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2015 Chemistry Semester 2 Final Exam Study Guide
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

The 2015 Semester 2 Cumulative Final Exam will be around 100 multiple choice questions.
A periodic table and calculator (chemical symbols – no names) will be provided.
NO GRAPHING CALCULATORS WILL BE ALLOWED.

Resources for studying
(1) Metacognitive Logs
(2) Old Quizzes
(3) Homework and Homework Videos
(4) Textbook
(5) Laboratories
(6) Handouts
(7) Study guides including your survival words

Study in a quiet place with few distractions over the next week. Read, write, and try sample problems.
Prequisite (Background) Knowledge from 1st semester:
Chapter
A.
B.
C.
6: The Periodic Table
Explain how elements are organized in a periodic table.
Describe the three broad classes of elements: metals, nonmetals, and metalloids.
Know the charges of elements in the s and p regions of the periodic table.
1.
Sulfur is a [metal / metalloid / non-metal] and it has a charge of ___________. It will form a [cation / anion ]
2.
Lithium is a [metal / metalloid / non-metal] and it has a charge of ___________. It will form a [cation / anion ]
3.
Hydrogen is a [metal / metalloid / non-metal] and it has a charge of ___________. It will form a [cation / anion ]
Chapter
A.
B.
C.
1.
2.
3.
4.
5.
6.
7.
8.
9.
Chapter
A.
B.
C.
D.
1.
2.
3.
7: Ionic and Metallic Bonding
Determine if an atom tends to gain or lose electrons.
Explain how cations and anions form.
Determine the chemical formula and formula unit for a compound.
How many valence electrons are in sulfur? _____ Lithium _____ Hydrogen _____
How does calcium obey the octent rule when reacting to form compounds>
a. It gains electrons
b. It gives up electrons
c. It does not change its number of electrons
d. Calcium does not obey the octet rule
What is the formula when potassium and sulfur combine?
What is the formula when magnesium and phosphorous combine?
What is the formula when Calcium and Sulfate combine?
What is the formula when Iron (III) and Nitrate combine?
What is the name of CuBr2?
What is the name of Mg3(PO4)2
Which of the following combinations would form an ionic bond?
a. Sodium and chlorine
b. Carbon and oxygen
c. Hydrogen and oxygen
d. Arsenic and bromine
8: Covalent (Molecular) Compounds
Determine what information a molecular formula provides.
Explain how molecular compounds are different from ionic compounds.
Define covalent bond, molecule, diatomic molecule, molecular compounds, and molecular formula.
Compare and contrast intermolecular attractions with ionic and covalent bonds.
How many atoms of each element are contained in Al2(SO4)3
Which of the following combinations would form a covalent bond?
a. Sodium and chlorine
b. Carbon and oxygen
c. Hydrogen and oxygen
d. Arsenic and bromine
Order the following ideas in terms of strength of attraction. (high to low)
1|Page
4.
Chapter
A.
B.
C.
D.
E.
F.
G.
H.
I.
J.
K.
a. Ionic Bond
b. Covalent Bond
c. Hydrogen Bond
d. Dispersion
Make a list of the diatomic molecules.
9 & 22.1: Naming Chemical Compounds
Name Monatomic and Polyatomic Ions
Determine how the charges of atoms relate to their position on the periodic table.
Determine how the charges of some transition metal ions are determined.
Name the most common suffixes used for polyatomic ions.
Explain how polyatomic anions differ from monatomic anions.
Determine how the names of binary ionic compounds are determined.
Write the formulas for binary ionic compounds.
Write the formulas and names of compounds containing polyatomic ions.
Explain what prefixes in the name of binary molecular compounds tell you about their composition.
Write the formula for a binary molecular compound.
Name a chemical compound using the proper guidelines.
____ 1. What type of ions have names ending in -ide?
a. only cations
c.
b. only anions
d.
only metal ions
only gaseous ions
____ 2. When Group 2A elements form ions, they ____.
a. lose two protons
c.
b. gain two protons
d.
lose two electrons
gain two electrons
____ 3. What is the correct name for the N
a. nitrate ion
b. nitrogen ion
nitride ion
nitrite ion
ion?
c.
d.
____ 4. When naming a transition metal ion that can have more than one common ionic charge, the numerical value of the charge is
indicated by a ____.
a. prefix
c. Roman numeral following the name
b. suffix
d. superscript after the name
____ 5. Aluminum is a group 3A metal. Which ion does A1 typically form?
a.
c.
Al
Al
b.
d.
Al
Al
____ 6. Which of the following correctly provides the name of the element, the symbol for the ion, and the name of the ion?
a.
c.
fluorine, F , fluoride ion
copper, Cu , cuprous ion
b.
d.
zinc, Zn , zincate ion
sulfur, S , sulfurous ion
____ 7. Which of the following correctly provides the names and formulas of polyatomic ions?
a.
carbonate: HCO ; bicarbonate: CO
b.
c.
d.
nitrite: NO ; nitrate: NO
sulfite: S
; sulfate: SO
chromate: CrO
; dichromate: Cr O
____ 8. An -ate or -ite at the end of a compound name usually indicates that the compound contains ____.
a. fewer electrons than protons
c. only two elements
b. neutral molecules
d. a polyatomic anion
____ 9. Which of the following compounds contains the Mn
ion?
a. MnS
c. Mn O
b.
MnBr
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d.
MnO
____ 10.Which of the following compounds contains the lead(II) ion?
a. PbO
c. Pb2O
b. PbCl4
d. Pb2S
____ 11.Which set of chemical name and chemical formula for the same compound is correct?
a. ammonium sulfite, (NH ) S
c. lithium carbonate, LiCO
b.
iron(III) phosphate, FePO
d.
magnesium dichromate, MgCrO
c.
d.
polyatomic ionic
binary molecular
____ 12.What type of compound is CuSO ?
a.
b.
monotomic ionic
polyatomic covalent
____ 13.Molecular compounds are usually ____.
a. composed of two or more transition elements
b. composed of positive and negative ions
c. composed of two or more nonmetallic elements
d. exceptions to the law of definite proportions
____ 19.What is the correct name for the compound CoCl ?
a.
b.
cobalt(I) chlorate
cobalt(I) chloride
c.
d.
cobalt(II) chlorate
cobalt(II) chloride
____ 20.What is the correct formula for barium chlorate?
a. Ba(ClO)
c. Ba(ClO )
b.
Ba(ClO )
d.
BaCl
____ 21.What is the correct formula for calcium dihydrogen phosphate?
a. CaH PO
c. Ca(H PO )
b.
Ca H PO
d.
Ca(H HPO )
____ 22.Which of the following is the correct name for N O ?
a.
b.
nitrous oxide
dinitrogen pentoxide
c.
d.
nitrogen dioxide
nitrate oxide
____ 23.How many double covalent bonds are in an alkane?
a. 0
c. 2
b. 1
d. 3
____ 24.How many carbons are in a molecule of hexane?
a. 3
c. 5
b. 4
d. 6
____ 25.What is the general formula for a straight-chain alkane?
a. C H
c. C H
b.
C H
____ 26.What is the name of BaF2?
a. Barium Fluorite
b. Barium Difluoride
____ 27.Ammonium Phosphide is which formula?
a. (NH4)3P
b. NH4P
c. (NH4)2PO4
d. (NO3)3P
e. none of these
3|Page
d.
C H
c.
d.
Barium Fluoride
Barium Fluorate
____ 28.Sodium Bromide is which of the following formulas?
a. Na3B
b. NaBr
c. NaBr2
d. NaBI
e. none of these
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
LiOH __________________________________________
PBr3 __________________________________________
Na2SO4 __________________________________________
(NH4)2S __________________________________________
CaCO3 __________________________________________
CF4 __________________________________________
NaNO3 __________________________________________
P2S3 __________________________________________
Al(NO3)3 __________________________________________
Mg(OH)2 __________________________________________
Write the formulas for the following compounds. Remember, they may be either ionic or covalent compounds, so make sure you use the
right method!
11)
12)
13)
14)
15)
16)
17)
18)
19)
20)
potassium oxide __________________________________________
phosphorus tribromide ________________________________________
calcium hydroxide __________________________________________
dinitrogen sulfide __________________________________________
carbon monoxide __________________________________________
diboron tetrahydride __________________________________________
phosphorus pentabromide _____________________________________
sulfur dichloride __________________________________________
sodium carbonate __________________________________________
aluminum acetate __________________________________________
Hydrocarbons

Because carbon has four valence electrons, carbon atoms always form four covalent bonds.

Molecules of hydrocarbons are nonpolar

Carbon atoms can form chains that are named using the following naming convention:
o
Prefix is determined by the number of carbons

1 = meth

2 = eth

3 = prop

4 = but

5 = pent

6 = hex

7 = hept

8 = oct

9 = non

10 = dec
o
Suffix is determined by the type of bond

Alkane CnH2n+2 (all bonds are single)

Alkene CnH2n (one bond is a double)

Alkyne CnH2n-2 (one bond is a triple)
Requisite Knowledge from 2nd semester:
Chapter
A.
B.
C.
D.
E.
F.
G.
H.
10: Chemical Quantities & the Mole
Describe methods of measuring the amount of something.
Define Avogadro’s number as it relates to a mole of a substance.
Distinguish between the atomic mass of an elements and its molar mass.
Describe how the mass of a mole of a compound is calculated
Describe how to convert the mass of a substance to the number of moles of a substance, and moles to mass. (Use the molar road)
Identify the volume of a quantity of gas at STP.
Describe how to calculate the percent by mass of an element in a compound.
Interpret an empirical formula.
4|Page
I.
Distinguish between empirical and molecular formulas.
1.
2.
What is the mass of 10.0 moles of Nitrogen molecules (N2)?
_____ Which one of the following contains the greatest number of moles of Sulfur atoms?
a. 1.0 mol S6
b. 3.0 mol SO3
c. 5.0 mol SO4
d. 3.0 mol P2S5
Adenine (C5H5N5) is an organic base found in the nucleotides of deoxyribonucleic acids (DNA). What is the molar mass of adenine?
How many Chlorine atoms are in 10 grams of Cl2? How many Hydrogen atoms are in 2 moles of Phosphine (PH3)?
Calculate the number of moles present in 2 grams of H2O.
How many Oxygen molecules are in 3.36 L of oxygen gas?
If there are 9.6 x 1015 molecules of sugar (C6H12O6) in a solution, then how many moles of sugar are there? What is the mass of
3.
4.
5.
6.
7.
8.
9.
the sugar in the solution?
Sucralose is a non-nutrative sweetener with the molecular formula C12H19Cl3O8. The molar mass of sucralose is 397.6 g/mol. What
is the percent composition by mass of Carbon in sucralose?
6 moles of an element have a mass of 312 grams. What is the element?
a. Zn
b. Cr
c. Fe
d. Al
A) Sodium Chloride
B) Diamond
C) Gold
D. Ice
10. Use the following classifications to label the four substances above:
a. Network Covalent Solid
b. Metallic Compound (solid)
c. Ionic Compound (solid)
d. Polar Covalent Compound with intermolecular forces of attraction (solid)
11. Which of the above substances has the highest melting points? Lowest melting points?
12. Which of the above substances has the highest boiling points? Lowest boiling points?
13. Which of the above substances is the most conductive of electricity?
14. Which of the above substances is the hardest? Softest?
15. [Solids / Liquids / Gases] exhibit the strongest intermolecular forces.
16. Arrange the following molecules in order of strongest to weakest intermolecular forces: C (s), Hg (l), I2 (g)
17. Which of these pairs of molecules forms a hydrogen bond?
a. NH3 & H2O
b. CH4 & H2O
c. C2H6 & C3H8
d. NaCl & H2O
e. O2 & H2
18. What two factors must you find in each polar molecule in order for hydrogen “bonds” to be present?
19. A mole of Magnesium Bromide contains 6.02 x 1023
a. atoms.
b. molecules.
c. formula units.
d. ions.
20. A mole of Copper contains 6.02 x 1023
a. atoms.
b. molecules.
c. formula units.
d. ions.
5|Page
21. A mole of methane contains 6.02 x 1023
a. atoms.
b. molecules.
c. formula units.
d. ions.
22. Nicotine, a compound found in cigarettes, is 74.04% carbon, 17.27% nitrogen, and 8.7% hydrogen by mass. What is the empirical
formula of nicotine?
23. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. One of them is a blue solid that comes
apart, reversibly, in the gas phase. It contains 36.84% N. What is the empirical formula of this oxide? What is the molecular
formula is the molecular mass is 456 grams/mole?
24. A sample of indium chloride weighing 0.5000 g is found to contain 0.2404 g of chlorine. What is the empirical formula of the
indium compound?
25. Given the structural formula to the right, what is the name of the following hydrocarbon?
26. What is the name of
C2H6? C6H12? C8 H14?
27. What is the chemical formula for hexyne? octane? ethene?
28. Hydrocarbons are [polar / nonpolar / unable to be determined]
29. The type of intermolecular forces of attraction that exist between hydrocarbons are
called
a. Dispersion
b. Dipole-dipole
c. Hydrogen bonds
30. H2, N2, O2, F2, I2, and Cl2 are all called ___________________ . All of these molecules are [polar / nonpolar / unable to be
determined].
31. The type of intermolecular forces of attraction that exist between diatomics are called
a. Dispersion
b. Dipole-dipole
c. Hydrogen bonds
32. Why would a hydrogen bond be formed between two molecules?
a. The positive end of one molecule is attracted to the negative end of another molecule
b. The positive end of one molecule repels the negative end of another molecule.
c. The nonpolar end of one molecule is attracted to the polar end of another molecule
d. The nonpolar end of one molecule is attracted to the nonpolar end of another molecule.
Chapter
A.
B.
C.
D.
E.
F.
G.
H.
I.
J.
K.
L.
33.
34.
35.
36.
11: Chemical Reactions
Describe and write a word equation.
Describe and write a skeleton equation (unbalanced equation).
Describe and write a balanced chemical equation (including proper labels for solids, liquids, gases, and aqueous solutions).
Describe, identify, and predict the products of the five general types of reactions (synthesis, decomposition, single replacement,
double replacement, and combustion)
Explain the law of conservation of mass.
Understand and determine the reactivity of halogens.
Understand and determine the activity (reactivity) of metals.
Describe the information found in a net ionic equation.
Predict the formation of a precipitate in a double replacement reaction using the solubility rules (provided). Identify and explain the
significance of spectator ions.
Explain why chemical reactions occur.
Know and use the activity series of metals
Describe exothermic and endothermic reactions.
What are the symbols for solid, liquid, gas, and aqueous in a chemical reaction?
What changes during a chemical reaction?
What remains the same during a chemical reaction?
Label the following reactions as Synthesis (S), Decomposition (D), Single-Replacement (SR), Double Replacement (DR), or
Combustion.
a. _____ PbO2  PbO + O2
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b.
_____ A + Z  AZ
c.
_____ Fe(OH)3  Fe2O3 + H2O
d.
_____ CxHy + O2  H2O + CO2
e.
_____ H2SO4 + BaCl2  BaSO4 + HCl
f.
_____ MQ + R  RQ + M
g.
_____ Hf + N2  Hf3N4
h.
_____ Mg + H2CO3  MgCO3 + H2
i.
_____ UT3  U2 + T
j.
_____ C12H22O11 + O2  CO2 + H2O
37. Balance the following equations
a. ___Mg + ___H2SO4  ___MgSO4 + ___H2
b.
___C2H6 + ___O2  ___CO2 + ___H2O
c.
___Pb(NO3)2 + ___NaI  ___PbI2 + ___ NaNO3
d.
___Fe + ___O2  ___Fe3O4
e.
___H2 + ___O2  ___H2O
f.
___C3H8 + ___O2  ___CO2 + ___H2O
38. Write the balanced chemical equations for the following reactions. Write NR if there is no reaction.
a.
Bubbling chlorine gas through a solution of potassium iodide gives elemental iodine and a solution of potassium chloride.
b.
Solid tetraphosphorus decoxide reacts with water to produce phosphoric acid.
c.
Octane gas reactions with oxygen gas in a combustion reaction.
d.
Zinc sulfide is added to sulfuric acid in an aqueous solution.
e.
Solutions of potassium fluoride and calcium nitrate are mixed.
f.
A piece of steel wool (iron) is placed in sulfuric acid (H2SO4).
g.
Mercury is poured into an aqueous solution of zinc nitrate.
h.
Bromine reacts with aqueous barium iodide.
i.
Aluminum is obtained from aluminum oxide with the addition of a large amount of electrical energy.
j.
Heating tin (IV) hydroxide gives tin (IV) oxide and water.
39. Predict the products and then balance the equations:
a.
___Ca (s) + ___S (s) 
b.
___Zn (s) + ___AgNO3 (aq) 
c.
___Al (s) + ___H2SO4 (aq) 
d.
___Al (s) + ___H2O (l) 
e.
___CdBr2 (aq) + ___Na2S (aq) 
f.
___Cl2 + ___NaF 
g.
___Cl2 + ___NaBr 
h.
___H2O2 
i.
___NH3 
40. Write a balanced net ionic equation for the following reactions:
a. ___HCl (aq) + ___ Ca(OH)2 (aq) 
b.
___AgNO3 (aq) + ___AlCl3 (aq) 
Chapter
A.
B.
C.
D.
12: Stoichiometry
Interpret balanced chemical equations in terms of moles, representative particles, mass, and gas volume at STP.
Identify the quantities that are always conserved in chemical reactions.
Construct mole ratios from balanced equations and apply these ratios in stoichiometric calculations. (switcheroo)
Calculate stoichiometric quantities from balanced chemical equations using units of moles, mass, representative particles, and
volumes of gases at STP.
E. Identify and use the limiting reagent in a reaction to calculate the maximum amount of product(s) produced and the amount of excess
reagent that remains unreacted.
F. Calculate theoretical yield, actual yield, percent yield, or percent error given appropriate information.
G. Identify and describe the function of a catalyst.
41. What changes during a chemical reaction?
42. What remains the same during a chemical reaction?
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43. Interpret the given equation in terms of relative numbers of representative particles, numbers of moles, and masses of reactants
and products.
2K (s) + 2H2O (l) 2KOH (aq) +H2 (g)
44. Balance the equation: C2H5OH (l) + O2 (g)  CO2 (g) + H2O (g). Show the balanced equation obeys the law of conservation of
mass.
45. Acetylene gas (C2H2) and Calcium hydroxide are produced by adding water to calcium carbide (CaC2). Write the balanced
chemical equation. How many grams of acetylene are produced by adding excess water to 5.00 grams of calcium carbide?
46. Using the equation you balanced in the problem above, determine how many moles of calcium carbide are needed to react
completely with 49.0 grams of H2O.
47. The last step in the production of nitric acid is the reaction of nitrogen dioxide with water.
3NO2 (g) + H2O (l)  2HNO3 (aq)
+ NO (g) How many grams of nitrogen dioxide must react with water to produce 5.00 x 1022 molecules of nitrogen monoxide?
48. How are mole ratios used in chemical calculations?
49. Write a sequential list of steps required to solve most typical stoichiometric problems.
50. Write the 12 mole ratios that can be derived from the equation for the combustion of isopropyl alcohol. 2C3H7OH (l) + 9O2 (g) 
6CO2 (g) + 8H2O (g)
51. Write the equation for the complete combustion of ethane. If 2.7 mol of ethane is reacted with 6.3 mole of oxygen gas, identify
the limiting reagent. Calculate the number of moles of water produced.
52. When 84.8 grams of iron (III) oxide reacts with an excess of carbon monoxide, iron and carbon dioxide are produced. Write the
balanced chemical equation. What is the theoretical yield of this reaction for each product?
53. If 50.0 grams of silicon dioxide is heated with an excess of carbon, 27.9 grams of silicon carbide is produced. SiO2 (s) + 3C (s) 
SiC (s) + 2CO (g) What is the percent yield of this reaction?
54. If you were given 12 grams of Mg and excess amounts of 6 molar hydrochloric acid, exactly how much hydrochloric acid would
you need to react all of the magnesium? What is the balanced chemical equation for this balanced chemical reaction?
Chapter
A.
B.
C.
D.
E.
F.
G.
H.
I.
J.
K.
17: Thermochemistry
Explain how energy, heat, and work are related.
Classify processes as either exothermic or endothermic.
Identify the units used to measure heat transfer. (Joules, KJ, calories, Kilocalories, Calories)
Distinguish between heat capacity and specific heat.
Describe how calorimeters are used to measure heat flow.
Construct thermochemical equations.
Solve for enthalpy changes in chemical reactions by using heats of reaction.
Classify the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves. (Heating curve)
State Hess’s law of heat summation and describe how it is used in chemistry.
Solve for enthalpy changes by using Hess’s law or standard heats of formation.
55. If a steak is removed from the freezer and then submerged in hot water, then the
a. temperature of the water will increase
b. water will begin to freeze
c. temperature of the steak will decrease and the water’s temperature will increase
d. temperature of the water will decrease and the steak’s temperature will increase
56. When a chemist holds a flask that has an exothermic reaction in it, the substance inside the flask will __________ and the chemist’s had
will feel _______________.
a. Lose heat; hot
b. Gain heat; hot
c. Lose heat; cold
d. Gain heat; cold
57. The heat lost by the surroundings is always
a. equal to the heat gained by the system
b. equal to the heat lost by the system
c. greater than the heat gained by the system
d. cannot be determined
58. What is the amount of heat required to raise the temperature of 500.0 g of aluminum by 20 C if the specific heat capacity is 0.2
59. What is the specific heat for Substance Z considering that 100 grams of Z absorbed 500 calories of heat resulting in a temperature
change of 30C?
60. How much energy is released with 24 g of Na2O reacts in the following reaction?
Na2O (s) + 2HI (g)  2NaI (s) + H2O (l) H -120 Cal
61. During a phase change
a. The temperature of a substance changes
b. The space between the particles changes
c. The speed of the particles remains the same
d. The energy in the particles remains the same
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?
62. [T/F] The heat of solution can be positive, negative, or zero.
63. When 50 grams of solid sodium hydroxide are dissolved in water, how much heat is released? (Hsolution = -445.1 Kj/mol)
64. Using a table that lists the standard heats of formation, you can calculate the change in enthalpy for a given chemical reaction. The
change in enthalpy is equal to
a. Hf of products minus Hf of reactants
b. Hf of products plus Hf of reactants
c. Hf of reactants minus Hf of products
d. Hf of reactants plus Hf of products
65. Calculate the H using the standard heats of formation for the following reactions:
a. 2H2O2 (l)  2H2O (l) + O2 (g)
b. 4NH3 (g) = 5O2 (g)  4NO (g) + 6H2O (g)
Substance
Hf (kJ/mol)
H2O (g)
-241.8
H2O (l)
-285.8
H2O2 (l)
-187.8
NH3 (g)
-46.19
NO (g)
90.37
O2 (g)
0
66. Using the average bond energies, calculate the enthalpy change in the following reactions
a. 2H2O2 (l)  2H2O (l) + O2 (g)
b. 4NH3 (g) = 5O2 (g)  4NO (g) + 6H2O (g)
Average Bond Energies (kJ/mol)
H– H
H– Cl
436 kJ/mol
431 kJ/mol
C– H
C– C
413 kJ/mol
348 kJ/mol
H–
N–
N–
O–
O–
F–
567
391
201
463
146
155
C– N
C– O
C– F
C– Cl
C– S
Cl– Cl
293
358
485
328
259
242
F
H
O
H
O
F
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
kJ/mol
C=C
C=C
C=O
O=O
C=O
C=N
N=N
N=N
C=N
614 kJ/mol
839 kJ/mol
799 kJ/mol
495 kJ/mol
1072 kJ/mol
615 kJ/mol
418 kJ/mol
941 kJ/mol
891 kJ/mol
67. Calculate the enthalpy change (H) for the formation of nitrogen monoxide from its elements.
N2 (g) + O2 (g)  2NO (g)
Use these thermochemical equations and Hess’s law:
4NH3 (g) + 3O2 (g)  2N2 (g0 + 6H2O (l) H = -1.53 x 103 kJ
4NH3 (g) = 5O2 (g)  4NO (g) + 6H2O (g) H = -1.17 x 103 kJ
68. According to the chart below, the material requiring the most energy to raise its temperature from 10C to 50C is _______? This
means its resistance to change temperature is [high / low / average]
Material
water
69.
Specific heat (J/kgºC)
4,184
aluminum
900
steel
silver
470
235
oil
concrete
1,900
880
glass
gold
wood
800
129
2,500
Consider the following heating curve:
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a.
b.
c.
d.
e.
f.
g.
h.
i.
Which section represents the solid?
Which section represents the gas?
Which section represents the liquid?
Which section represents the melting/freezing point?
Which section represents the boiling/condensing point?
In which sections is the temperature stays the same?
In which sections is the temperature changing?
In which section is the equation q=mcT used?
In which section is a conversion factor used?
70. How many calories are required to raise the temperature of a 50 gram sample of iron from 25C to 45C if iron has a specific heat
capacity of 0.11
?
71. How many calories are required to raise the temperature of 30 grams of ice from a temperature of -20C to 150C?
CH20 (g) =0.48
; CH20 (s) =0.50
; CH20 (l) =1.0
; Hvaporization = 540
; Hfusion = 800
72. It takes
a.
b.
c.
200 calories of heat to raise the temperature of 14 grams of canola oil by 30C.
Calculate the specific heat capacity of canola oil.
How does the specific heat capacity of canola oil compare to that of water (is it higher, lower, equal)?
If you have 100 g of canola oil and 100 g of water and they are heated by the same source (both are absorbing equal
amounts of heat), which substance will reveal the greatest change in temperature during 10 minutes of heating? Explain your
reasoning?
73. A 5 gram sample of Ben and Jerry’s Chubby Hubby Ice Cream is placed in a bomb calorimeter with 500.00 grams of water at an initial
temperature of 25C. After combustion of the ice cream the water has a temperature of 52C.
a. How much energy was absorbed by the water?
b. How much energy was released by the ice cream?
c. What is the caloric value in calories/gram?
74. A chunk of ice whose temperature is -40C is added to an insulated cup filled with water at 0C.
a. In what direction does heat flow?
b. Which of the following statements is true? Explain why!!
i. The ice melt until it reaches the temperature of the water.
ii. The liquid water cool until it reaches the temperature of the ice.
iii. The liquid water freezes so the chunk of ice gets larger.
Chapter
A.
B.
C.
D.
18: Reaction Rates and Equilibrium
Describe how to express the rate of a chemical reaction
Identify four factors that influence the rate of a chemical reaction and how they speed up or slow down a reaction.
Be able to describe the role of a catalyst or inhibitor in a chemical reaction.
Interpret the hills and valleys in a reaction progress curve.
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75. Which expression represents a reaction rate?
a. time/mass
c. energy/time
b. number/time
d. time/energy
76. At what stage of a reaction do atoms have the highest energy?
a. reactant stage
b. product stage
c. transition state stage
d. The stage of highest energy depends on the atom.
77. Why does a higher temperature cause a reaction to go faster?
a. There are more collisions per second only.
b. Collisions occur with greater energy only.
c. There are more collisions per second and the collisions are of greater energy.
d. There are more collisions per second or the collisions are of greater energy.
78. Why does a higher concentration make a reaction faster?
a. There are more collisions per second only.
b. Collisions occur with greater energy only.
c. There are more collisions per second and the collisions are of greater energy.
d. There are more collisions per second or the collisions are of greater energy.
79. What happens to a catalyst in a reaction?
a. It is unchanged.
c.
It is incorporated into the reactants.
b. It is incorporated into the products.
d.
It evaporates away.
80. A catalyst works by ____.
a. lowering the activation energy barrier
b. shifting the equilibrium position toward the products
c. changing the temperature of the reactants
d. changing the particle size of the reactants
81. The rate of a chemical reaction normally ____.
a. decreases as temperature increases
b. is slowed down by a catalyst
c. increases as reactant concentration increases
d. decreases as reactant concentration increases
82. T/F Surface are is high if the substance is crushed.
83. If the particles size is small, the surface area tends to be [high / low / average]
84. When the surface area is high, the speed of a reaction will be [fast / flow / average]
85. When the temperature of a reaction is high, the speed of the reaction will be [fast / flow / average]
86. When a lot of energy is provided to a reaction, the speed of the reaction will be [fast / flow / average]
87. When the concentration of the reactants is high, the speed of the reaction will be [fast / flow / average]
88. When the number of collisions between reactants is high, the speed of the reaction will be [fast / flow / average]
A
B
C
E.
F.
89.
Which of the activation energy graphs demonstrates a slow exothermic reaction at room temperature?
a.
A
c.
C
b. B
d.
none of the above
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You will be provided a clean copy of the following (unless otherwise specified) with your test:
Reactivity
Most Reactive
Least Reactive
Activity Series for METALS
Name
Symbol
Notes
Lithium
Li
Will replace H in
acids (compounds
Potassium
K
starting with H) and
Calcium
Ca
Water.
Sodium
Na
Magnesium
Mg
Will replace H in
acids only.
Aluminum
Al
Zinc
Zn
Iron
Fe
Lead
Pb
Hydrogen
H+
Copper
Cu
Mercury
Hg
Silver
Ag
Solubility Rules for Ionic Compounds
Salts of Alkali (Group I) metals and ammonium (NH4+1)
Nitrate (NO3-1) salts and Chlorate (ClO3-1) salts
Sulfate (SO4-2) salts
EXCEPT compounds with Pb+2, Ag+1, Hg2+2, Ba+2, Sr+2, and Ca+2
Chloride (Cl-1) salts
EXCEPT compounds with Ag+1, Pb+2, and Hg2+2
Carbonates (CO3-2), Phosphates (PO4-3), Chromate (CrO4-2), Sulfides
(S-2), and Hydroxide (OH-1) (Except compounds with group I
metals)
Activity Series for HALOGENS
Reactivity
Name
Symbol
Most
Reactive
Fluorine
F
Chlorine
Cl
Bromine
Br
Iodine
I
Least
Astatine
At
Reactive
You must memorize these
Soluble
(aq)
Soluble
(aq)
Soluble
(aq)
Soluble
(aq)
Insoluble
(s)
 actual yield
Percent Yield  
 theoretical yield

 x100

You must memorize the polyatomic ions
on your green sheet
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