Download 2016 Summer Assignment

AP Chemistry Summer Assignment
May 2016
Future AP Chemistry Student,
Welcome to AP Chemistry! I am eagerly anticipating a great year of Chemistry. In order to ensure the
best start for everyone next fall, I have prepared a Summer Assignment that reviews basic chemistry
concepts. Most of the material covered in the summer packet will be familiar to you, but is designed to
strengthen your foundation in chemistry and ensure that all students are on a relatively even plane. It
will be important for everyone to come prepared to class on the first day. While we will review,
extensive remediation is not an option as we work towards the goal of being prepared for the AP Exam
in early May. The Summer Assignment is DUE THE FIRST DAY OF SCHOOL! There will be a quiz
covering the basic concepts included in the summer packet during the first week of school.
I hope you are gearing up for an exciting, challenging, and rewarding academic course. The grades you
earn on tests in AP Chemistry may be lower than those you’ve earned in Honors Chem. That is the
nature of a difficult AP course, and this is one of the “tough” AP courses. But, I know you will learn a
lot in AP Chemistry, and it will be an excellent way for you to prepare for college. It will build your
knowledge base, enhance your work habits and your organizational skills, and you will grow as an
independent learner. You are all certainly fine students, and with motivation and hard work, you should
find AP Chemistry a successful and rewarding experience.
Finally, I recommend that you find a “study buddy” and spread out the summer assignment. Please do
not try to complete it all in one day. Chemistry takes time to process and grasp at a level necessary for
success in AP Chemistry. Remember, AP Chemistry is equivalent to a first-year chemistry college
course, and should be considered as such. College Board recommends that students taking AP
Chemistry put in a minimum of 5 hours of study per week outside the classroom. Taking a college
level course in high school is difficult, requires dedication, and is a great investment in your
education, so prepare yourself and arrive ready to learn.
Have a great summer and enjoy the chemistry.
Mrs. Heitmann
Mrs. Heitmann
Perry High School Chemistry
Materials for first day of class:
 Completed Summer Assignment
 Scientific Calculator
 3-Ring Binder
 Carbonless Duplicate Graph Paper AP Chem Lab Book (to be purchased from the PHS
Book Store)
1
PART ONE
Make flashcards and memorize the names, symbols and charges of the common ions listed below.
Names, Formulas, and Charges of Some Common Polyatomic Ions
NH4
+
Ammonium
C2H3O2
CO3
2-
C2O4
HSO4
NO2
NO3
-
Chlorate
2-
ClO3
Thiosulfate
ClO4
-
Perchlorate
S2O3
-
HPO4
O2
Nitrate
32-
H2PO4
-
Hydrogen sulfide
HS
Nitrite
-
-
BrO
Hydroxide
2-
CrO4
Hypochlorite
ClO
Hydrogen sulfite
Oxalate
OH
Hydrogen sulfate
-
2-
Thiocyanate
Hypofluorite
-
ClO2
HSO3
-
-
Sulfite
Hydrogen carbonate
Cyanide
SCN
-
FO
-
-
-
Sulfate
2-
SO3
-
PO4
Acetate
2-
Carbonate
HCO3
CN
-
SO4
BrO3
Peroxide
2-
-
BrO4
Chromate
IO
Hypobromite
-
Bromate
-
Perbromate
-
Hypoiodite
2-
Dichromate
-
IO3
Permanganate
2-
IO4
Manganate
AsO4
Phosphate
Cr2O7
Hydrogen phosphate
MnO4
Dihydrogen phosphate
MnO4
Chlorite
-
Iodate
-
Periodate
3-
Arsenate
Names, Formulas, and Charges of Some Common Ions
3+
Aluminum
Au
-
Hydride
Sn
Al
H
2+
Mn
2+
Ni
2+
Zn
Cd
2+
+
+
Ag
Au
3+
Gold (III)
Fe
2+
Tin (II)
4+
Manganese (II)
Sn
Nickel (II)
Zinc
Pb
2+
Pb
4+
Tin (IV)
Lead (II)
Lead (IV)
Iron (III)
Co
2+
Cobalt (II)
Co
3+
Cobalt (III)
Cu
+
Copper (I)
2+
Copper (II)
Cu
Cr
2+
Silver
Cr
3+
Chromium (III)
Gold (I) or aurous
Fe
2+
Iron (II)
Cadmium
3+
Chromium (II)
Hg2
2+
2+
Hg
Mercury (I)
Mercury (II)
PART TWO
More Memorization.
1. Memorize metric prefixes, and know how to convert from one unit to the next.
6
Mega
(M)
10
Kilo
(k)
10
Deci
(d)
10
Centi
(c)
Milli
(m)
10-2
-3
10
Micro
(μ)
10
Nano
(n)
10
3
-1
-6
-9
or
1,000,000
One million of
or
1000
One thousand of
or
1/10
One tenth of
or
1/100
One hundredth of
or
1/1000
One thousandth of
or
1/1,000,000
One millionth of
or
1/1,000,000,000
One billionth of
2
2. Know solubility rules.
a)
b)
c)
d)
e)
f)
All compounds containing alkali metal cations and the ammonium ion are soluble.
All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble.
All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb 2+, or Hg2+.
All sulfates are soluble except those containing Hg2+, Pb2+, Sr2+, Ca2+, or Ba2+.
All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+.
All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble except those that also contain alkali
metals or NH4+.
3. Know rules for assigning oxidation numbers.
a)
The oxidation number for an atom in its elemental form is always zero.
Examples:
S8: The oxidation number of S = 0
Fe: The oxidation number of Fe = 0
b) The oxidation number of a monoatomic ion = charge of the monatomic ion.
Examples:
Oxidation number of S2- is -2.
Oxidation number of Al3+ is +3.
c)
The oxidation number of all Group 1A metals = +1 (unless elemental).
d) The oxidation number of all Group 2A metals = +2 (unless elemental).
e)
Hydrogen (H) has two possible oxidation numbers:
o
o
f)
+1 when bonded to a nonmetal
-1 when bonded to a metal
Oxygen (O) has two possible oxidation numbers:
o
o
-1 in peroxides (O22-)....pretty uncommon
-2 in all other compounds...most common
g) The oxidation number of fluorine (F) is always -1.
h) The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.
i) The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion.
PART THREE
Start familiarizing yourself with the following special case examples of reaction types.
Synthesis:
Metal oxide + water  metal hydroxide
Nonmetal oxide + water  acid (polyatomic)
MgO + H2O  Mg(OH)2
CO2 + H2O  H2CO3
Decomposition:
Metal carbonate  metal oxide + carbon dioxide
Metal hydroxide  metal oxide + water
Metal chlorate  metal chloride + oxygen gas
Cs2CO3  Cs2O + CO2
2 LiOH  Li2O + H2O
2 LiClO3  2 LiCl + 3 O2
Acid/Base Reactions
acid + base  salt + water
H2SO4 + NaOH  Na2SO4 + H2O
metal + acid  salt + hydrogen gas
3Mg + 2H3PO4  Mg3(PO4)2 + 3H2
metal oxide + acid  salt + water
CaO + HCl  CaCl2 + H2O
carbonate + acid  salt + carbon dioxide + water
BaCO3 + H2SO4  BaSO4 + CO2 + H2O
(The carbonate ion breaks apart. The other negative ion becomes part of the salt)
nonmetal oxide + hydroxide  salt + water
CO2 + Ba(OH)2  BaCO3 + H2O
metal oxide + nonmetal oxide  salt
Na2O + SO3  Na2SO4
3
PART FOUR
Complete the following assignment. Show your work where relevant. Answers are in [ ] after selected
problems/questions.
Nomenclature and Formulas
1. Name these binary compounds of two nonmetals.
IF7_____________________________________
N2O5________________________________________
N2O4__________________________________
As4O10_____________________________________
XeF2 ________________________________
SF6___________________________________
2. Name these binary compounds with a fixed charge metal.
AlCl3 ______________________
MgO __________________________
KI__________________________
SrBr2__________________________
BaI2______________________
Na2S ____________________
3. Name these binary compounds of cations with variable charge.
CuCl2 ______________________
Fe2O3_________________________
SnO______________________
PbCl4_______________________
Cu2S__________________________
HgS______________________
4. Name these compounds with polyatomic ions.
Fe(NO3)3___________________
NaOH_____________________
Ca(ClO3)2___________________________
KNO2_____________________
Cu2SO4___________________________________
NaHCO3______________________
5. Name these binary acids
HCl _______________________
HF ___________________________
HI ________________________
6. Name these acids with polyatomic ions.
HClO4________________________________
H3PO4_____________________
H2SO4_________________________________
HNO2_______________________
HC2H3O2___________________________
H2CO3_____________________
7. Name these compounds appropriately.
CO________________________
NI3________________________
LiMnO4___________________
SO2________________________
FeF3_______________________
NH4CN ____________________
AlP ________________________
HClO(aq) ______________________
CuCr2O7___________________
KC2H3O2___________________
HIO3(aq)______________________
OF2____________________________
HF(aq)________________________
K2O____________________________
MnS___________________________
8. Write the formulas for the following compounds.
tin (IV) phosphide______________________
magnesium hydroxide__________________
sulfurous acid___________________________
potassium nitride _______________________
gallium arsenide_________________________
copper (II) cyanide______________________________
sodium peroxide_________________________________
lithium oxalate __________________________________
chromium (III) carbonate_______________________
cobalt (II) chromate___________________________
4
Solubility rules
9. Review solubility rules and identify each of the following compounds as soluble (S) or insoluble
(IS) in water.
Na2CO3 _____
K2S_____
AgI_____
FeS_____
Li2O______
AgClO3_____
CoCO3_____
BaSO4_____
Ni(NO3)2______
PbCl2______
Mn(C2H3O2)2_____
Sn(SO3)4 _____
Pb(NO3)2_____
(NH4)2S _____
KI ______
CuSO4_____
Cr(OH)3______
FeF2______
10. Predict whether each of these double replacement reactions will give a precipitate or not
based on the solubility of the products. If yes, identify the precipitate.
a) silver nitrate and potassium chloride
b) magnesium nitrate and sodium carbonate
c) strontium bromide and potassium sulfate
d) cobalt (III) bromide and potassium sulfide
e) ammonium hydroxide and copper (II) acetate
f) lithium chlorate and chromium (III) fluoride
Balancing Equations
11. Balance the following equations with the lowest whole number coefficients.
__S8 + __O2 → __SO3
__Fe + __O2 → __Fe2O3
__C7H6O2 + __ O2 → __CO2 + __H2O
__V2O5 + __HCl → __VOCl3 + __H2O
__Hg(OH)2 + __H3PO4 → __Hg3(PO4)2 + __H2O
Significant Figures
12. How many significant figures are in each of the following? [5,4,1,4]
a)
b)
c)
d)
0.030100 kJ
6.022 x1023 atoms
100
1001
5
13. Calculate the following to the correct number of significant figures.
a) 1.27 g / 5.296 cm3 [3 sf]
b) 12.2 mL + 0.38 mL [3 sf]
c) 7.355 g - 2.785 g [4 sf]
d) 0.1 x 3.21m [1 sf]
Moles, Stoichiometry, etc.
14. Na2O + H2O → 2NaOH
What mass of water would be needed to react with 10.0g of sodium oxide?
15. 2NaClO3 → 2NaCl + 3O2
What mass of sodium choride is formed along with 45.0g of oxygen gas?
16. 4NH3 + 5O2 → 4NO + 6 H2O
What mass of water will be produced when 100.0g of ammonia is reacted with excess oxygen?
17. If the reaction in #16 is done with 25.0g of each reactant, which would be the
limiting reactant?
18. The molecular formula of morphine, a pain-killing narcotic, is C17H19NO3.
a) What is the molar mass? [285 g/mol]
b) What fraction of atoms in morphine is accounted for by carbon? [17/40]
c) Which element contributes least to the molar mass? [N]
19. You have 1.023g of a CoSO4 hydrate. After heating and driving off all the water, the mass of the
anhydride is 0.603 g. Determine the hydrate’s chemical formula (number of waters attached).
[CoSO4 • 6 H 2O]
20. Determine the empirical formula given 76.0% iodine and the other element in the
compound is oxygen. [I2O5]
6
Review Questions – All Topics
1. Given the reaction:
What is the mole-to-mole ratio between nitrogen gas and hydrogen gas?
1. 1:2
2. 1:3
3. 2:2
4. 2:3
2. Given the balanced equation representing a reaction:
2H2 + O2 → 2H2O
What is the total mass of water formed when 8 grams of hydrogen reacts completely with 64 grams of oxygen?
1. 18 g
3. 56 g
2. 36 g
4. 72 g
3. A sample of water is being heated from 20oC to 30oC, and the temperature is recorded every 2 minutes. Which table would be
most appropriate for recording the data?
4. Expressed to the correct number of significant figures, the sum of two masses is 445.2 grams. Which two masses produce this
answer?
1. 210.10 g + 235.100 g
2. 210.100 g + 235.10 g
3. 210.1 g + 235.1 g
4. 210.10 g + 235.10 g
5. A student constructs a model for comparing the masses of subatomic particles. The student selects a small, metal sphere with a
mass of 1 gram to represent an electron. A sphere with which mass would be most appropriate to represent a proton?
1. 1 g
2. 1/2 g
3. 1/2000 g
4. 2000 g
6. A 100.00-gram sample of naturally occurring boron contains 19.78 grams of boron-10 (atomic mass - 10.01 atomic mass units)
and 80.22 grams of boron- 11 (atomic mass = 11.01 atomic mass units). Which numerical setup can be used to determine the atomic
mass of naturally occurring boron?
1. (0.1978)(10.01) + (0.8022)(11.01)
2. (0.8022)(10.01) + (0.1978)(11.01)
3. [(0.1978)(10.01)]/[(0.8022)(11.01)]
4. [(0.8022)(10.01)]/[(0.1978)(11.01)]
7. In Rutherford's gold foil experiments, some alpha particles were deflected from their original paths but most passed through
the foil with no deflection. Which statement about gold atoms is supported by these experimental observations?
Alpha particles and gold nuclei have opposite
1. Gold atoms consist mostly of empty space.
3. charges.
2. Gold atoms are similar to alpha particles.
4. Alpha particles are more dense than gold atoms.
8. The characteristic bright-line spectrum of an element occurs when electrons
1. move from lower to higher energy levels
3. are lost by a neutral atom
2. move from higher to lower energy levels
4. are gained by a neutral atom
7
3+
9. How many electrons are contained in an Au ion?
1. 76
2. 79
3. 82
4. 197
10. Equal volumes of 0.1 M NaOH and 0.1 M HCl are thoroughly mixed. The resulting solution has a pH closest to
1. 5
3. 3
2. 7
4. 9
11. Given the equation:
HCl(g) + H2O(l) --> X(aq) + Cl-(aq)
Which ion is represented by X ?
1. hydroxide
2. hydronium
3. hypochlorite
4. perchlorate
12. Which compound when dissolved in water is an Arrhenius acid?
1. CH3OH
2. HCl
3. NaCl
4. NaOH
13. An acid can be defined as an
+
1. H acceptor
+
2. H donor
3. OH acceptor
4. OH donor
14. Which statement is true about the charges assigned to an electron and a proton?
1. Both an electron and a proton are positive.
3. An electron is negative and a proton is positive.
2. An electron is positive and a proton is negative.
4. Both an electron and a proton are negative.
15. In the wave-mechanical model, an orbital is a region of space in an atom where there is
1. a high probability of finding an electron
3. a circular path in which electrons are found
2. a high probability of finding a neutron
4. a circular path in which neutrons are found
8