Introduction to Titration
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AP Chemistry › Introduction to Titration
A student titrates $\text{HCl}(aq)$ with $\text{NaOH}(aq)$ (strong acid–strong base). Which statement is correct about the solution composition before the equivalence point?
Only $\text{Na}^+$ and $\text{Cl}^-$ are present because neutralization is complete
Only water determines the pH because strong acids and bases cancel each other out at all points
A buffer of $\text{HCl}$ and $\text{Cl}^-$ controls the pH
Excess $\text{OH}^-$ is present because $\text{NaOH}$ is a strong base
Excess $\text{H}_3\text{O}^+$ is present because the acid has not yet been completely neutralized
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.
A weak base–strong acid titration is carried out by adding $\text{HCl}(aq)$ to $\text{NH}_3(aq)$. Which statement correctly compares the pH at the equivalence point to 7?
The pH is equal to 7 because equivalence always corresponds to neutral pH
The pH is greater than 7 because the solution contains excess $\text{NH}_3$ at equivalence
The pH is greater than 7 because $\text{Cl}^-$ ions hydrolyze to produce $\text{OH}^-$
The pH is less than 7 because the solution contains the conjugate acid $\text{NH}_4^+$ formed during the titration
The pH cannot be predicted qualitatively without calculating $K_a$
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.
A weak base–strong acid titration is performed by adding $\mathrm{HCl(aq)}$ to $\mathrm{NH_3(aq)}$. A simplified titration curve is divided into four labeled regions: initial solution (mostly $\mathrm{NH_3}$), buffering region (mixture of $\mathrm{NH_3}$ and $\mathrm{NH_4^+}$), equivalence region, and post-equivalence (excess $\mathrm{H_3O^+}$).
Which statement best describes what happens to the pH when a small amount of $\mathrm{HCl}$ is added while the titration is in the buffering region?
The pH decreases abruptly because the solution cannot resist pH change until equivalence.
The pH becomes neutral because the buffering region corresponds to pH 7.
The pH decreases, but only slightly, because both $\mathrm{NH_3}$ and $\mathrm{NH_4^+}$ are present.
The pH remains exactly constant because buffers prevent any pH change.
The pH increases because added acid converts $\mathrm{NH_4^+}$ into $\mathrm{NH_3}$.
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.
A student titrates a weak monoprotic acid, $\mathrm{HA(aq)}$, with a strong base, $\mathrm{NaOH(aq)}$. The student uses the following qualitative table to describe the curve.
| Point on curve | Qualitative description |
|---|---|
| I | Initial acidic solution; no base added |
| II | pH increases slowly as base is added |
| III | Rapid pH change over a small volume range |
| IV | High pH; additional base causes small changes |
Which point (I–IV) corresponds most closely to the equivalence region?
Points II and IV equally
Point II
Point IV
Point III
Point I
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.
A weak base–strong acid titration is performed by adding $\mathrm{HCl(aq)}$ to a solution of $\mathrm{B(aq)}$ (a weak base). A simplified titration curve is described with four labeled regions: initial solution, buffering region, equivalence region, and post-equivalence.
Which statement best describes the pH at the equivalence point compared with 7?
The pH is less than 7 because excess strong acid must be present at equivalence.
The pH is greater than 7 because the conjugate base formed at equivalence makes the solution basic.
The pH is greater than 7 because the weak base remains unreacted at equivalence.
The pH is approximately 7 because all titrations have neutral equivalence points.
The pH is less than 7 because the conjugate acid of the weak base is present 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.
A student titrates a weak monoprotic acid $\text{HA}(aq)$ with $\text{NaOH}(aq)$ (strong base). The student is at the equivalence point. Which statement best describes the major acid–base species in the solution at this point (ignoring water)?
Mostly $\text{H}_3\text{O}^+$ remains because the acid is still in excess
Equal amounts of $\text{H}_3\text{O}^+$ and $\text{OH}^-$ are present because equivalence always means neutral pH
Mostly excess $\text{OH}^-$ is present because the equivalence point occurs after the base is in excess
Mostly $\text{A}^-$ is present because all $\text{HA}$ has been converted to its conjugate base
Mostly $\text{HA}$ remains because weak acids are not fully neutralized by strong bases
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.
A student performs a strong acid–strong base titration by adding $\text{NaOH}(aq)$ to $\text{HCl}(aq)$ in a beaker. Which statement best describes what is happening in the equivalence region of the titration curve?
The pH changes linearly with added $\text{NaOH}$ because strong acids and bases dissociate completely
A small addition of titrant causes a large change in pH because the acid and base are nearly completely consumed in stoichiometric proportions
The pH remains nearly constant because the solution is buffered by $\text{Cl}^-$ ions
The pH decreases because adding base increases the concentration of $\text{H}_3\text{O}^+$ through hydrolysis
The solution contains significant amounts of both a weak acid and its conjugate base, so it resists pH change
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.
A strong acid–strong base titration is carried out by adding $\text{NaOH}(aq)$ to $\text{HNO}_3(aq)$. The student is slightly past the equivalence point. What happens to the pH when an additional small amount of $\text{NaOH}$ is added?
The pH increases sharply because the solution is in the equivalence region
The pH remains constant because nitrate ions form a buffer with $\text{HNO}_3$
The pH decreases sharply because the solution contains excess strong acid
The pH increases, but only gradually compared with the sharp change near the equivalence region
The pH decreases because additional $\text{NaOH}$ shifts water autoionization to produce more $\text{H}_3\text{O}^+$
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.
A weak acid–strong base titration is performed by adding $\mathrm{NaOH(aq)}$ to $\mathrm{HC_2H_3O_2(aq)}$ (acetic acid). A simplified titration curve is labeled with an initial solution region, a buffering region, an equivalence region, and a post-equivalence region.
Which qualitative statement best describes the solution at the equivalence point of this weak acid–strong base titration?
The solution contains equal amounts of strong acid and strong base, producing a neutral buffer.
The solution contains excess $\mathrm{H_3O^+}$ because the weak acid is not fully neutralized.
The solution contains a significant mixture of $\mathrm{HC_2H_3O_2}$ and $\mathrm{C_2H_3O_2^-}$, so it is a buffer.
The solution contains primarily $\mathrm{C_2H_3O_2^-}$ (the conjugate base) and spectator ions, with no excess strong base.
The solution contains excess $\mathrm{OH^-}$ because the equivalence point occurs only after base is in excess.
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.
A strong acid–strong base titration is carried out by adding $\mathrm{NaOH(aq)}$ to $\mathrm{HNO_3(aq)}$. The titration curve includes an initial low-pH region, a steep equivalence region, and a post-equivalence high-pH region.
Which region of the curve corresponds to the solution acting most like a buffer (resisting pH change upon small additions of titrant)?
Post-equivalence region, because excess $\mathrm{OH^-}$ neutralizes added base.
Initial solution region, because the strong acid and its conjugate base are both present.
Equivalence region, because pH changes very slowly there.
Buffering region before equivalence, because strong acid–strong base titrations form a buffer.
None of the regions; a strong acid–strong base titration does not have a true buffering region.
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.