Download AP Chemistry Summer Assignment

Material That You Should Have Already
Covered in Honors Chemistry
Chapter 1 Intro
The Scientific Method
Chapter 3 Scientific Measurement
Metric Conversions
Significant Figures
Scientific Notation
Chapter 4 Atomic Structure
Dalton, Thomson, Rutherford, Bohr experiments and models of the atom
Isotopes and average mass calculations
Ions
Chapter 5 Electrons in the Atom
Schrodingers model (quantum mechanics)
Shapes of s, p, d and f orbitals
Electron configurations
Electron configurations of ions
Electromagnetic Radiation (calculation of wavelength, frequency and energy)
Absorption and Emission of Light and Emission Spectra
Flame Test
Chapter 6 Periodic Table\
Trends in
Radius
Ionization energy (first, second, third, etc)
Electrogenativity
Ion size
Chapter 7 Ionic and Metallic Bonding
Lewis drawings of elements and ions
Characteristics of ionic and metallic bonds
Chapter 8 Covalent Bonding
Draw Lewis structures for molecules
Molecular Geometry : Name VSEPR shapes for molecules
Hybridization (sp, sp2, sp3, sp3d and sp3d2)
Resonance
Polar bonds
Polarity based on molecular geometry
Intermolecular forces: Hydrogen bonds, London dispersion, Dipole-Dipole, Ion- Dipole,
Covalent Network Solids
Chapter 9 Named and Formula
Memorize polyatomic ions
Write names and formula for A salts, B salts, and molecules
Write names and formula for acids
Chapter 10 Chemical Quantities
Dimensional Analysis
Mole/mass/volume/particle conversions
Empirical and Molecular Formula
Chapter 11 Reactions
Types of Chemical Reactions
Double Displacement
Single Displacement (oxidation and reduction concepts)
Synthesis
Decomposition
Combustion
Acid Base Reactions
Must be able to write balanced reactions from English!!!!
Chapter 12 Stoichiometry
Mole/mass/volume/particle conversions with balanced reactions
Limiting Reactants
Chapter 13 and 14 Matter and Gases
Standard Temperature and Pressure
Phase Diagrams
Boyles, Charles, Gay-Lussac, Avogardro’s Gas Laws
Ideal Gas Law
Gas Stoichiometry
Dalton’s Law
Graham’s Law
Kinetic Molecular Theory
Chapter 15 and 16 Solutions
Characteristics of water (specific heat, surface tension, hydrogen bonding)
Solute vs. Solvent
Factors that effect solubility (heat, agitation, particle size)
Vapor Pressures of Solutions
Colligative Properties (including the math!)
Molarity
Dilution Problems (M1V1 = M2V2)
Chapter 17 Thermochemistry
Specific Heat Capacity
Enthalpy or reactions
Hess’s Law
Calorimetry
Chapter 18 Reactions Rates and Equilibrium
Reaction Rates (temperature, concentration, particle size, catalysts)
Equilibrium Constant (including the math!)
Equilibrium Expression
Le Chatelier’s Principle
Thermodynamics – spontaneous reactions
Free Energy
Chapter 19 Acids and Bases
Concepts behind acid strength
Calculation of pH (including logarithmic problems)
Properties or acids and bases
Titrations (concentration, molar mass and unknown volume)
Chapter 22 Organic Chemistry
Alkanes, alkenes, alkynes
Functional Groups
Chapter 25 Nuclear Chemistry
Nuclear Stability and Radioactive Decay
Nuclear Reactions
Half life problems
Reaction Types
In balancing equations and predicting products it is always helpful to know what type of reaction it is. Also by
looking at the reactants you can predict which kind of product will form. Using oxidation states you can
determine the formula of the product.
1. Synthesis (Combination) Two reactants combine to give one product.
General Form:
A + B → AB (A and B can be either elements or compounds
a) Metal oxide + water  base
CaO(s) + H2O(l)  Ca(OH)2(aq)
Na2O(s) + H2O (l)  2 NaOH (aq)
b) Nonmetal oxide + water  acid
SO2(s) + H2O(l)  H2SO3(aq)
SO3(s) + H2O(l)  H2SO4(aq)
N2O5(s) + H2O(l) → 2 HNO3(aq)
P2O5(s) + 3 H2O(l) → 2 H3PO4(aq)
d) Metal + oxygen → metal oxide
2 Mg(s) + O2(g) → 2 MgO(s)
4 Al(s) + 3 O2(g) → 2 Al2O3(s)
(basic anhydrides)
(acid anhydrides)
e) Nonmetal + oxygen → nonmetal oxide
P4(s) + 3O2(g) → P4O10(s)
S(s) + O2(g) → SO2(g)
N2(g) + O2(g) → 2NO(g)
f) Metal + nonmetal → Salt
2Al(s) + 3 Cl2(g) → 2 AlCl3(s)
2 Al(s) + 3Br2(l) → 2 AlBr3(s)
3 Ca(s) + N2(g) → Ca3N2(s)
2. Decomposition: a single substance is decomposed or broken down to give two or more different products.
General Form AB → A + B
a) Hydrolysis of water
2H2O(l) → 2H2(g) + O2(g)
b) Hydrogen Peroxide
(with either a catalyst or heat)
MnO2
2 H2O2(g) → 2 H2O(g) + O2(g)
c) Chlorates
KClO3(s) + Δ → KCl(s) + O2(g)
d) Carbonates
(NH4)2CO3(s) → H2O(g) + CO2(g) + NH3(g)
MgCO3(s) → MgO(s) + CO2(g)
CaCO3(s) → CaO(s) + CO2(g)
NaHCO3(s) → Na2CO3(s) + H2O(g) + CO2(g)
e) Misc
CaSO3(s) → CaO(s) + SO2(g)
NH4NO3(s) → N2(g) + NO(g) + H2O(g)
Ca(OH)2(s)  CaO(s) + H2O(g)
3. Single Displacement (Redox) An element reacts with a compound totake the place of one of the elements of
that compound. A new element is formed along with a new compound.
General Form:
A + BC  B + AC
(A is a metal)
A + BC  C + BA
(C is a nonmetal)
a) Metal and Acid  hydrogen + salt
H2SO4(aq) + Fe(s) → FeSO4(aq) + H2(g)
2 Al(s) + 3 Pb(NO3)2(aq) → 2 Al(NO3)3(aq) + 3 Pb(s)
b) Metal and Water hydrogen + metal hydroxide OR metal oxide
2Na(s) + 2H2O (l) → NaOH(aq) + H2(g)
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
Mg(s) + H2O(l) → MgO(s) + H2(g)
(water must be hot for this one!)
c) Metal + salt  metal + salt
CuSO4(aq) + Fe(s) → FeSO4(aq) + Cu(s)
AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + Ag(s)
d) Halogens + Halogen salt  halogen + halogen salt
Elements higher up on the halogens will replace element lower on the table
Cl2(g) + 2 KBr(aq) → Br2(l) + 2 KCl(aq)
Cl2(g) + 2 KF(aq) → NO REACTION!
These will never be assigned…….all questions ARE reactions.
4. Double Discplacement (Metathesis) Two compounds exchange partners to produce two different compounds.
General Form: AB + CD →AD + CB
a) Neutralization
H3PO4(aq) + 3NH4OH(aq) → 3H2O(l) + (NH4)3PO4(aq)
b) Formation of a Precipitate: Know your Solubility Rules!
CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2NaCl(aq)
Fe2(SO4)3(aq) + 3Pb(ClO3)2(aq) → 3PbSO4(s) + Fe(ClO3)3(aq)
c) Metal Oxide and Acid
CuO(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2O(l)
d) Formation of a gas:
.
Na2S(aq) + 2HBr(aq) → 2NaBr(aq) + H2S(g)
e) Some unstable compounds formed will decompose into water and a gas:
1. Carbonic Acid
(NH4)2CO3(aq) + 2HCl(aq) → H2CO3(aq) + 2NH4Cl(aq) → H2O(g) + CO2(g) + 2NH4Cl
2. Ammonia
KOH(aq) + NH4Br(aq) → NH4OH(aq) + KBr(aq) → NH3(g) + H2O(g) + KBr
3. Sulfurous Acid
2HCl(aq) + K2SO3(aq)
+ 2KCl(aq)
2SO3(aq) + 2KCl(aq) → H2O + SO2
5. Combustion: a carbon compound is burned in oxygen to creat carbon dioxide and water.
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
2 CH3OH(l) + 3 O2(g) → 2 CO2(g) + 4 H2O(g)
C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(g)
Naming Compounds Test Review
Answer each of the following questions and give an example that helps with your explanation.
1. If a compound ends in –ide, what does it tell you about the compound?
2. If a compound ends in –ate what does it tell you about the compound?
3. If a compound ends in –ite what does it tell you about the compound?
4. How does a cation form?
What charge does it form?
What type of element forms these?
5. How does a anion form?
What charge does it form?
What type of element forms these?
6. What is a binary compound?
7. What is a ternary compound?
8. What is a polyatomic ion?
9. What is a molecular compound?
10. What is a compound?
11. What is an ionic compound?
12. What must be true about al ionic compounds?
13. What charge does each element in a group on the periodic table form?
14. What is a transition metal?
15. What type of ions have names that end in –ide?
16. If a polyatomic ion ends in –ate or –ite, what does it tell you?
17. What do metals do with their valence electrons?
18. What do nonmetals want to do to become stable?
19. What group does not react with other elements naturally?
20. How do you determine the charge of the transition element when you are given the chemical formula?
21. What are the formulas and charges of the polyatomic ions you have to memorize?
22. When do you have to use parentheses in a chemical formula?
23. When do you use prefixes in the name of a compound?
24. How do you name a binary acid?
25. How do you name an acid that involves a polyatomic ion?
26. In transition elements, which charge is given the –ic ending and which is given the –ous ending?
27. What does an acid do when it dissolves in water (Hint: how do you know a compound is an acid?)
28. What are the prefixes for 1-10?
29. How should an ionic compound always be written?(What is written first, then second?)
30. What do you do to the polyatomic ion when you add an “H+” to it? i.e. HCO331. What is the order for the 4 ions of chlorate, bromate or iodate polyatomic ions?
32. What is unique about the charges for the element mercury?
33. How do you name acids that involved cyanide and hydroxide ions?
34. How do you determine if a compound is ionic or moleculatr?
35. How are you going to do on this test? STUDY!
Practice Problems:
Name the following compounds and characterize the solubility of these compounds when dissolved in water.
Mark each as Soluble (S) or Insoluble (IN).
1. NH4NO3
16. BaCl2
2. CsClO4
17. KCl
3. NaOH
18. BaCO3
4. Na3PO4
19. K2CO3
5. Hg2I2
20. KOH
6. Ca(OH)2
21. H2O
7. BaO
22. K3PO4
8. Al(C2H3O2)2
23. SrBr2
9. SrCl2
24. (NH4)2CO3
10. MgSO3
25. SrCO3
11. SrCl2
26. NH4Br
12. Pb(ClO3)2
27. Ag2SO4
13. PbCl2
28. AlCl3
14. Sr(ClO3)2
29. AgCl
15. H3PO4
30. Al2(SO4)3
Solubility Rules for Ionic Compounds
Name the following compounds and characterize the solubility of these compounds when dissolved in water.
Mark each as Soluble (S) or Insoluble (IN).
31. a. Ammonium nitrate
i. Zinc oxide
b. Mercuric chloride
j. Cuprous bromide
c. Lead (II) Phosphate
k. Ammonium perchlorate
d. Strontium carbonate
l. Sodium chloride
e. Barium sulfate
m. Nickel (II) hydroxide
f. Calcium sulfide
n. Cadmium (II) nitrate
g. Silver iodide
o. Hydrogen oxalate
h. Sodium chromate
p. Ferric Sulfide
32. Name the following compounds, using both Roman numerals and –ic/-ous endings when necessary.
a. CO
m. H2S2O3
b. H2O
o. N2O
c. SO3
p. NiO
d. Cl2
q. P4O10
e. HCN
r. Cl2O8
f. Cu
s. NaSCN
g. CoO
t. HF
h. NaHSO4
u. S2Cl2
i. N2H4
v. HClO4
j. Mg(MnO4)2
w. PbCr2O7
k. PI3
x. HNO2
l. Fe(C2H3O2)2
y. HIO
m. SnO2
z. ICl
33. Write the chemical formula for each of these compounds.
a Aluminum hydrogen carbonate
n. Hydroiodic acid
b. Potassium perchlorate
o. nickel (II) nitrite
c. Calcium carbonate
p. Cuprous carbonate
d. Acetic acid
q. silver cyanide
e. Boron trichloride
r. Ammonium chlroate
f. Dichlorine heptachloride
s. nitrogen
g. Stannic phosphate
t. Magnesium sulfide
h. lithium fluoride
u. Ferric sulfate
i. Hydrochloric acid
v. Zinc oxide
j. Sulfur hexafluoride
w. Hydronium fluoride
k. Iron (III) sulfide
x. Cupric thiocyanate
i. Hypobromous acid
y. Manganese (V) phosphate
m. Mercuric iodide
z. Stannous oxalate
Naming Practice
For each of the following compounds give the correct formula name or chemical formula.
Make sure to include Roman numerals when working with a transition metal.
1. AgCl
2. Fe(NO3)3
3. Sr(HCO3)2
4. ZnSO4
5. K2CrO4
6. HBr
7. CuCl2
8. Mg3(PO4)2
9. (NH4)2SO3
10. NaClO2
11. Au(NO3)3
12. H2SO4
13. Cr(OH)3
14. Pb(MnO4)4
15. SnC2O4
16. H2CO3
17. N2O5
18. HClO2
19. Ba(HSO3)2
20. KCN
21. PH3
22. Co(IO4)2
23. Ni2S3
24. Al(BrO3)3
25. H2O
26. zinc sulfite
27. ammonium sulfide
28. nickel (III) chromite
29. lithium nitride
30. aluminum carbonate
31. nitrogen monoxide
32. lead (II) phophite
33. manganese (III) acetate
34. antimony (IV) hydroxide
35. silver hypobromite
36. sufurous acid
37. sulfur hexafluoride
38. barium chloride
39. nitric acid
40. ammonium phosphate
41. tetraphosphorus decoxide
42. hydrofluoric acid
43. calcium cyanide
44. silver chromate
45. aluminum fluoride
46. acetic acid
47. nickel (III) periodate
48. carbon disulfide
49. oxalic acid
50. ammonium acetate
Balancing Equations Review
Vocabulary: For each of the words below you need to be able to recognize them in an equation and be able to use
them in an explanation.
Coefficient
Synthesis/Combination
Products
Gas
Subscript
Decomposition
Reactants
Liquid
Precipitate
Single Displacement
Yields
Solid
Aqueous
Double Displacemenet
Soluble
Balanced equation
Diatomic elements
Combustion
Insoluble
Law of conservation of Mass
word equation
Chemical equation
1. Aluminum is added to hydrochloric acid to produce hydrogen has and aluminum chloride.
2. Calcium chlorate decomposes into calcium chloride and oxygen gas.
3. Diphosphorus pentoxide is added to water to form phosphoric acid.
4. Magnesium chloride is added to boron trifluoride to form magnesium fluoride and boron trichloride.
5. Fluorine gas reacts with aluminum oxide to form oxygen gas and aluminum fluoride.
6. Iron metal reacts with water for form iron (III) oxide and hydrogen gas.
7. Sodium carbonate decomposes into carbon dioxide gas and sodium oxide.
8. Lead metal is placed in a solution of silver nitate to form lead (II) nitrate and silver metal.
9. Solutions of aluminum hydroxide and sulfuric acid are mixed.
10. Zinc phosphate solution is mixed with tin (IV) nitrate solution.
11. Sodium oxide is mixed with water to form the metal base.
12. Sulfuric acid decomposes into its nonmetal oxide and water.
13. Ammonium hydroxide is mixed with copper (II) phosphate.
14. Combustion of sucrose yields carbon dioxide and water.
15. Gold (III) sulfide is mixed with hydrogen gas to produce gold and hydrogen sulfide gas.
16. Sodium bicarbonate decomposes into sodium carbonate, carbon dioxide and water.
17. Carbon disulfide liquid reacts with ammonia to produce hydrogen sulfide gas and solid ammonium
thiocyanate.
18. Sodium chloride salt is placed in water and electricity is applied. Chlorine gas, hydrogen gas and sodium
hydroxide is produced.
19. When heated solid potassium perchlorate produces potassium chloride and oxygen gas.
20. Lithium hydroxide solution is poured into nitric acid (HNO3).
21. Aluminum metal burns in oxygen to form a salt.
22. Zinc is dropped into hydrochloric acid.
23. Hydrogen peroxide is heated to form water vapor and oxygen gas.
24. Chlorine gas is bubbled through a solution of magnesium bromide displacing liquid bromine and a salt.
25. Metallic lithium and oxygen gas produces lithium oxide.
26. Barium carbonate is heated producing barium oxide and carbon dioxide.
27. Solid potassium oxide salt is dropped into water to produce potassium hydroxide solution.
28. When solid lime (calcium oxide) is heated, it decomposes into its constituent parts.
29. Potassium phosphate solution reacts with a solution of lead (II) nitrate to produce a lead solid.
30. A piece of nickel is place into silver nitrate forming nickel (II) nitrate.
Stoichiometry and Gas Laws Review Problems
Solve the following problems. Show all work, round off correctly to the number of significant digits. Use units in
your work and in your answer.
1. How many grams are in 0.420 moles of HF?
2. Bees release about 1 x 10-6 g C7H14O2. How many molecules are released in a typical bee sting?
How many atoms of carbon are present?
3. How many moles of hydrogen are in 0.684 grams of sulfuric acid?
4. What is the volume of 8.27 x 103 mole of ammonia gas at STP?
5. How many grams of phosphorus are in 8.45 x 1024 molecules of phosphoric acid?
6. a) What is the volume of a mixture of 0.813 moles of hydrogen sulfide and 0.439 mol of fluorine gas at STP?
b) What is the partial pressure of each of these gases?
7. How many atoms of hydrogen are in 10.586 grams of ammonia gas (NH3)?
8. How many grams of carbon tetrachloride are present in a sample of gas in a 2.50 liter container at 30 °C and
1.0 atm?
9. What is the mass in grams of one molecule of C145H293O168?
10. What is the volume of 1 molecules of ozone at STP?
11, What is the weight in grams of 1 U atom?
12. A sample of gas at 5 °C and 1.5 atm has a volume of 2.58 L. What volume will the gas occupy at 38 °C and
1.0 atm?
13. Determine the empirical formula of a compound that is 17 % Mg, 38% Al and 45%O.
14. Determine the empirical formula for 86% C and 14% H. What is its molecular formula if it has a molar mass
of 70.0 g?
15. Determine the empirical formula of a compound that is 27% K, 35% Cr, and 38% O.
16. Determine the percentage of iron in ferrous sulfate.
17. Determine the percentage of iron in a rock that has a mass of 10.0 g that is an equal mixture of 2 iron
compounds, ferrous carbonate and ferric sulfide.
18. If a rock has a mass of 5.98 grams and consists of 37.8% calcium carbonate and 62.2% ferric sulfide, how
many moles of each compound are in the rock?
19. Determine the percentage of carbon in acetic acid.
20. How many grams of phosphorus is in 250.0 grams of copper II phosphate?
21. What mass of sodium is in 405.3 grams of sodium carbonate?
22. How many oxygen atoms are found in 109.2 grams of HClO2?
23. There are 8.09 x 10 22 atoms of helium gas in a balloon at standard temperature and pressure. What is the
volume in milliliters?
24. How many grams of nickel II bicarbonate contains 607.0 grams of nickel?
25. Consider a 1.53 L sample of gaseous sulfur dioxide at a pressure of 5.6 kPa. If the pressure is changed to 150
kPa at a constant temperature what will be the new volume of gas?
26. Aureomycin is an antibiotic which contains 55.2 % C, 26.7 % O, 5.85 % N, 7.5 % Cl and 4.84 % H. What is
the true molecular formula if the molecular mass if 478.5 g/mol?
27. A white powder and sends it in to be analyzed. It is found to contain 43.64 % phosphorus and 56.36 %
oxygen by mass. Obtaining the empirical formula will be tricky because you will find a fraction in the ratios.
Multiply these numbers by 2 to get two whole numbers. The compound has a
molar mass of 283.88 g/mol. What is the molecular formula?
28. Caffeine, a stimulant found in coffer, tea, and chocolate, contain 49.18% carbon, 5.05 % hydrogen, 28.87 %
nitrogen and 16.49 % oxygen by mass. The molar mass is 194.2 g/mol. What is the molecular formula?
29. When a 1.50 g sample of an organic compound (Made of C, H, and O) is burned, 3.3 g of carbon dioxide and
1.8 g of water is formed. What is the empirical formula?
(Combustion Analysis: Page 91)
30. Combustion of 19.81 g of terephthalic acid produces 42.98 g of carbon dioxide and 6.45 grams of water.
The molar mass of terphtalic acid is 166.0 g/mol. Calculate the empirical and molecular formulas of the acid.
31. A sample of a compound of Cl and O reacts with the excess H2 to give 0.233 g of hydrochloric acid and
0.4003 g of water. Determine the empirical formula of the compounds.
32. Solid lithium hydroxide is used in space vehicles to remove exhaled carbon dioxide from the living
environment by forming solid lithium carbonate and liquid water. What mass of gaseous carbon dioxide can be
absorbed by 1.00 x 103 g of lithium hydroxide?
Limiting Reactants
33. DDT, an insecticide harmful to mammals is produced by the following reaction:
2 C6H5Cl + C2HOCl3 → C14H9Cl5 + H2O
(chlorobenzene)
(chloral)
(DDT)
In a government lab 1142 g of chlorobenzene were reacted with 485 g of chloral.
a) In grams how much DDT was formed?
b) What amount of excess reactant is leftover?
c) If the actual yield of DDT was 200.0 g what was the percent yield?
34. Nitrogen gas can be prepared by passing gaseous ammonia over solid copper (II) oxide at high temperatures.
The other products of this reaction are solid copper and water vapor. If a sample containing 18.1 g of ammonia
reacted with a 90.4 g of copper (II) oxide, how many grams of nitrogen will be formed?
2 NH3(g) + 3 CuO(s) → N2(g) + 3 Cu(s) + 3 H2O(g)
35. A sample of methane gas having a volume of 2.80 L at 25 °C and 1.65 atm was mixed with a sample of
oxygen gas having a volume of 35.0 L at 31 °C and 1.25 atm. The mixture was then ignited to form carbon
dioxide and water. Calculate the volume of carbon dioxide formed at a pressure of 2.50 atm and a temperature of
125 °C.
(R = 0.08206 l atm/mol K)