AP Unit 8 Problems

At a particular temperature, Kw = 4.00 × 10⁻¹⁴. Calculate [H₃O⁺], [OH⁻], pH, and pOH in pure water at that temperature. A student calls this water acidic because its pH is below 7. Evaluate the claim and state the correct pH + pOH relation.

At 25 °C, calculate pH and pOH for (a) 0.00300 M HNO₃ and (b) a fully dissolved 0.00300 M Ba(OH)₂ solution. Show the major solute ions and their concentrations for each. Explain why “strong” does not mean “concentrated.”

A 0.100 M solution of a monoprotic weak acid HA has Ka = 1.00 × 10⁻⁵ at 25 °C. Determine pH, percent ionization, the major acid-related species, [OH⁻], and Kb of A⁻. Show and check the small-x approximation. Predict qualitatively what dilution does to percent ionization and hydronium concentration.

A 0.0200 M weak base B has Kb = 2.00 × 10⁻⁵ at 25 °C. Estimate pOH, pH, percent protonation, and Ka of its conjugate acid HB⁺. Write the equilibrium equation and show that the approximation is reasonable. Describe the particle population.

At 25 °C, 25.0 mL of 0.120 M HCl is mixed with 35.0 mL of 0.100 M NaOH. Determine the excess reagent, [OH⁻], pH, and the major spectator-ion concentrations. Assume additive volumes. Explain why subtracting molarities would be wrong.

Three separate 25.0 mL samples of 0.100 M HA, Ka = 1.00 × 10⁻⁵, receive 0.100 M NaOH at 25 °C: sample I gets 10.0 mL, II gets 25.0 mL, III gets 30.0 mL. Identify major acid-related species and select the pH method for each. Calculate pH for II and III; do not compute a buffer pH change for I.

A weak base B has Kb = 1.00 × 10⁻⁵. Explain the dominant species and pH method when a 20.0 mL 0.100 M B sample receives 0.100 M HCl volumes of 5.0, 20.0, or 25.0 mL at 25 °C. Then consider a separate mixture HA + B ⇌ A− + HB+. Ka(HA) = 1.00 × 10⁻⁴ and Ka(HB+) = 1.00 × 10⁻⁸. Is quantitative neutralization always justified for arbitrary weak–weak mixtures? Use these constants to support the reasoning for this example.

At 25 °C, a 25.00 mL monoprotic weak-acid sample is titrated with 0.1000 M NaOH. Supplied data are: added volume 0.00, 10.00, 19.50, 20.00, 20.50, 25.00 mL; pH 2.92, 4.80, 6.39, 8.72, 11.04, 12.00. The equivalence volume from the steep region is 20.00 mL.

1. Determine the acid’s initial concentration and pKa. Identify major acid-related species at 10.00 and 20.00 mL, write the conjugate base’s reaction with water, and explain why the equivalence pH exceeds 7.
2. Predict the bias in calculated acid concentration if a delivery-tip air bubble fills during titration, so recorded volume exceeds delivered volume.
3. Two separate titrations at 25 °C are stopped exactly at their stoichiometric equivalence points: NH₃(aq) with HCl(aq), and HCl(aq) with NaOH(aq). For each, identify the major solute ions, write the relevant proton-transfer reaction with water or explain why there is no appreciable one, and predict whether pH is below, equal to, or above 7. Explain why the same definition of equivalence can give different pH values.

Assume dilute aqueous solutions, complete strong-acid/strong-base dissociation, no excess titrant, and no competing reactions. Chloride and sodium ions have negligible acid–base effects in this model.

A hypothetical diprotic acid H₂A is titrated with strong base. Its resolved curve has equivalence points at 15.0 and 30.0 mL; pH values at 7.50 and 22.5 mL are 3.00 and 7.00. Infer the number of distinguishable acidic protons, pKa1 and pKa2, and the major acid-related species at 0, 15.0, 22.5, and 30.0 mL. Explain a limitation of inferring proton count from an unresolved titration curve. No species-concentration calculations are required.

1. Explain why CH₃COOH is more acidic than CH₃CH₂OH and why ClCH₂COOH is more acidic than CH₃COOH. Identify the transferred proton and compare the resulting conjugate-base structures.
2. Explain how NH₃ acts as a weak base. Identify its proton-accepting site and draw or describe the product.
3. HNO₃ is a strong acid in dilute water. One valid Lewis contributor is H–O–N⁺(=O)–O⁻. Identify its transferred proton, draw one valid contributor of its conjugate base, and describe how the ion’s structure stabilizes negative charge and makes it a very weak base. The supplied strong-acid classification is not something you must derive from resonance alone.
4. NaOH and Ca(OH)₂ are ionic hydroxides. Identify the proton-accepting ion released into water, its accepting atom and conjugate acid. Explain why the dissolved portion gives strong-base behavior, and why this does not mean every added piece of Ca(OH)₂ must dissolve. Contrast that ion with ammonia’s molecular accepting site.

Use comparable dilute aqueous conditions. Relate strength to structure and proton transfer; hydrogen count alone is not evidence.

A weak acid indicator has pKa = 6.50. Its protonated HIn form is yellow and deprotonated In− form is blue. Predict the predominant form at pH 5.50 and 7.50 and the approximate ratio [In−]/[HIn] at each. Would this indicator be a good choice for a titration whose sharply rising equivalence region runs from pH 8.0 to 10.0? Explain using the curve region, not merely whether a color change exists.

Three equal-volume solutions contain (I) 0.100 M HA only, (II) 0.100 M HA and 0.100 M A−, or (III) 0.100 M HCl and 0.100 M NaCl. HA is a weak acid. Which is an effective buffer near pKa(HA)? Write the reactions that consume a small added quantity of strong acid and strong base in that solution. Explain why its pH changes slightly rather than remaining perfectly fixed, and why the others are less suitable.

A buffer has pKa(HA) = 4.80, [HA] = 0.200 M, and [A−] = 0.0500 M. Calculate its current pH and identify the more abundant form. To prepare a buffer at pH 5.40 using the same pair, what concentration ratio [A−]/[HA] is needed? State the condition under which using a mole ratio would give the same result. Do not calculate a pH change after an acid/base addition.

Equal 100 mL buffers A and B have the same HA/A− pair. A contains 0.100 M of each component; B contains 0.0100 M of each. Compare their initial pH and ability to resist the same added amount of acid. Buffer C instead contains 0.150 M HA and 0.0500 M A−. Compare C’s ability to consume added acid versus added base. If B was accidentally diluted further before an investigation, explain the expected bias in its apparent capacity while the ratio remains unchanged.

1. Compare adding acid to saturated solutions containing M(OH)₂(s), MCO₃(s), and MX(s), where X⁻ is the extremely weak conjugate base of a strong acid. Write a linked proton-transfer equation for the hydroxide or carbonate, explain the dissolution response, and state whether Ksp changes at fixed temperature.
2. A separate hypothetical sparingly soluble salt JHY(s) dissolves as JHY(s) ⇌ JH⁺(aq) + Y⁻(aq). JH⁺ is a weak acid; its neutral conjugate base J remains dissolved. Y⁻ is acid–base inactive. Write a proton-transfer reaction for JH⁺ with water or hydroxide. Predict the effect of raising pH on free JH⁺, the immediate dissolution quotient, and saturated total solubility. Distinguish free JH⁺ from total J-containing dissolved material.
3. For this JHY model, predict how lowering pH instead affects the JH⁺/J distribution and saturated total solubility. State one condition under which the simplified prediction would need to be reconsidered.

Assume constant temperature, negligible dilution, a macroscopic amount of the original solid present, and no common-ion addition, complexes, competing precipitates or other acid–base reactions over the specified pH range. These are supplied models, not chemical-handling instructions. No solubility-versus-pH calculation is required.