AP Chem Unit 4 Review: Chemical Reactions

7–9% of AP multiple choice
9 topics
AP exam review verified for 2027

Unit 4 covers chemical and physical changes, balanced molecular and net ionic equations, particulate representations, stoichiometry, titration, and the classification of reactions as acid-base, precipitation, or redox. It accounts for 7-9% of the AP exam score.

Fiveable's AP Chemistry Unit 4 review covers 9 topics, with guided practice questions for the unit.

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AP Chem unit 4 topics

4.1

Introduction to Reactions

Distinguish physical changes (phase, mixture) from chemical changes (new substances) using macroscopic evidence such as heat, light, gas, precipitate, or color change.

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4.2

Net Ionic Equations

Write balanced molecular, complete ionic, and net ionic equations; remove spectator ions; conserve mass and charge in all three forms.

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4.3

Representations of Reactions

Translate balanced equations into particulate diagrams where particle counts match stoichiometric coefficients and state symbols determine whether species appear as ions or intact formulas.

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4.4

Physical and Chemical Changes

Explain changes at the bond level: chemical processes break or form covalent or ionic bonds; physical processes change only intermolecular forces. Dissolving ionic salts is a borderline case involving both.

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4.5

Stoichiometry

Use mole ratios from balanced equations to convert between masses, moles, volumes of gases (PV = nRT), and solution volumes (molarity). Identify limiting reactants and calculate theoretical and percent yield.

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4.6

Introduction to Titration

Determine the equivalence point when moles of titrant exactly consume moles of analyte; use n = M x V and the stoichiometric mole ratio to calculate unknown concentrations.

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4.7

Types of Chemical Reactions

Classify reactions as acid-base (proton transfer), precipitation (insoluble solid from aqueous ions), or redox (electron transfer shown by oxidation number changes); combustion is a redox subtype.

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4.8

Introduction to Acid-Base Reactions

Apply the Bronsted-Lowry model to identify proton donors and acceptors, label conjugate acid-base pairs, and recognize water as amphoteric in aqueous solution.

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4.9

Oxidation-Reduction (Redox) Reactions

Balance redox equations by writing and combining oxidation and reduction half-reactions; balance atoms, then oxygen with H2O, then hydrogen with H+, then charge with electrons.

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Practice AP Chem unit 4 questions

Try AP-style questions on unit 4, then use the review notes below for anything you miss.

Example stimulus-based MCQs

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Stimulus-based practice question

$0.10$ mol of $\text{NaHCO}_3$ is added separately to $1.0\ \text{L}$ of a strong acid solution and to $1.0\ \text{L}$ of a strong base solution. The bar chart shows the final concentrations of the carbon-containing species in each mixture.

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Question

Which statement best explains the behavior of $\text{HCO}_3^-$ ?

$\text{HCO}_3^-$ is amphiprotic, donating a proton in Experiment 1 and accepting a proton in Experiment 2.

$\text{HCO}_3^-$ acts as a strong acid in both experiments because it completely ionizes in water.

$\text{HCO}_3^-$ acts as a strong base in both experiments because it completely ionizes in water.

$\text{HCO}_3^-$ is amphiprotic, accepting a proton in Experiment 1 and donating a proton in Experiment 2.

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Stimulus-based practice question

$0.10\ \text{M}\ NaOH(aq)$ is added to $20.0\ \text{mL}$ of $0.10\ \text{M}\ HCl(aq)$ . The scatter plot shows the solution conductivity as a function of the volume of $NaOH$ added.

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Question

Which statement best supports the claim that neutralization occurred?

The minimum at $20\ \text{mL}$ shows that an insoluble precipitate removed the ions.

The minimum at $20\ \text{mL}$ shows that highly mobile $H^+$ ions were consumed to form water.

The initial decrease in conductivity shows that $Na^+$ gained electrons from $Cl^-$ .

The final increase in conductivity shows that the $HCl$ completely evaporated.

Example FRQs

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SAQ

Redox titration iron content analysis

7. A student analyzes the iron content of a supplement tablet using a redox titration. The student dissolves the tablet in an acidic solution and titrates it with potassium dichromate, K₂Cr₂O₇(aq), using the experimental setup shown in Figure 1.

The reaction between the iron(II) ions and the dichromate ions is represented by the following balanced net ionic equation.

Table 1. Data recorded during the titration of the iron tablet

Measurement	Value
Mass of iron tablet	1.500 g
Molarity of K₂Cr₂O₇ solution	0.0150 M
Initial buret reading	0.50 mL
Final buret reading	24.50 mL

Figure 1. Redox titration setup for determining iron content using standard 0.0150 M K₂Cr₂O₇(aq)

Using the data in Table 1, calculate the number of moles of Cr₂O₇²⁻ used to reach the endpoint of the titration.

Calculate the number of moles of Fe²⁺ present in the dissolved tablet solution.

ii.

Using your answer to part B(i), calculate the mass percent of iron in the tablet. (The molar mass of Fe is 55.85 g/mol).

The student rinses the buret with distilled water but fails to rinse it with the standard K₂Cr₂O₇ solution before filling it. Would the calculated mass percent of iron in the tablet be greater than, less than, or equal to the actual mass percent? Justify your answer.

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SAQ

Isotope abundance, ionization energy, chloride formation

6. A scientist investigates a pure sample of Element M. The scientist obtains the mass spectrum for the element, shown in Figure 1, and determines the successive ionization energies for Element M, shown in Table 1.

Figure 1. Mass spectrum of Element M (relative abundance of three isotopes).

Table 1. Successive Ionization Energies of Element M

Ionization	Energy (kJ/mol)
First	738
Second	1450
Third	7730
Fourth	10500

Calculate the average atomic mass of Element M based on the data in Figure 1.

Based on the data in Table 1, identify Element M. Then, write the complete ground-state electron configuration for an atom of Element M.

A 1.215 g sample of Element M reacts completely with excess $Cl_2(g)$ to form $MCl_2$ . Calculate the mass, in grams, of $MCl_2$ produced. Element M reacts with chlorine gas to form the compound $MCl_2$ .

Element Z is found immediately to the left of Element M in the same period of the periodic table. A student claims that the first ionization energy of Element Z is less than the first ionization energy of Element M. Do you agree or disagree with the student's claim? Justify your answer using principles of atomic structure.

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SAQ

Limiting reactant and precipitation products

4. A student mixes an aqueous solution of silver nitrate, $AgNO_3$ , with an aqueous solution of potassium chromate, $K_2CrO_4$ . The student generates the particle diagram in Figure 1 to represent the specific amounts of ions in the solutions before they are mixed.

Figure 1. Particle diagram of reactant solutions

Identify the limiting reactant in the mixture represented in Figure 1.

Figure 2. Beaker for product mixture after reaction

Figure 2 represents the mixture after the reaction goes to completion. In Figure 2, draw the particulate representation of the reaction products, including the correct number of ions and precipitate particles.

Measurement	Mass (g)
Mass of dry filter paper	0.500
Mass of filter paper + dry precipitate	1.827

The student filters the mixture to collect the precipitate, dries it, and obtains the data shown in the table.

Propose a reason why the calculated mass of the precipitate would be higher than the true mass if the student forgot to wash the precipitate with distilled water before drying.

ii.

Calculate the number of moles of precipitate, $Ag_2CrO_4$ (molar mass 331.73 g/mol), that were collected in the experiment.

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AP Chem unit 4 practice test

A timed, AP-style test built only from unit 4 questions — multiple choice plus free response, scored with AI feedback. Use the standard format or build your own.

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Unit 4 review notes

4.1

Physical vs. Chemical Changes

A physical change alters properties or phase without changing chemical composition. A chemical change makes or breaks chemical bonds and produces new substances. At the macroscopic level, evidence of a chemical change includes heat or light production, gas evolution, precipitate formation, and color change. At the particle level, the distinction depends on which bonds or forces are involved: breaking or forming covalent or ionic bonds signals a chemical change, while changes only in intermolecular forces signal a physical change. Dissolving an ionic salt like NaCl is a borderline case because ionic bonds break and ion-dipole interactions form, so it can be argued either way.

Physical change: Change in phase or mixture state with no change in chemical composition; only intermolecular forces are disrupted, as in melting or boiling.
Chemical change: Transformation producing new substances through bond breaking and formation; evidenced by heat, light, gas, precipitate, or color change.
Ion-dipole interactions: Attractions between ions and polar water molecules that form when an ionic compound dissolves; these replace the broken ionic bonds.
Ionic bond breaking in dissolution: When NaCl dissolves, the ionic lattice breaks apart; this is why dissolution can be classified as either physical or chemical depending on the argument.

Given a scenario such as iron rusting, water boiling, or salt dissolving, identify whether it is a physical or chemical change and justify your answer at the particle level.

Feature	Physical Change	Chemical Change
Bonds affected	Intermolecular forces only	Covalent or ionic bonds broken/formed
Composition	Unchanged	New substances produced
Macroscopic evidence	Phase or shape change	Heat, light, gas, precipitate, color change
Example	Ice melting	Magnesium burning in oxygen
Reversibility (general)	Often reversible	Often not reversible

4.2

Net Ionic Equations and Particulate Models

Any chemical or physical process can be written as a balanced equation. Three forms exist: the molecular equation shows full formulas, the complete ionic equation splits all soluble strong electrolytes into their ions, and the net ionic equation removes spectator ions that appear identically on both sides. Both mass and charge must be conserved in every form. Particulate diagrams translate the balanced equation into a visual model showing individual atoms, molecules, and ions before and after the reaction. Coefficients in the equation correspond directly to the number of particles drawn. State symbols (aq, s, l, g) determine whether a species is shown as dissociated ions or as an intact formula unit.

Spectator ions: Ions that appear in the same form on both sides of a complete ionic equation and are removed to produce the net ionic equation.
Net ionic equation: Shows only the species that actually change during the reaction; charge and mass must balance.
Particulate representation: A visual depiction of a reaction at the atomic or ionic level; particle counts must match stoichiometric coefficients.
Conservation of charge: The total charge on the reactant side must equal the total charge on the product side in any balanced ionic equation.

Write the molecular, complete ionic, and net ionic equations for the reaction of aqueous lead(II) nitrate with aqueous potassium iodide, then sketch a particulate diagram of the net ionic equation.

Equation Form	What is shown	Spectator ions included?
Molecular	Full chemical formulas	Yes
Complete ionic	All soluble strong electrolytes split into ions	Yes
Net ionic	Only species that react	No

4.5

Stoichiometry

Stoichiometry uses the mole ratios from a balanced equation to convert between amounts of reactants and products. The standard pathway is: convert the given quantity to moles using molar mass or molarity, apply the mole ratio from the balanced equation, then convert to the desired unit. For gases, the ideal gas law (PV = nRT) connects moles to volume and pressure. For solutions, molarity (mol/L) connects moles to volume. The limiting reactant is the one that runs out first and sets the theoretical yield. Percent yield equals actual yield divided by theoretical yield, multiplied by 100.

Mole ratio: The ratio of coefficients from a balanced equation used to convert moles of one substance to moles of another.
Limiting reactant: The reactant completely consumed first; it determines the maximum amount of product that can form.
Theoretical yield: The maximum mass or moles of product calculated from the limiting reactant assuming complete reaction.
Percent yield: Actual yield divided by theoretical yield, multiplied by 100; measures reaction efficiency.
Dimensional analysis: Unit-tracking method used to chain conversions from a given quantity to the target unit without losing track of units.

Given that 5.00 g of hydrogen gas reacts with excess oxygen, calculate the theoretical yield of water in grams and the percent yield if 40.0 g of water is actually collected.

4.6

Introduction to Titration

A titration determines the unknown amount of an analyte by reacting it with a titrant of known concentration. The equivalence point is reached when the analyte is completely consumed by the titrant, based on the stoichiometric mole ratio of the reaction. The endpoint is the observable signal, usually an indicator color change, that approximates the equivalence point. To find the moles of analyte, multiply the titrant molarity by the volume used, then apply the mole ratio. Titration calculations are a direct application of solution stoichiometry: n = M x V.

Titrant: Solution of known concentration added from a buret to react with the analyte.
Analyte: The species of unknown amount or concentration being determined in the titration.
Equivalence point: The point at which moles of titrant added exactly consume all moles of analyte according to the stoichiometric ratio.
Endpoint: The observable event, such as an indicator color change, that signals the titration is complete.
Indicator: A substance that changes color at or near the equivalence point, used to signal when to stop adding titrant.

A 25.00 mL sample of HCl is titrated with 0.150 M NaOH. If 32.40 mL of NaOH is needed to reach the equivalence point, calculate the molarity of the HCl solution.

4.7

Types of Chemical Reactions

AP Chemistry classifies reactions into three types. Acid-base reactions transfer protons (H+) from acid to base. Precipitation reactions combine aqueous ions to form an insoluble solid; solubility rules determine which products precipitate. Redox reactions transfer electrons between species, tracked by changes in oxidation numbers. Combustion is a subtype of redox in which a substance reacts with oxygen; complete combustion of a hydrocarbon produces CO2 and H2O. To identify the reaction type, look for proton transfer, an insoluble product, or a change in oxidation numbers.

Precipitation reaction: Two aqueous ionic solutions combine to form an insoluble solid product; identified using solubility rules.
Combustion reaction: A redox reaction in which a substance reacts with O2; complete hydrocarbon combustion yields CO2 and H2O.
Oxidation number: A hypothetical charge assigned to an atom to track electron distribution; a change in oxidation number signals a redox reaction.
Solubility rules: Guidelines predicting whether an ionic compound dissolves in water; used to identify precipitates in double-displacement reactions.

Classify each of the following as acid-base, precipitation, or redox, and justify: (a) HCl + NaOH, (b) AgNO3 + NaCl, (c) Zn + CuSO4.

Reaction Type	What transfers	Key evidence	Example
Acid-base	Proton (H+)	Neutralization, pH change	HCl + NaOH → NaCl + H2O
Precipitation	Nothing transfers; ions combine	Insoluble solid forms	AgNO3 + NaCl → AgCl(s) + NaNO3
Redox	Electrons	Oxidation number change	Zn + CuSO4 → ZnSO4 + Cu

4.8

Bronsted-Lowry Acid-Base Reactions

A Bronsted-Lowry acid donates a proton (H+) and a Bronsted-Lowry base accepts one. Every acid-base reaction produces a conjugate acid-base pair: the conjugate base is the acid after it loses H+, and the conjugate acid is the base after it gains H+. Conjugate pairs differ by exactly one proton. Water is amphoteric and can act as either acid or base in aqueous solution. Lewis acid-base concepts are excluded from the AP exam; focus on proton transfer in aqueous solution. Note that strong acids and strong bases dissociate completely, while weak acids and bases ionize only partially.

Bronsted-Lowry acid: A proton donor; loses H+ in the reaction.
Bronsted-Lowry base: A proton acceptor; gains H+ in the reaction.
Conjugate acid: The species formed when a base accepts a proton; differs from the base by one H+.
Conjugate base: The species remaining after an acid donates its proton; differs from the acid by one H+.
Ionization: The partial dissociation of a weak acid or base in water to produce ions; strong acids and bases dissociate completely.

For the reaction CH3COOH + H2O ⇌ H3O+ + CH3COO-, identify the Bronsted-Lowry acid, base, conjugate acid, and conjugate base.

4.9

Balancing Redox Reactions with Half-Reactions

Redox equations are balanced by splitting the overall reaction into an oxidation half-reaction and a reduction half-reaction. In each half-reaction, balance atoms other than O and H first, then balance O by adding H2O, balance H by adding H+, and balance charge by adding electrons. Multiply each half-reaction by a factor so the electrons cancel when the two are added together. In basic solution, add OH- to neutralize any H+ after balancing in acid. The oxidizing agent is reduced (gains electrons) and the reducing agent is oxidized (loses electrons).

Oxidation half-reaction: Shows the species that loses electrons; electrons appear as products on the right side.
Reduction half-reaction: Shows the species that gains electrons; electrons appear as reactants on the left side.
Oxidizing agent: The species that accepts electrons and is itself reduced in the reaction.
Reducing agent: The species that donates electrons and is itself oxidized in the reaction.

Balance the following redox reaction in acidic solution using half-reactions: MnO4- + Fe2+ → Mn2+ + Fe3+.

Key terms

Term	Definition
Chemical Change	A transformation that produces new substances through the breaking and forming of chemical bonds; evidenced by heat, light, gas, precipitate, or color change.
Conservation of Mass	The principle that atoms cannot be created or destroyed in a chemical reaction; the total mass of reactants equals the total mass of products.
Conservation of Charge	The total electric charge must be equal on both sides of a balanced chemical equation; essential for writing correct ionic and redox equations.
Spectator Ions	Ions that appear in identical form on both sides of a complete ionic equation and do not participate in the reaction.
Mole Ratios	Conversion factors derived from the coefficients of a balanced equation, used to relate moles of one substance to moles of another in stoichiometry calculations.
Limiting Reactant	The reactant completely consumed first in a reaction; it determines the theoretical yield of product.
theoretical yield	The maximum amount of product calculated from the limiting reactant, assuming complete reaction with no losses.
Percent Yield	Actual yield divided by theoretical yield, multiplied by 100; measures how efficiently a reaction produces its expected product.
Equivalence Point	The point in a titration at which moles of titrant have exactly consumed all moles of analyte according to the stoichiometric mole ratio.
Oxidation numbers	A hypothetical charge assigned to an atom to track electron distribution; a change in oxidation number from reactants to products identifies a redox reaction.
Precipitation Reaction	A reaction in which two aqueous ionic solutions combine to form an insoluble solid product, identified using solubility rules.
Redox Reaction	A reaction involving electron transfer between species; the oxidizing agent is reduced and the reducing agent is oxidized, tracked by changes in oxidation numbers.
particulate representation	A visual depiction of a reaction at the atomic or ionic level showing individual particles before and after the reaction; particle counts must match stoichiometric coefficients.

How to review unit 4

Topics students miss most

Miss rates come from Fiveable multiple-choice practice.

4.9

Oxidation-Reduction (Redox) Reactions

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4.8

Introduction to Acid-Base Reactions

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4.2

Net Ionic Equations

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Common unit 4 mistakes

Forgetting to check charge balance in net ionic equations

Students often balance atoms but leave the net charge unequal. After canceling spectator ions, verify that the total charge on the left equals the total charge on the right.

Using the wrong mole ratio in stoichiometry

The mole ratio must come from the balanced equation coefficients, not from the subscripts in the formulas. Always write the balanced equation first before setting up any conversion.

Confusing equivalence point with endpoint

The equivalence point is defined by stoichiometry (moles of titrant equal moles of analyte per the mole ratio). The endpoint is the observable indicator color change, which approximates but is not identical to the equivalence point.

Misassigning oxidation numbers in polyatomic ions

Apply the rules systematically: oxygen is usually -2, hydrogen is usually +1, and the sum of oxidation numbers must equal the ion charge. Errors here lead to incorrectly identifying which species is oxidized or reduced.

Not dissociating strong electrolytes in complete ionic equations

Soluble ionic compounds and strong acids must be written as separated ions in the complete ionic equation. Leaving them as molecular formulas produces incorrect spectator ion cancellation and a wrong net ionic equation.

How this unit shows up on the AP exam

Justify claims with particle-level evidence

AP Chemistry free-response questions frequently ask you to explain a macroscopic observation at the particle level. For Unit 4, this means connecting visible evidence of a reaction (precipitate, color change, temperature change) to bond breaking or formation, ion interactions, or electron transfer rather than simply naming the observation.

Multi-step quantitative calculations

Stoichiometry and titration calculations on the AP exam often chain multiple conversions: for example, converting a solution volume and molarity to moles, applying a mole ratio, then converting to grams or liters of gas using PV = nRT. Showing dimensional analysis with units at every step is expected and earns method credit even if a numerical error occurs.

Classify and represent reactions in multiple forms

The AP exam tests whether you can identify a reaction type, write the correct net ionic equation, and represent the same reaction as a particulate diagram. Questions may give you one form and ask you to produce another, or ask you to identify the reaction type from a diagram alone. Practicing all three representations together is more efficient than studying them separately.

More ways to review

Topic study guides

Open the individual guides for Unit 4 when you want a closer review of one topic.

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Practice questions

Use AP-style practice after you review the notes so you can check what you understand.

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FRQ practice

Practice free-response reasoning and compare your answer with scoring guidance.

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Cram archive videos

Watch past review streams filtered to Unit 4 when you want a video walkthrough.

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Official unit cheatsheet

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Score calculator

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Unit 4 printables

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Frequently Asked Questions

What's on the AP Chem Unit 4 progress check (MCQ and FRQ)?

The AP Chem Unit 4 progress check includes both MCQ and FRQ parts drawn from all 9 topics in the unit. MCQ questions test stoichiometry calculations, net ionic equations, identifying types of chemical reactions, and physical vs. chemical changes. The FRQ portion typically asks you to write or interpret reactions, balance equations, or work through a titration or redox problem. Practicing those same topics before you take the progress check in AP Classroom is the best prep move. Find matched practice at /ap-chem/unit-4.

How do I practice AP Chem Unit 4 FRQs?

The best way to practice AP Chem Unit 4 FRQs is to focus on the topics that generate free-response questions most often: stoichiometry calculations, titration problems, net ionic equations, and oxidation-reduction (redox) reactions. FRQ prompts in this unit usually ask you to write a balanced equation, calculate moles or concentrations, or justify whether a change is physical or chemical. Practice by writing out full solutions and checking your work step by step, not just the final answer. Past FRQs from College Board and topic-specific practice sets at /ap-chem/unit-4 are both solid starting points.

Where can I find AP Chem Unit 4 practice questions?

For AP Chem Unit 4 practice questions, including multiple-choice and practice test sets, head to /ap-chem/unit-4. You'll find MCQ practice covering stoichiometry, types of chemical reactions, net ionic equations, and titration, plus FRQ sets that mirror the format of the real exam. Mixing MCQ drills with full FRQ write-outs gives you the best coverage of all 9 topics in the unit.

How should I study AP Chem Unit 4?

Start AP Chem Unit 4 by locking in stoichiometry first, since mole calculations run through almost every other topic in the unit. From there, work through net ionic equations and types of chemical reactions together, since both require you to recognize what's actually happening in a reaction. Then move into titration and acid-base reactions as a pair, and finish with redox. A few concrete steps that help:
- Practice balancing equations by hand until it's automatic.
- For net ionic equations, always cancel spectator ions before checking your answer.
- For titration problems, write out the mole ratio before plugging in numbers.
- Do at least one timed FRQ per topic so you know how to show your work under pressure. All 9 topics and practice sets are at /ap-chem/unit-4.

Ready to review Unit 4?Start with the notes, check the topic cards, and use the practice or resource links when they are available for this course.

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