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AP Chemistry (REA) - The Best Test Prep for: 10th Edition

AP Chemistry (REA) - The Best Test Prep for: 10th Edition

by P. E. Dumas, R. M. Fikar, J. W. Samples, Jay M. Templin, W. C. Uhland


REA … Real review, Real practice, Real results.


Get the college credits you deserve.


AP CHEMISTRY, 10th Edition

Completely aligned with today’s AP exam


Are you prepared to excel on the AP exam? 

* Set up a study schedule by



REA … Real review, Real practice, Real results.


Get the college credits you deserve.


AP CHEMISTRY, 10th Edition

Completely aligned with today’s AP exam


Are you prepared to excel on the AP exam? 

* Set up a study schedule by following our results-driven timeline

* Take the first practice test to discover what you know and what you

   should know

* Use REA's advice to ready yourself for proper study and success

Practice for real

* Create the closest experience to test-day conditions with 6 full-length practice tests

* Chart your progress with full and detailed explanations of all answers

* Boost your confidence with test-taking strategies and experienced advice

Sharpen your knowledge and skills

* The book’s full review features coverage of all AP Chemistry main topic areas, such as solutions, stoichiometry, kinetics, and thermodynamics, as well as all subject areas found on the official exam, including the structure of matter, the states of matter, chemical reactions, and descriptive chemistry.

* Smart and friendly lessons reinforce necessary skills

* Key tutorials enhance specific abilities needed on the test

* Targeted drills increase comprehension and help organize study

Ideal for Classroom, Family, or Solo Test Preparation!

REA has provided advanced preparation for generations of advanced students who have excelled on important tests and in life. REA’s AP study guides are teacher-recommended and written by experts who have mastered the course and the test.



Product Details

Research & Education Association
Publication date:
Test Preps
Edition description:
Product dimensions:
6.60(w) x 9.90(h) x 1.60(d)
Age Range:
16 - 17 Years

Read an Excerpt



This book provides a thorough review for the Advanced Placement (AP) Chemistry Examination written in a way that high school students will readily grasp and appreciate. REA’s mission is to translate chemistry into terms the student can understand and benefit from.

Six full-length practice exams are included to get you ready for the actual exam. Use them, along with the detailed explanations of answers, to help determine your strengths and weaknesses, and to prepare you to score well on exam day.


The Advanced Placement Chemistry Examination is offered each May at participating schools and testing centers throughout the world. The Advanced Placement Program is designed to allow high school students to pursue college-level studies while attending high school. Participating colleges, in turn, grant credit and/or advanced placement to students who do well on the examinations.

The Advanced Placement Chemistry course is designed to be the equivalent of a college introductory chemistry course, often taken by chemistry majors in their first year of college. Since the test covers a broad range of topics, no student is expected to answer all of the questions correctly. (Consequently, it is important for students to not feel defeated when confronting a question that appears unanswerable.)

The AP Chemistry exam is divided into two sections. The first section is composed of 75 multiple-choice questions designed to test recall of a broad range of concepts and calculations. The score students earn on the multiple-choice test composes 50% of the total score. Calculators are not permitted on the multiple-choice portion of the exam, although simple arithmetic may be needed to answer some questions.

The second section is a free-response exam and constitutes the other 50% of the final grade. Calculators are allowed on Part A, during which the student has 55 minutes to complete three problems. Calculators are not allowed on Part B, during which students have 40 minutes to complete chemical reaction questions and two essay questions. Most hand-held calculators are allowed in the examination. However, calculators with

typewriter-style (QWERTY) keypads are not allowed. If you are unsure if your calculator is permitted, check with your teachers or Educational Testing Service.

A detailed outline of the topics on the examination and specific strategies for both portions of the Advanced Placement Chemistry examination follows.


The following is an outline of the general breadth of topics that the College Board identifies as being on the AP Chemistry Examination.

I. Structure of Matter (20% of the AP test)

A. Atomic theory and structure

1. Evidence for atomic theory

2. Atomic mass

3. Atomic number, mass number, isotopes

4. Electron energy levels, quantum numbers, atomic orbitals

5. Periodic relationships: atomic radii, ionization energy, electron affinity, oxidation states.

B. Chemical Bonding

1. Binding forces

a. Types of forces: ionic, covalent, network covalent, metallic,

hydrogen bonding, van der Waals.

b. Relationships to states, structure, and properties of matter.

c. Polarity of bonds, electronegativities.

2. Molecular Models

a. Lewis structures

b. Hybridization of orbitals, resonance, sigma and pi


c. VSEPR (valence shell electron pair repulsion)

3. Geometry of molecules and ions, structural isomerism of simple organic molecules and coordination compounds; dipole moments of molecules, relation of properties to structure.

C. Nuclear chemistry: nuclear equations, half-lives, and radioactivity; chemical applications.

II. States of Matter (20% of the AP test)

A. Gases

1. Laws of ideal gases

a. Equation of state for an ideal gas

b. Partial pressures

2. Kinetic-molecular theory

a. Interpretation of ideal gas laws on the basis of theory

b. Avogadro’s hypothesis and the mole concept

c. Dependence of kinetic energy on temperature

d. Deviations in the ideal gas laws

B. Liquids and Solids

1. Liquids and solids from the kinetic-molecular viewpoint

2. Phase diagrams

3. Changes of state, including critical points and triple points

4. Structure of solids; lattice energies

C. Solutions

1. Types of solutions and factors affecting solubility

2. Methods of expressing concentration

3. Raoult’s law and colligative properties; osmosis

4. Behavior of non-ideal solutions

III. Reactions (35-40% of the AP test)

A. Reaction types

1. Acid-base reactions, concepts of Arrhenius, Brønsted-Lowry, and Lewis; coordination complexes, amphoterism.

2. Precipitation reactions

3. Oxidation-reduction reactions

a. Oxidation number

b. The role of the electron in oxidation-reduction

c. Electrochemistry: electrolytic and galvanic cells; Faraday’s laws; standard half-cell potentials; Nernst equation; prediction of the direction of redox reactions

B. Stoichiometry

1. Ionic and molecular species present in chemical systems; net ionic reactions

2. Balancing of equations, including oxidation-reduction reactions

3. Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants

C. Equilibrium

1. Concept of dynamic equilibrium, physical and chemical; Le Chatelier’s principle; equilibrium constants

2. Quantitative treatment

a. Equilibrium constants for gaseous reactions: Kp, Kc

b. Equilibrium constants for reactions in solution

i. Constants for acids and bases; pK; pH

ii. Solubility product constants and their application to

precipitation and dissolution of slightly soluble compounds

iii. Common ion effect; buffers; hydrolysis

D. Kinetics

1. Concept of reaction rate

2. Use of experimental data and graphical analysis to determine reaction order, rate constants, and rate laws

3. Effect of temperature on reaction rates

4. Energy of activation; the role of catalysts

5. Relationship between the rate-determining step and mechanism of reaction

E. Thermodynamics

1. State functions

2. First law: change in enthalpy; heat of formation; heat of reaction; Hess’s law; heats of vaporization and fusion; calorimetry

3. Second law: entropy; free energy of formation; free energy of reaction; dependence of change in free energy of enthalpy and entropy changes

4. Relationship between change in free energy, equilibrium constants, and electrode potentials

IV. Descriptive Chemistry (10-15% of the AP test)

A. Chemical reactivity and products of chemical reactions

B. Relationships in the periodic table; horizontal, vertical, and diagonal with examples of alkali metals, alkaline earth metals, halogens, and the first series of transition metals

C. Introduction to organic chemistry: hydrocarbons and functional groups (structure, nomenclature, chemical properties)

V. Laboratory (5-10% of the AP test)

A. Making observations of chemical reactions

B. Recording data

C. Calculating and interpreting results based on observed quantitative data

D. Effectively communicating experimental results, including error analysis


Format and Scoring

Each correct answer in the multiple-choice section earns one point, while each wrong answer causes one-quarter (1/4) point to be taken off the score earned on this section. Omitted questions receive neither credit nor deduction.

The multiple-choice section is subtly broken into two subsections, Part A and Part B. Part A questions give information and a set of answers that can be used for three to five questions. Part B presents standard multiple-choice questions that provide information, and then offer five options for answers, options (A) through (E).Strategies

The following strategies will help you study for and take the multiple-choice section of the AP Chemistry examination.

• To prepare for the multiple-choice section, use the content review in this book. If you know each item in this content review, you will get a great score on your multiple-choice section. After studying the content review, use the practice exams to test your knowledge. When you miss a problem, go back to the content review to be sure that you understand the material.

• When taking the exam, keep a one-question-per-minute pace and go directly through the test. Quite often, easy questions toward the end of the multiple-choice test are not answered because students do not get to them. Mark the unanswered questions that you skip on your answer sheet, so you may easily return to them. Always be sure that you fill in the bubble that corresponds to the question that you are answering.

• Once you have answered most of the multiple-choice questions during the initial 60-minute pass-through of the exam, you now have 30 minutes to answer the tougher or more time-consuming questions. Once again, start at the beginning; try not to skip

around. You will probably fi nd that you are more warmed up; some of these answers will now come to you and you will wonder why you skipped the question during the first pass-through of the exam.

• For questions you don’t know, it is worth guessing IF you can eliminate at least two possible answers. If you can’t eliminate at least two answers, then leave the answer blank; the odds are stacked against you to get the question correct. Answers for

conceptual questions that contain extreme words like “never” or “always” are likely candidates for elimination. Once you guess on an answer, move on and forget about the question. Studies show that students tend to change guesses from right answers to

wrong answers more often than the other way around. Trust your intuition on these.



• You will be given 95 minutes to complete Section II, which is 50% of the value of the entire test.

• You will be given 55 minutes for Part A and 40 minutes for Part B, for a total of 95 minutes for Section II.

• In Part A, you will complete three problems, each with multiple parts (usually a through e). The fi rst question involves equilibrium topics. One of the other problems may came from a laboratory example.

• In Part B, you will complete three questions: one based on a group of three chemical reactions, and two essay questions. One of the essay questions will use a laboratory example, if one was not used in one of the Part A problems.

• Clearly demonstrate the steps you use to arrive at your answers. The curve for the test is steep, so partial credit can make the difference between a very good score and a poor score.

• Pay attention to significant figures.

• Be sure to write your answers in the space provided following each question.

• Data necessary for the solution of the problems may be found in the tables at the beginning of the test. Be familiar with the type of information that is available on this data sheet. An example of the data sheet can be found with the practice examinations in

this book.


Format and Scoring

In the first question of Part A, you will be given one equilibrium question with multiple parts. The equilibrium question may focus on gaseous systems, the solubility product constant, the base dissociation constant, or the acid dissociation constant. This part of the free-response test is worth 20 percent of the free response section of the test.


This is the most predictable section of the AP Chemistry examination,

and it can sometimes be the most difficult. However, if a student is comfortable

with all the topics mentioned under “What to Study,” then this section

will feel easy.

Nine points are distributed over the different parts of this question. Usually, the distribution of these points is easily understood by normal student intuition. For example, when given past Part A questions, current students are just as successful as experienced teachers in guessing how College Board readers will distribute points. In other words, if it seems important to you, the student, put it down to be sure that you earn credit. A few pointers will maximize the grade that you earn on this section:

• Always show your work. Even if you can’t answer the question, write down the formula that you think would be involved. You may still get two of the three points for that question without ever solving the problem. When the average score is about 4.5 points

out of 9 points, picking up two-thirds of the available points still puts you in the top half.

• Include units and the correct number of significant digits in your answer when appropriate. A general rule of thumb is to use three significant figures if you don’t know what is expected of you.

Forgetting the units of the answers may cost a point for each of the four parts of the question.

• Graders typically do not count off more than once for any mistake. Therefore, if you make a mistake in Part A, and then carrying that answer forward to solve Part B also leads to an incorrect answer for B, even though you were otherwise correct, they might

still give you full credit. However, this is unlikely to happen if you don’t show your work, so, once again, show your work.

• Circle or underline your answer so the reader can find it. This seems like a funny thing to get in the way, but it does happen that a student works out the correct answer and the reader simply isn’t sure which answer the student intends to submit.

• Practice equilibrium problems in this book and your text. In particular, be sure that you know how to find the pH of different types of solutions.

Examples of Past Questions

• A acid dissociation reaction, along with the corresponding Ka, is given.

Question a: Write the equilibrium expression for the reaction.

Question b: Determine the pH of a 0.1 M solution of the acid.

Question c: Titrate the weak acid with a strong base and determine the resulting pH.

Question d: Calculate the pH of the solution at the equivalence point.

Question e: Use the Henderson-Hasselbach equation to determine the pH of a buffer involving the weak acid.

Question f: Given the pKa of several indicators, determine the best indicator to use when titrating this particular weak acid.

• The percent ionization was given for a reaction in which a weak base is added to water.

Question a: Write the equilibrium expression for the reaction.

Question b: Calculate the Kb from the percent ionization of the weak base in water.

Question c: Calculate the pH that results when some of the weak base is titrated with a given number of moles of strong acid.

Question d: Combine one of the above equations with another equation, such as Kw, to calculate the Ka of the conjugate acid of the weak base.

• A reaction that involves gases in the reactants and/or products, and the corresponding equilibrium constant, is given.

Question a: Write the equilibrium expression for the equilibrium constant.

Question b: Given partial pressures under non-equilibrium conditions, determine Q and the direction of the reaction.

Question c: If Kp were originally given, calculate Kc.

Question d: Given other equilibrium constants for related reactions, use the combination of constants and reactions to calculate the equilibrium constant for a new reaction.

• A reaction that involves the dissolving of a slightly soluble salt, and its corresponding Ksp , is given.

Question a: Write the equilibrium expression for a dissolving process.

Question b: Determine the molar solubility of one of the ions involved in the solubility reaction.

Question c: Determine the molar solubility of one of the ions under conditions in which the other ion is already present (common ion effect), as it would be for hydroxide ions if the pH were not neutral.

Question d: Given a volume and concentration of two solutions which, when combined, cause the precipitation reaction to occur, determine whether or not the precipitation will take place.

Question e: Determine the quantity of precipitate that formed under one of the previous conditions.

What to Study

• Be sure to study Chapter 10 in this review book.

• Understand Le Chatelier’s principle; in particular how a shift in reaction concentration can affect the direction of the reaction.

• Know well the chapter on equilibrium. In particular, be sure that you understand the following about weak acid/base systems since most of these questions will involved acids and bases.

1. Determining the pH of a weak acid or basic solution.

2. Determining the pH of a buffer solution.

3. Determining the pH of a the solution that results from a partial titration—particularly adding a weak acid and strong base together, or adding a weak base and strong acid together.

• Understand the reactions associate with different types of equilibrium constants, and how those equilibrium constants might be related to one another. (For example, how to convert between Kw, Ka, and Kb; or how to convert between Kp and Kc.)

• Gaseous equilibrium may involve equilibrium vapor pressure, so be sure to study the chapter on the phases of matter (Chapter 6).

• Solubility product constant problems, including the following:

1. Determine the molar solubility knowing the Ksp, and visa versa.

2. Determine the molar solubility at various pH levels using the common ion effect, or in a buffered solution.

3. Determine whether or not a precipitate will form when two solutions are combined.

• Relate the equilibrium constant to other values, such as cell voltage, the reaction quotient, or free energy.

• Understand the relationship between equilibrium constants and multiple reactions.


Format and Scoring

For the remainder of Part A, you must solve two further problems. These problems usually involve calculations from more than one topic regarding reactions, such as thermodynamics, calorimetry, kinetics, oxidationreduction (voltaic or electrolytic cell) stoichiometry or colligative properties.

One of these problems may be drawn from the set of recommended laboratories.

Like the equilibrium question, students can earn up to nine points on each of these problems. Also like the equilibrium question, each problem is worth 20 percent of the free-response section of the test. Remember, partial credit is available for writing formulas and correct set-up, even if you can’t reach the correct solution. Quite often, one of the points is attributed to getting the units correct, so just writing the correct formula and the units of the answer may earn two-thirds of the points for a problem. This is a lot of points given the fact that the average score on these problems is typically about 50% (about 4.5 out of 9 points).


All the same problem-solving hints mentioned about the Part A problem apply to this problem as well. Show your work. Write correct formulas even if you can’t solve the problem. Include units in your answer. Take significant figures into consideration.

Examples of Past Problems

• The mass of two reactants in an oxidation-reduction reaction is given.

Question a: Identify the limiting reactant; show all calculations.

Question b: Determine the concentration of one of the products in solution.

Question c: Determine the standard electrode potential for the galvanic cell based on this reaction.

Question d: Given a similar oxidation-reduction reaction, determine the change in free energy for the reaction under standard conditions.

Question e: Given non-standard concentrations, calculate the electric potential for the cell.

• The molecular formula of a small hydrocarbon molecule is given.

Question a: Write the balanced equation for the combustion of the hydrocarbon molecule.

Question b: Determine the volume of gas that would be collected at a specific temperature and pressure if a given mass of the organic molecule were completely combusted.

Question c: The amount of heat given off during the above question is given; calculate the molar change in enthalpy for the reaction.

Question d: The hydrocarbon molecule is a gas at room temperature; compare its effusion rate with that of another gas of known molar mass.

Question e: The structure of the hydrocarbon is given; write a structure for an isomer of the hydrocarbon.

• The percent composition is given for a gaseous hydrocarbon.

Question a: Determine the empirical formula.

Question b: Determine the molar mass, given the density of the gas at a given temperature and pressure.

Question c: Compare the effusion rate of this gas with that of another, given the molar mass of the second gas.

Question d: The gas reacts in a confined chamber at a given temperature and pressure; calculate the number of moles of gas produced.

Question e: Heats of formation of water and carbon dioxide are given; calculate the heat of formation for the hydrocarbon gas.

Question f: With the heat of formation known, calculate the maximum amount of temperature change of a bomb calorimeter with a given heat capacity.

• The mass of an oxide is combined with a given volume of gas at a specific temperature and pressure.

Question a: Determine the number of moles of gas available for the reaction.

Question b: Determine the limiting reactant for the reaction.

Question c: Determine the amount of one of the products formed during the reaction.

Question d: Another reaction is given, which occurs entirely in solution. Calculate the mass of product formed given the concentrations of the reactants.

Question e: One of the products is a strong acid. Determine the pH of the solution when the reaction is complete.

• A chemical reaction is given, along with a chart of initial concentrations of each reactant and the initial reaction rate that

corresponds to the concentrations, for each of four different experiments.

Question a: Determine the order of the reaction with respect to each reactant.

Question b: Write the rate law for the overall reaction.

Question c: Calculate the value and units for the specific rate constant.

Question d: The reaction is an oxidation-reduction reaction; calculate the standard cell potential given the reduction potential for each half reaction.

Question e: Determine the total number of electrons transferred in the reaction.

What to Study

• Be sure to study the calculations involved with the following chapters in this review book; each chapter has examples that simulated the kind of questions you will find on these problems:

– Chapter 5: Gases

– Chapter 7: Solutions

– Chapter 9: Stoichiometry

– Chapter 11: Kinetics

– Chapter 12: Thermodynamics

– Chapter 14: Laboratories


Format and Scoring

In the first question of Part B, you will be given three sets of chemical reactants written in words. You must write the balanced net ionic equation (reactants and products) for each of the three reactions, then answer a question about the reaction.

Example: Solutions of ammonia and nitric acid are mixed.

i. Write the balanced, net ionic reaction.

ii. How many moles of acid will be used for every mole of ammonia?


i. NH3 + H+ ? NH4+

ii. One mole of acid will be used up for every mole of ammonia that reacts in this reaction.

You will be given a periodic table and a table of standard reduction potentials, but no other information. A calculator is not allowed or needed on this section. This section is worth 10 percent of the points on the Free-Response section of the test. This review will give you a set of steps so that it shouldn’t be difficult to earn most, if not all, of the 15 available points.


The most effective strategy in succeeding on question 4 is to recognize that most of the reactions that will be given to you will be one of four easily recognizable types of reactions that are summarized below. The following steps will help you succeed on question 4 of the free response portion of the test.

1. Write down the formulas for as many reactants as you can. If you don’t know the formula, eliminate that question and move on to another one.

2. Is it a precipitation reaction?

If you know the solubility rules, identifying these reactions and writing the resulting net ionic equations should be straightforward. Remember to not include spectator ions in the net ionic reaction. Only those ions that come together to form the insoluble inorganic compound should be written.

Question: Solutions of silver nitrate and potassium iodide are mixed.

Answer: Ag+ + I– ? AgI

Question: A solution of lead nitrate is added to a solution of ammonium sulfide.

Answer: Pb++ + S– – ? PbS

3. Is it an acid-base reaction?

Remember that most acid-base reactions simply move a proton from one reactant to another.

a. Does a strong acid neutralize a strong base?

The net ionic reaction for this type of acid-base neutralization is always the same.

Question: Strong hydrochloric acid is added to a solution of sodium hydroxide.

Answer: H+ + OH– ? H2O

b. Does a strong acid neutralize a weak base?

The net ionic equation should depict a proton combining with the basic molecule.

Question: Strong hydrochloric acid is added to an ammonia solution.

Answer: H+ + NH3 ? NH4+

c. Does a strong base react with a weak acid?

A proton from the weak acid combines with the hydroxide ion to form an anion and water.

Question: Solutions of sodium hydroxide and acetic acid are mixed.

Answer: HC2H3O2 + OH– ? C2H3O2– + H2O

d. Does a weak acid react with a weak base?

Keep an eye out for Lewis acid-base reactions, in which the acid accepts an electron pair from and combines with a weak base. Question: Solutions of boron trifluoride and ammonia are mixed.

Answer: BF3 + NH3 ? BF3NH3

e. Is a coordination-complex formed?

This is another example of a Lewis acid-base reaction, where a transition metal serves as an electron-pair acceptor (Lewis acid). You would recognize this if a transition metal is placed in a solution with soluble ammonia, cyanide, hydroxide, or thiocyanate ions. You may combine the metal with as many polyatomic anions as you wish; just be sure that the total charge on the ion is correct. The charge on the metallic atom does not change.

Question: A solution of iron (III) chloride is added to a solution of potassium thiocyanate.

Answer: Fe3+ + SCN– ? FeSCN2+

4. Is it an oxidation-reduction reaction?

One element increases in oxidation state while another is reduced in oxidation state.

a. Are two uncombined elements coming together?

Combine the two elements to form a compound with reasonable oxidation states for each element.

Question: Magnesium metal is burned in oxygen gas.

Answer: Mg + O2 ? MgO

b. Is a carbon compound combusting with oxygen?

An alkane, alkene, or alkyne is oxidized by oxygen gas to form carbon dioxide and water.

Question: Butane gas ignites in the presence of oxygen gas.

Answer: C4H10 + O2 ? CO2 + H2O

c. Is a single reactant decomposing?

Decomposition usually occurs because an uncommon oxidation state in one element gives way to a more common oxidation state.

Question: Hydrogen peroxide solution is exposed to bright light.

Answer: H2O2 ? O2 + H2O

d. Is a solid transition metal placed in a solution of metallic ions?

Use the chart of standard reduction potentials. The change with the highest reduction potential is reduced in charge. Voltaic or galvanic cells are an example of this type of reaction.

Question: Copper metal is placed in a solution of silver nitrate.

Answer: Cu + Ag+ ? Cu++ + Ag

e. Does an electrical current pass through a solution?

If so, this reaction takes place in an electrolytic cell. Only a limited number of possible reactions are possible.

Question: An electrical current runs between two electrodes in molten sodium chloride.

Answer: NaCl ? Na + Cl2

f. Is a solid metal placed into an acid?

The metal is oxidized, and hydrogen gas is formed. Remember, water can also be an acid; check to be sure the dissociation of water on the reduction potential chart shows a higher

tendency to be reduced, such as the case with calcium in water.

Question: Magnesium metal is placed into a weak solution of hydrochloric acid.

Answer: Mg + H+ ? Mg++ + H2

5. Be sure that reactions in solution are written as net ionic reactions.

Also, since you will not be given any points for balancing the reaction or writing the phases down, don’t waste time doing either of these.

Summary of Steps

1. Write down the formulas for as many reactants as you can.

2. Is it a precipitation reaction?

3. Is it an acid-base reaction?

4. Is it an oxidation-reduction reaction?

5. Be sure that reactions in solution are written as net ionic reactions.

What to Study

• Use the Periodic Chart to discern common oxidation states in Chapter 2

• Memorize the strong acids and strong bases in Chapter 7.

• Memorize the solubility rules in Chapter 7.

• Know how to write net ionic reactions as outlined in Chapter 8.

• Understand the different types of reactions from Chapter 8.

• Know how to solve stoichiometry problems from Chapter 9.

• Know the names of common ions in Chapter 13.

• Know the names of organic compounds identified in Chapter 13.

• Know traits about each reaction you learn. For example, what does the product look like? What is oxidized and reduced? Which species acts as the acid, or the base? Is it a Lewis acid-base reaction, or a Bronsted reaction? Is the resulting solution acidic or

basic? What oxidation numbers are involved?


Format, Scoring, and Strategy

• Questions 5 and 6 are multiple-part questions that address some general aspect of chemistry.

• Each question is worth 15% of the total score on the Free-Response section of the AP Chemistry Exam.

• These question may involve the reading of graphs, laboratory data, decisions about what to measure, and the interpretation of data, if a previous problem in this part did not involve laboratory observations.


For registration bulletins or other information about the AP Chemistry exam, contact:

AP Services

Educational Testing Service

P.O. Box 6671

Princeton, NJ 08541-6671

Phone: (609) 771-7300 or (888) 225-5427

E-mail: apexams@ets.org

Website: www.collegeboard.com

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