neutralizes some of the excess acid. (correct answer)
3. The pH remains constant because a buffer is present in all strong acid–strong base titrations.
4. The pH increases abruptly by several units because any addition of base causes a steep rise.
5. The pH becomes neutral because pre-equivalence solutions are always at pH 7.

Explanation: This question assesses the introduction to titration. Titration curves for strong acid-strong base show composition evolving from excess acid (low, slowly rising pH) to neutral at equivalence (sharp rise) to excess base (high, slow rise). Before equivalence, adding base neutralizes some excess acid, increasing pH gradually as H3O+ concentration decreases logarithmically. This qualitative pre-equivalence behavior lacks buffering, so changes are steady but not resistant like in weak systems. A common misconception is that pH remains constant due to a buffer (choice C), but strong-strong titrations have no buffer regions. Before calculating pH, determine if the system is pre- or post-equivalence to predict change direction.

Question 19

1. A mixture of $\mathrm{HA}$ and $\mathrm{A^-}$ is present in comparable amounts. (correct answer)
2. Only $\mathrm{HA}$ is present; no conjugate base has formed yet.
3. Only $\mathrm{A^-}$ is present because all $\mathrm{HA}$ has been neutralized.
4. Excess $\mathrm{OH^-}$ from the titrant predominates.
5. Excess $\mathrm{H_3O^+}$ from the titrant predominates.

Explanation: This question assesses the introduction to titration. Titration curves plot pH against the volume of titrant added, reflecting how the solution's composition evolves from predominantly weak acid to a mixture of acid and conjugate base, then to mostly conjugate base and excess base. In the buffering region, the presence of comparable amounts of HA and A- allows the solution to resist pH changes, resulting in a relatively flat curve segment where pH increases gradually. As more base is added, the curve steepens near equivalence because the buffer capacity is exhausted, and pH becomes more sensitive to additional titrant. A common misconception is that only A- is present in the buffering region (choice C), but buffering requires both species to effectively neutralize added acid or base. To understand titration behavior, identify what remains in excess before worrying about pH.

Question 20

1. $\text{NH}_3$ and $\text{NH}_4^+$ (correct answer)
2. $\text{HCl}$ and $\text{Cl}^-$
3. $\text{H}_3\text{O}^+$ and $\text{OH}^-$
4. $\text{NH}_4^+$ and $\text{Cl}^-$
5. $\text{NH}_3$ and $\text{Cl}^-$

Explanation: This problem focuses on introduction to titration, specifically identifying the buffer components in a weak base-strong acid titration. When HCl is added to NH₃, the reaction produces NH₄⁺: NH₃ + H⁺ → NH₄⁺. Early in the titration, both NH₃ (weak base) and NH₄⁺ (its conjugate acid) are present in significant amounts, creating a buffer that resists pH changes. Option B (HCl and Cl⁻) is incorrect because HCl is a strong acid that completely dissociates, and Cl⁻ has no basic properties. The strategy is to identify the weak base and its conjugate acid formed during the titration reaction.