CHEMISTRY 11 REVIEW
Appendix D
D1 Matter and Chemical Bonding
(a)
SUMMARY
(b)
most active
most active
lithium
mass number (number of
protons plus neutrons)
(Z + N)
A
Z
potassium
X
fluorine
barium
chlorine
calcium
bromine
sodium
iodine
magnesium
least active
aluminum
zinc
iron
atomic number
(number of protons)
nickel
tin
Figure 1
Symbolism representing an individual atom of an element
lead
hydrogen
copper
Ionization energy, electron affinity, and electronegativity increase.
silver
gold
least active
Ionization energy,
electron affinity,
and electronegativity
increase.
Atomic radius
and metallic
properties
increase.
Atomic radius and metallic properties increase.
Figure 2
Trends in periodic properties
Figure 3
(a) In the activity series of metals, each metal
will displace any metal listed below it.
Hydrogen is usually included in the series,
even though it is not a metal, because
hydrogen can form positive ions, just like
the metals.
(b) The halogens can also be ordered in an
activity series.
Table 1 Summary of Bonding Characteristics
Intramolecular force Bonding model
ionic bond
• involves an electron transfer, resulting in the formation of cations
and anions
• cations and anions attract each other
polar covalent bond
• involves unequal sharing of pairs of electrons by atoms of two
different elements
• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest),
double, or triple (strongest) bonds
nonpolar
covalent bond
• involves equal sharing of pairs of electrons
• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest),
double, or triple (strongest) bonds
806 Appendix D
NEL
Appendix B
Table 2 Summary of Reaction Type Generalizations
Reaction type
Reactants
Products
combustion
metal + oxygen
nonmetal + oxygen
fossil fuel + oxygen
metal oxide
nonmetal oxide
carbon dioxide + water
synthesis
element + element
element + compound
compound + compound
compound
more complex compound
more complex compound
decomposition
binary compound
complex compound
element + element
simpler compound + simpler compound
or
simpler compound + element(s)
single displacement
A + BC
B + AC
double displacement
AB + CD
AD + CB
D
Table 3 Classical and IUPAC Names of Common Multivalent Metal Ions
Metal
Ion
Classical name
IUPAC name
iron
Fe2
Fe3
ferrous
ferric
iron(II)
iron(III)
copper
Cu
Cu2
cuprous
cupric
copper(I)
copper(II)
tin
Sn2
Sn4
stannous
stannic
tin(II)
tin(IV)
lead
Pb2
Pb4
plumbous
plumbic
lead(II)
lead(IV)
antimony
Sb3
Sb5
stibnous
stibnic
antimony(III)
antimony(V)
cobalt
Co2
Co3
cobaltous
cobaltic
cobalt(II)
cobalt(III)
gold
Au
Au2
aurous
auric
gold(I)
gold(II)
mercury
Hg+
mercurous
mercury(I)
Hg2+
mercuric
mercury(II)
Table 4 Prefixes Used When Naming Binary
Molecular Compounds
Subscript in
chemical formula
NEL
Prefix in chemical
nomenclature
1
mono
2
di
3
tri
4
tetra
5
penta
6
hexa
7
hepta
8
octa
9
nona
10
deca
Chemistry 11 Review 807
Table 6 Solubility of Ionic Compounds at SATP
Anions
Cations
Cl–, Br–, I–
S2–
OH–
+
most
SO42–
CO32–, PO43–, SO32–
+
C2H3O2–
NO3–
+
High solubility (aq)
0.1 mol/L
(at SATP)
Group 1, NH4
Group 1, NH4
most
Group 1, NH4
most
all
Group 2
Sr2+, Ba2+, Tl+
All Group 1 compounds, including acids, and all ammonium compounds are assumed to have high solubility in water.
Low Solubility (s)
0.1 mol/L
(at SATP)
Ag+, Pb2+, Tl+,
Hg22+ (Hg+),
Cu+
most
most
Ag+, Pb2+, Ca2+, most
Ba2+, Sr2+, Ra2+
Ag+
none
Practice
1. Write the chemical name and symbol corresponding to
each of the following theoretical descriptions:
(a) 3 protons, 4 neutrons, and 3 electrons
(b) 20 protons, mass number 40, and 18 electrons
(c) 10 electrons, net charge of 2
(d) 6 protons, 8 neutrons, no charge
2. When a gas is heated, the gas will emit light. Use the Bohr
model of the atom to explain why this phenomenon
occurs.
3. Use the periodic table to predict the most common
charges on ions of chlorine, potassium, and calcium.
Provide a theoretical explanation of your answer.
4. Are the following pairs of atoms more likely to form ionic
or covalent bonds? Give reasons for your answer.
(a) chlorine and chlorine
(b) potassium and iodine
(c) carbon and oxygen
(d) magnesium and fluorine
5. Draw a Lewis structure and a structural formula for each
of the following:
(a) O2
(b) CH4
(c) NH3
(d) PF3
(e) CO2
(f)
(g)
(h)
(i)
(j)
N2H4
HCN
H2S
OH
H3O
6. Identify the more polar bond in each of the following pairs:
(a) CH; OH
(b) CO; NO
(c) CC; CH
(d) SH; OH
(e) HCl; HI
7. Predict whether carbon tetrachloride, CCl4, is a polar or
nonpolar substance. Give reasons for your answer.
8. Write the formula, including state of matter, for each of the
following compounds.
(a) aluminum chloride
(b) copper(II) sulfate
(c) calcium hydroxide
(d) lead(II) nitrate
(e) sulfuric acid
(f) ferrous iodide
(g) ammonium nitrate
(h) sodium phosphate
(i) stannic bromide
(j) iron(III) carbonate
808 Appendix D
(k) potassium dichromate
(l) cobalt(III) sulfate
9. Write the IUPAC name for each of the following:
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
(k)
(l)
CuCl(s)
Fe2O3(s)
plumbic iodide
SF6(l)
NH4ClO3(s)
Cu(NO3)2(s)
hydrochloric acid
pentaphosphorus decaoxide
SnH4(g)
Ca(HCO3)2(s)
KMnO4(s)
CuSO4•5H2O(s)
10. For each of the following reactions, write a balanced equa-
tion and classify the reaction as synthesis, decomposition,
combustion, single displacement, or double displacement:
(a) iron + copper(I) nitrate → iron(II) nitrate + copper
(b) phosphorus + oxygen → diphosphorus pentoxide
(c) calcium carbonate → calcium oxide + carbon dioxide
(d) propane + oxygen → carbon dioxide + water
(e) lead(II) hydroxide → lead(II) oxide + water
(f) ammonia + sulfuric acid → ammonium sulfate
(g) potassium phosphate + magnesium chloride →
magnesium phosphate + potassium chloride
11. For each of the following, use an activity series to deter-
mine which single displacement reactions will proceed.
For the reactions that do occur, predict the products and
complete and balance the equation. Note reactions that do
not occur with NR.
(a) Cu(s) HCl(aq) →
(b) Au(s) + ZnSO4(aq) →
(c) Pb(s) + CuSO4(aq) →
(d) Cl2(g) + NaBr(aq) →
(e) Fe(s) + AgNO3(aq) →
12. Predict the products potentially formed by double displace-
ment reactions in aqueous solutions of each of the following pairs of compounds. In each case, write a balanced
chemical equation indicating the physical state of the products formed, and predict whether the reaction will proceed.
(a) copper(II) chloride and magnesium nitrate
(b) ammonium sulfate and silver nitrate
(c) barium hydroxide and potassium sulfate
NEL
Appendix D
D2 Quantities in Chemical Reactions
Determining the Limiting Reactant
SUMMARY
1. Write a balanced equation for the reaction.
2. Select one of the reactants and calculate the amount
in moles available.
Table 7 Stoichiometry, Symbols and Units
Quantity
Unit
3. Use mole ratios in the balanced equation to calculate
n
amount
mol
the amount in moles needed of the other reactants.
m
mass
mg, g, kg
M
molar mass
g/mol
N
number of entities
atoms, ions, formula
units, molecules
NA
Avogadro’s constant,
6.02 1023/mol
Symbol
4. Calculate the available amount in moles of the other
reactants. If the available amount of a reactant is more
than sufficient, it is in excess. If the available amount is
insufficient, it is limiting. (See example, Figure 5)]
(a)
CH4(g)
+
2 O2(g)
CO2(g) + 2 H2O(g)
—
we have
6.0 mol
we have
2.5 mol
Calculating Mass of Reactants
and Products
Begin with a balanced chemical equation, with the measured
mass of reactant or product written beneath the corresponding formula.
need
5.0 mol
1. Convert the measured mass into an amount in moles.
2. Use the mole ratio in the balanced equation to predict
have more
than enough
the amount in moles of desired substance.
3. Convert the predicted amount in moles into mass
(See example, Figure 4).
measured mass
of substance
CH4(g) is
limiting
reagent
mass of required
substance
step 1
(b)
step 3
moles of measured
substance
step 2
CH4(g)
O2(g) is in
excess
+
we have
2.5 mol
2 O2(g)
CO2(g) + 2 H2O(g)
we have
6.0 mol
moles of required
substance
need
3.0 mol
Fe2O3(s) + 3 CO(g)
2 Fe(s) + 3 CO2(g)
143.0 g
100.0 g
step 1
step 3
0.8952 mol
1.790 mol
step 2
Figure 4
Steps showing calculations
NEL
have less
than needed
CH4(g) is
limiting
reagent
O2(g) is in
excess
Figure 5
Steps showing limiting reagent
Chemistry 11 Review 809
D
Practice
1. Calculate the molar mass of each of the following. Express
your answers in g/mol.
(a) nitrogen gas
(b) C8H18(6)
(c) oxygen gas
(d) nickel(II) nitrate
(e) zinc hydrogen carbonate
(f) CuSO4•5H2O(s)
(g) helium gas
(h) sulfur trioxide liquid
(i) ammonia gas
(j) hydrochloric acid
2. What is the amount (in moles) of each type of atom in
each of the following samples?
(a) 3.0 mol of chlorine gas
(b) 2.0 mol of iron(III) nitrate
(c) 4.5 mol of potassium dichromate
(d) 1.5 mol of liquid nitrogen
(e) 5.0 mol of ammonium sulfate
3. Calculate the mass of each of the following:
(a)
(b)
(c)
(d)
(e)
2.5 mol of Mg(OH)2(s)
0.25 mol of glucose, C6H12O6(s)
6.75 mmol of oxygen molecules
1.20 1024 atoms of copper
3.01 1022 molecules of methane, CH4(g).
4. Calculate the amount (in moles) of each of the following
samples:
(a) 10.00 g of H2O(l)
(b) 1.50 kg of aluminum oxide
(c) 2.35 mg of sodium phosphate
(d) 1.20 105 g of hydrogen
(e) 1.00 1025 molecules of CO2(g)
5. Calculate the percentage composition of each of the
following:
(a) H2SO4(l)
(b) 2.50 g of AgNO3(s)
(c) NH4NO3(s)
6. An oxide of nitrogen was found to contain 36.8% nitrogen
by mass.
(a) Find the empirical formula for this compound.
(b) The molar mass of this compound was found to be
76.02 g/mol. What is the molecular formula of this
compound?
7. A gaseous compound contains 16.0 g of hydrogen and
96.0 g of carbon. If the molar mass of this compound is
28.06 g/mol, what is its molecular formula?
8. Balance the following equations. (You can use whole or
fractional coefficents.)
(a) NH3(g) O2(g) → NO(g) H2O(l)
(b) NO2(g) H2O(l) → HNO3(aq) NO(g)
(c) C12H22O11(s) O2(g) → CO2(g) H2O(l)
(d) KClO3(s) → KCl(s) O2(g)
810 Appendix D
(e)
(f)
(g)
(h)
MnO2(s) HCl(aq) → MnCl2(aq) Cl2(g) H2O(l)
Al2O3(s) → Al(s) O2(g)
Ni(s) AgNO3(aq) → Ag(s) Ni(NO3)2(aq)
KOH H3PO4 → K3PO4 H2O
9. Write a balanced chemical equation for each of the fol-
lowing reactions:
(a) phosphorus oxygen → diphosphorus pentoxide
(b) aluminum sulfate calcium hydroxide →
aluminum hydroxide calcium sulfate
(c) ammonia oxygen → nitrogen water
(d) calcium chloride nitric acid →
calcium nitrate hydrochloric acid
(e) ammonium sulfide lead(II) nitrate →
ammonium nitrate lead(II) sulfide
(f) aluminum sulfate ammonium bromide →
aluminum bromide ammonium sulfate
(g) sodium nitrate → sodium nitrite oxygen
(h) potassium phosphate magnesium chloride →
magnesium phosphate potassium chloride
(i) ammonia sulfuric acid → ammonium sulfate
(j) mercury(II) hydroxide phosphoric acid →
mercury(II) phosphate water
10. Methanol, CH3OH(l), burns in excess oxygen to produce
carbon dioxide and water, according to the following
equation:
2 CH3OH(l) 3 O2(g) → 2 CO2(g) 4 H2O(g)
(a) What amount of oxygen is required to completely
burn 5 mol of methanol?
(b) What amount of carbon dioxide is produced when
12.5 mol of methanol is completely burned?
11. Magnesium metal reacts with chlorine gas to produce
magnesium chloride.
(a) Write a balanced equation for the reaction.
(b) What mass of magnesium metal is needed to
completely react with 15.00 g of chlorine gas?
(c) What mass of magnesium metal is required to
produce, in excess chlorine, 8.00 g of magnesium
chloride?
12. Calcium hydroxide reacts with aqueous sodium carbonate
to produce sodium hydroxide and calcium carbonate.
(a) Write a balanced equation for this reaction.
(b) What mass of sodium carbonate is needed to
completely react with 175.0 g of calcium hydroxide?
(c) What mass of sodium hydroxide is produced when
175.0 g of calcium hydroxide is completely reacted in
an excess of sodium carbonate?
13. A single displacement reaction occurs when zinc metal is
immersed in lead(II) nitrate solution.
(a) Predict the products of the reaction.
(b) Write a balanced equation for the reaction.
(c) Predict the mass of lead formed when 4.55 g of zinc is
completely reacted in an excess of lead(II) nitrate.
(d) What mass of zinc metal is required to produce 50.0 g
of lead in this reaction, in an excess of lead(II) nitrate?
NEL
Appendix D
14. Propane, C3H8(g), burns in oxygen to produce carbon
dioxide and water, according to the following equation:
C3H8(g) 5 O2(g) → 3 CO2(g) 4 H2O(g)
Which is the limiting reagent if:
(a) 1 mol of propane and 1 mol of oxygen are available.
(b) 5 mol of propane and 5 mol of oxygen are available.
(c) 2 mol of propane and 5 mol of oxygen are available.
(d) 2 mol of propane and 12 mol of oxygen are available.
(e) 0.36 mol of propane and 1.60 mol of oxygen are available.
15. In a blast furnace, iron(III) oxide reacts with carbon
monoxide to produce iron and carbon dioxide.
(a) Write a balanced equation for the reaction.
(b) Identify the limiting reagent if 2.50 mol of iron(III)
oxide and 6.50 mol of carbon monoxide are available.
(c) Identify the limiting reagent if 200.0 g of iron(III) oxide
and 100.0 g of carbon monoxide are available.
(d) Predict the mass of iron produced in the reaction
when 200.0 g of iron(III) oxide and 100.0 g of carbon
monoxide are available.
16. When a solution containing 15.0 g of aluminum chloride is
mixed with a solution containing 15.0 g of sodium
hydroxide, a double displacement reaction occurs.
(a) Predict the mass of aluminum hydroxide produced.
(b) What mass of the excess reagent remains unreacted?
17. Silicon tetrafluoride is produced from the reaction of
silicon dioxide and hydrofluoric acid, with water as the
other product.
(a) What mass of silicon tetrafluoride can be produced
from 15.00 g of silicon dioxide in excess hydrofluoric
acid?
(b) If the actual yield of silicon tetrafluoride is 17.92 g,
what is the percentage yield?
18. When 8.40 g of zinc metal is placed in a solution in which
11.6 g of HCI(g) is dissolved, hydrogen gas and zinc
chloride are produced.
(a) Calculate the expected yield of hydrogen gas.
(b) If 0.19 g of hydrogen gas is produced, what is the
percentage yield?
D3 Solutions and Solubility
SUMMARY
Molar Concentration (mol/L)
molar concentration n
C ,
v
Hydrogen Ion Concentration and pH
amount of solute (in moles)
}}}
n vC,
n
v C
Preparing Standard Solution by Diluting Stock Solution
viCi vfCf
where
vi initial volume (volume of stock solution used)
Ci initial concentration
(concentration of stock solution used)
vf final volume (volume of dilute solution)
Cf final concentration (concentration of dilute solution)
pH is the negative power of ten of the hydrogen ion
concentration.
pH log[H
(aq)]
solution:
pH
[H
(aq)] 10
or
acidic
>107
neutral
[H+(aq)]:
107
<107
basic
pH:
<7
7
>7
Note the inverse relationship between [H+(aq)] and pH. The
higher the hydrogen ion molar concentration, the lower the pH.
Practice
1. Write equations to represent the dissociation of the fol-
lowing ionic compounds when they are placed in water:
(a) sodium chloride
(b) potassium sulfate
(c) ammonium nitrate
2. Calculate the molar concentration (mol/L) of each of the
(a) 0.174 mol of sodium hydroxide dissolved in water to a
final volume of 0.250 L of solution
(b) 60.0 g of NaOH(s) dissolved in water to a final volume
of 750.0 mL of solution
(c) 15.0 g of glucose, C6H12O6(s), dissolved in water to a
final volume of 125.0 mL of solution
following solutions:
NEL
Chemistry 11 Review 811
D
3. What volume of a 0.36-mol/L solution of KCl(aq) contains
0.09 mol of the solute?
4. What mass of sodium carbonate is required to make
0.500 L of a 0.12 mol/L solution?
5. The solubility of NaCl in water at 0°C is 31.6 g/100 mL.
What mass of NaCl(s) can be dissolved in 375 mL of
solution at 0°C?
6. Calculate the molar concentration of a solution that con-
tains 13.8 g of potassium bicarbonate in 354 mL of solution.
7. A 0.500-L sample of a sodium sulfate solution contains
0.320 mol of the solute. Calculate the molar concentration
of
(a) sodium sulfate
(b) sodium ions
(c) sulfate ions
8. Calculate the amount, in moles, of solute in 24.9 mL of a
0.200 mol/L solution of NaOH(aq).
9. What mass of copper(II) sulfate pentahydrate is needed to
prepare 150.0 mL of a 0.125 mol/L solution?
10. A 15.0-mL sample of 11.6 mol/L HCl(aq) is added to water to
make a final volume of 500.0 mL. Calculate the concentration of the final HCl(aq) solution.
11. What volume of concentrated 17.8 mol/L stock solution of
sulfuric acid would you need in order to prepare 2.00 L of
0.215 mol/L sulfuric acid?
12. The density of water is 1.00 g/mL.
(a) Calculate the mass of H2O in 1.00 L of water.
(b) Calculate the amount, in moles, of H2O(l) in 1.00 L of
water.
(c) What is the molar concentration of water?
(d) Does the concentration of water change?
13. Write the net ionic reaction for each of the following
reactions:
(a) aqueous barium chloride and aqueous sodium sulfate
(b) aqueous copper(II) sulfate and aluminum
(c) aqueous lead(II) nitrate and aqueous potassium
iodide
14. A 27.5-mL sample of 0.112 mol/L CuSO4(aq) solution is
added to 45.0 mL of 0.088 mol/L Na2CO3(aq). A precipitate
is formed.
(a) Write a balanced equation for the reaction.
(b) Identify the limiting reagent in the reaction.
(c) Calculate the mass of CuCO3 that is produced in the
reaction.
15. When 5.00 mL of a solution of KCl(aq) is added to an excess
of 1.00 mol/L Pb(NO3)2(aq), a precipitate of PbCl2(s) is
formed. The mass of the precipitate is found to be 0.075 g.
(a) Write a balanced equation for the reaction.
(b) Calculate the molar concentration of the KCl(aq) solution.
16. Write a sentence to distinguish between the terms in each
of the following pairs:
(a) dissociation and ionization
(b) a strong acid and a weak acid
812 Appendix D
(c) a strong base and a weak base
17. Calculate the pH of each of the following solutions:
(a) a vinegar solution with [H+(aq)] 1 102 mol/L
(b) an antacid solution with a hydrogen ion concentration
of 4.5 1011 mol/L
(c) orange juice with [H+(aq)] 5.5 103 mol/L
(d) a household cleaner with [H+(aq)] 7.2 1010 mol/L
18. Calculate the concentration of hydrogen ions in solutions
with the following pH values:
(a) pH 5.00
(b) pH 2.1
(c) pH 9.88
(d) pH 7.00
19. The pH of a hydrochloric acid solution was measured to
be 1.1.
(a) Write an ionization equation for hydrochloric acid.
(b) What is the concentration of hydrogen ions in the
solution?
(c) What is the concentration of the HCl(aq) solution?
20. How do acids differ from bases
(a) according to the Arrhenius definitions?
(b) according to the BrØnsted-Lowry definitions?
21. Identify the two acid–base conjugate pairs in each of the
following reactions:
(a) H3O+(aq) NH3(aq) → H2O(l) NH4+(aq)
2
(b) OH
(aq) HSO3(aq) → H2O(l) SO3(aq)
2
(c) HPO42
(aq) HSO4(aq) → H2PO4(aq) SO4(aq)
2
(d) HS
(aq) HCO3(aq) → CO3(aq) H2S(aq)
22. A 25.0-mL portion of 0.125 mol/L hydrochloric acid
requires 21.4 mL of potassium hydroxide solution for
neutralization. Calculate the molar concentration of the
potassium hydroxide solution.
23. A 20.0-mL portion of sulfuric acid solution requires
16.8 mL of 0.250 mol/L sodium hydroxide solution for
neutralization. Calculate the molar concentration of the
sulfuric acid solution.
24. A 10.0-mL portion of calcium hydroxide solution neutral-
izes 15.5 mL of 0.100 mol/L nitric acid. Calculate the molar
concentration of the barium hydroxide solution.
25. Calculate the molar concentration of a solution of phos-
phoric acid if 17.8 mL of it neutralizes 20.0 mL of
0.050 mol/L calcium hydroxide.
26. A solution of KOH is prepared by dissolving 2.00 g of KOH
in water to a final volume of 250 mL of solution. What
volume of this solution will neutralize 20.0 mL of 0.115
mol/L sulfuric acid?
27. Oxalic acid dihydrate, (COOH)2• 2H2O, reacts with sodium
hydroxide according to the following equation:
(COOH)2•2 H2O(s) 2 NaOH(aq) →
(COONa)2(aq) 4 H2O(l)
If a 0.118-g sample of oxalic acid dihydrate is dissolved in
water and exactly neutralized with 10.4 mL of a NaOH
solution, what is the molar concentration of the NaOH
solution?
NEL
Appendix D
D4 Gases and Atmospheric Chemistry
p1v1
p2v2
T1
T2
combined gas law: SUMMARY
(for constant amount of gas)
Gas Laws
STP: 0°C and 101.325 kPa (exact values)
SATP: 25°C and 100 kPa (exact values)
101.325 kPa 1 atm 760 mm Hg (exact values)
or 101 kPa (for calculation)
absolute zero 0 K
or 273.15°C, or 273°C (for calculation)
T (K) t (°C) 273 (for calculation)
Boyle’s law: p1v1 p2v2
(for constant temperature and amount of gas)
v1
v2
T1
T2
Charles’s law: (for constant pressure and amount of gas)
p1
p2
T1
T2
pressure–temperature law: Ideal gas law:
pv nRT
where n amount (in moles)
R 8.31 kPa•L/(mol•K)
Other Concepts
Dalton’s law of partial pressures the total pressure of a mixture of nonreacting gases is equal to the sum of the partial
pressures of the individual gases.
ptotal p1 + p2 + p3 + ...
Avogadro’s theory equal volumes of gases at the same
temperature and pressure contain equal numbers of molecules
molar volume the volume that one mole of a gas occupies at
a specified temperature and pressure
VSTP 22.4 L/mol; VSATP 24.8 L/mol
(for constant volume and amount of gas)
Practice
1. A balloon filled to 2.00 L at 98.0 kPa is taken to an altitude
at which the pressure is 82.0 kPa, the temperature
remaining the same. What is the new volume of the
balloon?
2. What volume will a sample of gas occupy at 88°C if it
occupies 1.50 L at 32°C?
3. A sample of gas in a metal cylinder has a pressure of 135.0
kPa at 298 K. What is the pressure in the cylinder if the gas
is heated to a temperature of 398 K?
4. A balloon has a volume of 2.75 L at 22.0°C and 101.0 kPa.
What is its volume at 37.0°C and 90.0 kPa?
5. A sample of gas occupies 1.00 L at 22°C and has a pressure
of 700.0 kPa. What volume would this gas occupy at STP?
6. Calculate the volume occupied by 2.50 mol of nitrogen gas
at 58.6 kPa and 40.0°C.
7. Calculate the pressure exerted by 6.60 g of carbon dioxide
gas at 25°C in a 2.00-L container.
8. What amount of chlorine gas is present in a sample that
has a volume of 500.0 mL at 20°C and exerts a pressure of
450.0 kPa?
9. Calculate the volume of 240.0 g of hydrogen gas when it is
at STP.
10. 1.00 L of an unknown gas has a mass of 1.25 g at STP.
Calculate the molar mass of the gas.
NEL
11. A sample of a mixture of gases contains 80.0% nitrogen
gas and 20.0% oxygen gas by volume. Calculate the mass
of 1.00 L of this mixture at STP.
12. Hydrogen gas reacts with nitrogen gas to produce
ammonia gas. In an experiment, 75.0 L of hydrogen gas is
reacted with an excess of nitrogen gas. All gases are at the
same temperature and the pressure is kept constant.
(a) What volume of nitrogen gas is required to react
completely with the hydrogen gas?
(b) What volume of ammonia gas is produced?
13. In a laboratory, hydrogen gas was collected by water dis-
placement at an atmospheric pressure of 98.2 kPa and a
temperature of 22.0°C. Calculate the partial pressure of the
dry hydrogen gas. (The vapour pressure of water at 22.0°C
is 2.64 kPa.)
14. Hydrogen gas is produced when zinc metal is added to
hydrochloric acid. What mass of zinc is necessary to
produce 250.0 mL of hydrogen at STP?
15. Ammonium nitrate, a solid, can decompose rapidly to
produce nitrogen gas, oxygen gas, and water vapour.
(a) Write a balanced equation for the decomposition of
ammonium nitrate.
(b) What is the total volume of the gases, measured at
SATP, produced from the decomposition of 1.00 kg of
ammonium nitrate?
Chemistry 11 Review 813
D
D5 Hydrocarbons and Energy
SUMMARY
Hydrocarbons
Table 8 Prefixes in Naming Alkanes,
Alkenes,and Alkynes
organic compounds
hydrocarbon
derivatives
hydrocarbons
aromatic
(e.g., benzene)
aliphatic
acyclic
alkanes
cyclic
alkenes
cycloalkanes
alkynes
C
C
C
C
C
C
C
C
C
C
C
Number of carbon atoms
meth-
1
eth-
2
prop-
3
but-
4
pent-
5
hex-
6
hept-
7
oct-
8
non-
9
dec-
10
cycloalkenes
C
C
Prefix
C
C
C
C
C
Figure 6
This classification system helps scientists organize their
knowledge of organic compounds.
Isomers
Structural isomers: chemicals with the same molecular formula, but with different structures and different names.
Geometric (cis-trans) isomers: organic molecules that differ
in structure only by the position of groups attached on either
side of a carbon–carbon double bond. (A cis isomer has both
groups on the same side of the molecular structure; a trans
isomer has groups on opposite sides of the molecular
structure.)
The quantity of heat energy, q, transferred to or from a
sample can be calculated:
q mc∆T
Thermochemical Equations
endothermic reaction: reactants energy (kJ) → products
exothermic reaction: reactants → products energy (kJ)
Specific Heat Capacity
A measure of the quantity of heat required to change the
temperature of a unit mass of a substance by one degree
Celsius (represented by c).
The specific heat capacity for water, c 4.18 J/(g•°C)
814 Appendix D
NEL
Appendix D
Practice
1. Draw a structural diagram for each of the following
hydrocarbons:
(a) 3-ethyl-2-methylhexane
(b) 2,2,3-trimethyloctane
(c) 1,3-dimethylcyclopentane
(d) 4-ethyl-2-hexene
(e) 3,4-dimethyl-2-pentene
(f) 1-butyne
2. Write IUPAC names for the following hydrocarbons:
(a) CHC—CH2—CH2—CH2—CH3
(b) CH3—CH—CHCH2—CH3
|
|
CH3 CH3
(c) CH3—CHCH2—CH—CH—CH3
|
CH3
(d) CH2—CH2
|
|
CH2—CH2
(e) CH3(CH2)7CH3
(f)
CH2—CH2—CH3
|
CH2—CH—CH2—CH2—CH2—CH3
|
CH3
3. Draw structural diagrams and write the IUPAC names for
the five structural isomers of C4H8(g).
4. Draw structural diagrams and write the IUPAC names for
the geometric isomers of 2-pentene.
5. Write a balanced equation for the complete combustion of
butane.
NEL
6. Classify each of the following hydrocarbons as saturated
or unsaturated:
(a) cyclohexane
(b) ethyne
(c) C3H8(g)
(d) a compound containing only single covalent bonds
(e) a hydrocarbon that reacts rapidly with bromine water
or potassium permanganate solution
7. Calculate the quantity of heat required to raise the tem-
perature of 1.50 L of water from 15.0°C to 75.0°C. The
specific heat capacity of water is 4.18 J/(g•°C).
8. Calculate the quantity of heat required to raise the tem-
perature of 500.0 g of water in a 325.0 g copper pot, from
12.0°C to 60.0°C. The specific heat capacity of copper is
0.385 J/(g•°C).
9. When 5.0 g of urea, NH2CONH2(s), is completely dissolved
in 150.0 mL of water, the temperature of the water changes
from 22.0°C to 18.3°C.
(a) Is the dissolving of urea in water endothermic or
exothermic?
(b) Calculate the specific heat of solution of urea (the
energy change in dissolving 1.0 g of urea).
(c) Calculate the molar heat of solution of urea (the
energy change in dissolving 1.0 mol of urea).
10. When methanol, CH3OH(l), burns in air, the products
formed are carbon dioxide gas and water vapour. When
10.0 g of methanol is completely combusted, 227.0 kJ of
heat is transferred.
(a) Is the combustion of methanol endothermic or
exothermic?
(b) Calculate the molar heat of combustion of methanol.
(c) Write a thermochemical equation for the combustion
of 1.0 mol of methanol.
(d) Write a thermochemical equation for the combustion
of 3.0 mol of methanol.
Chemistry 11 Review 815
D