AP Chemistry Unit 4

AP Chemistry – Answer Key (Questions 1–35)

Chemical & Physical Changes • Equations • Particulate Models • Stoichiometry • Titrations • Redox

I. Chemical vs. Physical Changes (Q1–5)

1. Dissolution as both a physical and chemical change

Dissolving an ionic compound such as NaCl in water appears physical because the overall chemical identity of Na⁺ and Cl⁻ remains the same—they are simply separated and dispersed through the solvent. However, because the dissolution process requires breaking ionic bonds in the solid lattice and forming new ion–dipole interactions with water molecules, it can also be argued as a chemical change at the microscopic level. No new substances are formed, but the fundamental interactions holding ions together are replaced with different interactions in solution. Chemically, the species remain ions, but energetically and structurally, meaningful changes occur.

2. Evidence of chemical change from precipitation and heat

The formation of a cloudy solid (precipitate) indicates that a new substance with different solubility properties formed from the reaction of ions in solution. A chemical reaction has occurred because ions reorganized into an insoluble compound. The slight temperature increase reflects an exothermic process where new bonds or lattice structures form, releasing energy. Together, these observations show that composition changed, not just state or mixture.

3. Why phase change does not alter composition

In boiling water, molecules of H₂O remain chemically identical before, during, and after vaporization; only the spacing and intermolecular interactions (hydrogen bonds) change. No covalent bonds are broken. Heat energy is absorbed to overcome intermolecular attractions, increasing kinetic energy until the molecules transition into gas phase, but the chemical structure remains unchanged—thus, this is a physical change.

4. Why heat alone is insufficient evidence

Heat production can accompany either a chemical reaction (e.g., combustion) or a physical process (e.g., dissolving CaCl₂). However, the evolution of gas bubbles when metal reacts with acid indicates formation of a new substance (H₂ gas), which necessitates bond breaking and forming. Additionally, if the metal dissolves, the metal atoms are oxidized into ions—clear chemical change indicators. The combination of bubbles and temperature change strongly indicates chemical change.

5. Melting ice vs. combustion of methane

Melting ice is a physical change involving only the disruption of hydrogen bonding between water molecules; the molecules themselves remain H₂O. In contrast, combustion of methane breaks C–H bonds and O=O bonds and forms new C=O and O–H bonds, producing CO₂ and H₂O. The atoms are rearranged into new substances, so this is a chemical change. Melting involves no new substances, while combustion fundamentally alters molecular structure.

II. Chemical, Complete Ionic, & Net Ionic Equations (Q6–10)

6. CaCl₂(aq) + Na₂CO₃(aq)

(a) Molecular:
 CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)

(b) Complete ionic:
 Ca²⁺(aq) + 2Cl⁻(aq) + 2Na⁺(aq) + CO₃²⁻(aq) → CaCO₃(s) + 2Na⁺(aq) + 2Cl⁻(aq)

(c) Net ionic:
 Ca²⁺(aq) + CO₃²⁻(aq) → CaCO₃(s)

Spectator ions (Na⁺, Cl⁻) are removed because they do not participate in formation of the precipitate and undergo no chemical change.

7. Conservation of mass & charge in acid–base neutralization

In HCl(aq) + NaOH(aq), the ionic equation is:
H⁺ + Cl⁻ + Na⁺ + OH⁻ → H₂O + Na⁺ + Cl⁻
Net ionic: H⁺ + OH⁻ → H₂O.

The total number of atoms is identical on both sides. Charge is also conserved: +1 + (–1) = 0 on both sides. This demonstrates that chemical equations represent rearrangement of atoms, not their formation or destruction.

8. Determining correct net ionic equation

To validate a net ionic equation:

  1. Check that strong electrolytes are written as ions.

  2. Confirm atoms and charge balance.

  3. Ensure only species that chemically change remain.

  4. Verify that the reaction forms a precipitate, gas, or weak electrolyte.
    Only the equation that satisfies all these conditions accurately represents the reaction.

9. When to use each equation form

  • Molecular: used when discussing the overall process, particularly in labs where solutions are identified by formula rather than ions.

  • Complete ionic: used to show dissociation behavior of strong electrolytes.

  • Net ionic: used when focusing on species undergoing chemical change.
    Example: HCl + Mg(OH)₂ → ionic shows spectator ions; net ionic shows only H⁺ + OH⁻.

10. Why some reactions yield “no net ionic equation”

If all species remain aqueous ions before and after mixing, no chemical change occurs. For example, mixing NaCl(aq) and KNO₃(aq) results in no precipitate, gas, weak acid, or weak base. Ions remain dissociated and unchanged at the particle level, so there is no net reaction.

III. Particulate Representations (Q11–15)

11. Particulate model for Mg + HCl

Solid Mg atoms appear as a lattice or sheet of neutral atoms. When contacting H⁺ ions in solution, electrons transfer from Mg to H⁺, forming Mg²⁺ ions that enter solution while H₂ molecules form and leave as gas. Cl⁻ ions remain free spectator ions in solution.

12. Converting balanced equations to particulate models

Each coefficient in the balanced equation corresponds to the ratio of molecules or ions drawn. For example, 2H₂ + O₂ → 2H₂O means two H₂ molecules must be shown for every one O₂ molecule, and the product must show twice as many water molecules. Particle counts must reflect stoichiometric ratios.

13. Differences between strong and weak electrolytes

Strong electrolytes (e.g., HCl) fully dissociate, so the particulate model shows only free ions. Weak electrolytes (e.g., HC₂H₃O₂) partially dissociate, showing a mixture of intact molecules and a small number of ions. This difference explains why strong electrolytes conduct electricity more effectively.

14. Physical vs. chemical particulate models

A physical change model shows particles rearranged in spacing or phase but with the same connectivity (e.g., closer together in solids).
A chemical change shows atoms bonded differently in products than in reactants, with new molecular structures. Particle identity changes only in chemical processes.

15. Limiting reactant in particulate models

Particles of the limiting reactant are entirely consumed first, shown by their disappearance while particles of the excess reactant remain unreacted. This visual representation matches stoichiometric calculations.


IV. Bond Interactions & Processes (Q16–20)


16. Vaporization vs. decomposition energies

Vaporization only breaks intermolecular forces such as hydrogen bonds, which require far less energy than breaking covalent bonds. Decomposition requires rearranging atoms by breaking intramolecular bonds, which requires significantly greater energy input because covalent bonds are much stronger.

17. Dissolving LiBr as physical or chemical

LiBr(s) breaks ionic bonds as it dissolves, which suggests chemical change. However, no new chemical species are created; Li⁺ and Br⁻ simply become solvated by water. Since solvated ions have different interactions but unchanged identity, dissolution straddles physical/chemical classification.

18. Bond changes in H₂ + O₂ → H₂O

Reactants undergo bond breaking (H–H and O=O) which requires energy. Product formation (O–H bonds) releases more energy than is consumed. The large net release of energy manifests as heat and light macroscopically. The macroscopic exothermic behavior reflects net bond formation energy.

19. Plateaus on heating curves

Plateaus occur during melting or boiling when temperature remains constant while energy is used to disrupt intermolecular forces instead of raising kinetic energy. These plateaus represent physical changes because molecular identity remains constant.

20. Simultaneous chemical + physical processes

Burning a candle involves melting wax (physical) and combustion of vaporized hydrocarbons (chemical). The solid–liquid transition occurs without changing composition, while combustion forms CO₂ and H₂O, demonstrating chemical change at the molecular level.

V. Stoichiometry (Q21–25)

21. Mg + HCl stoichiometric reasoning

Balanced equation: Mg + 2HCl → MgCl₂ + H₂
HCl moles = 3.0 M × 0.050 L = 0.150 mol
Mg is in excess if more than 0.075 mol is present, because HCl requires double the moles. The limiting reactant is HCl.
Moles of H₂ = ½ × moles of HCl = 0.075 mol H₂.
The limiting reactant determines maximum product formed.



22. Coefficients vs. subscripts

Coefficients show the number of molecules; subscripts show the number of atoms within each molecule. Changing a subscript changes identity (H₂O → H₂O₂). Changing a coefficient changes the amount. Stoichiometric mole ratios rely solely on coefficients.

23. Reasons for <100% yield

  1. Side reactions can consume reactants.

  2. Incomplete reaction may leave some reactant unreacted.

  3. Loss during transfer such as spillage or sticking to glassware.
    Each reason reflects imperfect conversion of reactants to product at the particulate level.

24. Linking stoichiometry + ideal gas law

Use PV = nRT to find moles of CO₂ formed, which can be converted back to moles of reactant via mole ratios. Gas volume allows quantification of reactant usage.

25. Using molarity for solution stoichiometry

Moles = M × V. Use balanced equation to convert between reactant and product moles.
Example: 25.0 mL of 2.0 M NaOH → 0.050 mol.
Use mole ratio to find amount of acid neutralized.

VI. Titrations (Q26–30)

26. Equivalence vs. endpoint

Equivalence point: moles of titrant = moles of analyte based on reaction stoichiometry.
Endpoint: observable signal (e.g., phenolphthalein turning pink).
The endpoint approximates the equivalence point; ideally, they match closely.

27. Determining concentration of monoprotic acid

  1. Measure volume of acid.

  2. Titrate with standardized strong base.

  3. Record volume of base at endpoint.

  4. Use:
    M_acid × V_acid = M_base × V_base

  5. Solve for unknown concentration.
    Accurate titration assumes reaction goes to completion and has 1:1 stoichiometry.

28. Effect of overshooting endpoint

Excess base makes the titrant volume artificially high, causing the calculated concentration of acid to be too low, since the student appears to have needed more titrant to neutralize it.

29. Comparing strong vs. weak acid titration curves

A strong acid titration begins at very low pH and has a steep equivalence region.
A weak acid titration begins at higher pH, has a buffer region, and has a higher pH at equivalence because conjugate base hydrolysis occurs. Structure influences degree of ionization and curve shape.

30. Why titrations require completion

If the reaction does not go to completion, the titrant won’t accurately reflect analyte amount because not all analyte molecules react. Stoichiometric predictions assume full conversion.

VII. Redox (Q31–35)


31. Zn + CuSO₄

Oxidation numbers:
Zn(s) = 0
Cu²⁺ = +2
SO₄²⁻ = –2 overall (internal S=+6, O=–2)
In products: Zn²⁺ = +2, Cu(s) = 0.

Zn is oxidized (0 → +2).
Cu²⁺ is reduced (+2 → 0).
Particulate: Zn atoms lose electrons; Cu²⁺ ions gain electrons and plate out as solid copper.

32. Why redox must be simultaneous

Electrons cannot exist freely in solution. When one species loses electrons (oxidation), another must gain them (reduction). Charge must be conserved, so every oxidation event couples with a reduction event.

33. Correct oxidation number method

  1. Assign to elemental forms: 0.

  2. Use rules for oxygen (–2), hydrogen (+1), halogens (–1).

  3. Apply sum-of-oxidation-numbers = charge on ion/molecule.
    This systematic approach prevents misassignments and reveals electron transfer correctly.


34. Oxidation number changes in combustion vs. single replacement

Combustion: carbon increases oxidation state (oxidized), oxygen decreases (reduced).
Single replacement: metals typically oxidize (Zn → Zn²⁺) while ions reduce (Cu²⁺ → Cu).
Both show electron movement through changes in oxidation states.

35. How oxidation numbers reveal electron flow

Balanced molecular equations rarely show electrons explicitly. Oxidation numbers allow identification of which atoms lose or gain electrons by tracking their oxidation states. Increases indicate oxidation; decreases indicate reduction.