Download AP Chemistry Summer Assignment Summer 2015 Ms. Osquist

AP Chemistry Summer Assignment
Summer 2015
Ms. Osquist
Severn School
Purpose: The purpose of this summer assignment is to help you prepare for a fast-paced
course in AP Chemistry. We have a lot of material to cover in a short amount of time, so
you will be expected to have certain background knowledge at the start of the course.
You will be expected to be familiar with the content covered in the first four chapters of
our AP chemistry textbook. You do NOT need to purchase the textbook to complete this
assignment however. Electronic resources are provided to help you review the content of
the summer assignment. The majority of the summer assignment will be review from first
year chemistry.You will be assessed on the content of the assignment in the early
weeks of the school year.
Objectives (what you should know):
• Distinguish between accuracy and precision.
• Record a measurement with the proper number of significant figures.
• Maintain the correct number of significant figures when performing
computations.
• Convert between metric prefixes and mass and volume using density.
• Distinguish between elements, compounds, molecules, mixtures, and solutions.
• Determine the number of protons, electrons, or neutrons in an atom or ion.
• Explain how the average atomic mass is calculated from mass numbers and
relative abundance of isotopes.
• Determine if an element is a metal or nonmetal and the charge of its most
common ion (for A-group elements) using the periodic table.
• Write the formula of an ionic compound given its name (including polyatomic
ions).
• Explain how a cation or anion would form from a neutral atom.
• Identify the most common state of an element (metal, nonmetal, diatomic
molecule).
• Name an ionic or molecular compound.
• Covert between grams, moles, and liters using molar mass, stoichioimetric
relationships, or concentration (molarity).
• Explain why the average atomic mass can be used as estimation of the mass of
each individual atom in a large sample of matter.
• Determine the limiting reactant in a chemical reaction given the masses or
volumes and concentrations of reactants.
• Determine the empirical or molecular formula of a compound given the percent
composition of the elements and molar mass.
• Determine an empirical formula of a compound using stoichiometric relationships
and lab data of mass of reactants or products.
• Use mass percent data to verify if a sample is a pure substance (single compound)
or to determine if two samples contain identical substances.
• Create a model (drawing) of a substance or chemical reaction that accurately
depicts the particles involved.
• Classify a chemical reaction as a combination, decomposition, combustion,
double replacement, single replacement, or acid-base neutralization.
• Predict the products, including states, of a precipitation reaction.
• Identify the substance oxidized and reduced in a redox reaction.
• Determine the oxidation state of an element in a compound.
Resources:
The UC Davis ChemWiki is written like a textbook with information and sample
problems. The following link will take to you the ChemWiki aligned to your textbook.
http://chemwiki.ucdavis.edu/Textbook_Maps/General_Chemistry_Textbook_Maps/Map
%3A_Zumdahl's_%22Chemistry%22
The Chem Team is another website where you can find useful videos, explanations, and
practice problems organized by topic. The primary downside is content is not directly
aligned to your textbook. However, you will find more sample problems and
explanations, particularly if your stoichiometry is a bit rusty.
http://www.chemteam.info/ChemTeamIndex.html
Instructions:
The problems below are organized to correspond with the chapters of your textbook.
Before completing the problems, you should either read your book (if you have
purchased it) or read the notes for the chapter on the ChemWiki site linked above.
Complete the problems below on a separate piece of paper. Your work will be collected
and graded for accuracy, so please take time with this assignment and show your work
when necessary. You do not need to get every problem correct to earn a “good” grade,
but you do need to take the assignment seriously. I am going to use your work to
determine the amount of review needed during the early weeks of the school year.
Chapter 1. Classifications and Properties of Matter & Measurement
1.1 Classify each of the following as a pure substance or a mixture. If a mixture, indicate whether
it is homogeneous or heterogeneous. (a) chocolate chip cookie, (b) seawater, (c) magnesium, (d)
gasoline.
1.2 In the process of attempting to characterize a substance, a chemist makes the following
observations: the substance is a silvery white, lustrous metal. It melts at 649°C and boils at
1105°C. Its density at 20°C is 1.738 g/cm3. The substance burns in air, producing an intense
white light. It reacts with chlorine to give a brittle white solid. The substance can be pounded
into thin sheets or drawn into wires. It is a good conductor of electricity. Which of these
characteristics are physical properties, and which are chemical properties?
1.3 Label each of the following as either a physical process or a chemical process: (a) corrosion
of aluminum metal, (b) melting of ice, (c) pulverizing of aspirin, (d) digesting a candy bar, (e)
explosion of nitroglycerin.
1.4 What exponential notation do the following abbreviations represent in the metric system: (a)
d, (b) c, (c) f, (d) µ (e) M, (f) k, (g) n, (h) m, (i) p?
1.5 Make the following conversions:
(a) 233°C to K, (b) 100 K to °C, (c) 0°C to K, (d) 1088 °C to K, (e) 0 K to °C.
1.6 (a) A sample of carbon tetrachloride, a liquid once used in dry cleaning, has a mass of 39.73 g
and a volume of 25.0 mL at 25°C. What is its density at this temperature? Will carbon
tetrachloride float on water? (b) The density of platinum is 21.45 g/cm3 at 20°C. Calculate the
mass of 75.00 cm3 of platinum at this temperature. (c) The density of magnesium is 1.738 g/cm3
at 20°C. What is the volume of 87.50 g of this metal at this temperature?
1.7 To identify a liquid substance, a student determined its density. Using a graduated cylinder,
she measured out a 45 mL sample of the substance. She then measured the mass of the sample,
finding that its mass was 38.5 g. She knew the substance had to be either isopropyl alcohol
(density 0.785 g/mL) or toluene (density 0.866 g/mL). What are the calculated density and the
probable identity of the substance? (b) An experiment requires 45.0 g of ethylene glycol, a liquid
whose density is 1.114 g/mL. Rather than weigh the sample on a balance, a chemist chooses to
measure the liquid using a graduated cylinder. What volume of liquid should he use? (c) A cubic
piece of metal measures 5.00 cm on each edge. If the metal is nickel, whose density is 8.90 g/cm3,
what is the mass of the cube?
1.8 Indicate which of the following are exact numbers: (a) the mass of a paper clip, (b) the surface
area of a dime, (c) the number of inches in a mile, (d) the number of ounces in a pound, (e) the
number of microseconds in a week, (f) the number of pages in this book.
1.9 What is the number of significant figures in each of the following measured quantities? (a)
358 kg, (b) 0.054 s, (c) 6.3050 cm, (d) 0.0105 L, (e) 7.0500 × 10-3m3.
1.10 Round each of the following numbers to four significant figures, and express the result in
standard exponential notation: (a) 102.53070, (b) 656,980, (c) 0.008543210, (d) 0.000257870, (e)
– 0.0357202
1.11 Perform the following conversions: (a) 0.076 L to mL, (b) 5.0 × 10-8 m to nm, (c) 6.88 × 105
ns to s, (d) 1.55 kg/m3 to g/L.
1.12 The density of air at ordinary atmospheric pressure and 25°C is 1.19 g/L. What is the mass,
in kilograms, or the air in a room that measures 12.5 × 15.5 × 8.0 ft?
Chapter 2. Atoms, Molecules, and Ions
2.1 How does Dalton’s atomic theory account for the fact that when 1.000 g of water is
decomposed into its elements, 0.111 g of hydrogen and 0.889 g of oxygen are obtained regardless
of the source of the water?
2.2 A chemist finds that 30.82 g of nitrogen will react with 17.60 g, 35.20 g, 79.40 g, or 88.00 g
of oxygen to form four different compounds. (a) calculate the mass of oxygen per gram of
nitrogen in each of the following. (b) How do the numbers in part (a) support Dalton’s atomic
theory?
2.3 Summarize the evidence used by J.J. Thompson to argue cathode rays consist of negatively
charged particles.
2.4 (a) What are the three subatomic particles that make up the atom? (b) What is the relative
charge on each of the particles? (c) Which of the particles is the most massive? (d) Which is the
least massive?
2.5 (a) Define atomic number and mass number. (b) Which of these can vary without changing
the identity of the element?
2.6 How many protons, neutrons and electrons are in each of the following atoms: (a) 40Ar, (b)
65
Zn, (c) 70Ga, (d) 80Br, (e) 184W, (f) 243Am
2.7 Fill in the gaps in the following table; assuming each column represents a neutral atom:
Symbol
52
Cr
Protons
25
Neutrons
30
Electrons
Mass
No.
82
64
48
86
222
207
2.8 Write the correct symbol, with both superscript and subscript, for each of the following. (a)
the isotope of platinum that contains 118 neutrons, (b) the isotope of krypton with mass number
84, (c) the isotope of arsenic with mass number 75, (d) the isotope of magnesium that has an
equal number of protons and neutrons.
2.9 For each of the following elements, write its chemical symbol, determine the group to which
it belongs, and indicate whether it is a metal, metalloid, or nonmetal: (a) potassium, (b) iodine,
(c) magnesium, (d) argon, (e) sulfur.
2.10 Write the empirical formula corresponding to each of the following molecular formulas: (a)
Al2Br6, (b) C8H10, (c) C4H8O2, (d) P4O10, (e) C6H4Cl2, (f) B3N3H6
2.11 How many hydrogen atoms are in each of the following: (a) C2H5OH, (b) Ca(CH3COO)2, (c)
(NH4)3PO4?
2.12 Fill in the gaps in the following table:
Symbol
59
Co3+
Protons
34
76
80
Neutrons
46
116
120
Electrons
36
Net
Charge
2+
2.13 Each of the following elements is capable of forming an ion in chemical reactions. By
referring to the periodic table, predict the charge of the most stable ion of each: (a) Mg, (b) Al, (c)
K, (d) S, (e) F.
2.14 Predict the chemical formula for the ionic compound formed by (a) Ca2+ and Br-, (b) K+ and
CO32-, (c) Al3+ and CH3COO-, (d) NH4+ and SO42-, (e) Mg2+ and PO43-.
2.15 Complete the table by filling in the formula for the ionic compound formed from each pair
of cations and anions, as shown for the first pair.
Ion
K+
Cl-
KCl
NH4+
Mg2+
Fe3+
OHCO32PO43-
2.16 Predict whether each of the following compounds is molecular or ionic: (a) B2H6, (b)
CH3OH, (c) LiNO3, (d) Sc2O3, (e) CsBr, (f) NOCl, (g) NF3, (h) Ag2SO4.
2.17 Give the names and charges of the cation and anion in each of the following compounds: (a)
CaO, (b) Na2SO4, (c) KClO4, (d) Fe(NO3)2, (e) Cr(OH)3
2.18 Name the following ionic compounds: (a) MgO, (b) AlCl3, (c) Li3PO4, (d) Ba(ClO4)2, (e)
Cu(NO3)2, (f) Fe(OH)2, (g) Ca(C2H3O2)2, (h) Cr2(CO3)3, (i) K2CrO4, (j) (NH4)2SO4
2.19 Write the chemical formulas for the following compounds: (a) aluminum hydroxide, (b)
potassium sulfate, (c) copper (I) oxide, (d) zinc nitrate, (e) mercury (II) bromide, (f) iron (III)
carbonate, (g) sodium hypobromite.
Chapter 3. Stoichiometry
3.1. What is the difference between adding a subscript 2 to the end of the formula for CO to give
CO2 and adding a coefficient in front of the formula to give 2CO?
3.2 Is the following equation, as written, consistent with the law of conservation of mass? Why
or why not. 3Mg(OH)2(s) + 2H3PO4(aq)  Mg3(PO4)2 (s) + 6H2O(l)
3.3 Balance the following equations:
(a) CO(g) + O2(g)  CO2(g)
(b) N2O5(g) + H2O(l)  HNO3(aq)
(c) CH4(g) + Cl2(g)  CCl4(l) + HCl(g)
(d) Al4C3(s) + H2O (l)  Al(OH)3(s) + CH4(g)
3.4 Determine the formula weights (molar masses) of the following compounds: (a) nitric acid,
HNO3; (b) KMnO4; (c) Ca3(PO4)2; (d) quartz, SiO2; (e) gallium sulfide, (f) chromium (III) sulfate,
(g) phosphorous trichloride.
3.5 Use the equation: Fe2O3(s) + 2Al (s)  2Fe (l) + Al2O3 (s) to answer the following questions:
(a) How many moles of aluminum, Al, are needed to react with 10 moles of iron (III) oxide,
Fe2O3? (b) How many moles of iron (III) oxide are needed to react with 0.2 moles of aluminum?
(c) How many grams of liquid iron can be produced from the reaction of 4.0 moles of iron (III)
oxide? (d) How many grams of aluminum would be needed to react with 1.2 moles of Fe2O3?
3.6 A student heats solid magnesium, which reacts with oxygen in the air to form a magnesium,
oxide compound. Mg(s) + O2(g)  MgxOy(s). (a) predict the formula of the compound based on
the most commonly formed ions of magnesium and oxygen. (b) if a student heats 5.65 grams of
magnesium in the presence of oxygen, what is the theoretical yield of the magnesium oxide based
on the formula you expect? (c) if a student heats 5.65 grams of magnesium in the presence of
oxygen then measures the mass of the product formed and finds it to be 8.97 grams, how many
grams of oxygen are in the product? (d) What is the empirical formula of the compound based on
the mass of magnesium and mass of oxygen calculated in part ‘c’? (e) what is the percent yield of
this experiment?
3.7 The reaction of methane and water is one way to prepare hydrogen for use as a fuel:
CH4(g) + H2O(g)  CO(g) + 3H2(g). If you begin with 995 g of CH4 and 2510 g of water, (a)
which is the limiting reactant? (b) What is the maximum mass of H2 that can be prepared? (c)
What mass of excess reactant remains when the reaction is completed?
3.8 When limestone, CaCO3, is heated, it decomposes according to the equation:
CaCO3(s)  CaO(s) + CO2 (g). A 1.506 gram sample of a solid material containing an
UNKNOWN amount of limestone is heated and 0.558 grams of carbon dioxide were produced.
(a) How many grams of CaCO3 from the sample must have decomposed (reacted) to produce the
0.558 grams of CO2? (b) what is the mass of percent of CaCO3 in the 1.506 gram sample?
3.9 Use the equation: KClO3(s) + P4(s)  P4O10(s) + KCl(s)
to answer the following questions: (a) How many moles of phosphorous, P4, are needed to react
with 13.5 grams of potassium chlorate, KClO3? (b) How many grams of phosphorous, P4, are
needed to produce 12.3 moles of P4O10? (c) How many grams of potassium chloride, KCl, are
produced when 52.3 grams of KClO3 react? (d) How many grams of potassium chlorate, KClO3,
are needed to react with 300 grams of phosphorous, P4?
3.10 Use the reaction: 3Al(s) + 3NH4ClO4(s)  Al2O3(s) + AlCl3(s) + 3NO(g) + 6H2O(g)
to answer the following questions. (a) How many moles of aluminum chloride, AlCl3, are
produced by the reaction of 14 moles of aluminum metal, Al? (b) How many grams of
ammonium perchlorate, NH4ClO4, are needed to react with 1.3 moles of aluminum metal, Al? (c)
How many grams of water are produced by the reaction of 50 grams of ammonium perchlorate,
NH4ClO4? (d) If a reaction produces 20 grams of aluminum oxide, Al2O3, how many grams of NO
are also produced in the reaction?
Chapter 4. Chemical Reactions
4.1 Write a balanced chemical equation for each of the following reactions.
(a) magnesium solid reacts with Cl2 gas.
(b) barium carbonate decomposes into barium oxide and carbon dioxide gas when heated.
(c) the hydrocarbon styrene (C8H8(l)) is combusted in air.
(d) dimethylether (CH3OCH3(g)) is combusted in air.
4.2 Balance the following equations and indicate whether they are combination, decomposition,
or combustion reactions.
(a) Al(s) + Cl2(g)  AlCl3(s)
(b) C2H4(g) + O2(g)  CO2(g) + H2O(g)
(c) Li(s) + N2(g)  Li3N(s)
(d) PbCO3(s)  PbO(s) + CO2(g)
(e) C7H8O2(l) + O2(g)  CO2(g) + H2O(g)
4.4 Write a balanced chemical equation and net-ionic equation for the precipitation reaction that
occurs when the following aqueous solutions are combined (a) AgC2H3O2 (aq) and NaCl (aq),
(b) aqueous barium chloride and aqueous sodium sulfate, (c) aqueous magnesium chloride and
aqueous sodium hydroxide.
4.5 The products formed by a reaction (either double replacement, single replacement, or
neutralization) are given below. Determine the formula and state of the two reactants.
(a) _____ + ______  Ba(NO3)2 (aq) + AgCl (s)
(b) _____ + ______  Fe(NO3)3 (aq) + H2O (l)
(c) _____ + ______  Cu (s) + Zn(NO3)2 (aq)
4.6 Assign oxidation numbers to each element in the following ions or compounds.
(a) BrO3-, (b) C2O42-, (c) H4SiO4, (d) PF6-, (e) POCl3, (f) N2O5
4.7 In the following reactions, determine which element is oxidized and which is reduced.
(a) C2H4 + 3O2  2CO2 + 2H2O
(b) Si + 2Cl2  SiCl4
4.8 A student is precipitating iron ions from 100-mL of an aqueous Fe2+ solution of unknown
concentration using excess hydroxide ions. Hydroxide and iron (II) react according to the netionic equation: Fe2+(aq) + 2OH-(aq)  Fe(OH)2(s). The student combines 100 mL of a 5.0M
sodium hydroxide solution with the 100-mL of aqueous iron solution. A precipitate forms, which
is then collected and dried; the final mass of precipitate is found to be 15.23 grams. (a) how many
moles of iron (II) ion were in the water sample? (b) what was the concentration of iron (II) ion in
the 100-mL sample? (c) how many moles of hydroxide remain unreacted after the two solutions
are combined?
4.9 A student is neutralizing 25 mL of hydrochloric acid with unknown concentration by reacting
it with a 0.5 M solution of barium hydroxide. Hydrochloric acid reacts with barium hydroxide
according to the following equation:
2HCl (aq) + Ba(OH)2 (aq)  BaCl2 (aq) + 2H2O (l).
(a) if it requires 13.5 mL of the 0.5 M barium hydroxide to completely neutralize the HCl, how
many moles of HCl were neutralized? (b) what is the concentration of the HCl?