AP Chemistry II
Summer Work 2014
Contact information: bizraganem@warwickschools.org
General Information:
Over the summer it is important to review concepts from chemistry I. The AP Chemistry Exam in May
covers topics from Chemistry I and Chemistry II. The AP Chemistry Exam is scheduled for Monday, May 4th
2015.
Directions:
1.)
2.)
3.)
4.)
Read the Chemistry I Review Outline
Read the assigned chapters in the book and take good notes
a. Since this is a college level course you should think of your note-taking as good practice for how
you will need to study in college.
Complete the practice problems. Show all work. Always use units. Use sig fig rules for rounding.
Answers key is in my shared document on First Class.
Complete the online practice problems. You will receive directions via email in order to log on the site.
Take advantage of the “Show me” and “Guided Solution” options if you are stuck on a problem.
Chemistry I Review Outline:
I. Matter & Measurement
a) Definitions & Examples for the following chemistry terms:
1. Intensive and extensive physical properties, Chemical vs physical changes, Chemical
properties, Mixtures: Heterogeneous and homogeneous mixtures, Pure Substances:
Elements and compounds, Law of Conservation of Mass, Units of Measurement: The SI
System, Metric System, Accuracy and Precision, Sig Fig Rules of Rounding, Dimensional
Analysis, Scientific Notation, Temperature Scales (oF, oC, K)
b) Compare and contrast particles in a solid, liquid, and gas
c) Be able to answer questions about heating curves
II.
Atomic Structure and the Periodic Table
a) For any element on the Periodic Table you should be able to:
1. Determine the number of protons, neutrons and electrons in one atom of the element
i. Ex. One atom of aluminum has 13 protons, 13 electrons, and 14 neutrons
2. Determine the charge of the ion for that element using the octet rule
i. Ex. Aluminum forms Al+3 as the aluminum cation
3. Describe how one atom of the element forms an ion and why the atom forms an ion
i. Ex. One aluminum atom gets an octet by giving away its three valence electrons.
4. Write an electron configuration, electron dot notation, Noble gas configuration, and orbital
notation
i. Ex. Electron configuration: Al: 1s22s22p63s23p1
.
ii. Ex. Electron dot notation: Al:
iii. Ex. Noble gas configuration: [Ne] 3s23p1
iv. Ex. Orbital notation:
5. Know the name of the group to which the element belongs
i. Ex. Fluorine is a halogen
6. Compare the electronegativity, ionization energy, and atomic radius of one atom of the
element compared to other elements on the Periodic Table and be able to define the terms
i. Ex. Compare Fluorine to aluminum: Fluorine has the highest electronegativity of any
element but is smaller than aluminum. Fluorine has a higher ionization energy
7. Determine if the element is a metal, nonmetal, or a metalloid/semimetal
i. Ex. Fluorine is a nonmetal, Aluminum is a nonmetal
8. Know the state of the element at room temperature
(Assume room temperature is 250C)
i. Ex. Fluorine is a gas at room temperature
9. Discuss the three types of bonds: metallic, ionic, covalent
i. Discuss polar covalent vs nonpolar covalent
ii. Discuss the types of atoms that form certain types of
bonds
10. Predict the type of bond that will form between atoms of two elements.
i. Use electronegativity differences to discuss the polarity of a bond
1. Ex. Aluminum will form an ionic bond with chlorine.
2. Ex. Hydrogen and oxygen will form a polar covalent bond.
3. Ex. In diatomic oxygen, the oxygen atoms have a nonpolar
covalent bond
11. Compare and contrast the terms atom, isotope, and ion
12. Be able to solve problems and discuss the mathematical relationship between wavelength,
frequency, and energy
13. Be able to discuss the terms absorption and emission, how spectral lines are created, and how
the Bohr Model of the atom describes the production of spectral lines.
III. Atomic History
a) Be able to describe the experiments and major contributions to atomic structure for the following
scientists:
1. Rutherford: Gold Foil Experiment, discovered
the nucleus
2. Thomson: Cathode Rays, discovered the
electron
3. Bohr: Spectral Lines, electrons orbit the
nucleus.
4. Schrodinger: Orbitals, probability of finding
an electron in an orbital (s, p, d, f)
IV. Compounds & Bonding
a) Two types of compounds: ionic and covalent (molecular)
b) Three types of bonds: ionic, covalent, and
metallic
1. Be able to compare and contrast the
three types of bonds and describe
how the electrons are involved in the
bond and how the electrostatic force
plays a part in the bond
i. Compare and contrast polar
covalent and nonpolar covalent
bonds
V.
Chemical Reactions
a) Evidence of a chemical reaction (Examples: change in pH, change in temperature, color change,
formation of a precipitate, formation of gas bubbles, etc.)
b) Five types of reactions: synthesis, decomposition, single replacement, double replacement,
combustion
c) Balancing equations
d) Predicting products of reactions using the activity series or the solubility rules
1. MEMORIZE: Nitrates (NO3- ), and salts containing H+, Na+ or K+ are always aqueous
e) Using information (temperature change, energy diagrams) to determine if a reaction is exothermic or
endothermic
f) REDOX reactions: be able to identify the species that is oxidized or reduced in
a REDOX reaction
g) Identify oxidation numbers in a chemical compound
VI. Moles
a)
b)
c)
d)
e)
f)
Use dimensional analysis to convert from grams to moles to particles
Use the correct unit for particles: atoms, molecules, formula units
Be able to calculate an empirical or molecular formula
Name and write formulas for hydrates (ex. CuSO4 . 5H20 copper II sulfate pentahydrate)
Use lab data to determine the empirical formula of a hydrated ionic compound
Be able to solve molarity problems ie determine the number of grams of a salt needed to make a
solution of a specific concentration
g) Use the dilution equation to solve problems: M1V1 = M2V2
VII. Thermodynamics:
a) Be familiar with the terms enthalpy (H), entropy (S) and free energy (G)
b) Understand that all chemical reactions involve a change in energy
c) Bond breaking is always endothermic (requires energy) whereas bond forming is always exothermic
(releases energy)
d) Be able to use q=mcT to solve calorimetry problems
VIII. Stoichiometry
a) Use a balanced equation to convert between masses, moles, and particles of two different substances
b) Limiting Reactants
1. Determine a limiting reactant
2. Determine the amount of excess reactant
c) Yields
1. Calculate a theoretical yield, percent yield, or actual yield
IX. Gases
a) Identify Boyle’s and Charles’ Law
b) Discuss the relationship between pressure, temperature, volume, and moles of a gas when a constant
is identified
c) Use the Ideal Gas Law (PV=nRT) to solve problems
d) Use molar volume (1 mole = 22.4 L at STP) to solve problems
e) Use the Kinetic Molecular Theory and the properties of gases to explain gas behavior
f) Solve problems using gas laws and stoichiometry
g) Dalton’s Law of Partial Pressures
h) Graham’s Law of Effusion
X.
Lab Work
a) Lab work is an integral part of AP Chemistry II and it is important that you remember the lab skills
that you learned in chemistry I so you can be successful in lab during AP Chemistry II.
1. You should be familiar with each of the following pieces of lab equipment: balance,
graduated cylinder, beaker, Bunsen Burner, spectroscope, metric ruler, funnel.
2. You should keep your labs from chemistry I since the AP Chemistry Exam includes
questions from lab experiments that you have performed in Chemistry I and will perform in
Chemistry II.
Book Chapters/Reading:
Chemistry I Chapters 1-5, 7-9, 10, 15.1-15.3
Skip these Sections: 3.4, 4.7, 4.8, 5.8, 7.6, 7.7, 9.3, 9.6-9.9
- Take good notes
- Do not skip diagrams and pictures! They are in the book to help
explain concepts.
Name:___________________________________________________________
AP Chemistry II
Summer Work 2014: Practice Problems
1.) Cyclopentane is a highly flammable, colorless liquid used in refrigerators and freezers and in the
manufacture of rubber adhesives. Selected properties of cyclopentane are given in the table below. Use
the information in the table to answer questions a-d.
PROPERTIES OF CYCLOPENTANE
Structural Formula
Melting Point
Boiling Point
Density
-94oC
49oC
0.751 g/cm3
a. A refrigerator contains 3.54 grams of cyclopentane. How many molecules of cyclopentane are in
the refrigerator?
b. A sample of cyclopentane is kept in a sealed container with a constant temperature of -210oC.
i. Identify the state of cyclopentane. Justify your answer.
ii. Describe the motion of the particles in the sample.
c. If 0.50 moles of cyclopentane are used a chemical reaction to manufacture rubber, what is the
volume of cyclopentane used?
d. Cycloheptane is another molecule containing carbon and hydrogen atoms. The formula for
cycloheptane is C7H14. If you were given samples of cycloheptane and cyclopentane each
containing exactly 1.5 moles, which sample would have the larger mass? Justify your answer.
2.) Draw a Bohr Model for Oxygen.
a. Draw a new Bohr Model for oxygen and show an electron transition from n=4 to n=2
b. Draw a new Bohr Model for oxygen and show an electron transition from n=3 to n=2
c. Which electron transition produces photons with a higher energy, the electron transitioning from
n=4 to n=2 or the electron transitioning from n=3 to n=2? Justify your answer.
d. The n=4 to n=2 transition produces a photons with a wavelength of 567 nm.
i. Determine the frequency of the photon.
ii. Determine the energy of the photon.
iii. Determine the type of radiation.
e. Describe why oxygen has a higher electronegativity than nitrogen using their atomic structures to
explain the difference. (It is not enough to say that oxygen is farther to the right on the Periodic
Table)
f. Describe why oxygen has a higher electronegativity than sulfur using their atomic structures to
explain the difference. (It is not enough to say that oxygen is higher up on the Periodic Table.)
3.) Write a balanced chemical equation for: Solid magnesium metal is burned in the presence of oxygen gas
to produce a solid, white, ionic compound. (Remember states!)
a. If 0.02465 grams of magnesium metal are burned, determine the theoretical yield of the white
ionic compound.
b. Determine the percent yield if 0.03988 grams of ionic compound are produced.
c. What is the limiting reactant in the experiment? Justify your answer.
4.) Write balanced chemical equations for the following if a reaction takes place. If no reaction takes place
explain why not.
a. Aqueous solutions of lead II nitrate and sodium sulfate are combined.
b. Calcium metal is added to a beaker containing a 150 mL of a 1.0 M solution of aluminum nitrate.
The beaker is heated and the solution is stirred continuously for 15 minutes.
c. Aqueous solutions of sodium hydroxide and iron III sulfate are mixed.
d. Zinc metal is mixed with a solution of sulfuric acid (H2SO4).
e. Tin metal is mixed with 55.0 mL of a 1.0 M solution of zinc sulfate.
5.) Write a balanced chemical equation for: Nitrogen monoxide gas reacts in the atmosphere with oxygen to
form nitrogen dioxide gas.
a.
If 3.50 grams of nitrogen monoxide gas is placed into a sealed container with exactly 0.510
grams of oxygen gas, what mass of nitrogen dioxide gas will be formed?
b. A sample of nitrogen monoxide gas is placed into a rigid 3.0 L vessel at a pressure of 0.55 atm
and a temperature of 23oC. Determine the new pressure of the gas if the temperature is doubled.
c. Determine how many moles of oxygen gas are in a 250 mL sample with a pressure of 1.10 atm at
a temperature of 18.5oC.
d. What is the volume of a sample of nitrogen dioxide gas with a mass of 3.45 g at STP?
6.) Determine the empirical formula and molecular formula of an organic compound containing 76.5%
carbon, 12.2% hydrogen, and 11.3% oxygen with a molar mass of 705 g/mol.
7.) Discuss the atomic theories of the following scientists using labeled diagrams of each scientists’ atomic
model. (Dalton, Thomson, Rutherford, Bohr).
8.) Write an electron configuration for Phosphorus.
a. Write a noble gas configuration for Phosphorus.
b. How many electrons are in the 3 p orbital?
c. How many protons, neutrons, and electrons are in one atom of phosphorus?
d. How many protons, neutrons, and electrons are in a phosphide ion?
e. What is the name and formula of the ionic compound formed when phosphorus reacts with
calcium?