Chapter 1 Study Questions
1. Indicate the metric unit for: a) mass, b) length, and c) volume
2. The time is recorded from three different clocks as indicated below. The “true” time is exactly 8:30 A.M.
a) Which of the three clocks is the most precise? b) Which clock is the most accurate?
Measurement / Clock A / Clock B / Clock C1 / 8:25.20 AM / 8:29 AM / 8:36 AM
2 / 8:25.00 AM / 8:31 AM / 8:36 AM
3 / 8:25.10 AM / 8:30 AM / 8:36 AM
c) Which clock(s) show a systematic error?
3. Indicate the number of significant figures in the following numbers:
a) 2,348b) 7.0001 c) 0.0023
d) 24,500e) 0.1060
4. Perform the following operations and express the answers in significant figures:
a) 1.24 x 8.2 = b) 6.78 - 3.3 =
c) 9.999 + 0.22 =d) (5.67 x 103) x (2.1 x 102)
5. Express the following numbers in scientific notation:
a) 650 (2 sig fig)b) 0.0005 (1 sig fig)c) 207,000 (3 sig fig)
6. Bozo determined the density of a sample of aluminum. For his sample, he found the volume was 0.350 cm3 and the mass was 0.822 g. Given that the density of aluminum is 2.70 g/cm3, calculate Bozo’s percent accuracy error.
NOTE: Use dimensional analysis (conversion factors) to answer the problems below. Answers must be in significant figures, include units and show work. Use the table on the inside cover of the back of the text as needed.
7. Find the mass in pounds (lbs) of a 275-gram sample of sugar.
8. Find the number of cm in 0.286 miles.
9. Find the volume in quarts of 10.7 kg of iron. The density of iron is 7.87 g/cm3.
10. Convert the density of ethanol (0.789 g/cm3) into units of pounds/liter.
11. Give two common examples of each of the following: a pure substance, a mixture, a solution, an element, a compound. Are your mixtures homogeneous or heterogenous?
12. List two chemical properties and two physical properties of the element magnesium. (You may use your textbook.)
13. (OPTIONAL) Assuming each ant is 5.0 mm long, how many ants would it take to make a line, single file, from one end to the other of a 100-yard football field? (2 sig fig)
Chapter 2 Study Questions
1. Which law is supported by each of the following statements? (The laws are: conservation of mass, definite proportion, and multiple proportions.)
a) In hydrogen peroxide there are 15.9 grams of oxygen per 1.00 g of hydrogen and in water there are 7.94 grams of oxygen per 1.00 g of hydrogen.
b) The total mass of reactants (starting materials) is the same as the total mass of products when a chemical reaction is carried out in a closed system.
c) In any sample of a given compound, the mass proportion of each element is the same.
2. What are two conclusions supported by Rutherford’s experiment?
3. Fill in the following table:
Nuclear Atomic Mass Number of Number of Number of Charge
Symbol Number Number Protons Electrons Neutrons
______
______39 19 18 ______
______16 ______20 -2
4. Write the nuclear symbols for the isotopes of neon which contain 10 neutrons and 12 neutrons.
5. For each of the following elements, indicate whether it is a main group element (MG), transition metal (TM), or inner transition metal (ITM). If the element is a main group element, indicate the group number and whether it is a metal, a nonmetal or a metalloid. Also indicate the Period of each element.
a) Sr (atomic # 38)b) Br (atomic # 35)c) Mo (atomic # 42)
d) P (atomic # 15)e) B (atomic # 5)f) U (atomic # 92)
g) Sn (atomic # 50)h) Hg (atomic # 80)
6. Provide the common names of Groups 1, 2, 17 and 18.
7. Give an example of at least one element made up of molecules and one compound made up of molecules.
8. For each of the following atoms, indicate whether it forms a positive or a negative ion, and include the ion charge.
a) Nab) Bac) Cld) Se) Ag
9. Which of the following are ionic compounds? Which are covalent compounds? Name each compound.
a) N2Ob) K2Oc) PCl3d) AlPO4
e) HClf) NH4Fg) Pb(NO2)2h) H2SO3
10. Name the following ionic compounds:
a) CaCO3b) ZnSc) CuOHd) Mg(ClO4)2
11. Give the formulas for the following ionic compounds:
a) potassium phosphateb) ammonium sulfate
c) cobalt(II) hydroxided) iron(III)nitride
12. Provide the formulas for the following covalent compounds:
a) phosphorus triiodideb) dinitrogen pentoxidec) chloric acid
Chapter 3 Study Questions
1. Glycerol (C3H8O3) is sold in drug stores as glycerine and is commonly found in soaps and shampoos.
a) What is the molar mass of glycerol?
b) What is the mass in grams of 1.00 mole of glycerol?
c) How many molecules are in one mole of glycerol?
d) How many grams are in 0.217 moles of glycerol?
2. Ammonia (NH3) is the active ingredient in many kitchen cleansers. How many atoms are in
a) one molecule of ammonia?
b) one mole of ammonia?
c) 3.40 grams of ammonia?
3. Sodium nitrite is a controversial food preservative added to processed meat and thought to form cancer-causing compounds when heated. What are the mass percentages of each element in sodium nitrite?
4. A compound consists of 40.7% C, 5.1% H, and 54.2% O?
a) What is its empirical formula?
b) The molar mass of this compound is 118 grams/mole. What is the molecular formula of this compound?
5. A 25.0 gram sample of a compound made up of magnesium, carbon and oxygen contains 7.20 grams magnesium and 3.55 grams carbon.
a) Find the empirical formula of this compound.
b) Find the mass percentage of each element in this compound.
c) What is the mass of magnesium in a 13.9 gram sample of this compound?
d) What is the mass of this compound that contains 0.290 moles of carbon?
6. A sample of zinc is heated in air to form zinc oxide. Assuming all of the zinc is converted to the oxide, use the data table below to calculate the empirical formula of zinc oxide.
mass of crucible= 32.00 g
mass of crucible + zinc (before heating)= 33.64 g
mass of crucible + oxide (after heating)= 34.04 g
7. Balance the following equations:
a) the combustion of the rocket fuel diborane,
B2H6(l) + O2(g) B2O3(s) + H2O(l)
b) the combustion of the poisonous gas, PH3,
PH3(g) + O2(g) H2O(l) + P4O10(s)
8. Write a balanced equation for each of the following reactions:
a) the reaction of solid lithium with nitrogen to form solid lithium nitride.
b) the reaction between aqueous solutions of cobalt(III) nitrate and sodium hydroxide to form aqueous sodium nitrate and solid cobalt(III) hydroxide.
c) the reaction between solid zinc and aqueous hydrochloric acid in a single replacement reaction.
d) classify the reactions in (a) and (b).
9. Hydrogen sulfide, given off by decaying organic matter, is converted to sulfur dioxide in the atmosphere by the reaction:
2 H2S(g) + 3 O2(g) 2 SO2(g) + 2 H2O(l)
a) How many moles of H2S are required to form 8.20 moles of SO2?
b) How many grams of O2 are required to react with 1.00 mole of H2S?
c) How many grams of water are produced from 6.82 g H2S?
d) If 12.0 grams of SO2 are formed from 7.98 g of H2S, what is the percent yield?
e) How many grams of SO2 are produced starting from 2.66 g H2S and 3.00 g O2?
Which reactant is limiting?
10. A gaseous mixture containing 7.50 mol H2(g) and 9.00 mol Cl2(g) reacts to form hydrogen chloride (HCl) gas.
a) Write a balanced equation for the reaction.
b) Which reactant is limiting?
c) If all the limiting reactant is consumed, how many moles of hydrogen chloride are formed?
d) How many moles of the excess reactant remain unreacted?
Chapter 4 Study Questions
1. Classify each of the following substances as: 1) acid, base, or neutral, and 2) strong or weak. Then 3) write a balanced equation for the ionization of the substance in water:
a) HNO3b) HClO c) NH3d) NaNO3e) Ba(OH)2
2. A common method of preparing solutions is to make up a concentrated solution and then dilute it to the desired concentration.
a) What is the molarity of a solution prepared by dissolving 29.2 g NaCl in enough water to make 0.250 liters of solution?
b) What volume of the above solution is needed to make 125 ml of a 0.350 M NaCl solution?
3. What mass of glucose (C6H12O6) is needed to prepare 200.0 ml of a 2.50 M glucose solution?
4. Solid magnesium is added to 125 ml of 2.00 M hydrochloric acid to produce dissolved magnesium chloride and hydrogen gas.
a) Write a balanced equation for this reaction.
b) If excess magnesium is added, how many moles of hydrogen gas are produced?
5. Indicate whether a precipitate will form when the following solutions are mixed. If a precipitate forms, write a net ionic equation for the reaction.
a) iron(III) nitrate and potassium hydroxide
b) ammonium chloride and lithium carbonate
c) sodium sulfide and nickel(II) sulfate
6. Name two solutions which could be mixed to form strontium sulfate.
7. Write a balanced netionic equation for the acid-base reaction between HNO3 with KOH.
8. How many ml of 2.00 M NaOH would be required to neutralize 12.5 ml of 0.0800 M HBr?
9. When solutions of lead(II) nitrate and aluminum chloride are mixed, a precipitate forms.
a) Write a balanced formula equation for the reaction.
b) What volume of a 0.200 M lead(II) nitrate solution is needed to completely form a precipitate when added to 2.48 mL of 0.300 M aluminum chloride?
c) What is the mass of precipitate formed in (b)?
10. What mass of precipitate is formed when 71.3 mL of 0.500 M iron(III) nitrate are mixed with 112 mL of 0.800 M sodium carbonate?
11. How many ml of 2.50 M HNO3 contain enough nitric acid to dissolve an old copper penny with a mass of 3.94 grams?
3 Cu(s) + 8 HNO3(aq) 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O
Chapter 5 Study Questions
1. A sample of air collected at STP contains 0.039 moles of N2, 0.010 moles of O2, and 0.001 moles of Ar. (Assume no other gases are present.)
a) Find the partial pressure of O2.
b) What is the volume of the container?
2. A sample of hydrogen gas (H2) is collected over water at 19C.
a) What are the partial pressures of H2 and water vapor if the total pressure is 756 mm Hg?
b) What is the partial pressure of hydrogen gas in atmospheres?
3. If 600. cm3 of H2 at 25C and 750. mm Hg is compressed to a volume of 480. cm3 at 41C, what does the pressure become?
4. Find the density of helium gas at STP.
5. a) Write a balanced chemical equation for the reaction of butane gas with oxygen gas to form carbon dioxide and water vapor.
b) How many liters of oxygen are required to produce 2.0 liters of CO2?
c) How many liters of CO2 are produced from 11.6 g of butane at STP?
d) How many molecules of water vapor are produced from 5.6 liters of butane gas at STP?
6. Find the molar volume of a gas at 68C and 2.00 atmospheres pressure.
7. How many liters of methane are there in 8.00 grams at STP?
8. Calculate the density of carbon dioxide at 546 K and 4.00 atmospheres pressure.
9. What volume of O2 at 710. mm Hg pressure and 36C is required to react with 6.52 g of CuS?
CuS(s) + 2 O2(g) CuSO4(s)
10. What is the molar mass of a gas if 7.00 grams occupy 6.20 liters at 29C and 760. mm Hg pressure?
11. At a particular temperature and pressure, 15.0 g of CO2 occupy 7.16 liters. What is the volume of 12.0 g of CH4 at the same temperature and pressure?
12. To prepare a sample of hydrogen gas, a student reacts 7.78 grams of zinc with acid:
Zn(s) + 2 H+ (aq) Zn2+(aq) + H2(g)
The hydrogen is collected over water at 22C and the total pressure of gas collected is 750. mm Hg. What is the partial pressure of H2? What volume of wet hydrogen gas is collected?