Download 432 Final Exam Study Guide

432 Final Exam Study Guide
Ch. 9:
Name ionic and covalent compounds
Write the formulas for compounds
Examples:
The correct name for the N─3 ion is the:
a. nitride ion
b. nitric ion
c. nitrate ion
d. nitrite ion
An -ite or –ate ending on the name of a compound indicates that the compound:
a. is a binary ionic compound
b. is a binary molecular compound
c. contains a polyatomic anion
d. contains a polyatomic cation.
The nonmetals in Groups 5A, 6A, and 7A:
a. lose electrons when they form ions
b. form positively charged ions
c. form ions with charges of 3─, 2─, and 1─, respectively.
d. Form ions with a numerical charge equal to their group number.
Name these compounds:
K3PO4 __________________________________
N2O5 ___________________________________
Write the formulas for these compounds:
potassium sulfide ____________________________
tin(IV) chloride _____________________________
hydrosulfuric acid ___________________________
calcium oxide ______________________________
dichlorine monoxide _________________________
Ch. 10:
Calculate moles, number of atoms, volumes, and masses from chemical formulas and
equations.
Calculate percent by mass
Determine empirical and molecular formulas
Examples:
.____1. How many moles of SO3 are in 2.4 x 1024 molecules of SO3?
a. 0.25
b. 3.4 x 1022
c. 4.0
d. 2.9 x 10─23
___2. The molar mass of molecular oxygen is:
a. equal to the mass of one mole of oxygen atoms
b. 16.0g
c. 32.0g
d. None of the above
___3. A large weather balloon filled with helium has a volume of 7.00 x 102 L at STP.
Which expression should be used to find the mass of helium in the balloon?
a. 22.4 L | 4g He
c. 22.4 L/mol | 4g He
mol | mol
7.0 x 102 L | mol
b. 7.00 x 102 | 4g He
L
| mol
d. 7.0 x 102 L | 1 mol | 4g___
| 22.4 L | mol
___4. What is the percent by mass of carbon in acetone, C3H6O?
a. 1.60%
b. 30.0%
c. 62%
d. 20.7%
___5. The empirical formula of a compound is CH2F. The molar mass of this
compound is 66.0g. The molecular formula of the compound is:
a. C4H8F4
b. C4H4F4
c. C2H4F2
d. CH2F
____6. The density of a gaseous compound is 1.623 g/L at STP. Determine the molar
mass of the compound.
a. 13.80 g
b. 7.246 x 10─2 g
c. 1.623 g
d. 36.36 g
____7. The representative particle for nitrogen is:
a. an atom
b. a molecule
c. a formula unit
d. none of these
___8. How many oxygen atoms are in 10 formula units of Al2(SO4)3?
a. 10 atoms O
c. 70 atoms O
b. 7 atoms O
d. 120 atoms O
__9. A compound has the empirical formula C2H3O and a molar mass of 172g.
What is its molecular formula?
a. C6H9O3
b. C8H12O4
c. C9H16O3
d. C7H8O5
____10. What is the volume in liters at STP of 2.50 mol of carbon monoxide?
a. 0.112 L
b. 3.10 L
c. 56.0 L
d. 8.96 L
___11. What amount in grams of Krypton (Kr) gas occupies 89.6 L?
Ch. 11
Balance chemical equations
Classify reactions, predict products
Examples:
____1. A double-replacement reaction takes place when aqueous K2SO4 reacts with
aqueous Pb(NO3)2. You would expect one of the products of this reaction to be:
a. K2S
b. NaPb
c. PbSO4
d. CNO3
____2. A chemical formula written above or below the yield sign indicates:
a. that a gas is formed.
b. That the substance is a catalyst.
c. That heat must be supplied.
d. A reversible reaction.
___3. In a double-replacement reaction:
a. one of the products is always gas.
b. One of the products must be an element.
c. Positive ions are exchanged between ionic compounds.
d. All of the above.
___4. The equation H3PO4 + KOH → K3PO4 + 3H2O is an example of which type of
reaction?
a. double-replacement
c. decomposition
b. combination
d. single-replacement
____5. Based on the activity series, the reaction most likely to occur from among the
following is:
a. Cu + HCL →
c. Pb + Na2SO4 →
b. Zn + HNO3 →
d. Ag + Cu(NO3)2 →
____6. If butane (C4H10) undergoes complete combustion:
a. 8 CO2 is one product
c. 9 O2 is one product
b. 8 CO is one product
d. 10 O2 is one product
___7. The complete combustion of octane (C8H18) would:
a. require 25 O2
c. Produce 18 H2O
b. produce 16 CO2
d. All of these
____8. When the reaction Fe + O2 → Fe2O3 is balanced, what is the coefficient
for Fe2O3 ?
a. 1
b. 2
c. 3
d. 4
____9. In the chemical equation 2H2O2 (aq) MnO2 → 2H2O (l) + O2 (g), the MnO2 is a:
a. reactant
c. spectator ion
b. product
d. catalyst
___10. In order for the reaction Cu + 2AgNO3 → Cu(NO3)2 + 2Ag to occur, which of
the following must be true?
a. Cu must be above Ag in the activity series.
b. Ag must be above Cu in the activity series.
c. Cu must be above H in the activity series
d. Ag must be above H in the activity series.
____Ca
+
____ NO2
____H3PO4
+
O2
→
____ AgNO3 + H2S →
____ Fe(NO3)3
→
Ca3(PO4)2
+ ____H2
____ N2O5
Ag2S + ____HNO3
+ ____ Na2CO3
→
Fe2(CO3)3
+ ____ NaNO3
Ch. 12
Determine stoichiometric conversions (mole-mole-, mass-mole, mass-volume, etc.)
Calculate percent yield
Examples:
____1. Which of these expressions is an incorrect interpretation of the
balanced equation?
2S(s) + 3O2(g) → 2SO3(g)
a. 2 atoms S + 3 molecules O2 → 2 molecules SO3
b. 2 g S + 3 g O2 → 2 g SO3
c. 2 mol S + 3 mol O2 → 2 mol SO3
d. None of the above
___2. In a chemical reaction, the mass of the products
a. is less than the mass of the reactanats
b. is greater than the mass of the reactants
c. is equal to the mass of the reactants
d. has no relationship to the mass of the reactants
____3. How many molecules of NO2 are produced when 2.0 x 1020
molecules of N2O4 are decomposed according to the following
equation?
N2O4 → 2 NO2
a. 4
b. 1.0 x 1020
c. 2.0 x 1020
d. 4.0 x 1020
____4. A reaction that has been calculated to produce 60.0g of CuCl2
actually produces 50.0g of CuCl2. What is the percent yield?
a. 0.833%
b. 96.1%
c. 83.3%
d. 120%
Questions 5 and6 refer to the following equation:
3 Cu + 8 HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
____5. Calculate the number of moles of water produced when 3.3 mol of
Cu(NO3)2 are formed in the reaction.
a. 4.4 mol
b. 6.6 mol
c. 4.9 mol
d. 8.8 mol
___6. How many grams of Cu would be needed to react with 2.0 mol HNO3 ?
a. 95.3 g
b. 63.5 g
c. 47.6 g
d. 1.50 g
___7. Given the reaction: Zn + 2HCL → ZnCl2 + H2
If 2.0 mol Zn and 5.0 mol HCl are allowed to react
a. Zn is the limiting reagent
b. HCl is the limiting reagent
c. 1.0 mol ZnCl2 is produced
d. 5.0 mol H2 is produced
8. What is the limiting reagent when 49.84 g of nitrogen react with 10.7 g
of hydrogen according to this balanced equation?
N2 + 3H2 → 2NH3
Ch. 13
Classify states of matter / properties of particles in different states
Examples:
____1. Sublimation can occur because solids have
A. a melting point
B. a vapor pressure
C. a boiling point
D. a triple point
____2.A decrease in the temperature of a contained liquid causes
A. an increase in the vapor pressure
B. a decrease in the vapor pressure
C. more particles to evaporate
D. an increase in the average kinetic energy of the particles
___3. Evaporation is type of
A. boiling
B. vaporization
C. sublimation
D. condensation
___4. Compared with gases, liquids:
A. have stronger intermolecular attractions.
B. have more space between their particles.
C. are much less dense.
D. are more easily compressed.
____5.A pressure of 2.5 atm is equivalent to:
A. 41 kPa
B. 1900 mm Hg
C. 3.0 x 102 mm Hg
D. 2.5 kPa
____6. According to kinetic theory, gases consist of particles that
A. occupy considerable volume
B. are relatively close together
C. exert attractive and repulsive forces on other particles
D. have motion that is constant, random, and rapid
___7. Water could be made to boil at 105°C by
A. applying a great deal of energy.
B. increasing the air pressure above the water.
C. heating the water more gradually.
D. decreasing the air pressure above the water.
___8. An increase in the temperature of a contained liquid
A. causes the vapor pressure above the liquid to increase
B. decreases the vapor pressure above a liquid.
C. causes fewer particles to escape the surface of the liquid.
D. has no effect on the kinetic energy of the liquid.
Ch. 14
Gas Laws and calculations
Examples:
___1. Which of these changes will not cause an increase in the pressure of a contained
gas?
A. Another gas is added to the container.
B. Additional amounts of the same gas are added to the container.
C. The temperature is increased.
D. The gas is moved to a larger container.
___2. The volume of a gas is doubled while the temperature is held constant. The
pressure of the gas ______.
A. remains unchanged
C. is doubled
B. is reduced by half
D. depends on the kind of gas
___3. What is the mass, in grams, of 0.125 L of CO2 at STP?
A. 0.246 g
B. 2.80 g
C. 181 g
D. 4.11 g
___4. How many moles of H2 would be contained in 4.0 L of the gas at 202.6 kPa and
127° C?
A. 89.6 mol
B. 6.38 mol
C. 0.24 mol
D. 0.77 mol
___5. If a sample of oxygen occupies a volume of 2.15 L at a pressure of 58.0 kPa and a
temperature of 25° C, what volume would this sample occupy at 101.3 kPa and
0°C?
A. 1.35 L
B. 1.13 L
C. 4.10 L
D. 3.44 L
___6. Which of the following would double the pressure on a contained gas at constant
temperature.
A. doubling the volume of the container
B. halving the number of particles in the container.
C. Doubling the number of particles in the container.
D. None of the above.
___7. A gas occupies a volume of 0.2 L at 25 kPa. What volume will the gas occupy at
2.5 kPa?
A. 4 L
B. 20 L
C. 2 L
D. 0.02 L
___8. The temperature of 6.24 L of a gas is increased from 25.0°C to 55.0°C at constant
pressure. The new volume of gas is ____.
A. 13.7 L
B. 5.67 L
C. 6.87 L
D. 2.84 L
___9. If the temperature of a gas in a closed container increases
A. the pressure of the gas decreased.
B. The average kinetic energy of the molecules decreases.
C. The molecules collide with the walls of the container less frequently.
D. The pressure of the gas increases.
10. A gas occupies a volume of 360 mL at 35.0° C and 191.8 kPa. What is the volume of
the gas at conditions of STP?
Ch. 15
Properties of water (polarity, surface tension, etc.)
Examples:
___1. What is the percent by mass of water in MgSO4•5H2O?
A. 51.1%
B. 195%
C. 56.0%
D. 42.8%
___2. Which of the following would dissolve readily in water?
A. CH4
B. H2
C. NaOH
D. CBr4
__3. An electric current is conducted by:
A. a solution of NaCl
B. a sugar solution
C. solid NaCl
D. solid sugar
___4. A liquid that has strong intermolecular attractions has:
A. a high surface tension.
B. An intermediate surface tension.
C. A low surface tension.
D. No surface tension.
___5. Which of these would you expect to be soluble in the nonpolar solvent carbon
disulfide, CS2?
A. MgCl2
B. CaCO3
C. CBr4
D. H2O
___6. The attractions between adjacent water molecules are called______.
A. hydrogen bonds
C. nonpolar covalent bonds
B. ionic bonds
D. polar covalent bonds
___7. How many water molecules are in two formula units of barium hydroxide
octahydrate, Ba(OH)2 • 8H2O?
A. 2
B. 8
C. 16
D. 20
___8. If salt is dissolved in water, water serves as the:
A. solute
B. solvent
C. dissolved medium
D. none of the above
___9. The density of ice is less than the density of water because:
A. ice has a lower molecular mass than water.
B. The same mass occupies a smaller volume.
C. The molecules are more closely packed.
D. Hydrogen bonding in ice produces an open framework
Ch. 16
Calculate the molality and molarity of solutions
Calculates the mass of solute in a molar solution
Examples:
________ 1. If the solubility of AgNO3 at 40°C is 311 g per 100 g of water, what mass
of this solute can be dissolved in 350 g of water at the same temperature?
a. 0.89 g
c. 1.1 g
b. 1100 g
d. 110 000 g
________ 2. Which of the following would increase the solubility of a gas in a liquid?
a. stirring the solution
b. increasing the temperature of the solvent
c. increasing the pressure of the gas above the solution
d. adding more solvent
________ 3. If the solubility of a gas in water is 1.22 g/L at 2.75 atm, what is its
solubility (in g/L) at 1.0 atm?
a. 0.44 g/L
c. 2.25 g/L
b. 3.97 g/L
d. 3.36 g/L
________ 4. The molarity of a solution that contains 14 g KOH per 150 mL of solution is:
a. 93M.
c. 0.093M.
b. 1.7M.
d. 11M.
________ 5. How many moles of solute are present in 1.25 L of a 0.75M NaNO3 solution?
a. 1.7 mol
c. 0.75 mol
b. 0.60 mol
d. 0.94 mol
________ 6. What volume of 1.25M HCl would be required to prepare 180 mL of a 0.500M HCl solution?
a. 450 mL
c. 0.014 mL
b. 72 mL
d. 2.2 x 10_3 mL
________ 7. What is the percent (m/m) of a solution containing 25 g of NaCl in 175 g of solution?
a. 700%
c. 7.0%
b. 0.14%
d. 14%
________ 8. What is the molality of a solution prepared by dissolving 13.0 g of Ba(NO 3)2 in 450 g of
water?
a. 0.029 molal
c. 0.11 molal
b. 29 molal
d. 1.1 x 10_4 molal
________ 9. The mole fraction of ethanol in a solution containing 1.50 moles of ethanol and 3.25 mol of
water is:
a. 0.316
c. 0.462
b. 0.217
d. 0.681
________ 10. What is the molality of an aqueous solution of a molecular solute if the boiling point of the
solution is 101.4°C? (Kb H2O = 0.512°C/m)
a. 2.70 molal
c. 1.04 molal
b. 0.515 molal
d. 0.556 molal
Ch. 19
Calculate pH and pOH
___1. A solution where the [ H+ ] = 10─4 M is ________:
a. basic
b. neutral
c. acidic
___2. According to the Bronstad-Lowry definition, a base is:
a. a substance that can donate an electron pair to the formation of a covalent
bond.
b. A substance that can accept a H+ ion from an acid.
c. A substance that increases the hydroxide ion concentration in water.
___3. If a solution has a pH 3, this solution is said to be:
a. acidic
b. basic
c. neutral
___4. Which of the following must be true if a solution is to be considered acidic?
a. [ H+ ] < [ OH─ ]
b. [ H+ ] > [ OH─ ]
c. [ H+ ] = [ OH─ ]
___5. If a solution has a pH 8.4, then the pOH is:
a. 14
b. 8
c. 8.4
d. 5.6
___6. What is the pH of a solution which has [ H+ ] of 1 x 10−10 ?
a. 1
b. 4
c. 10
d. 14
Terms to Know
Acid
Alkaline
Amphoteric
Anion
Aqueous solution
Arrhenius acid and base
Base
Binary compound
Boiling-point elevation
Bronsted-Lowry acid and base
Cation
Colligative properties
Combustion reaction
Concentrated solution
Concentration
Decomposition reaction
Dilute solutions
Diprotic acid
Double-replacement reaction
Electrolyte
Freezing-point depression
Henry’s Law
Immiscible
Law of Conservation of Matter
Lewis acid and base
Miscible
Mixture
Molality
Molarity
Molecule
Monoprotic acid
pH of solutions
Physical state of ionic compounds
Polyatomic ion
Polyatomic molecule
Precipitate
Precipitation reaction
Saturated solution
Single-replacement reaction
Spectator ion
Strong acid
Strong base
Supersaturated solution
Synthesis reaction
Unsaturated solution