AP Unit 4 Problems

Two clear aqueous samples are mixed. The temperature falls 3 °C, but no solid or gas appears. A student says a chemical reaction is proved. Evaluate the claim and propose one additional type of evidence that would distinguish chemical change from physical mixing.

A sealed container holds liquid water and water vapor. Heating increases the vapor amount while the total mass stays fixed. State a particle prediction and classify the change. Contrast it with decomposition of water into H₂ and O₂.

Given that CaCO₃ is sparingly soluble, write molecular, complete ionic, and net ionic equations for CaCl₂(aq) mixed with Na₂CO₃(aq). Identify spectators and verify charge.

Write (a) the dissolution of K₂SO₄(s) in water, and (b) neutralization of weak acid HF(aq) by NaOH(aq), as a net ionic equation. Explain why HF is treated differently from dissolved potassium sulfate.

For the model reaction 2A₂ + B₂ → 2A₂B, a closed whole-sample box starts with seven A₂ and three B₂ molecules. The reaction proceeds completely until one reactant is exhausted; A and B denote different neutral atoms.

1. Construct a labeled before-and-after particle diagram. Use the same atom key and icon scale on both sides. Show bonds within each molecule, every product molecule and every unreacted molecule. Counts are the complete model sample, not equal-volume windows. Particle positions need not represent measured coordinates.
2. Identify the limiting reactant and verify both A-atom and B-atom counts before and after. Explain why coefficients give reacting proportions rather than forcing all starting particles to react.

Draw on paper or in the workspace. An accessible alternative is a coordinate construction identifying each bonded molecular group and its atom labels in the two panels. A final inventory alone does not replace constructing the representation.

A whole-sample solution model initially contains two Mg²⁺, four NO₃⁻, six Na⁺ and six OH⁻ ions. Mg(OH)₂ is sparingly soluble. For this particle-count model, assume precipitation consumes all available Mg²⁺; ignore the small equilibrium dissolved Mg²⁺ amount and any competing reaction.

1. Write the balanced net ionic equation. Construct labeled before-and-after particle diagrams with one unchanged ion key. In the before panel show separate aqueous ions. In the after panel show a single labeled ionic-solid array plus every remaining aqueous ion, including spectators and any excess reagent. Keep OH⁻ and NO₃⁻ intact. Water may be omitted if the legend says so.
2. Verify Mg, O, H, N and Na atom counts and the total charge in each panel. Identify spectators and explain why cancellation from a net ionic equation does not erase them from the solution.

Draw on paper or in the workspace. An accessible alternative specifies coordinates and labels for each aqueous ion plus a connected solid-region layout containing its constituent ions; include phase labels and the same before/after inventory. A list of counts without a represented solid region and dispersed aqueous region is insufficient.

A student says “evaporation and NaCl dissolution are physical, so no bonds or attractions change.” Evaluate each process at the particle level and explain the limitation of the student’s rule.

For N₂ + 3H₂ → 2NH₃, 0.400 mol N₂ and 0.900 mol H₂ react as far as possible. Find NH₃ produced and excess reactant remaining. Then predict the product if H₂ alone is increased to 1.80 mol.

30.0 mL of 0.120 M BaCl₂ is mixed with 20.0 mL of 0.150 M Na₂SO₄. BaSO₄ is sparingly soluble; assume complete precipitation and additive volumes. Find solid moles and the concentration of excess Ba²⁺.

A 0.0150 mol sample of CaCO₃ reacts completely with excess acid: CaCO₃(s) + 2H⁺(aq) → Ca²⁺(aq) + CO₂(g) + H₂O(l). CO₂ is collected over water at 298 K and total pressure 1.000 atm; water vapor pressure is 0.0313 atm. With R =0.08206 L atm mol⁻¹ K⁻¹, find gas volume and explain the bias if water vapor pressure is ignored.

A 25.00 mL sample of monoprotic acid HA is titrated by 0.1250 M NaOH. Buret readings are 1.35 mL initially and 19.75 mL at the indicator endpoint, assumed to match equivalence. Calculate [HA]. On a graph of indicator signal versus added base volume, state the x-axis quantity and where the equivalence estimate belongs.

A titration of an unknown acid uses known base concentration. Three endpoints agree closely, but all are reached after adding a little excess base. Explain the effect on calculated acid concentration and why repeated trials do not remove this error.

Classify and justify: (i) NH₄⁺ + OH⁻ → NH₃ + H₂O; (ii) 2Mg + O₂ → 2MgO; (iii) Ba²⁺ + SO₄²⁻ → BaSO₄(s), given low BaSO₄ solubility. Include oxidation-number changes and total electron transfer where applicable.

For C₂H₆ + O₂ → CO₂ + H₂O, balance complete combustion and find the oxidation-number change per carbon and electrons transferred per two ethane molecules.

In CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺, label acids, bases and conjugate pairs. Given that H₃O⁺ is a stronger acid than CH₃COOH, predict which side is favored and compare the bases.

Two diagrams show water reacting: model I, H₂O becomes H₃O⁺ while HX becomes X⁻; model II, H₂O becomes OH⁻ while B becomes BH⁺. Explain water’s role in each, write both equations and account for charge.

Use half-reactions to balance Cr₂O₇²⁻ + I⁻ → Cr³⁺ + I₂ in acidic solution. Verify total charge.

Balance Cl₂ → Cl⁻ + ClO⁻ in basic aqueous solution using half-reactions. Explain why one element can be both oxidized and reduced.