Unit 15: Equilibria of Other Reaction Classes
How a solution can be built to resist pH change, how a titration curve reveals it happening, and what equilibrium says about a solid that won't dissolve.
10 topics · 37 questions
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15.1 Buffers: How They Resist pH Change
- How it worksHow does a buffer's pair of a weak acid and its conjugate base each neutralize an added acid or base without a big pH shift?
- How it worksWalk me through what happens inside a buffer solution the moment a small amount of strong acid is added to it.
- Common mistakeWhy isn't a buffer solution completely immune to any pH change at all?
- Going furtherWhat happens to a buffer's pH when both the weak acid and conjugate base concentrations are diluted by the same factor?
Your turn1 of 2
A buffer resists large pH swings when small amounts of acid or base are added. Why isn't a buffer solution completely immune to any pH change at all?15.1 The Henderson-Hasselbalch Equation
- How it worksHow does the Henderson-Hasselbalch equation turn the ratio of conjugate base to weak acid directly into a pH?
- Worked exampleFind the pH of a buffer made from 0.30 M acetic acid and 0.20 M sodium acetate, given a Ka of 1.8 times 10−5.
- Common mistakeWhy isn't the Henderson-Hasselbalch equation valid for calculating the pH of just any acidic solution?
Your turn1 of 2
Find the pH of a buffer made from 0.363 M acetic acid and 0.342 M sodium acetate, given a Ka of 1.8 times 10−5.15.1 Buffer Capacity
- How it worksHow does the absolute amount of weak acid and conjugate base present, not just their ratio, determine a buffer's capacity?
- Worked exampleFind which of two buffers has the greater capacity: one made from 0.10 mol each of acid and base, or one made from 1.0 mol each.
- Common mistakeWhy doesn't a buffer's pH by itself tell you anything about how much stress it can actually absorb?
- Going furtherWhat happens to a buffer's effectiveness once enough strong acid is added to use up all of its conjugate base?
Your turn1 of 2
One buffer is made from 0.10 mol each of a weak acid and its conjugate base, and another from 1.0 mol each of the exact same acid and base. Both have the identical pH. Which buffer has the greater capacity to resist added acid or base?15.2 Titration Curves: Strong Acid with Strong Base
- How it worksHow does the sharp vertical jump in a strong acid-strong base titration curve mark the exact point where moles of acid and base match?
- How it worksWalk me through how a strong acid-strong base titration curve is built point by point as base is added from a burette.
- Common mistakeWhy does the equivalence point of a strong acid-strong base titration land at exactly pH 7?
- Going furtherWhat happens to the steepness of the vertical jump once the titration is run with more concentrated acid and base?
Your turn1 of 2
Why does the equivalence point of a strong acid-strong base titration land at exactly pH 7?15.2 Titration Curves: Weak Acid with Strong Base
- How it worksHow does the flattened buffer region of the curve form from the weak acid and its conjugate base coexisting in solution?
- Worked exampleFind the pH at the half-equivalence point of a weak acid titration, given that the acid's Ka is 4.5 times 10−4.
- Common mistakeWhy isn't the equivalence point of a weak acid-strong base titration at pH 7?
- Ties togetherHow does the half-equivalence point on a titration curve connect back to the Henderson-Hasselbalch equation?
Your turn1 of 2
Find the pH at the half-equivalence point of a weak acid titration, given that the acid's Ka is 7 times 10−4.15.2 Choosing an Indicator for a Titration
- How it worksHow does an indicator's own color-change pH range need to overlap with a titration's steep vertical jump for an accurate result?
- Worked exampleFind whether phenolphthalein, with a range of about 8.2 to 10, is a good indicator choice for titrating a weak acid with a strong base.
- Common mistakeWhy isn't the equivalence point of a titration the same thing as the endpoint marked by an indicator's color change?
Your turn1 of 2
Phenolphthalein changes color over a pH range of about 8.2 to 10. Is it a good indicator choice for titrating a weak acid with a strong base?15.3 The Solubility Product Constant, Ksp
- How it worksHow does the solubility product constant, Ksp, describe an equilibrium between an undissolved solid and its dissolved ions?
- Worked exampleFind the molar solubility of AgCl given a Ksp of 1.8 times 10−10.
- Common mistakeWhy isn't Ksp itself the same number as a salt's molar solubility?
- Going furtherWhat happens to Ksp calculations when comparing two salts with different ion ratios, like AgCl versus Ag2CrO4?
Your turn1 of 2
Find the molar solubility of a salt AX, which dissociates into one A+ ion and one X- ion, given a Ksp of 7 times 10−10.15.3 Predicting Precipitation from Ksp and Q
- How it worksHow does comparing the ion product Q to Ksp reveal whether a solution is saturated, unsaturated, or ready to precipitate?
- Worked exampleFind whether a precipitate forms when solutions are mixed to give 0.010 M Pb2+ and 0.010 M Cl-, given a Ksp for PbCl2 of 1.7 times 10−5.
- Common mistakeWhy isn't Q calculated any differently from Ksp, aside from using concentrations that may not yet be at equilibrium?
Your turn1 of 2
Find the ion product Q when solutions are mixed to give 0.0363 M Pb2+ and 0.0342 M Cl- (for PbCl2, Q = [Pb2+][Cl-]2).15.3 The Common Ion Effect on Solubility
- How it worksHow does Le Chatelier's principle, applied to a solubility equilibrium, explain why a shared ion suppresses a salt's solubility?
- Worked exampleFind the molar solubility of AgCl in a solution that already contains 0.10 M NaCl, given a Ksp of 1.8 times 10−10.
- Common mistakeWhy doesn't the common ion effect change the value of Ksp itself, even though it changes how much solid dissolves?
- Ties togetherHow does the common ion effect on solubility connect back to the same Le Chatelier's principle that governs buffers and gas equilibria?
Your turn1 of 2
Find the molar solubility of AgCl in a solution that already contains 0.351 M NaCl, given a Ksp of 1.8 times 10−10 (assume the dissolved AgCl contributes a negligible amount of Cl- compared to 0.351).15.4 Complex Ion Equilibria and Solubility
- How it worksHow does a metal ion binding to several ligands to form a complex ion pull dissolved ions out of the original solubility equilibrium?
- Worked exampleFind the effect on AgCl's solubility of adding excess ammonia, given that Ag+ forms the complex ion Ag(NH3)2+.
- Common mistakeWhy doesn't forming a complex ion violate the solubility product constant for the original solid?
- Going furtherWhat happens to a metal hydroxide precipitate's solubility once the solution becomes basic enough to form a soluble hydroxo complex?
Your turn1 of 2
AgCl, normally considered essentially insoluble, dissolves noticeably once excess ammonia is added, forming the complex ion Ag(NH3)2+. What effect does this have on AgCl's solubility?