AP Chemistry Review Preparing for the AP Chemistry Exam

Preparing to load PDF file. please wait...

0 of 0
100%
AP Chemistry Review Preparing for the AP Chemistry Exam

Transcript Of AP Chemistry Review Preparing for the AP Chemistry Exam

AP Chemistry Review Preparing for the
AP Chemistry Exam
1

Table of Contents

Title

Pg

Top 25 Things to Know before you take the AP Chemistry Exam

3

About the AP Chemistry Exam

5

AP Chemistry Study Guide

8

Types of Reactions

19

AP Subject Review

24

Multiple Choice Practice test

40

Practice Free Response Questions

58

Practice Free Response Scoring Guide

63

2

Top 25 Things to know before you take the AP chemistry Exam.
Assume the person grading your exam is an idiot. Make it clear for them to understand your process and grade your exam with ease. It is not acceptable to lose points because of not showing work/units or messy handwriting. If you need to cross something out do it like this “The reaction is spontaneous because.” . No need to scribble and make a mess.
Focus on your weakest areas; it is doubtful you can do/know everything. The AP Chemistry Exam is designed so that it is impossible to know absolutely everything on it (in case you haven’t noticed).
I might as well place the biggest two at the start – You need to review your incorrect MC from the Practice Exam(s) and Princeton exam(s) and understand the concepts. It can really increase your MC score.
Know the 6 strong acids HCl, HI, HBr, H2SO4, HClO4, HNO3 and the one weak by formula acetic acid CH3COOH, everything else is weak. Remember that strong acids/bases don’t make buffers!!! You should be 100% confident what ionizes and what doesn’t
Know the strong bases: Group 1 hydroxides, Ba(OH)2, Sr(OH)2, Ca(OH)2
Know how to determine which molecule has the largest dipole moment (difference in electronegativity).
Know hybridization on a carbon atom (Remember Double Bonds and Triple bonds don’t count) Lone Pairs do count!!! Single bonds are sigma, double and triple are pi-bonds.
Solubility Rules are a must. If you still don’t know them get cracking!!!
Know some basic geometry’s of molecules. Do not expect complicated ones like See-Saw etc. Stick with the basics. If any really hard ones are on there you will prob. need to guess.
Know the bond angles on a Bent Geometry, Trigonal Planer, Trigonal Pyrimidal, and Tetrahedral. Also know why bond angles shrink as lone pairs are added (b/c if increased repulsion amongst the electrons causing the bond angles to squeeze)
It would be a safe bet to assume that when a metal by itself it placed in acid you will get H2 gas and some aqueous salt and the negative ion is the spectator ion. Cu does not react in HCl, but in AP chemistry it is safe to assume that there is always a reaction. Also groups 1 and group 2 metal by themselves placed in water always give you H2 gas and a base (which usually ionizes).
Know the carbonate reactions! You always get CO2 on the other side….
Be ready for a complex ion. You should be confident in determining the answer. Remember to “cheat” on the reaction and determine the product first (using the double it rule) and work backwards to the get the reactants.
Know the units on the rate constant k for kinetics.
Know the kinetics graphs.
Lighter atoms (molar mass) are faster than heavier atoms (molar mass) at the same temperature. Also lighter atoms “effuse” (leak) out of small hole in a container faster than heavier ones.
3

Know the periodic table trends. It would useful to know the E.N values for F =4.0, O=3.5, N=3.0 C=2.5, H=2.1, and what type of bonds occur as a result.

Know your intermolecular forces, how to I.D. than and what they mean in terms of boiling points. Know about covalent network bonds and what they mean!

Know what the signs on Delta H, S and G mean and be able to explain each.

Know how to use Ksp

Be sure to look over your Labs, there will a question on the part two this topic.

The Chart. You will need to know it. No doubt. Be Ready

G

Ecell

Keq

+

--‐

<1

--‐

+

>1

0

0

1

OIL RIG and REDCAT. The only time reduction does not occur at the cathode is when it is an electrolysis reaction in which energy is added. Usually you are given the reduction or the oxidation on this type of diagram/question.
Decide what else to study from the exam memorization guide.
25)The test will be hard, but don’t get frustrated, just keep plugging through, and don’t give up. You might surprise yourself. You are prepared. Good luck and Congratulations on finishing the course!!!

4

About the AP Chemistry Exam
Section I: Multiple-Choice Section: 90 minutes (50% of your grade) Section I consists of 60 multiple-choice questions, either as discrete questions or question sets, that represent the knowledge and science practices outlined in the AP Chemistry Curriculum Framework, which students should understand and be able to apply. Question sets are a new type of question: They provide a stimulus or a set of data and a series of related questions. Calculators are not permitted on the Multiple Choice section.
 Examine each question for a maximum of thirty seconds (on the average, some will take less time allowing more time for others).
 Quickly determine the subject of the question.  By the end of the thirty seconds either:  Mark the correct answer.  Mark a “Y” next to the questions that you know how to work but need more time.  Mark a “N” next to the questions that you don’t have any idea how to work.
Force yourself to move through twenty questions each ten minutes and the full seventy-five questions in forty minutes. Now make a second pass concentrating on the “Y” questions only. Do not spend any time on the “N” questions. If you don’t know the correct answer see if some key piece of knowledge will allow you eliminate two or three of the choices. Complete this pass in forty minutes. Now make your third pass. Focus only on the “N” questions. Attempt to eliminate at least two choices. If you can, then make an intelligent guess. If not, guess. You are not penalized for the wrong answer. Any correct choices on this pass are bonus points. You have only ten minutes, so make it count! Before time expires, count the number that you have answered. You should answer at least sixty (60) questions.
Section II: Free-Response Section 90 minutes (50% of your grade) Section II contains two types of free-response questions (short and long), and each student will have a total of 90 minutes to complete all of the questions. This section also contains questions pertaining to experimental design, analysis of authentic lab data and observations to identify patterns or explain phenomena, creating or analyzing atomic and molecular views to explain observations, articulating and then translating between representations, and following a logical/analytical pathway to solve a problem. Beginning with the May 2014 administration of the AP Chemistry Exam, multiple-choice questions will contain four answer options, rather than five. This change will save students valuable time without altering the rigor of the exam in any way. A student's total score on the multiple choice section is based on the number of questions answered correctly. Points are not deducted for incorrect answers or unanswered questions.
What to Bring to the Exam
 2-3 Sharpened #2 pencils  Acceptable calculator  A positive attitude  Get a good night’s rest and eat breakfast that morning!
5

AP Chemistry Exam Format Section I

Question Type

Number of Questions

Multiple Choice

60

Section II

Long Free Response 3

Short Free Response 4

Timing 90 minutes
90 minutes

Calculators are allowed on the free-response section for the first 55 minutes. During that time, students will work on three required problems. For the last 40 minutes, calculators must be put away as students work on the remaining free-response questions.

Section II Free Response

 Questions require you to apply and explain chemical concepts and solve multiple step problems.  You do not have to answer in essay form and may save time using one of the following methods: bullet format,
chart format or outline format.  Write your answers in the space provided and number your answer clearly.  There is a slight penalty for incorrect sig figs.  Stating a Principle, Law, Theory or stating the name of the Principle law or theory is not an explanation or
justification for an answer. Stating a trend is also not an explanation. State and apply the Principle , Law or Theory to the specific situation in the questions to explain your answer and use the reason for trends as explanations for trends not just stating the trend.  Answer everything, no matter what  If you don’t know the answer, substitute in a number to complete the remaining parts of the question. You will be given credit for following the correct procedure even if the answer may be wrong

Part A (Question 1) Equilibrium

Read all of the question before doing any work. Items later in the problem may provide keys to earlier sections.

Part A is always equilibrium. Determine which type (Gaseous equilibrium, acid/base, buffer, or precipitation). Look for

key words and clues.

Acid/Base: Look for the words acid or base, Ka or Kb, [H+], [OH-], or [H3O+]. Any of these indicate an

acid/base problem.

Buffer: Look for the word buffer. Also, check for a weak acid and its conjugate base.

Precipitation: Look for Ksp or the word solubility.

Gas Equilibrium: Look for (g) on most of the reactants and products.

After determining the type of reaction, write a reaction if one is not provided. Use the general forms given below:

Acid HA +

H2O -----> H3O+

or HA ----- H+ + A>

6

Base A-

+

H2O -----> OH-

Precipitation

HA + MA(s)

--- +

A-(aq)

Write an equilibrium constan-t-e>xpressMio+n.(aLqea)ve out solids and liquids.

Solve the problem. THINK! Put in all of the given quantities in the equilibrium constant expression and solve for the

unknown allowing the units to direct the problem.

Part A (Questions 2 and 3) Read both problems all the way through before doing any work. Determine which type of problem each is. Select the problem you know the most about and solve it. Remember that if you cannot solve an earlier part you may still get some credit for a later section by showing how you could use the earlier answer in succeeding parts of the problem. Question 2 & 3 generally cover the following:
Lab procedure Kinetics Electrochemistry Stoichiometry Thermochemistry

Part B (Question 4) Consists of three reactions and usually a lab question about the reaction. Write the reactant in symbol form for all reactions showing each reactant in net ionic form as follows: Strong Acids, bases, and soluble salts written as ions. Weak acids, bases, and insoluble salts written as molecules.
Classify the reactions as:  Acid/Base - Look for H or OH or salts which could act as a weak acid or base.  Precipitation - Look for insoluble salts which could form as products according to the solubility rules.  Redox-Ifitisnotacid/baseorprecipitation,itisprobablyoxidation-reduction. Checkfor elements which could change oxidation states. Pay particular attention to the common oxidizing agents (NO3-, MnO4-, Cr2O72-, H2O2) and reducing agents (Cl-, Br-, I- and elemental metals).  Other - Anything else which doesn’t fit above (usually either organic or complexation).
Remember you score one point for getting the reactants in the correct form and two points for each product. At least get all of the reactants correct and possibly two or three products.
Part B (Questions 5-7) Typically includes the following topics:
Bonding/intermolecular forces/hybridization Electrochemistry Lewis dot structures/Periodic Trends Be as specific as possible in your answer. Look for clues in the question as to what is really important. Answer the question. State exactly what you are asked not what you would like to answer. Do not simply restate the question. Remember that you will be getting partial credit. Answer any part about which you have any knowledge.

7

AP Chemistry Study Guide
2 • Atoms and Elements
The Development of the Atomic Theory:  Define the three theories that Dalton explained in terms of atoms:
o Law of Conservation of Matter o Law of Definite/Constant Proportions o Law of Multiple Proportions  Give examples and solve calculation problems related to each of the three theories.  Sketch a cathode ray tube as demonstrated in class and state how J.J. Thomson’s experiments led to the idea that atoms have positive and negative parts, the negative parts are all the same, and the negative parts (called electrons) have a certain charge/mass ratio.  Define cathode rays.  State the factors that determine how much a moving charged particle will be deflected by an electric or magnetic field.  Explain Millikan’s oil drop experiment & how it added to the atomic theory.  Sketch the set-up used by Ernest Rutherford (the gold-foil experiment), show what he observed, and explain how these observations led to the idea that most of the mass of the atom is concentrated into a tiny, amazingly massive, positively-charged nucleus. Parts of the Atom:  State the three particles that make up an atom, their symbol, their charge, their mass, and their location.  State the number of protons, neutrons, and electrons in any atom or ion.  Explain that isotopes are two atoms with the same atomic number (number of protons) but different mass numbers (number of nucleons— protons + neutrons).
220  Represent the nucleus with isotopic notation, such as: 86 Rn  Recognize when two nuclei are isotopes of each other.
Molar Mass Calculations:
 Calculate the isotopic mass of an atom given the resting mass of protons and neutrons.  Explain that a mole of any element is actually made up of various isotopes in a constant percentage abundance.  Calculate the average atomic mass of an element using the percent abundance and mass of each isotope.  Calculate the percent abundance of isotopes given the average atomic mass and isotopic masses of an element. The Families of the Periodic Table:  List the common families of the periodic table and recognize to which family any element belongs.  Recognize metals, non-metals, and metalloids (semi-metals) on the periodic table.  State and define the terms conductivity, malleability, and ductility.  State some element facts such as which elements are too radioactive to exist, which is the largest non-radioactive element, which element has
the greatest density, and which element has the highest melting point.  Explain how Dmitri Mendeleev put together the periodic table and why we give him credit for the table even though others were working along
the same lines.  List the three elements that Mendeleev predicted and where they are located on the periodic table. Nuclear Chemistry:  State that Henri Becquerel discovered radioactivity and Marie Curie studied it.  List the three “Becquerel rays” (alpha, beta, and gamma) and state why alpha particles were the perfect tool for Ernest Rutherford to study the
structure of atoms.  State that the alpha particle is the same as a helium nucleus, a beta particle is a high-speed electron, and a gamma ray is a high-energy form of
light.
3  Chemical Formulas
Formulas
 Look at a formula and state how many elements and atoms are in that compound.  Calculate the molecular mass or molar mass of any compound.  State that the mass of a molecule is measured in amu’s and the mass of a mole is measured in grams.  Give examples of empirical formulas, molecular formulas, and structural formulas.  Identify a formula as empirical, molecular, or structural.
8

Ionic Compounds
 State whether a compound is an ionic compound or a nonmetal compound.  Write the formula of an ionic compound given the two ions or its name. Know when to use parentheses.  Name an ionic compound given the formula.  Determine the charge on an ion from information in an ionic formula.
Nonmetal Compounds (aka Molecular Compounds)
 Write the formula of a binary nonmetal compound (molecular compound) given its name.  Name a binary nonmetal compound (molecular compound) given its formula.
Percent Composition
 Calculate the percent composition (by mass) for any compound.  Calculate the empirical formula from percent composition data.  Determine the molecular formula of a compound given its empirical formula and molar mass.
Hydrates
 Give examples of hydrates and anhydrous compounds.  Calculate the formula of a hydrate from dehydration data.
The Mole
 State the significance of the mole.  State the three mole facts for any substance (molar volume, molar mass, Avogadro’s number)
o 1 mole = 22.4 Liters @ STP (gases only) o 1 mole = 6.02 x 1023 particles o (particles = molecules or atoms) o 1 mole = gram molecular mass of chemical  Use dimensional analysis to convert between moles, mass, volume, and number of particles for a chemical.  Use density as a conversion factor in mole problems.  Use gas density to calculate molar mass.
4  Chemical Equations and Stoichiometry
Chemical Equations  Give examples of products and reactants in a chemical equation.  State that Antoine Lavoisier introduced the law of conservation of matter. Combustion  State that combustion is another name for burning.  Write an equation for a combustion reaction given only the fuel that is burned.  Correctly label substances in an equation as solid (s) , gas (g), liquid (l), or aqueous (aq) Balancing Equations  Balance equations by adding coefficients.  Recognize when an equation is balanced.  State that the formulas of reactants and products should not be changed in order to balance equations. Stoichiometry Problems  Use the stoichiometric factor ( of the problem) to convert from moles of one substance to moles of a different substance.
(i.e. In the equation: N2 + 3H2  2NH3, 3 mol H2  2 mol NH3)  Convert between the quantities of mass, volume, molecules and moles using dimensional analysis
(i.e. use 1 mol = 22.4 L, 1 mol = 6.02 x 1023 molecules, and 1 mol = gram molecular mass)  Show the units of molar mass as grams/mol or g·mol-1. Limiting Reactant Problems  Recognize that a problem with two “given values” is a limiting reactant problem.  Determine the limiting reactant and excess reactant in a problem.  Solve problems involving Limiting Reactants  Calculate how much excess chemical is left over after a reaction. Percent Yield Problems  Use stoichiometry to calculate the theoretical yield (mass of a product) in a problem.  State that actual yields are usually given in a problem.  Use the theoretical yield and actual yield to calculate the percent yield. Chemical Analysis Problems
9

 Calculate the mass of each element in a given compound given data such as the masses of CO2 and H2O formed in a combustion reaction.
 Use mass and mole information to calculate the empirical formula of an unknown substance.  Use percent composition to equalize mass and mole information derived from different samples.
5  Reactions in Aqueous Solution____________________________________________________________
Properties of Aqueous Solutions  Define solute, solvent, and solution. Give examples.  Define electrodes. Give operational and theoretical definitions of electrolytes.  Know that soluble ionic compounds and strong acids are strong electrolytes. Ionic compounds of low solubility [e.g. Mg(OH)2]
and weak acids/bases are weak electrolytes.  Know that molecular compounds (except acids) are non-electrolytes.  Know that alcohols (e.g. CH3OH )are not ionic hydroxides. Bases are usually metallic hydroxides.  Know the solubility rules. State whether an ionic compound is soluble in water. Precipitation Reactions  Know that ppt reactions are double replacement reactions that produce an insoluble product.  Given two ionic compounds in solution, correctly determine the products. (Know your ions).  Determine which product(s) is/are precipitates. Use (aq) and (s) symbols correctly.  Correctly write the ions in a soluble ionic compound. [e.g. CaCl2(aq) becomes Ca2+ + 2Cl ]  Identify spectator ions.  Write molecular, detailed ionic, and net ionic equations for a ppt reaction. Acids and Bases  Give operational (cabbage juice) and theoretical (ions) definitions of acids and bases.  Know that acids increase the H+ ion concentration in an aqueous solution. (Theoretical definition)  Memorize the 8 strong acids.  Know that acids are molecular compounds that form ions when in aqueous solution.  Be able to name acids according to their anion.
[ide  hydro__ic acid; ate  __ic acid; ite  __ous acid; sulfur: add “ur”; phosphorus: add “or”]  Know that bases increase the OH ion concentration in an aqueous solution. (Theoretical definition)  Memorize the soluble hydroxides (except NH4OH) that are the strong bases.  Understand that ammonia(aq), NH3 + H2O NH4+ + OH forms a weak basic solution.  Know that metal oxides form bases [CaO + H2O  Ca(OH)2] while nonmetal oxides form acids [CO2 + H2O  H2CO3]  Know that acids react with bases to form H2O and a salt. (Neutralization)  Write equations for acid-base reactions including NH3 (example on page 199) as the base.  Know that strong acids and strong bases are written as ions in the ionic equations. Gas Forming Reactions  Recognize the six products that turn into gases. Memorize the gases formed. Organizing Reactions in Aqueous Solution  Double Replacement reactions (text calls them exchange reactions) (Fred-Wilma/Barney-Betty reactions) also have the old
fashioned name: metathesis reactions.  Know the three examples of double replacement reactions and the “driving force” for each.
Precipitate reactions form an insoluble product. Acid-Base reactions form water (a very weak electrolyte therefore, a very stable product). Gas-forming reactions form a gas.  Know that a driving force is something that keeps the new combinations of ions from reforming the old combinations of ions.  Oxidation-Reduction is a fourth type of reaction driven by the transfer of electrons. Oxidation-Reduction Reactions  Know that an important type of reaction gets its name from atoms that combine with oxygen. During the refining of iron, carbon monoxide combines with oxygen (from the iron ore), CO  CO2 and is oxidized. Large masses of iron ore (Fe2O3) are reduced to a smaller amount of iron metal.  Understand that since CO helps the iron ore to be reduced, CO is called the reducing agent. Since Fe2O3 causes the C to be oxidized, iron ore is called the oxidizing agent. What ever is oxidized acts as the reducing agent. What ever is reduced acts as the oxidizing agent.
10
MassAnswerQuestionsCompoundFormula