1 s-block 18 H The Periodic Table of the Elements 1

Hydrogen

1H1.01

±1

Lithium

3 +1

Sodium

11

22.99

+1

Potassium

19

K39.10

+1

Rubidium

37 +1

Cesium

55

132.91

+1

Francium

87

(223)

+1

Beryllium

4 +2

Magnesium

12

24.31

+2

Calcium

20

40.08

+2

Strontium

38 +2

Barium

56

137.33

+2

Radium

88

(226)

+2

Iron

26

55.85

+3,2,6,-2

Ruthenium

44

101.07

+2,3,4,6,8,-2

Osmium

76

190.23

+4,8,6,3,±2

Hassium

108

(277)

Copper

29

63.55

+2,1

Silver

47

107.87

+1,2

Gold

79

196.97

+3,1,5,4,2,-1

Roentgenium

111 (+3)

Zinc

30

65.38

+2

Cadmium

48

112.41

+2

Mercury

80

200.59

Copernicium

112

(285)

Boron

5

10.81

+3

Aluminum

13

26.98

+3

Gallium

31

69.72

+3

Indium

49

114.82

+3

Thallium

81

204.38

Nihonium

113

(286)

Carbon

6

12.01

±4,3,2,1,0

Silicon

14

28.09

+4,-4

Germanium

32

72.63

+4

Tin

50

118.71

+4,2

Lead

82

207.20

+2,4

Flerovium

114

(289)

(+2)

Nitrogen

7

14.01

±3,5,4,2

Phosphorus

15

30.97

±3,5,4

Arsenic

33

74.92

±3,5

Antimony

51

121.76

±3,5

Bismuth

83

208.98

+3,5

Moscovium

115

(290)

Oxygen

8

16.00

-2,-1

Sulfur

16

32.06

+6,±2,4

Selenium

34

78.96

+4,6,-2

Tellurium

52

127.60

+4,6,-2

Polonium

84

(209)

+4,2,6

Livermorium

116

(293)

Fluorine

9

19.00

-1

Chlorine

17

35.45

±1,7,5,3

Bromine

35

79.90

±1,7,5,3

Iodine

53

126.90

±1,7,5

Astatine

85

(210)

±1,7,5,3

Tennessine

117

(294)

Helium

2

4.00

Neon

10

20.18

Argon

18

39.95

Krypton

36

83.80

+2,1

Xenon

54

131.29

+2,4,6

Radon

86

(222)

+2

1

2

3 5 6 8 11 12

13 14 15 16 17

18

Titanium

22

47.87

+4,3

4

1.5

Zirconium

40

91.22

+4

1.3

Hafnium

72 +4

1.3

Scandium

21 +3

1.4

Yttrium

39

88.91

+3

1.2

Rutherfordium

104

(267)

+4

Bohrium

107

(270)

7

Manganese

25

54.94

+2,3,4,6,-1

1.6

Technetium

43

(98)

+7

2.1

Rhenium

75

186.21

+7,6,4,2,-1

1.9

2.2

1.0

0.9

0.8

0.8

0.8

0.7

1.6

1.3

1.0

1.0

0.9

0.9

1.9

1.9

2.4

1.7

1.7

1.9

2.6

Oganesson

118

Darmstadtium

110

Ds

10

Nickel

28

58.69

+2,3

1.9

Palladium

46

106.42

+2,4

Molybdenum

42

95.95

+6,5,4,3,2

2.2

Tungsten

74

183.84

+6,5,4,3,2

Seaborgium

106

(269)

Chromium

24

52.00

+3,6,2

1.7

1.7

Meitnerium

109

Mt

9

Cobalt

27

58.931.9

Rhodium

45

102.91

+2,3,4,1

Iridium

77

192.22

+4,2,6,3,-1

2.2

Niobium

41

92.91

+5,3

1.6

Tantalum

73

180.95

+5

Dubnium

105

(262)

+5

Vanadium

23

50.94

+5,4,3,2

1.6

1.5

1.8

2.2

2.0

1.6

1.8

1.8

1.8

2.6

1.9

2.0

2.0

1.8

3.0

2.2

2.2

2.1

1.9

3.4

2.6

2.6

2.1

2.0

4.0

3.2

3.0

2.7

2.2

1

2

7

3

4

5

6

Period

Platinum

78

195.08

2.2

s-block

d-block

p-block

(+6) +7 +8

+2,3,-1

(+8,2) (+3,1)(+6,3,1) (+2) (+1,3)

Lanthanum

57

138.91

+3

1.1

Actinium

89

(227)

+3

1.1

+4,2

2.2

+2,1

(+2) (+3,1)

+3,1

(+2,4)

44.96

86.6285.47

6.94 9.01

178.49

2.22.3

(280)(281)(278)

Cn

O F

Te I1.7

Lv Ts(294)

Neodymium

60

144.24

+3

Uranium

92

238.03

1.1

Praseodymium

59

140.91

+3,4

Protactinium

91

231.04

1.1

Cerium

58

140.12

+3,4

1.1

Thorium

90

232.04

+4

Promethium

61

(145)

+3

Neptunium

93

(237)

+5,6,4,3

Europium

63

151.96

+3,2

Americium

95

(243)

Samarium

62 +3,2

Plutonium

94

(244)

+4,6,5,3

1.2+3,6,5,4

Gadolinium

64

157.25

+3

1.2

Curium

96

(247)

+3,4

Dysprosium

66

160.50

+3

Californium

98

(251)

1.2

Terbium

65

158.93

+3,4

Berkelium

97

(247)

+2,3,4 +2,3,4

Holmium

67

164.93

+3

1.2

Einsteinium

99

(252)

+3

Thulium

69

168.93

+3,2

Mendelevium

101

(258)

1.3

Erbium

68

167.26

+3

Fermium

100

(257)

+31.2

+3

Ytterbium

70

173.05

+3,2

Nobelium

102

(259)

+2,3

Lutetium

71

174.97

+3

Lawrencium

103

(262)

+31.0

1.3 1.5 1.7 1.3 1.3 1.3

f-block

+6,5,4,3+5,4

150.36

Pr

Pa

Cs

Fr

Rb

Ra

Ti VSc

Ta

Cr Mn

RuMo Tc

OsW

HsBhSgRf Db

Hf

ZrY

La

Ac

Ba

Sr

Ca

BeLi

Na

Ir Pt

Rh

Co ZnNi Cu

CdPd

Au

B C N

Al Si P

Ga Ge As

In Sn Sb

Tl Pb Bi

Nh Fl Mc

S Cl

Se Br

Po At

He

Ne

Kr

Ar

Xe

Rn

Th

Nd

U

Pm

Np Pu

Eu Gd

Am Cm

DyTb

CfBk

Ho

Es MdFm Lr

Ce Er YbTm Lu

No

Re

Fe

Nb

Sm

Ag

Mg

OgRg

Hg

Name of Element

Atomic Number

Atomic Mass

Atomic Symbol

*if in brackets mass number

for most stable isotope

Oxidation State (Bold most stable)

Electronegativity Pauling Scale;

Rounded

Platinum

78

195.08

+4,2

2.2

Pt

The Periodic Table of the Elements

Physical Constants

Constant Symbol Value

Avagadro number NA 6.022 x 1023 mol−1

Speed of light in a vacuum c 3.00 x 108 m/s

Planck constant h 6.626 x 10−34 Js

Molar volume of a gas at STP Vm 22.4 L/mol

Gas constant R 0.0821 Latm/molK

Ionization constant of water at 25°C Kw 1.0 x 10−14

STP Conditions

Pressure = 1 atm = 760 torr = 760 mmHg = 101.325 kPa = 1.01325 bar

Temperature = 0°C = 273.15 K

General Solubility Guidelines

1. Most sodium, potassium, and ammonium

compounds are soluble in water.

2. Most nitrates, acetates, and chlorates are soluble.

3. Most chlorides are soluble, except those of silver,

mercury(I), and lead. Lead(II) chloride is soluble in

hot water.

4. Most sulfates are soluble, except those of barium,

strontium, and lead.

5. Most carbonates, phosphates, and silicates are

insoluble, except those of sodium, potassium, and

ammonium.

6. Most sulfides are insoluble, except those of calcium,

strontium, sodium, potassium, and ammonium.

St. Joseph Chemistry 30 Course Outline

J. Nicholson, Room 117

Email: [email protected]

Website: http://nicholsonchem.weebly.com

Text: Davis, R. E., Metclafe, H. C., Williams, J. E., & Castka, J. F. (2002). Modern

chemistry.Toronto, Ontario: Holt, Rinehart, and Winston.

Evaluation

Outcome Exams (6) 65 %

Laboratory 5 %

Final Exam 30 %

Total 100 %

Chemistry 30 Outcomes Outcome Exams 1. MS1 Examine the role of valence electrons in the formation of chemical bonds. 2. MS2 Investigate how the properties of materials are dependent on their underlying intermolecular and intramolecular forces. 3. MS3 Explore the nature and classification of organic compounds, and their uses in modern materials. 4. EC1 Investigate the chemistry of oxidation and reduction reactions. 5. EQ1 Consider the characteristics and applications of equilibrium systems in chemical reactions. 6. EQ2 Analyze equilibrium of aqueous solutions. 7. EQ3 Observe and analyze phenomena related to acid-base reactions. SDS Research Project SDS1 Create and carry out a plan to explore one or more topics of personal interest relevant to Chemistry 30 in depth. Completion Assignment outcomes MS4 Determine the suitability of materials for use in specific applications. EC2 Examine the applications of electrochemistry and their impact on society and the environment.

Classroom Expectations

Materials

Each student should bring to class with them each day:

a binder with dividers, a pen (blue or black), a pencil, an eraser, a ruler, a calculator, and a text if

applicable.

Attendance

Regular attendance is essential for success. Research has shown that your mark decreases with

each absence by about 1 % on average. An excessive number of absences may result in you

being asked to discontinue the class. If you are absent, have a parent or guardian phone the

attendance desk - in advance if possible. (659-7656) You may not make up exams or class work

done on a day for which you have an unexcused absence. If you are late for class, come in

discretely and catch up on what you missed. Persons with habitual lates will be dealt with on an

individual basis.

Assignments

If you are absent from class, you are responsible for catching up in your notes and assignments.

Ask a friend (or me) what you have missed, and if possible, see me in advance - I may be able to

give you an assignment early. Assignments are due at the beginning of the period and late

assignment deductions are 5 % per day. Homework checks will be performed at random and will

count for marks. If you have an excused absence on a day I perform a homework check, you can

show me the completed work at a later time, or I can omit that particular check.

Exams

It is important that you write the exams at the time scheduled. If you have a legitimate reason to

miss an exam, you must see me before the day of the exam (the earlier the better) to make

arrangements to write it. If you miss an exam without prior consultation, you will need to

complete a make up assignment before you may write the exam. The requirements of the

assignment will depend on the circumstances of your absence. Skipping an exam may result in a

grade of 0 %.

There are no re-writes for exams.

General

1. No food or drinks are allowed in the classroom. (water bottles are an exception)

2. Cell phones need to be off and away at all times.

3. Be respectful of other people and their right to learn. Listen attentively when someone else is

speaking.

4. Be respectful of your classroom - use the garbage/recycling bin and be responsible for the

cleanliness of your station and desk.

5. You are expected to actively participate in classroom activities and discussions. Always give

your best effort.

6. You should use the washroom before or after class. If you must leave during class, use your

judgment as to whether it appropriate to go. Abuse of this responsibility will result in it being

taken away.

7. You will be treated with respect and you will learn all the things.

Selected Polyatomic Ions:

+1/ +2

H3O+ hydronium

NH4+ ammonium

Hg2 2+ mercury (I)

−1

ClO − hypochlorite

ClO2− chlorite

ClO3− chlorate

ClO4− perchlorate

NO3− nitrate

NO2− nitrite

OH− hydroxide

CH3COO− acetate (also C2H3O2−)

HCO3− bicarbonate (hydrogen carbonate)

HSO4− bisulfate (hydrogen sulfate)

HSO3− bisulfite (hydrogen sulfite)

CN− cyanide

SCN− thiocyanate

MnO4− permanganate

H2PO4− dihydrogen phosphate

−2

CrO4 2− chromate

Cr2O7 2− dichromate

SO4 2− sulfate

SO3 2− sulfite

CO3 2− carbonate

HPO4 2− hydrogen phosphate

C2O4 2− oxalate

O2 2− peroxide

−3

PO4 3− phosphate

PO3 3− phosphite

Chem 30: Chemistry Review

Equations:

n = m c = n n = #p a

M V 6.022 x 1023

% yield = actual yield x 100%

theoretical yield

1.Give the formula for the following compounds: 2. Name the following formulas:

a) sulfur trioxide a) SO2

b) dinitrogen pentoxide b) CS2

c) carbon tetraiodide c) N2O4

d) diphosphorous pentachloride d) Cl2O

e) carbon dioxide e) P4O10

f) beryllium nitrate f) Co(C2H3O2)2

g) phosphorous pentafluoride g) NH4HCO3

h) cobalt(III) sulfate h) Zr3P4

i) manganese(VII) oxide i) H2SO4

j) potassium hydroxide j) BF3

3. Balance and identify the type of these equations:

a) ___ Zn + ___ HCl → ___ ZnCl2 + ___ H2 _________________

b) ___ Al4C3 + ___ H2O → ___ CH4 + ___ Al(OH)3 _________________

c) ___ C3H6 + ___ O2 → ___ CO2 + ___ H2O _________________

4. Predict Products

a) Na + Cl2 → _________________

b) KCl → _________________

c) Mg + 2HCl → _________________

d) Mg + NaCl → _________________

e) FeS + 2HCl → _________________

f) C3H8 + 5O2 → _________________

5. Write and balance the chemical reaction:

a) Solid sodium reacts with liquid water to form aqueous sodium hydroxide and gaseous hydrogen. b) Solid iron(II) sulfide reacts with aqueous hydrochloric acid to form aqueous iron(II) chloride and

dihydrogen monosulfide gas.

6. How many atoms of gold are there in 6.25 mol of Au?

7. How many atoms of oxygen are there in 3.29 mol of H2SO4?

8. How many moles of Na atoms are there in 4.21 x 1024 atoms of Na?

9. How many moles of atoms are there in 126 g of silver?

10. What is the mass of 2.54 mol of aluminum?

11. Find the molar mass of MgCl2

12. How many moles of N2 are found in 6.8 L of the gas at STP?

13. How many litres of N2 are found in 2.77 mol of the gas at STP?

14. What is the mass of 2.59 x 1024 molecules of carbon dioxide?

15. Find the percent composition of a 42 g sample that contains 32.8 g iron and 9.2 g of oxygen.

16. Find the percent composition of ammonia, NH3, to the nearest percent.

17. A compound is analyzed and is found to contain 47.9 g of carbon and 127.7 g of oxygen. Find its

empirical formula.

18. Find the empirical formula of an iron-oxygen compound that is 70.0 % Fe and 30.0 % O by mass.

19. Find the empirical formula of a compound that contains 6.61 g of hydrogen, 105 g of sulfur and 209 g

of oxygen.

20. Find the molecular formula of a compound if its molar mass is 70.15 g/mol and its empirical formula

is CH2.

21. Find the molecular formula for a hydrocarbon if a 20.15 g sample contains 16.66 g of carbon, and its

molar mass is 87.21 g/mol.

22. Find the concentration of a solution made by dissolving 0.591 mol of H2SO4 in enough water to make

451 mL of solution.

23. What mass of sodium chloride would you add to water in order to make 500.0 mL of a 3.00 M

solution?

24. When sodium reacts completely with 6.00 mol of chlorine, how many moles of sodium chloride will

you get?

25. The complete combustion of 9.29 mol of octane would yield how many moles of carbon dioxide?

26. Excess of calcium reacts with 227 g of oxygen to form how many moles of calcium oxide?

27. What volume of 2.00 M AgNO3 is required to produce 4.00 L of 2.00 M Cu(NO3)2 when reacted with

copper?

28. If 4.00 mol of Na react with 3.50 mol of Cl2, find the limiting reactant.

29. Which reactant is limiting if 21.5 g of antimony react with 34 g of iodine? How many grams of

antimony(III) iodide would you expect?

30. When 158 g of zinc reacts with 640 mL of 2.0 M hydrochloric acid, what mass of zinc chloride is

formed? Find how many grams of the excess reactant is left over.

31. When 159.7 g of propene is burned with 145.2 g of oxygen, 89.4 g of carbon dioxide is actually

produced. Calculate the percent yield of CO2.

Solutions 1. a) SO3 2.a) sulfur dioxide

b) N2O5 b) carbon disulfide

c) CI4 c) dinitrogen tetroxide

d) P2Cl5 d) dichlorine monoxide

e) CO2 e) tetraphosphorous decoxide

f) Be(NO3)2 f) cobalt(II) acetate

g) PF5 g) ammonium bicarbonate

h) Co2(SO4)3 h) zirconium phosphide

i) Mn2O7 i) hydrogen sulfate/sulfuric acid

j) KOH j) boron trifluoride

3. a) Zn + 2HCl → ZnCl2 + H2 single replacement

b) Al4C3 + 12H2O → 3CH4 + 4Al(OH)3 double replacement

c) 2C3H6 + 9O2 → 6CO2 + 6H2O combustion

4. a) 2Na + Cl2 → NaCl

b) 2KCl → 2K + Cl2

c) Mg + 2HCl → MgCl2 + H2

d) Mg + NaCl → no reaction

e) FeS + 2HCl → FeCl2 + H2S

f) C3H8 + 5O2 → 3CO2 +4H2O

5. a) 2Na(s) + 2H2O(l) → NaOH(aq) + H2(g) b) FeS(s) + 2HCl(aq) → FeCl2(aq) + H2S(g)

6. 3.76 x 1024 atoms Au

7. 7.92 x 1024 atoms O 20. C5H10

8. 6.99 mol 21. C6H15

9. 1.17 mol 22. 1.31 M

10. 68.5 g 23. 87.7 g

11. 95.21 g/mol 24. 12.0 mol

12. 0.30 mol 25. 74.3 mol

13. 62.0 L 26. 14.2 mol

14. 189 g 27. 8.00 L

15. 78 % Fe 22% O 28. Na

16. 82% N 18% H 29. I2 is lim, 45g

17. CO2 29. I2 is lim, 45g

18. Fe2O3 30. 87 g ZnCl2, 116 g

19. H2SO4 31. 67.2%

Chem 30 MS1 Practice Problems 1

1. Determine the frequency of light with a wavelength of 4.257 x 10−7 cm.

2. Determine the energy in joules of a photon whose frequency is 3.55 x 1017 Hz.

3. Using the two equations E = hf and c = f, derive an equations expressing E in terms of h, c,

and .

4. How long would it take a radio wave with a frequency of 7.25 x 105 Hz to travel from Mars to

Earth if the distance between the two planets is approximately 8.00 x 107 km?

5. If the wavelength of the gamma radiation from a cobalt-60 source is 1.00 x 10−3 nm, calculate

the energy of a photon of this radiation.

6. What is the max kinetic energy of an electron that is ejected from sodium caused by light with

a frequency of 6.12 x 1016 Hz if its binding energy is 3.69 x 10−19 J?

7. What frequency of light causes the ejection of an electron with a KE max of 3.96 x 10−20 J from

a piece of copper, which has a binding energy of 7.69 x 10−19 J?

8. What is the binding energy of platinum if light with a frequency of 1.89 x 1015 Hz ejects an

electron with KE max of 2.42 x 10−19 J?

9. What is the longest wavelength of that that will eject an electron from tungsten, which has a

binding energy of 7.24 x 10−19 J? Can you see this light?

10. Given the binding energies of the following metals, which could produce the most energetic

electrons given the same light source? Note 1 eV = 1.60218 x 10−19 J.

11. What are the frequency and wavelength ranges of visible light?

12. List the colours of light in the visible spectrum in order of increasing frequency.

13. In the early 20th century, what two experiments involving light could not be explained by the

wave theory of light?

14. Which theory of light (wave or particle) best explains the following phenomena?

a) the interference of light b) the photoelectric effect c) the emission spectrum of an atom

15. Distinguish between the ground state and an excited state of an atom.

Ag Cu Fe Li K Mg Na Pt W

4.3 eV 4.8 eV 4.7 eV 2.4 eV 2.2 eV 3.7 eV 2.3 eV 6.3 eV 4.8 eV


16. Describe two major shortcomings of the Bohr model of the atom.

17. a) What is the principal quantum number?

b) How is it symbolized?

c) What are shells?

d) How does n relate to the number of orbitals per main energy level and the number of electrons

allowed per main energy level?

18. a) What information is given by the angular momentum quantum number?

b) What are sublevels or subshells?

19. For each of the following values of n, indicate the numbers and types of sublevels possible

for that main energy level.

a) n = 1 b) n = 2 c) n = 3 d) n = 4 e) n = 7

20. a) What information is given by the magnetic quantum number?

b) How many orbital orientations are possible in each of the s, p, d, and f sublevels?

c) Explain and illustrate the notation for distinguishing among the different p orbitals in a

sublevel.

21. a) What is the relationship between n and the total number of orbitals in a main energy level?

b) How many total orbitals are contained in the third main energy level? In the fifth?

22. a) What information is given by the spin quantum number?

b) What are the possible values for this quantum number?

23. How many electrons could be contained in the following main energy levels with n =

a) 1 b) 3 c) 4 d) 6 e) 7

24. a) In your own words, explain the Aufbau principle.

b) Explain the meaning of this principle in terms of an atom with many electrons.

25. a) In your own words, state Hund’s rule.

b) What is the basis for this rule?

26. a) In your own words, state the Pauli exclusion principle.

b) What is the significance of the spin quantum number?

c) Compare the values of the spin quantum number for two electrons in the same orbital.


27. a) What is meant by the highest occupied energy level in an atom?

b) What are inner shell electrons?

28. Write the electron configuration notation for an unidentified element that contains the

following number of electrons:

a) 3 b) 6 c) 8 d) 13

29. Write the orbital notation for the following elements:

a) P b) B c) Na d) O

30. Write the noble gas notation for the electron configuration of the following elements:

a) Cl b) Ca c) Se

31. Identify the following atoms based on its electron configuration:

a) 1s22s22p1 d) [Ne] 3s23p2 g) [Ar] 4s23d6

b) 1s22s22p5 e) [Ne] 3s23p5

c) [Ne] 3s2 f) [Ar] 4s1

32. Neon has 3 isotopes. Calculate the average atomic mass of neon given the following: 20Ne 19.992440 amu 90.48 % 21Ne 20.993847 amu 0.27 % 22Ne 21.991386 amu

33. Gallium has 2 isotopes. What are the percent abundances of each isotope of gallium if its

average atomic mass is 69.723 amu? 69Ga 68.925581 amu 71Ga 70.924705 amu

34. Magnesium has 3 isotopes. What are the masses of the isotopes of magnesium (24.3050 amu)

if their abundances and mass components are: 24Mg _________________ 78.99 % 18.94578468 amu 25Mg _________________ 26Mg _________________ 11.01 % 2.860683489 amu

35. Iron has 4 isotopes. Without doing any calculations, what is the approximate average atomic

mass of iron? 54Fe 53.940 5.845 % 56Fe 55.935 91.754 % 57Fe 56.935 2.119 % 58Fe 57.933 0.282 %


36. Nickel has 5 isotopes. What is the mass of 62Ni to 3 decimal places given: 58Ni 57.935348 amu 68.0769 % 60Ni 59.930791 amu 26.2231 % 61Ni 60.931060 amu 1.1399 % 64Ni 63.927970 amu 0.9256 %

37. a) What is meant by atomic radius?

b) What trend is observed among the atomic radii of main group elements across a period?

c) How can this trend be explained?

38. a) What trend is observed among the atomic radii of main-group elements down a group?

b) How can this trend be explained?

39. Define

a) ion c) first ionization energy

b) ionization energy d) second ionization energy

40. a) How do the first ionization energies of main-group elements vary across a period and

down a group?

b) Explain the basis for this trend.

41. a) What is electron affinity?

b) What signs are associated with electron affinity values, and what is the significance of each

sign?

42. a) What are valence electrons?

b) Where are such electrons located?

43. a) What is electronegativity?

b) Identify the 3 most and 3 least electronegative elements.

44. Of Cs, Hf and Au, which element has the smallest atomic radius? Explain your answer in

terms of trends in the periodic table.

45. a) Distinguish between the first, second, and third ionization energies of an atom.

b) How do the values of successive ionization energies compare?

c) Why does this occur?

46. Without looking at an electron affinity table, arrange the following elements in order of

decreasing electron affinities: C, O, Li, Na. Rb, and F


47. a) Without looking at an ionization energy table, arrange the following elements in order of

decreasing first ionization energies: Li, O, C, K, Ne, and F

b) Which of the elements listed would you expect to have the highest second ionization energy

and why?

48. a) Which of the following cations is least likely to form: Sr2+, Al3+, K2+?

b) Which of the following anions is least likely to form: I−, Cl−, O2−?

Solutions

1. 7.05 x 1016 Hz

2. 2.35 x 10−16 J

3. E = hc/

4. 267 s

5. 1.99 x 10−13 J

6. 4.02 x 10−17 J

7. 1.22 x 1015 Hz

8. 1.01 x 10−18 J

9. <2.75 x 10−7 m, no: UV light

K

11. f range: 8 x 1014 to 4 x 1014 Hz range:= 400 to 700 nm

12. red, orange, yellow, green, blue, indigo, violet (ROYGBIV)

13. The photoelectric effect and hydrogen’s line-emission spectrum.

14. a) wave theory b) particle theory c) particle theory

15. The ground state of an atom is the atom’s lowest energy state. An excited state is any energy

state higher than the ground state.

16. The Bohr model was only valid for a single-electron atom and it did not explain the chemical

nature of atoms.

17. a) The number used to specify the main energy level of an atom b) by the letter n

c) all the orbitals within the main energy level d) The number of orbitals per main energy

level is equal to n2. The number of electrons allowed per main energy level is equal to 2n2.

18. a) The angular momentum quantum number indicates an orbital’s shape.

b) A sublevel or subshell consists of the orbitals within a main energy level that share the same

value of l. For example, an atom’s 3d subshell consists of 5 d orbitals.

19. a) 1; s b) 2; s, p c) 3; s, p, d d) 4; s, p, d, f e) 7; s, p, d, f, g, h, i

20. a) The orientation of an orbital about the nucleus. b) 1, 3, 5, and 7

c) Subscripts x, y, and z may be used to designate the 3-dimensional orientation of each orbital.

21. a) The number of possible orbitals is equal to n2.

b) There are 9 orbitals in n = 3 and 25 orbitals in n = 5.

22. a) The spin quantum number indicates the electron’s spin state. b) +½, −½

23. a) 2 b) 18 c) 32 d) 72 e) 98


24. a) An electron occupies the lowest energy orbital that can receive it.

b) In a multi-electron atom, the lowest energy orbital is filled first. Electrons are then added to

the orbital with the next lowest energy until all of the electrons have been placed in orbitals.

25. a) Orbitals of equal energy are each occupied by one electron before any orbital is occupied

by a second electron.

b) By placing electrons in separate orbitals, electron-electron repulsion is minimized and lower

energy arrangements result.

26. a) No two electrons in the same atom can have the same four quantum numbers.

b) The values show that two electrons in the same orbital spin in opposite directions.

c) One electron is +½ and the other is −½.

27. a) The highest occupied energy level is the level with the highest n that contains electrons.

b) Inner-shell electrons are not in the highest occupied energy level.

28. a) 1s2 2s1

b) 1s2 2s2 2p2

c) 1s2 2s2 2p4

d) 1s2 2s2 2p6 3s2 3p1

29. a)

1s 2s 2p 3s 3p

b)

1s 2s 2p

c)

1s 2s 2p 3s

d)

1s 2s 2p

30. a) [Ne] 3s2 3p5

b) [Ar] 4s2

c) [Ar] 4s2 3d10 4p4

31. a) boron b) fluorine c) magnesium d) silicon e) chlorine f) potassium g) iron

32. 20.18 amu

33. 69Ga = 60.11 %, 71Ga = 39.89 %

34. 24Mg 23.99 amu 25Mg 24.99 amu 26Mg 25.98 amu

35. Around 55 because 92 % of the mass is 55.

36. 61.930 amu

37. a) Atomic radius is half of the distance between the nuclei of two identical bonded atoms.

b) They decrease. c) As electrons are added to the s and p sublevels in the same main energy

level, the increasing positive charge of the nucleus pulls electrons closer to the nucleus, resulting

in decreasing atomic radii.

38. a) They generally increase. b) Down a group, the outer electrons of each element

occupy comparable sublevels in successively higher main energy levels farther from the nucleus.

39. a) A charged atom or a charged group of bonded atoms.

b) Any process that results in forming an ion.


c) The energy required to remove one electron from a neutral atom to form an ion.

d) The energy required to remove one electron from a +1 ion.

40. a) They increase across a period and decrease down a group.

b) Across a period, the increasing nuclear charge more strongly attracts electrons in the same

energy level and makes them more difficult to remove. Down a group, the outermost electrons

become farther from the nucleus and thus more easily removed.

41. a) The energy taken in or given off when an electron is added to an atom

b) Positive or negative. A negative sign means that energy is given off; Positive means taken in.

42. a) Valence electrons are the electrons able to be lost, gained, or shared to form bonds.

b) They are the s and p electrons in the highest (outermost) energy level.

43. a) The ability of an atom in a compound to attract electrons from other atoms.

b) Most: 1) F 2) O 3) Cl, N (tie) Least: 1) Fr, Cs (tie) 2) Rb, K (tie) 3) Ra, Ba, Na (tie)

44. Gold - all 3 are in period 6, gold is furthest right and atomic radius decreased left to right.

45. a) First ionization energy is the energy required to remove an electron from a neutral atom;

Second ionization energy is the energy needed to remove an electron from a +1 ion;

Third ionization energy is the energy needed to remove an electron from a +2 ion.

b) IE1 < IE2 < IE3. c) Each successive electron must be removed from a more positive particle.

46. Decreasing electron affinity: F, O, C, Li, Na, Rb

47. a) Decreasing IE1: Ne, F, O, C, Li, K

b) Li and K – in both cases, the second electron must come from a completely filled noble-gas

configuration. Of the two, Li would have the higher IE2 because the Li+ ion is smaller than the

K+ ion.

48. a) K2+ b) O2− because second electron affinity values are always positive – it is always

more difficult to add a second electron to an already negatively charged ion.

Compound Lewis Structure Electron Group

Geometry Molecular

Geometry Bond Angle () Polarity

(P/NP) Resonance

(Y/N)

(level 5)

1. AlH3Cl−

2. BF2I2−

3. SF6

4. AlHFI

5. NBrI2

6. SCl2

7. NF4+

8. PBr3


Geometry Molecular


(P/NP) Resonance

(Y/N)

(level 7)

1. ClCN

2. COAt2

3. POF3

4. C2Cl6

5. HCAs

6. HCl

7. NAtBr4

8. AlH2FI−

9. PH2I3


Geometry Molecular


(P/NP) Resonance

(Y/N)

(level 9)

1. C2S2

2. NFO

3. AtO4−

4. HCO2−

5. IOOAt

6. PO2CN

O

N C P

O

7. Ge2OTe66−

Te Te

Te Ge O Ge Te

Te Te

8. BrSSI

9. S2Br2


Geometry Molecular


(P/NP) Resonance

(Y/N)

(level 10)

1. ArClF

2. XeF4

3. BiH2I2−

4. AsH2At2−

5. ClF5

6. XeOBr3−

7. CO32−

8. SO42−

VSEPR Theory

Level 5

E.G.G. M.G. Bond Angle Polarity Resonance

1 tetrahedral tetrahedral 109.5 (charged) N


3 octahedral octahedral 90 NP N

4 trigonal planar trigonal planar 120 P N

5 tetrahedral trigonal pyramidal <109.5 P N

6 tetrahedral bent <109.5 P N


8 tetrahedral trigonal pyramidal <109.5 P N

Level 7


1 linear linear 180 P N

2 trigonal planar trigonal planar 120 P N

3* tetrahedral tetrahedral 109.5 P N

4 tetrahedral tetrahedral 109.5 NP N

5 linear linear 180 P N

6 tetrahedral linear n/a P N

7 trigonal bipyramidal trigonal bipyramidal 120 + 90 P N


9 trigonal bipyramidal trigonal bipyramidal 120 + 90 P N

Level 9


1 linear linear 180 NP N

2 trigonal planar bent <120 P N

3* tetrahedral tetrahedral 109.5 (charged) N

4 trigonal planar trigonal planar 120 (charged) Y


6* C = linear P = trigonal planar

C = linear P = trigonal planar

C = 180 P = 120

P Y

7* Ge = tetrahedral O = tetrahedral

Ge = tetrahedral O = bent

Ge = 109.5 O = <109.5

(charged) N



Level 10


1 trigonal bipyramidal linear 180 P N

2 octahedral square planar 90 P N

3* trigonal bipyramidal see-saw <120 + <90 (charged) N

4* trigonal bipyramidal see-saw <120 + <90 (charged) N

5 octahedral square pyramidal <90 P N

6 octahedral square planar 90 (charged) N

7 trigonal planar trigonal planar 120 (charged) Y



1. Find the difference in electronegativity and use it to determine the types of bonds present.

a) RbCl

b) CO2

c) MgF2

d) SO2

e) CI4

f) NH3

g) CaO

h) CH4

i) Fe2O3

2. Use orbital notation to show

bonding in a) Cl2 b) O2 c) HF

3. What are hybrid orbitals?

b) What determines the number of

hybrid orbitals produced by an atom?

4. Identify the type of hybridization found in the following molecules:

a) CCl4 d) H2O

b) HCN e) XeF4

c) BH3 f) PCl5

5. What is a chemical bond?

6. In general, what determines whether atoms will form chemical bonds?

7. a) What determines bond length? b) In general, how are bond energies and bond lengths

related?

8. Describe the types of bonding in HCN in terms of sigma and pi bonds.

9. What property of metals contribute to their tendency to form metallic bonds?

10. a) What is metallic bonding?

b) How can the strength of metallic bonding be measured?

11. Aluminum’s heat of vapourization is 284 kJ/mol and beryllium is 224 kJ/mol. Which

element has stronger metallic bonding?


12. a) What is lattice energy?

b) In general, what is the relationship between lattice energy and the strength of ionic bonding?

13. The lattice energy of NaCl is – 787.5 kJ/mol and KCl is –715 kJ/mol. Which compound

would you expect to have a higher melting point and why?

14. In general, how do ionic and molecular compounds compare in terms of melting points,

boiling points, and ease of vapourization?

15. Arrange the following pairs from strongest to weakest attraction:

a) polar molecule and polar molecule

b) nonpolar molecule and nonpolar molecule

c) polar molecule and ion

d) ion and ion

16. What are dipole-dipole forces?

17. What is hydrogen bonding?

18. What are London dispersion forces?

19. What is surface tension?

20. List seven properties of solids.

21. List four common examples of amorphous solids.

22. a) List and describe the four types of crystals in terms of their component particles and the

type of bonding between them.

b) What physical properties are associated with each type of crystal?

23. What is a fluid?

24. What is a liquid?

25. What is a gas?

26. Convert the following into torr:

a) 1.25 atm b) 2.48 x 10–3 atm c) 4.75 x 104 atm d) 7.60 x 106 atm


27. Convert:

a) 125 mm Hg into atm b) 3.20 atm into Pa c) 5.38 kPa into torr

28. Convert from Celsius to Kelvin:

a) 0.00C b) 27C c) –50.0C d) – 273C

29. Convert from Kelvin to Celsius:

a) 273 K b) 350.0 K c) 100.00 K d) 20 K

30. What is the volume in litres of each of the following at STP?

a) 1.00 mol O2 b) 3.50 mol F2 c) 0.0400 mol CO2 d) 1.20x 10 –6 mol He

31. How many moles are contained in each of the following at STP?

a) 22.4 L N2 b) 5.60 L Cl2 c) 0.125 L Ne d) 70.0 mL

32. Find the mass in grams of each of the following at STP

a) 11.2 L H2 b) 2.80 L CO2 c) 15.0 mL SO2 d) 3.40 cm3 F2

33. Find the volume in litres of each of the following at STP

a) 8.00 g O2 b) 3.50 g CO c) 0.0170 g H2S d) 2.25 x 105 kg NH3

34. The pressure exerted on a 240. mL sample of hydrogen gas at constant temperature is

increased from 0.428 atm to 0.724 atm. What will be the final volume of the gas?

35. A flask containing 155 cm3 of hydrogen gas was collected under a pressure of 22.5 kPa.

What pressure would have been required for the volume to have been 90.0 cm3, assuming the

same temperature?

36. A gas has a volume of 450.0 mL. If the temperature is held constant, what volume would the

gas occupy if the pressure were a) doubled? b) reduced to one-fourth its original value?

37. A sample of oxygen that occupies 1.00 x 106 mL at 575 mm Hg is subjected to a pressure of

1.25 atm. What is the final volume of the sample if the temperature is held constant?

38. A sample of air has a volume of 140.0 mL at 67C. At what temperature will its volume be

50.0 mL at constant pressure?

39. At standard temperature, a gas has a volume of 275 mL. The temperature is then increased to

130.C, and the pressure is held constant. What is the new volume?


40. A sample of hydrogen at 47C exerts a pressure of 0.329 atm. The gas is heated to 77C at

constant volume. What is its new pressure?

41. To what temperature must a sample of nitrogen at 27C and 0.625 atm be taken so that its

pressure becomes 1.125 atm at constant volume?

42. The pressure on a gas at –73C is doubled, but the volume is held constant. What will be the

final temperature in degrees Celsius?

43. A sample of gas at 47C and 1.03 atm occupies a volume of 2.20 L. What volume would this

gas occupy at 107C and 0.789 atm?

44. A 350. mL air sample collected at 35C has a pressure of 550. torr. What pressure will the air

exert if it is allowed to expand to 425 mL at 57C?

45. A gas has a volume of 1.75 L at –23C and 150. kPa. At what temperature would the gas

occupy 1.30 L at 210. kPa?

46. A sample of oxygen at 40.C occupies 820. mL. If this sample later occupies 1250 mL at

60.C and 1.40 atm, what was its original pressure?

47. A gas at 7.75 x 104 Pa and 17C occupies a volume of 850. cm3. At what temperature, in

degrees Celsius, would the gas occupy 720. cm3 at 8.10 x 104 Pa?

48. Calculate the pressure in atm exerted by 7.50 x 102 mL of CO2 containing 2.15 mol at 57C.

49. Calculate the volume in litres occupied by 4.00 g of O2 at 57C and 0.888 atm.

50. Determine the moles of gas contained in 0.80 L at 27C and 0.925 atm.

51. Find the mass of 125 mL of SO2 at 0.822 atm at –53C.

52. Find the molar mass of a gas if a 1.05 g sample occupies 2.35 L at 37C and 0.840 atm.

53. If the density of an unknown gas is 3.20 g/L at –18C and 2.17 atm, what is the molar mass

of the gas?


54. Using the phase diagram for ammonia, describe the phase changes that occur as it is heated

from 100 K to 400 K at 2 atm.

55. What are point C, D and E called?

56. What phase change would occur if the pressure on ammonia at 185 K is reduced from 4 atm

to 0.000 000 5 atm?

Solutions

1. a) RbCl 2.2 ionic

b) CO2 1.0 polar covalent

c) MgF2 2.8 ionic

d) SO2 1.0 polar covalent

e) CI4 0.0 nonpolar covalent or pure covalent

f) NH3 0.9 polar covalent

g) CaO 2.5 ionic

h) CH4 0.4 polar covalent

i) Fe2O3 1.7 polar covalent

2. See figures to right

3. a) Hybrid orbitals are identically shaped orbitals of equal energy that are

produced by mixing two or more atomic orbitals of similar, but not identical

energies on the same atom.

b) The number of hybrid orbitals is equal to the number of orbitals that have combined.

4. a) sp3 b) sp c) sp2 d) sp3 e) sp3d2 f) sp3d

5. a) A chemical bond is a link between atoms resulting from the mutual attraction of their nuclei

and electrons.


6. In general, atoms will form a chemical bond if their potential energy is lowered by doing so.

7. Bond length is the distance at which potential energy is at a minimum. It is the point at which

there is a balance between attraction and repulsion between atoms in a covalent bond.

b) In general, the higher the bond energy, the shorter the bond.

8. The H-C bond is a sigma bond. The C-N triple bond consists of one sigma bond and two pi

bonds.

9. Most metals contain sparsely populated outermost orbitals, they have low ionization energies,

and they have low electronegativities.

10. a) Metallic bonding results from the attraction between metal atoms and a sea of surrounding

electrons. b) A metal’s heat of vapourization is a measure of the bond strength.

11. Al has a higher heat of vapourization therefore Al has stronger bonding.

12. a) Lattice energy is the energy released when one mole of an ionic compound is formed from

gaseous ions. b) The greater the lattice energy, the stronger the ionic bonding.

13. NaCl has a greater lattice energy therefore has stronger bonding and as a result has a higher

m.p. (801 vs 770C)

14. Ionic compounds have higher melting and boiling points than molecular compounds. Ionic

compounds do not vapourize at room temperature as do molecular compounds.

15. Strongest to weakest: d, c, a, b

16. a) Dipole-dipole forces are forces of attraction between polar molecules

17. Hydrogen bonding is a particularly strong dipole-dipole force that occurs among molecules

containing H-N, H-O, and H-F bonds. Because of the great difference in electronegativity

between H and N, O, and F, H has a positive charge of nearly a proton. This, combined with the

small size of the atoms involved, makes the dipole-dipole attraction very strong.

18. London dispersion forces are intermolecular forces of attraction resulting from the creation

of instantaneous dipoles.

19. Surface tension is a force that tends to pull adjacent parts of a liquid’s surface together,

decreasing surface area.

20. Solids maintain a definite shape, have a definite volume, are non-fluid, have definite melting

points, have high densities, are incompressible, and have a low rate of diffusion.

21. rubber, glass, plastic, synthetic fibres, wax etc.

22. a) 1) Ionic: ionically bonded positive and negative ions in a regular pattern;

b) 1. Covalent network crystals are hard, brittle, have high m.p., and are usually nonconducting

or semiconducting. Covalent molecular crystals are soft, have low m.p., easily vapourized, and

are good insulators. Metallic crystals are good conductors, have varying m.p.

2) Covalent network: single atoms covalently bonded to nearest neighbours in extensive pattern;

3) Covalent molecular: covalently bonded molecules held together by London dispersion forces,

dipole-dipole forces or hydrogen bonding;

4) Metallic: positive metal cations surrounded by negative valence electrons that belong to the

crystal as a whole and are attracted to the cations.

b) Ionic crystals are hard, brittle, have high m.p., and are good insulators.


23. A fluid is a substance that flows and conforms to the shape of the container.

24. A liquid is a fluid that has a definite volume but no definite shape.

25. A gas is a fluid that has neither definite shape nor definite volume – it expands to fill the

container.

26. 26. a) 9.50 x 102 torr b) 1.88 torr c) 3.61 x 107 torr d) 5.78 x 109 torr

27. a) 0.164 atm b) 3.24 x 105 Pa c) 40.4 torr

28. a) 273.15 K b) 3.00 x 102 K c) 223.2 K d) 0 K

29. a) 0C b) 76.9C c) –173.15C d) –250C

30. a) 22.4 L b) 78.4 L c) 0.896 L d) 2.69 x 10–5 L

31. a) 1.00 mol b) 0.250 mol c) 5.58 x 10–3 mol d) 3.12 x 10–3 mol

32. a) 1.01 g b) 5.50 g c) 0.0429 g d) 5.77 x 10–3 g

33. a) 5.60 L b) 2.80 L c) 0.0112 L d) 2.96 x 108 L

142 mL

35. 38.8 kPa

36. a) 225.0 mL b) 1.800 x 103 mL

37. 6.05 x 105 mL

38. 121 K or –152C

39. 406 mL

40. 0.360 atm

41. 540. K or 267C

42. 127C

43. 3.41 L

44. 485 torr

45. 260. K or –13C

46. 2.01 atm

47. –16C

48. 77.7 atm

49. 3.82 L

50. 0.030 mol

51. 0.364 g

52. 13.5 g/mol

53. 30.9 g/mol

54. At 100 K ammonia is a solid and continues to be solid until the temperature reaches about

198 K, where it melts. Ammonia remains liquid unit it reaches about 250 K, where it boils.

55. C is the triple point; D is the normal melting/freezing point; E is the normal boiling point.

56. Ammonia will sublimate.

Draw: 1 3,4,5-trimethylheptane 2 5-methylhex-2-ene 3 3-ethylhexa-2,4-diene 4 5-methylhex-2-yne 5 1,4-dimethylbenzene 6 1,3-dimethylcycloheptane 7 6-ethyl-5-methyloct-3-yne 8 4-methylhepta-2,5-diyne 9 3-bromo-2-chloro-5-fluoro-4-iodoheptane 10 3,6-diethyl-5,5,7-trimethyl-4-propyldecane 11 1,2,3,4,6,7-hexaethylcyclooctene

Chem 30 Functional Groups Worksheet

Draw: 1) heptan-4-one 6) methyl butanoate

2) 2-ethyl-3-methylbutanoic acid 7) ethoxyethane

3) 3-hydroxypropanamide 8) 2-chloro-5-ethyl-4-methylheptan-3-ol

4) 2-phenylbutanal 9) pentan-2-amine

5) hexan-1-amine 10) 1-chloro-3-hydroxypentan-2-one

N-ethyl-N-methylethanamine 3-ethyl-2-methylpentanal

3,4-dimethylpentan-2-one 2-ethylbutanoic acid

2-methylhexan-3-ol hexan-3-one

3,3-dimethylbutanal 1-ethoxyhexane

2,2,3-trimethylhexan-3-ol ethyl heptanoate

FUNCTIONAL GROUPS II - WORKSHEET

Name or draw the following compounds: Chemical structure IUPAC Name

1

2

3

4

N-butylhexanamide

5

6

7

8

9

2-methylhexanal

10

N-(1-methylethyl) ethanamide

11

12

13 2,3-dibromononan-5-amine

14

2-chloropropyl butanoate

15

16

N-ethyl-N-methyl-4,5-difluorooct-4-

enamide

For each of the following, state the highest priority functional group and

name the organic compound

17

18

19

20

21

22

Functional Groups II Solutions

*not necessary in this class

Name or draw the following compounds: Chemical structure IUPAC Name

1

1-butoxyheptane

2

butyl propanoate (or but-1-yl propanoate)*

3

7-fluorononan-4-amine

4

N-butyl hexanamide

5

N-propylethanamide

6

2-methoxypentane

7

5-methylhexan-3-amine

8

N-ethyl-N-methyl pentanamide

9

2-methylhexanal

10

N-(1-methylethyl) ethanamide

11

pentyl 2-methylbutanoate

12

N-ethyl-N-propylbutan-1-amine

13

2,3-dibromononan-5-amine

14

2-chloropropyl butanoate

15

2-chloro-1-ethoxypropane

16

N-ethyl-N-methyl-4,5-difluorooct-4-

enamide

For each of the following, state the highest priority functional group and

name the organic compound

17

FG: ketone

4-methoxy-4-methylpentan-2-one

18

FG: amine

pentan-1-amine

19

FG: alkane

1-ethoxy-2-methylpropane

20

FG: amide

N-ethyl-N-methyl butanamide

21

FG: amine (tertiary)

N-ethyl-N-methyl-2-methylpropan-1-amine

22

FG: ester

butyl pent-2-enoate

Chem 30 EC1 Practice Problems 1

1. Arrange the following in order of increasing oxidation number of the xenon atom:

CsXeF8, Xe, XeF2, XeOF2, XeO3, XeF

2. Identify each of the following reactions as a redox or

non redox reaction:

a) 2Na + Cl2 → 2NaCl

b) C + O2 → CO2

c) 2H2O → H2 + O2

d) NaCl + AgNO3 → AgCl + NaNO3

e) NH3 + HCl → NH4+ + Cl−

f) 2KClO3 → 2KCl + 3O2

g) H2 + Cl2 → 2HCl

h) H2SO4 + 2KOH → K2SO4 + 2H2O

i) Zn + CuSO4 → ZnSO4 + Cu

3. For each redox reaction in the previous question, identify what is oxidized and what is

reduced.

4. Identify the most active reducing agent among all the common elements.

b) Why are all of the elements in its group in the periodic table very active reducing agents?

c) Identify the most active oxidizing agent among the common elements.

5. Based on table, identify the strongest and weakest reducing

agents among the substances listed within each of the following:

a) Ca, Ag, Sn, Cl−

b) Fe, Hg, Al, Br−

c) F−, Pb, Mn2+, Na

d) Al3+, Cu2+, NO3−, K+

e) Cl2, S, Zn2+, Ag+

f) Li+, F2, Ni2+, Fe3+

6. Use the table to respond to the following:

a) Would Al be oxidized by Ni2+ ?

b) Would Cu be oxidized by Ag+ ?

c) Would Pb be oxidized by Na+ ?

d) Would F2 be reduced by Cl− ?

e) Would Br2 be reduced by Cl− ?

7. What is the difference between a voltaic cell and an electrolytic

cell in terms of the nature of the reaction involved?


8. a) What is electroplating?

b) Distinguish between the anode and the cathode in such a process.

9. a) Explain what is meant by potential difference between two electrodes in an electrochemical

cell.

b) How, and in what units, is the potential difference measured?

10. For each of the following pairs of half-cells, determine the overall electrochemical reaction

that proceeds spontaneously and its E: (Use the chart of reduction potentials)

a) Cu2+/Cu, Ag+/Ag

b) Cd2+/Cd, Co2+/Co

c) Na+/Na, Ni2+/Ni

d) I2/I−, Br2/Br−

11. Suppose chemists had chosen to make I2 + 2e− 2I− half cell the standard electrode and had

assigned it a potential of zero volts.

a) What would be the E value for the Br2 + 2e− Br− half cell?

b) What would be the E value for the Al3+ + 3e− Al half cell?

c) How much change would be observed in the E value for the reaction involving Br2 + I− using

the I2 half cell as standard?

12. If a strip of Ni were dipped into a solution of AgNO3, what would be expected to occur?

Explain using E values and equations.

13. a) What would happen if an aluminum spoon were used to stir a solution of Zn(NO3)2?

b) Could a strip of Zn be used to stir a solution of Al(NO3)3? Explain using E values.

14. Can a solution of Sn(NO3)2 be stored in an aluminum container? Explain using E values

15. A voltaic cell is made from a cadmium electrode in a solution of CdSO4 and a zinc electrode

in a solution of ZnSO4. The two half cells are separated by a porous barrier.

a) Identify the anode and cathode.

b) In which direction are the electrons flowing?

c) Write a balanced redox equation for this cell.

d) Determine the Eof the cell.

16. Consider the following electrochemical cell: 3Ag + Cr3+ → 3Ag+ + Cr.

Calculate the voltage of this cell and is it spontaneous?



Solutions

1. Xe, XeF, XeF2, XeOF2, XeO3, CsXeF8

2. a) redox

b) redox

c) redox

d) non redox

e) non redox

f) redox

g) redox

h) non redox

i) redox

3. a) Na is oxidized

Cl2 is reduced

b) C is oxidized

O2 is reduced

c) O is oxidized

H is reduced

f) O is oxidized

Cl is reduced

g) H2 is oxidized

Cl2 is reduced

i) Zn is oxidized

Cu is reduced

4. a) Li

b) They have a weak attraction for

valence electrons and readily

lose e−s to reduce substances

5. a) Strongest: Ca, Weakest: Cl−

b) Str: Al Wk: Br−

c) Str: Na Wk: F−

d) Str: K+ Wk: NO3−

e) Str: Zn2+ Wk: Cl2

f) Str: Li+ Wk: F2

6. a) yes

b) yes

c) no

d) yes

e) no

7. voltaic reactions produce e−

and are spontaneous; electrolytic

reactions require e− and are non

spontaneous

8. a) an electrolytic process that

deposits metal on a surface

b) object plated = cathode

plating metal = anode

9. a) a measure of the energy

required to move a certain

electric charge across the cell

b) measured in volts by a

voltmeter across two electrodes

10. a) 2Ag+ + Cu → 2Ag + Cu2+

E = +0.46V

b) Co2+ + Cd → Co + Cd2+

E = +0.12 V

c) 2Na + Ni2+ → 2Na+ + Ni

E = +2.45 V

d) Br2 + 2I− → 2Br− + I2

E = +0.53 V

11. a) +0.53 V

b) −2.20 V

c) no change

12. Ni + 2Ag+ → Ni2+ + 2Ag

E = +1.06 V

nickel dissolves, solid silver forms

13. a) 2Al + 3Zn2+ → 2Al3+ + 3Zn

E = +0.90 V, spoon disintegrates

solid Zn forms

b) 2Al3+ + 3Zn → 2Al + 3Zn2+

E = −0.90 V, yes, no reaction

14. E = +1.52 V, no, container

dissolves, solid Sn forms

15. a) Cd = cathode, Zn = anode

b) e− flow from Zn to Cd

c) Zn + Cd2+ → Zn2+ + Cd

d) +0.36 V

16. −1.54 V, not spontaneous

Chemistry 30 Redox Balancing a. SnCl2 + HgCl2 SnCl4 + HgCl b. HNO3 + H2S NO + S + H2O c. NaClO + H2S NaCl + H2SO4 d. CdS + I2 + HCl CdCl2 + HI + S e. I2 + HNO3 HIO3 + NO2 + H2O f. MnO4

+ H+ + Cl Mn2+ + Cl2 + H2O g. Sn + HNO3 + H2O H2SnO3 + NO h. K2Cr2O7 + SnCl2 + HCl CrCl3 + SnCl4 + KCl + H2O More Practice a. Br2 + I I2 + Br b. Fe2+ + SeO4

2 + H+ Fe3+ + H2SeO3 + H2O c. Mn2+ + MnO4

+ H2O MnO2 + H+ d. IO3

+ H2S + OH I2 + SO32 + H2O

e. Fe2+ + Cr2O7

2 + H+ Fe3+ + Cr3+ + H2O f. MnO4

+ H3AsO3 + H+ Mn2+ + H2O + H3AsO4 g. I + O2 + H+ I3

+ H2O h. Tl3+ + Mn2+ + H2O Tl+ + MnO4

+ H+ i. VO2+ + I + H+ I3

+ V3+ + H2O Balance the following reactions that occur in acidic conditions. a. I2 + ClO IO3

+ Cl b. H2SO3 + Cr2O7

2 Cr3+ + SO42

c. Cr2O7

2 + C2O42 Cr3 + CO2

Balance the following reactions that occur in basic conditions. a. Cr(OH)3 + ClO3

CrO42 + Cl

b. SO3

2 + CrO42 Cr(OH)4

+ SO42

c. S2 + I2 SO4

2 + I

1

Chem 30 Unit 4 EC1 Practice Test

Multiple Choice

Identify the letter of the choice that best completes the statement or answers the question.

____ 1. What are the oxidation numbers in the compound H2O2?

a. H = +1, O = –2 c. H = +2, O = –2

b. H = –1, O = –2 d. H = +1, O = –1

____ 2. What are the oxidation numbers in the ion SO32–?

a. S = +6, O = –2 c. S = +4, O = –2

b. S = +1, O = –1 d. S = 0, O = –1

____ 3. In an oxidation, atoms or ions

a. increase their oxidation number.

b. decrease their oxidation number.

c. do not change their oxidation number.

d. have a zero oxidation number after the reaction.

____ 4. How does the number of electrons lost in an oxidation compare with the number gained in the simultaneous

reduction?

a. The two numbers are always equal.

b. The number lost is always greater than the number gained.

c. The number lost is always less than the number gained.

d. No relationship exists between the two numbers.

____ 5. A species whose oxidation number decreases in a reaction is

a. oxidized. c. electrolyzed.

b. reduced. d. autooxidized.

____ 6. In the reaction O2 + 4e– → 2O2–, the species O2 is

a. oxidized. c. electrolyzed.

b. reduced. d. autooxidized.

____ 7. In the reaction F2 + Mg → 2F– + Mg2+, which species is oxidized?

a. F2 only c. both Mg and F2

b. Mg only d. neither Mg nor F2

____ 8. Which of the following is an oxidation-reduction reaction?

a. H2 → 2H

b. 2O– → O22–

c. H2 + Cl2 → 2HCl

d. HCl + NaBr → HBr + NaCl

____ 9. Which of the following substances could be produced from SO3 only by an oxidation-reduction reaction?

a. H2SO4 c. SF6

b. H2SO3 d. None of the above

____ 10. In a redox reaction, MnO4– is changed to MnO4

2–. How many electrons must be lost or gained by Mn?

a. none lost or gained c. one gained

b. one lost d. two gained

2

____ 11. In redox reactions,

a. the oxidizing agent is the substance reduced.

b. the reducing agent is the substance reduced.

c. the oxidizing agent is the substance oxidized.

d. both oxidizing and reducing agents are reduced.

____ 12. When hydrogen peroxide decomposes, usually the only products are

a. hydrogen and oxygen. c. oxygen and water.

b. hydrogen and water. d. hydrogen, oxygen, and water.

____ 13. Which of the following reactions is an autooxidation?

a. 2NO2 → N2O4 c. 2NO + O2 → 2NO2

b. N2O3 → NO2 + NO d. N2 + O2 → 2NO

____ 14. If the reactants in a spontaneous energy-releasing redox reaction are in direct contact, the energy is released

in the form of

a. light. c. heat.

b. electrical energy. d. mechanical energy.

____ 15. Where does reduction take place in an electrochemical cell?

a. the anode c. the anode or the cathode

b. the cathode d. the half-cell

____ 16. The voltage of a voltaic cell is determined by the E0 value(s) of the

a. half-reaction at the anode. c. half-reaction at the cathode.

b. half-reactions at the cathode and anode. d. standard hydrogen electrode.

____ 17. In which cell does a current drive a nonspontaneous redox reaction?

a. electrolytic cell c. electrochemical cell

b. dry cell d. voltaic cell

____ 18. In an electrolytic cell, oxidation occurs

a. at the cathode. c. at either the cathode or the anode.

b. at the anode. d. between the cathode and the anode.

____ 19. In an electrolytic cell, the cathode

a. can be either positively or negatively charged.

b. is not charged.

c. is positively charged.

d. is negatively charged.

____ 20. In an electroplating cell, the object to be plated is the

a. external circuit. c. anode.

b. electrolyte. d. cathode.

____ 21. In a cell used to electroplate silver onto an object, Ag+ is

a. oxidized at the anode. c. oxidized at the cathode.

b. reduced at the anode. d. reduced at the cathode.

____ 22. Electrical energy is provided to a rechargeable cell from an outside source when

a. it is charging. c. its external circuit is not closed.

b. it is discharging. d. the porous barrier is in place.

3

Standard Reduction Potentials

Half-cell reaction Standard

electrode

potential,

E0

(in volts)

Half-cell reaction Standard

electrode

potential,

E0

(in volts)

F2 + 2e– á F– +2.87 Fe3+ + 3e– á Fe –0.04

MnO4– + 8H+ + 5e– á Mn2+ + 4H2O +1.50 Pb2+ + 2e– á Pb –0.13

Au3+ + 3e– á Au +1.50 Sn2+ + 2e– á Sn –0.14

Cl2 + 2e– á 2Cl– +1.36 Ni2+ + 2e– á Ni –0.26

Cr2O72– + 14H+ + 6e– á 2Cr3+ + 7H2O +1.23 Co2+ + 2e– á Co –0.28

MnO2 + 4H+ + 2e– á Mn2+ + 2H2O +1.22 Cd2+ + 2e– á Cd –0.40

Br2 + 2e– á 2Br– +1.07 Fe2+ + 2e– á Fe –0.45

Hg2+ + 2e– á Hg +0.85 S + 2e– á S2– –0.48

Ag+ + e– á Ag +0.80 Cr3+ + 3e– á Cr –0.74

Hg22+ + 2e– á 2Hg +0.80 Zn2+ + 2e– á Zn –0.76

Fe3+ + e– á Fe2+ +0.77 Al3+ + 3e– á Al –1.66

MnO4– + e– á MnO4

2– +0.56 Mg2+ + 2e– á Mg –2.37

I2 + 2e– á 2I– +0.54 Na+ + e– á Na –2.71

Cu2+ + 2e– á Cu +0.34 Ca2+ + 2e– á Ca –2.87

Cu2+ + e– á Cu+ +0.15 Ba2+ + 2e– á Ba –2.91

S + 2H+(aq) + 2e– á H2S(aq) +0.14 K+ + e– á K –2.93

2H+(aq) + 2e– á H2 0.00 Li+ + e– á Li –3.04

____ 23. Calculate E0 for the reaction 3Sn2+ + 2Cr → 3Sn + 2Cr3+. Is the reaction spontaneous?

a. –0.88 V; yes c. +0.60 V; yes

b. +0.88 V; no d. +0.60 V; no

____ 24. Calculate E0 for the spontaneous reaction when a Co2+/Co half-cell is joined to a Cu2+/Cu half-cell. Name the

neutral metal produced.

a. +0.62 V; Cu c. +0.06 V; Cu

b. +0.62 V; Co d. +0.06 V; Co

____ 25. In a balanced redox equation, how does the total number of reactant molecules compare with the total

number of product molecules?

a. The two numbers are always equal.

b. Reactant molecules are always more numerous.

c. Product molecules are always more numerous.

d. No relationship exists between the two numbers.

____ 26. In a balanced redox equation, how does the total charge of reactants compare with the total charge of

products?

a. The two totals are always equal.

b. Total reactant charge is always greater.

c. Total product charge is always greater.

d. No relationship exists between the two totals.

____ 27. In a zinc-copper dry cell, oxidation of

a. zinc occurs at the anode. c. zinc occurs at the cathode.

b. copper at the anode. d. copper occurs at the cathode.

4

____ 28. Which is the most active oxidizing agent among the elements?

a. cesium c. fluorine

b. iodine d. lithium

____ 29.

In the figure above, which ion is reduced by Zn but reduces Ag+ to Ag?

a. Mg2+ c. Fe2+

b. F2 d. NO3–

____ 30. In the figure above, which element displaces Cu2+ ions from solution but is displaced by Ni metal when it is

in ionic form?

a. Al c. Fe

b. Ag d. Pb

Short Answer: answer of paper provided (5 marks each)

31. Balance the following redox reaction:

FeCl2 + KMnO4 + HCl → FeCl3 + MnCl2 + H2O + KCl

32. Balance the following redox reaction that takes place in acidic conditions:

Cr2O72– + SO2 → Cr3+ + HSO4

–

33. Balance the following redox equation that takes place in basic conditions:

Bi3+ + SnO22– → Bi + SnO3

2–

ID: A

1

Chem 30 Unit 4 EC1 Practice Test

Answer Section

MULTIPLE CHOICE

1. D

2. C

3. A

4. A

5. B

6. B

7. B

8. C

9. B

10. C

11. A

12. C

13. B

14. C

15. B

16. B

17. A

18. B

19. D

20. D

21. D

22. A

23. C

24. A

25. D

26. A

27. A

28. C

29. C

30. D

SHORT ANSWER

31. 31. 5FeCl2 + KMnO4 + 8HCl → 5FeCl3 + MnCl2 + 4H2O + KCl

32. 5H+ + Cr2O72– + 3SO2 → 2Cr3+ + 3HSO4

– + H2O

33. 6OH− + 2Bi3+ + 3SnO22– → 2Bi + 3SnO3

2– + 3H2O

Chem 30 EQ1 Practice Problems 1

(1)

(2) (3)

4. In general, which reaction is favoured (forward, reverse, or neither) if the value of Keq at a

specified temperature is

a) equal to 1?

b) very small?

c) very large?

5. Predict whether each of the following pressure changes would favour the forward or reverse

reaction. 2NO(g) + O2(g) 2NO2(g)

a) increased pressure

b) decreased pressure

6. In heterogeneous reaction systems, what types of substances do not appear in the equilibrium

constant expression? Why?

7. Explain the effect of a catalyst on an equilibrium system.


8. Predict the effect of each of the following on the equilibrium system in terms of which reaction

would be favoured (forward, reverse, or neither) H2(g) + Cl2(g) 2HCl(g) + 184 kJ

a) addition of Cl2

b) removal of HCl

c) increased pressure

d) decreased temperature

e) removal of H2

f) decreased pressure

g) addition of a catalyst

h) increased temperature

i) decreased system volume

9. How would parts a) through i) of the last question affect the new equilibrium concentration of

HCl and the value of Keq at the new equilibrium?

10. Changes in the concentrations of the reactants and products have no impact on the value of

the equilibrium constant. Explain why.

11. What relative pressure (high or low) would result in the production of the maximum level of

CO2 according to the following? Explain. 2CO (g) + O2(g) 2CO2(g)

12. What relative conditions (reactant concentrations, pressure, and temperature) would favour a

high equilibrium concentration of the underlined substance in each of the following equilibrium

systems?

a) 2CO (g) + O2(g) 2CO2(g) + 167 kJ

b) Cu2+(aq) + 4NH3(aq) Cu(NH3)4

2+(aq) + 42 kJ

c) 2HI(g) + 12.6 kJ H2(g) + I2(g)

d) 4HCl(g) + O2(g) 2H2O(g) + 2Cl2(g) + 113 kJ

e) H2O(l) + 42 kJ H2O(g)

13. Determine the value of the equilibrium constant for each reaction given the following

equilibrium concentrations in mol/L:

a) A + B C [A] = 2.0, [B] = 3.0, [C] = 4.0

b) D + 2E F + 3G [D] = 1.5, [E] = 2.0, [F] =1.8, [G] = 1.2

c) N2(g) + 3H2(g) 2NH3(g) [N2] = 0.45, [H2] = 0.14, [NH3] = 0.62

14. An equilibrium mixture at some temperature is found to consist of 1.2 x 10−3 mol/L HCl,

3.8 x 10−4 mol/L O2, 5.8 x 10−2 mol/L H2O, and 5.8 x 10−2 mol/L Cl2 according to the following:

4HCl(g) + O2(g) 2H2O(g) + 2Cl2(g) Determine the Keq for this system.


15. At 450C the value of the equilibrium constant for the following system is 6.59 x 10−3. If

[NH3] = 1.23 x 10−4 M and [H2] = 2.75 x 10−3 M at equilibrium, determine the concentration of

N2 at that point. N2(g) + 3H2(g) 2NH3(g)

16. The value of the equilibrium constant for the reaction below is 40.0 at a specified

temperature. What would be the value of the constant for the reverse reaction under the same

conditions? H2(g) + I2(g) 2HI(g)

17. Consider the equilibrium: H2(g) + Br2(g) 2HBr(g)

If the initial concentration of hydrogen is 0.020 M and the concentration of hydrogen bromide is

0.0048 M at equilibrium, what is the concentration of hydrogen at equilibrium?

18. For the equilibrium P4(g) + 10Cl2(g) 4PCl5(g), given the following initial concentrations,

[P4] = 0.17 M [Cl2] =0.40 M [PCl5] = 0.02 M what are the equilibrium

concentrations of P4 and PCl5 if the equilibrium concentration of chlorine is 0.08 M?

19. For the equilibrium H3PO4(aq) + H2O(l) H3O+

(aq) + H2PO4−

(aq), if the concentration of a 0.800

M phosphoric acid solution decreases to 0.750 M at equilibrium, determine the Keq.

20. Consider the equilibrium: 2Fe(g) + 3SnCl4(g) 2FeCl3(g) + 3SnCl2(g). If the equilibrium

concentrations of Fe and SnCl2 are 0.48 M and 0.13 M respectively, what is the initial

concentration of Fe? If the concentration of SnCl4 is 0.0032 M at equilibrium, find the Keq.

21. Starting with 0.500 M N2, 0.800 M H2, and no products, a reaction produces NH3 with a

concentration of 0.150 M at equilibrium. The reaction takes place at 250C and all components

are gases. Determine the Keq for this reaction.

22. Consider the equilibrium CH4(g) + H2O(g) CO(g) + 3H2(g).

If the reactants have the same initial concentration, the equilibrium concentration of CO is

0.472 M, and the Keq is 2.01, find the equilibrium concentration of CH4, H2O and H2.

Solutions

1. 12.20 a) left

b) right

c) left

d) left

e) left

f) right

g) no shift

h) left

i) right

2. 12.21 a) left, more pink

b) right, more blue

c) right, more red

d) right, more solid forms

e) left, less solid forms

f) left, more blue


g) no shift, no change

h) left, more violet

i) right, less violet

3. 12.22 a) right, H2O, CO2; H+

b) left, H2O; H2CO3

c) right, H2CO3; H2O, CO2

d) right, LiHCO3; LiOH

e) right, Na2O, CO2, H2O ; NaHCO3

f) right, CaO, CO2; CaCO3

g) left, CO2; solids don’t change

4. a) neither

b) reverse

c) forward

5. a) forward

b) reverse

6. Pure substances in the solid and liquid

phases; because their concentrations are

not changed by the addition or removal

of quantities of these materials.

7. It increases the rates of forward and

reverse reactions equally, so the relative

equilibrium amounts are not affected.

8. a) forward

b) forward

c) neither

d) forward

e) reverse

f) neither

g) neither

h) reverse

i) neither

9. [HCl] and Keq respectively would:

a) increase, remain the same

b) decrease, remain the same

c) both remain the same

d) both increase

e) decrease, remain the same

f) both remain the same

g) both remain the same

h) both decrease

i) both remain the same

10. Such changes have an equal effect on

the numerator and the denominator, so

Keq, the ratio is not affected.

11. High pressure because the forward

reaction converts three molecules into

two, relieving the stress imposed by the

pressure increase.

12. a) high reactant concentrations, high

pressure, low temperature

b) high reactant concentrations, pressure

not relevant, low temperature

c) high reactant concentrations, pressure

not relevant, high temperature

d) high reactant concentrations, high

pressure, low temperature

e) high reactant concentrations, low

pressure, high temperature

13. a) 0.67

b) 0.52

c) 310

14. 1.4 x 1010

15. 1.10 x 102 M

16. 0.0250

17. 0.018 M

18. P4 = 0.14 M, PCl5 = 0.15 M

19. Keq = 3.3 x 10−3

20. Fe = 0.57 M, Keq = 2200

21. Keq =0.278

22. CH4 = 0.817 M, H2O = 0.817 M,

H2 = 1.416 M


1. a) What is a saturated solution?

b) What visible evidence indicates that a solution is saturated?

c) What is an unsaturated solution?

2. What is meant by the solubility of a substance?

b) What conditions must be specified when expressing the solubility of a substance?

3. a) What rule is useful for predicting whether one substance will dissolve in another?

b) Describe what the rule means in terms of various combinations of polar and nonpolar solutes

and solvents.

4. Based on the graph of Temperature vs. Solubility,

determine the solubility of each of the following in

grams per 100 g of water:

a) NaNO3 at 10C

b) KNO3 at 60C

c) NaCl at 50C

5. Based on the same graph, at what temperature would

each of the following solubility levels be observed?

a) 40 g KCl in 100 g H2O

b) 100 g NaNO3 in 100 g H2O

c) 50 g KNO3 in 100 g H2O

6. Write the equation for the dissolution of each of the following ionic compounds in water.

a) KI

b) NaNO3

c) MgCl2

d) Na2SO4

7. For the compounds listed in the previous problem, determine the number of moles of each ion

produced as well as the total number of moles of ions produced when 1 mol of each compound

dissolves in water.

8. Write the equation for the dissolution of each of the following in water and indicate the total

number of moles of solute ions formed.

a) 0.50 mol strontium nitrate

b) 0.50 mol sodium phosphate

c) 3.28 mol aluminum sulfate


9. Write the balanced chemical equation, the overall ionic equation, and the net ionic equation for

the following reactions:

a) mercury(II) chloride (aq) + potassium sulfide (aq) →

b) sodium carbonate (aq) + calcium chloride (aq) →

c) copper(II) chloride (aq) + ammonium phosphate (aq) →

10. Identify the spectator ions in the reaction between KCl and AgNO3 in aqueous solution.

11. Copper(II) chloride and lead(II) nitrate react in aqueous solutions to form copper(II) nitrate

and lead(II) chloride. Write the balanced chemical equation, the overall ionic equation, and the

net ionic equation for this reaction. If 13.45 g of copper(II) chloride react, what is the maximum

amount of precipitate that could be formed?

12. The ionic substance EJ dissociates into E2+ and J2− ions. The solubility of EJ is 8.45 x 10−6

mol/L. What is the value of the solubility product constant?

13. Calculate the Ksp for each of the following, based on the solubility information provided:

a) BaSO4 = 2.4 x 10−4 g/100.0 mL of water at 20C

b) Ca(OH)2 = 0.173 g/100.0 mL of water at 20C

14. Calculate the molar solubility of substance MN that ionizes into M2+ and N2− ions, given that

Ksp = 8.1 x 10−6.


15. Use the Ksp values in table 18-3 to evaluate the solubility of each of the following in moles

per litre.

a) AgBr

b) CoS

c) Mg(OH)2

16. The Ksp of CaSO4 is 9.1 x 10−6. What is the molar concentration of CaSO4 in a saturated

solution (molar solubility)?

17. A salt has the formula X2Y and its Ksp is 4.25 x 10−7. Calculate the molar solubility of X2Y.

18. The ionic substance T3U2 ionizes to form T2+ and U3− ions. The solubility of T3U2 is 3.77 x

10−20 mol/L. What is the value of the Ksp?

19. A solution of AgI contains 2.7 x 10−10 mol/L Ag+. What is the maximum concentration I− that

can exist in this solution at 25C? (hint – what is the Ksp of AgI?)

20. Will a precipitate of Ca(OH)2 form when 320.0 mL of a 0.046 M solution of NaOH mixes

with 400.0 mL of a 0.085 M CaCl2 solution at 25C? Show the ion product.

21. A 25.0 mL solution of 0.0500 M Pb(NO3)2 is combined with 25.0 mL of 0.0400 M Na2SO4.

a) Write the solubility equilibrium for the potential precipitate.

b) Write the solubility product constant expression for the precipitate. Include the Ksp.

c) Find the ion product of the precipitate ions for this solution.

d) Does a precipitate form in this solution?

22. 20.00 mL of a 0.077 M solution of silver nitrate is mixed with 30.00 mL of a 0.043 M

solution of sodium acetate. Does a precipitate form at 37C if the Ksp of AgC2H3O2 is 2.5 x 10−3?

Show the ion product.

23. If you mix 100.0 mL of 0.036 M Pb(C2H3O2)2 with 50.0 mL of 0.074 M NaCl, will a

precipitate of PbCl2 form at 42C if its Ksp is 1.9 x 10−4? Show the ion product.

24. If 20.00 mL of a 0.0090 M solution of ammonium sulfide is mixed with 120.00 mL of a

0.0082 M solution of aluminum nitrate, does a precipitate form at a given temperature if the Ksp

of aluminum sulfide is 2.00 x 10−7? Show the ion product.

25. Calculate whether a precipitate will form if 0.35 L of 0.0044 M Ca(NO3)2 and 0.17 L of

0.000 39 M NaOH are mixed at 25C. Show the ion product.


26. Determine whether a precipitate will form if 1.70 g of solid AgNO3 and 14.5 g of solid NaCl

are dissolved in enough water to make 200.0 mL of solution at 25C. Show the ion product.

27. If 2.50 g of solid Fe(NO3)3 is added to 100.0 mL of a 1.0 x 10−20 M NaOH solution, will a

precipitate form at 25C? Show the ion product. (assume the solution volume remains 100.0 mL)

University level challenge:

Tooth enamel is contains the mineral hydroxyapatite, Ca5(PO4)3OH, which has a Ksp of 6.4 x

10−37. The molar solubility of hydroxyapatite is 2.7 x 10−5 mol/L. When hydroxyapatite is reacted

with the fluoride ion, it forms fluorapatite, Ca5(PO4)3F, which is harder and less susceptible to

caries and cavities. The Ksp of fluorapitite is 1.0 x 10−60. Calculate the molar solubility of

fluorapatite in water. Given your calculations, can you support the fluoridation of drinking water?

Solutions

1. a) A saturated solution is one that contains

the maximum possible amount of dissolved

solute at solution equilibrium under the

existing conditions.

b) A residual quantity of undissolved solute

remains in contact with a saturated solution.

c) An unsaturated solution contains less

solute than does a saturated solution under

the existing conditions.

2. a) The solubility of a substance is the amount

of that substance that is dissolved at solution

equilibrium at a specified temperature.

b) Temperature and pressure (for gases) must


be specified.

3. a) The rule of thumb for predicting solubility

is “like dissolves like.”

b) In general, the rule means that polar

substances dissolve in polar solvents, and

nonpolar substances dissolve in nonpolar

solvents.

4. a) about 84 g/100 g solvent

b) about 105 g/100 g solvent

c) about 36 g/100 g solvent

5. a) 25C

b) 40C

c) 32C

6. (for all reactions H2O goes over the arrow)

a) KI(s) → K+(aq) + I−(aq)

b) NaNO3(s) →Na+(aq) + NO3

−(aq)

c) MgCl2(s) → Mg2+(aq) + 2Cl−(aq)

d) Na2SO4(s) → 2Na+(aq) + SO4

2− (aq)

7. a) 1 mol K+ , 1 mol I−, 2 mol ions

b) 1 mol Na+ , 1 mol NO3−, 2 mol ions

c) 1 mol Mg2+ , 2 mol Cl−, 3 mol ions

d) 2 mol Na+ , 1 mol SO42−, 3 mol ions

8. a) Sr(NO3)2(s) →Sr2+(aq) + 2NO3

−(aq)

1.50 mol ions

b) Na3PO4(s) →3Na+(aq) + PO4

3−(aq)

2.00 mol ions

c) Al2(SO4)3(s) →2Al3+(aq) + 3SO4

2−(aq)

16.40 mol ions

9. a) HgCl2(aq) + K2S(aq) → HgS(s) + 2KCl(aq)

Hg2+(aq) + 2Cl−(aq) + 2K+

(aq) + S2−(aq) →

HgS(s) + 2K+(aq) + 2Cl−(aq)

spectator ions: K+ and Cl−

precipitate: HgS

Hg2+(aq) + S2−

(aq) → HgS(s)

b) Na2CO3(aq) + CaCl2(aq) →

CaCO3(s) + 2NaCl(aq)

2Na+(aq) + CO3

2− (aq) + Ca2+

(aq) + 2Cl−(aq) →

CaCO3(s) + 2Na+(aq) + 2Cl−(aq)

spectator ions: Na+ and Cl−

precipitate: CaCO3

Ca2+(aq) + CO3

2− (aq) → CaCO3(s)

c) 3CuCl2(aq) + 2(NH4)3PO4(aq) →

Cu3(PO4)2(s) + 6NH4Cl(aq)

3Cu2+(aq) + 6Cl− (aq) + 6NH4

+(aq) + 2PO4

3−(aq) →

Cu3(PO4)2(s) + 6NH4+

(aq) + 6Cl−(aq)

spectator ions: NH4+ and Cl−

precipitate: Cu3(PO4)2

3Cu2+(aq) + 2PO4

3−(aq) → Cu3(PO4)2(s)

10. K+ and NO3−

11. CuCl2(aq) + Pb(NO3)2(aq) →

Cu(NO3)2(aq) + PbCl2(s)

Cu2+(aq) + 2Cl−(aq) + Pb2+

(aq) + 2NO3−

(aq) →

Cu2+(aq) + 2NO3

−(aq + PbCl2(s)

Pb2+(aq) + 2Cl−(aq) → PbCl2(s)

27.82 g PbCl2

12. 7.14 x 10−11

13. a) 1.1 x 10−10

b) 5.09 x 10−5

14. 2.8 x 10−3 mol/L

15. a) 7.1 x 10−7 mol/L

b) 6.3 x 10−11 mol/L

c) 1.7 x 10−4 mol/L

16. 3.0 x 10−3 mol/L

17. 4.74 x 10−3 mol/L

18. 8.22 x 10−96

19. 3.1 x 10−7 mol/L

20. ppt occurs, ion prod = 2.0 x 10−5

21. a) PbSO4(s) Pb2+ (aq) + SO4

2− (aq)

b) Ksp = [Pb2+] [SO42−] = 1.6 x 10−8

c) ion prod = 5.00 x 10−4

d) ppt forms, ion prod > Ksp

22. no ppt, ion prod = 7.9 x 10−4

23. no ppt, ion prod = 1.5 x 10−5

24. no ppt, ion prod = 1.1 x 10−13

25. no ppt, ion prod = 4.8 x 10−11

26. AgCl will ppt, ion prod = 6.21 x 10−2

27. no ppt, ion prod = 1.0 x 10−61

Challenge

Solubility is 6.1 x 10−8 M. The fluoridation

produces a lowered solubility, which protects

tooth enamel.


1. Name each of the following binary acids:

a) HCl

b) H2S

c) HBr

2. Name each of the following oxyacids:

a) HNO3

b) H2SO3

c) HClO3

d) HNO2

3. Write the formulas for each of the following binary acids:

a) hydrofluoric acid

b) hydriodic acid

4. Write the formulas for each of the following oxyacids:

a) chlorous acid

b) phosphoric acid

c) acetic acid

d) chromic acid

e) perchloric acid

f) perbromic acid (bromate is BrO3−)

g) oxalic acid

5. Identify and describe the characteristic properties of five common acids: sulfuric acid, nitric

acid, phosphoric acid, hydrochloric acid, and acetic acid. Give examples of uses of each. (p.456)

6. Although HCl(aq) exhibits Arrhenius acidic properties, pure HCl gas and HCl dissolved in a

nonpolar solvent exhibit no acidic properties in the Arrhenius sense. Explain why.

7. a) What distinguishes strong acids from weak acids?

b) Give two examples of each.

8. H3PO4, which contains three hydrogen atoms per molecule, is a weak acid, whereas HCl,

which contains only one hydrogen per molecule, is a strong acid. Explain how this is possible.

9. a) What determines the strength of an Arrhenius base?

b) Give one example each of solutions that are strongly and weakly basic.

10. Distinguish among monoprotic, diprotic, and triprotic acids. Give an example of each.


11. Which of the three acid definitions is the broadest? Explain.

12. Define and give an equation to illustrate each of the following:

a) a conjugate base

b) a conjugate acid

13. a) What is the relationship between the strength of an acid and that of its conjugate base?

b) What is the relationship between the strength of a base and that of its conjugate acid?

14. a) What trend is there in the favoured direction of proton-transfer reactions?

b) What determines the extent to which a proton-transfer reaction occurs?

15. a) What is meant by the term amphoteric?

b) Give an example of a substance or ion (other than water) with amphoteric characteristics.

16. For each reaction listed, identify the proton donor or acid and the proton acceptor or base.

Label each conjugate acid-base pair.

a) CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO−

(aq)

b) HCO3−

(aq) + H2O(l) H2CO3(aq) + OH−(aq)

c) HNO3(aq) + SO42−

(aq) HSO4−

(aq) + NO3−

(aq)


17. Based on the information given in table 15-6, determine the following relative to HF, H2S,

HNO3 and CH3COOH:

a) the strongest acid.

b) the weakest acid.

c) the strongest conjugate base among the four produced by the acids listed.

d) the weakest conjugate base among the four produced by the acids listed.

18. Explain why the conjugate base of a strong acid is a weak base and the conjugate acid of a

strong base is a weak acid.

19. a) Write the balanced equations that describe the two-stage ionization of sulfuric acid in a

dilute aqueous solution.

b) How do the degrees of ionization in the two steps compare?

20. Dilute HCl(aq) and KOH(aq) are mixed in chemically equivalent quantities. Write the

following:

a) the formula equation for the reaction

b) the overall ionic equation

c) the net ionic equation

21. Repeat the previous question with H3PO4(aq) and NaOH(aq).

22. Write the formula equation and net ionic equation for each of the following reactions:

a) Zn(s) + HCl(aq) →

b) Al(s) + H2SO4(aq) →

23. Write the formula equation and net ionic equation for the reaction between Ca(s) and HCl(aq).

24. Complete the following neutralization reactions. Balance each reaction, then write the overall

ionic and net ionic equation for each.

a) HCl(aq) + NaOH(aq) →

b) HNO3(aq) + KOH(aq) →

c) Ca(OH)2(aq) + HNO3(aq) →

d) Mg(OH)2(aq) + HCl(aq) →

25. Write the formula equation, the overall ionic equation, and the net ionic equation for the

neutralization reaction involving aqueous solutions of H3PO4 and Mg(OH)2. Assume the

solutions are too dilute to form precipitates.


26. Write the balanced chemical equation for each of the following reactions between an acid and

a carbonate:

a) BaCO3(s) + HCl(aq) → c) Na2CO3(s) + H2SO4(aq) →

b) MgCO3(s) + HNO3(aq) → d) CaCO3(s) + H3PO4(aq) →

27. Write the formula equation, the overall ionic equation, and the net ionic equation for the

neutralization reaction that would form each of the following salts:

a) RbClO4 c) CaCl2

b) BaSO4 d) K2SO4

28. Zinc reacts with 100.0 mL of 6.00 M cold, aqueous sulfuric acid through single replacement.

a) How many grams of zinc sulfate are produced?

b) How many litres of hydrogen gas would be released at STP?

29. A 211 g sample of barium carbonate is placed in a solution of nitric acid. Assuming that the

acid is present in excess, what mass and volume of dry carbon dioxide gas at STP will be

produced?

30. A seashell, composed largely of calcium carbonate, is placed in a solution of HCl. As a result,

1500 mL of dry CO2 gas at STP is produced.

a) How many grams of calcium carbonate are consumed in the reaction?

b) What volume of 2.00 M HCl solution is used in this reaction?

31. Calculate the [H3O+] and [OH−] for each of the following:

a) 0.03 M HCl c) 5 x 10−4 M H2SO4

b) 1 x 10−4 M NaOH d) 0.01 M Ca(OH)2

32. Determine the pH of each of the following solutions.

a) 1.0 x 10−2 M HCl c) 1.0 x 10−5 M HI

b) 1.0 x 10−3 M HNO3 d) 1.0 x 10−4 M H2SO4

33. Given the following [OH−] values, determine the pH of each solution.

a) 1.0 x 10−6 M c) 1.0 x 10−2 M

b) 1.0 x 10−9 M d) 1.0 x 10−7 M

34. Determine the pH of each solution.

a) 1.0 x 10−2 M NaOH c) 1.0 x 10−4 M LiOH

b) 1.0 x 10−3 M KOH d) 1.0 x 10−5 M Sr(OH)2


35. Determine the pH of solutions with each of the following [H3O+].

a) 2.0 x 10−5 M c) 3.8 x 10−3 M

b) 4.7 x 10−7 M d) 4.0 x 10−1 M

36. Given the following pH values, determine the [H3O+] for each solution.

a) 3.0 d) 5.47

b) 7.00 e) 3.28

c) 11.0 f) 10.64

37. Given the following pH values, determine the [OH−] for each solution.

a) 7.00 d) 6.63

b) 11.00 e) 13.25

c) 4.00 f) 4.88

38. A nitric acid solution has a pH of 2.70. Deternine:

a) [H3O+]

b) [OH−]

c) the mass of the solute in 5.50 L of this acid solution.

39. If 1.00 L of a potassium hydroxide solution with a pH of 12.90 is diluted to 2.00 L, what is

the pH of the resulting solution?

40. Suppose that 15.0 mL of 2.50 x 10−2 M aqueous H2SO4 is required to neutralize 10.0 mL of

an aqueous solution of KOH. What is the molarity of the KOH?

41. In a titration, 12.5 mL of 1.75 x 10−2 M Ba(OH)2 neutralized 14.5 mL of HNO3. Calculate the

molarity of the HNO3 solution.

42. A student titrates a 20.00 mL sample of HBr with unknown molarity. The titration requires

20.05 mL of a 0.1819 M solution of NaOH. What is the molarity of the HBr solution?

43. A 20.00 mL sample of a solution of Sr(OH)2 is titrated to the equivalence point with 43.03

mL of 0.1159 M HCl. What is the molarity of the Sr(OH)2 solution?

44. A 35.00 mL sample of ammonia solution is titrated to the equivalence point with 54.95 mL of

a 0.400 M sulfuric acid solution. What is the molarity of the ammonia solution?

45. At 25C, a 0.025 M solution of formic acid, a monoprotic acid with the formula HCOOH, is

found to have a hydronium ion concentration of 2.03 x 10−3 M. Calculate the acid ionization

constant ( Ka ) of formic acid.


46. The pH of a 0.400 M solution of iodic acid, HIO3, is 0.726 at 25C. What is the Ka at this

temperature?

47. The pH of a 0.150 M solution of hypochlorous acid, HClO, is found to be 4.55 at 25C.

Calculate the Ka for HClO at this temperature?

48. The compound propylamine, CH3CH2CH2NH2 is a weak base. At equilibrium, a 0.039 M

solution of propylamine has a hydroxide ion concentration of 3.74 x 10−3 M. Calculate the pH of

this solution and the Kb for propylamine.

49. (Bonus Challenge) The Ka of nitrous acid is 4.6 x 10−4 at 25C. Calculate the [H3O+] of a

0.0450 M nitrous acid solution.

Solutions

1. a) hydrochloric acid

b) hydrosulfuric acid

c) hydrobromic acid

2. a) nitric acid

b) sulfurous acid

c) chloric acid

d) nitrous acid

3. a) HF

b) HI

4. a) HClO2

b) H3PO4

c) CH3COOH

d) H2CrO4

e) HClO4

f) HBrO4

g) H2C2O4

5. H2SO4 is the most common acid, a

dehydrating agent, used to make fertilizers,

in petroleum refining, in metallurgy, and in

car batteries. H3PO4 is used in making

fertilizers, cleaners, and animal feed. It is

also used as a flavouring agent. HNO3 is a

volatile, unstable liquid that stains proteins

yellow. It is used to make fertilizers,

explosives, rubber, plastics, dyes, and drugs.

HCl, also called muriatic acid, is used as a

cleaning agent, in food processing, to pickle

iron, and in activating oil wells. HCl is also

produced in the stomach to aid digestion.

CH3COOH is a clear, colourless organic acid

produced by the fermentation of malt, barley,

and fruit juices. It is found in household

vinegar and is used to make plastics, food

supplements, and fungicides.

6. HCl(g) consists of covalently bonded

molecules that do not ionize. Nonpolar

solvent molecules do not attract HCl

molecules to cause them to be ionized.

7. a) strong acids ionize completely in dilute

aqueous solution; weak acids ionize much

less.

b) Examples - strong acids: HCl, HNO3;

weak acids: HF, H3PO4

8. The strength of an acid depends on the

degree of ionization, not the amount of

hydrogen in the molecule. HCl ionizes

completely, H3PO4 only ionizes slightly.

9. a) Strength is determined by the

concentration of OH− ions in solution.

b) Answers vary – NaOH is strongly basic,

NH3 is weakly basic

10. Monoprotic acids, such as HCl, can donate

only one proton per molecule; diprotic acids,

such as H2SO4, can donate two protons per

molecule; triprotic acids, such as H3PO4, can

donate three protons per molecule.

11. The Lewis definition is the broadest. Any


substance defined as an Arrhenius or

Bronsted-Lowry acid is also a Lewis acid.

The reverse is not always true.

12. a) A conjugate base remains after an acid

has given up a proton. Example:

HF + H2O → H3O+ + F−

acid: HF; conjugate base: F−

b) A conjugate acid is what forms when a

proton is added to a base. Example:

CH3NH2 + H2O → CH3NH3+ + OH−

base: CH3NH2; conjugate acid: CH3NH3+

13. The stronger an acid, the weaker its

conjugate base.

b) The stronger a base, the weaker its

conjugate acid.

14. a) Production of the weaker acid and

weaker base is favoured.

b) It depends on the relative strengths of

the acids and bases involved. For an

acid-base reaction to approach

completion, the reactants must be much

stronger acids and bases than the

products.

15. a) Amphoteric describes a species that

can react as either an acid or a base.

b) Example: H2PO4− , the conjugate base of

H3PO4, can act as an acid to further ionize to

HPO4− .

16. a) acid: CH3COOH, conjugate base:

CH3COO− ; base: H2O, conjugate acid:

H3O+

b) acid: H2O, conjugate base:

OH− ; base: HCO3−, conjugate acid:

H2CO3

c) acid: HNO3, conjugate base: NO3− ;

base: SO42−, conjugate acid: HSO4

−

17. a) HNO3

b) H2S

c) HS−

d) NO3−

18. Strong acids and bases are readily ionized in

solution. For them to remain ionized, the

corresponding conjugate base and acid,

respectively, must be too weak to

successfully compete with them.

19. H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4

−(aq)

HSO4−

(aq) + H2O(l) → H3O+(aq) + SO4

2− (aq)

b) The degree of ionization in the first stage

is much greater than in the second.

20. a) HCl(aq) + KOH(aq) → KCl(aq) + H2O(l)

b) H3O+

(aq) + Cl−(aq) + K+(aq) + OH−

(aq) →

K+(aq) + Cl−(aq) + 2H2O(l)

c) H3O+

(aq) + OH−(aq) → 2H2O(l)

21. a) H3PO4(aq) + 3NaOH(aq) →

Na3PO4(aq) + 3H2O(l)

b) 3H3O+

(aq) + PO43−

(aq) + 3Na+(aq) + 3OH−

(aq)

→ 3Na+(aq) + PO4

3−(aq) + 6H2O(l)

c) H3O+

(aq) + OH−(aq) → 2H2O(l)

22. a) Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2 (g)

((Zn(s) + 2H3O+

(aq) + 2Cl−(aq) →

Zn2+ (aq) + 2Cl−(aq) + H2 (g) + 2H2O(l)))

Zn(s) + 2H3O+

(aq) + → Zn2+(aq) + H2(g) + 2H2O(l)

b) 2Al(s)+ 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2 (g)

((2Al(s) + 6H3O+

(aq) + 3SO42−

(aq) →

2Al3+(aq) + 3SO4

2−(aq) + 3H2(g) + 6H2O(l)))

2Al(s) + 6H3O+

(aq)→ 2Al3+(aq)+ 3H2(g) + 6H2O(l)

23. a) Cas) + 2HCl(aq) → CaCl2(aq) + H2 (g)

((Ca(s) + 2H3O+

(aq) + 2Cl−(aq) →

Ca2+ (aq) + 2Cl−(aq) + H2 (g) + 2H2O(l)))

Ca(s) + 2H3O+

(aq) + → Ca2+(aq) + H2(g) + 2H2O(l)

24. a) HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

H3O+

(aq) + Cl−(aq) + Na+(aq) + OH−

(aq) →

Na+(aq) + Cl−(aq) + 2H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

b) HNO3(aq) + KOH(aq) →KNO3(aq) + H2O(l)

H3O+

(aq) + NO3−

(aq) + K+(aq) + OH−

(aq) →

K+(aq) + NO3

−(aq) + 2H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

c) Ca(OH)2(aq) + 2HNO3(aq) →

Ca(NO3)2(aq) + 2H2O(l)

Ca2+(aq) + 2OH−

(aq) + 2H3O+

(aq) + 2NO3−

(aq) →

Ca2+(aq) + 2NO3

−(aq) + 4H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)


d) Mg(OH)2(aq) + 2HCl(aq) →

MgCl2(aq) + 2H2O(l)

Mg2+(aq) + 2OH−

(aq) + 2H3O+

(aq) + 2Cl−(aq) →

Mg2+(aq) + 2Cl−(aq) + 4H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

25. 2H3PO4(aq) + 3Mg(OH)2(aq) →

Mg3(PO4)2(aq) + 6H2O(l)

6H3O+

(aq)+ 2PO43−

(aq)+ 3Mg2+(aq)+ 6OH−

(aq) →

3Mg2+(aq) + 2PO4

3−(aq) + 12H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

26. a) BaCO3(s) + 2HCl(aq) →

BaCl2(aq) + H2O(l) + CO2(g)

b) MgCO3(s) + 2HNO3(aq) →

Mg(NO3)2(aq) + H2O(l) + CO2(g)

c) Na2CO3(s) + H2SO4(aq) →

Na2SO4(aq) + H2O(l) + CO2(g)

d) 3CaCO3(s) + 2H3PO4(aq) →

Ca3(PO4)2(s) + 3H2O(l) + 3CO2(g)

27. a) RbOH(aq)+ HClO4(aq) → RbClO4(aq)+ H2O(l)

Rb+(aq) + OH−

(aq) + H3O+

(aq) + ClO4−

(aq) →

Rb+(aq) + ClO4

−(aq) + 2H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

b) Ba(OH)2(aq) + H2SO4(aq) →

BaSO4(s) + 2H2O(l)

Ba2+ (aq) + 2OH−

(aq) + 2H3O+

(aq) + SO42−

(aq) →

BaSO4(s) + 4H2O(l)

Ba2+ (aq) + 2OH−

(aq) + 2H3O+

(aq) + SO42−

(aq) →

BaSO4(s) + 4H2O(l) (no spec. ions)

c) Ca(OH)2(aq) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)

Ca2+ (aq) + 2OH−

(aq) + 2H3O+

(aq) + 2Cl−(aq) →

Ca2+ (aq) + 2Cl−(aq) + 4H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

d) 2KOH(aq)+ H2SO4(aq) → K2SO4(aq) + 2H2O(l)

2K+(aq) + 2OH−

(aq) + 2H3O+

(aq) + SO42−

(aq) →

2K+(aq) + SO4

2− (aq) + 4H2O(l)

H3O+

(aq) + OH−(aq) → 2H2O(l)

28. a) 96.9 g ZnSO4

b) 13.4 L H2

29. 47.1 g, 24.0 L CO2

30. a) 6.7 g CaCO3

b) 0.067 L

31. a) H3O+ = 3 x 10−2 M, OH− = 3 x 10−13 M

b) H3O+ = 1 x 10−10 M, OH− = 1 x 10−4 M

c) H3O+ = 1 x 10−3 M, OH− = 1 x 10−11 M

d) H3O+ = 5 x 10−13 M, OH− = 2 x 10−2 M

32. a) 2.00 c) 5.00

b) 3.00 d) 3.70

33. a) 8.00 c) 12.00

b) 5.00 d) 7.00

34. a) 12.00 c) 10.00

b) 11.00 d) 9.30

35. a) 4.70 c) 2.42

b) 6.33 d) 0.40

36. a) 1 x 10−3 M

b) 1.0 x 10−7 M

c) 1 x 10−11 M

d) 3.4 x 10−6 M

e) 5.2 x 10−4 M

f) 2.3 x 10−11 M

37. a) 1.0 x 10−7 M

b) 1.0 x 10−3 M

c) 1.0 x 10−10 M

d) 4.3 x 10−8 M

e) 1.8 x 10−1 M

f) 7.6 x 10−10 M

38. a) 2.0 x 10−3 M

b) 5.0 x 10−12 M

c) 0.69 g HNO3

39. pH = 12.60

40. 7.50 x 10−2 M KOH

41. 3.02 x 10−2 M HNO3

42. 0.1824 M HBr

43. 0.1247 M Sr(OH)2

44. 1.26 M NH3

45. Ka = 1.8 x 10−4

46. Ka = 0.167

47. Ka = 5.3 x 10−9

48. pH = 11.573

Kb = 4.0 x 10−4

49. [H3O+] = 4.3 x 10−3 M

Chem 30 Final Review: Topics

-Democritus' model -Bohr's model -periodic table and e- configuration

-Dalton's model -the quantum model -valence electrons

-J. J. Thomson's model -quantum numbers -hadrons

-Rutherford's model -orbital notation -isotopes

-light as waves + particles -electron configuration -periodic trends: atomic radii, IE,

- c = f -Aufbau principle electron affinity + electronegativity

- E = hf -Hund's rule -Lewis structures

- KE = hf – -Pauli exclusion principle -VSEPR theory and geometry

-Spectra -noble gas notation -molecular polarity

-types of bonds -dipole-dipole forces -Charles’ Law

-nonpolar-covalent bonds -hydrogen bonding -Gay-Lussac's Law

-polar covalent bonds -London dispersion forces -Ideal Gas Law

-ionic bonds -surface tension, capillary -vapour pressure

-covalent bonds orb. not. action + viscosity -changes of state

-hybridization -solids: crystals, amorphous -evaporation and boiling

-sigma + pi bonds -liquids and fluids -phase diagrams

-metallic bonds -gases ideal and real -P1 V1 = P2V2

-bond + lattice energy -Avogadro’s Law T1 T2

-intermolecular forces -Boyle’s Law -m = DV

-PV = nRT

-types of formulas -isomers: structural and geometric

-naming and drawing: -functional groups

alkanes, alkenes, alkynes -polymers: addition and condensation

-oxidation numbers -electrochemistry

-oxidation-reduction reactions -voltaic cells -wet cells, dry cells

-LEO GER -An Ox, Red Cat

-half reactions -electrolytic cells

-redox and covalent bonds -electroplating

-balancing redox reactions -electrode potentials

-hydrogen peroxide -Eº = Eºcathode – Eºanode

-equilibrium -calculating Keq

-Le Châtelier's Principle -equilibrium shift: press. conc. temp.

-favoured reactions -factors affecting Keq

-the equilibrium constant Keq -ICE method

-solubility -precipitation reactions

-common ion effect -calculating Ksp

-electrolytes/nonelectrolytes -using Ksp to find solubility

-strong/weak electrolytes -predicting precipitates

-dissociation/ionization

-Acid names -ionization constant of water:

-Arrhenius acids and bases -Kw = [H3O+] [OH–] = 1.0 10–14 M2

-strength of acids and bases -calculating concentrations

-Brønsted-Lowry acids and bases -pH, pOH

-polyprotic acids -titration

-Lewis acids and bases -calculations with pH and titration

-conjugate acids and bases -Ka and pKa

-amphoteric/amphiprotic -Kb and pKb

-neutralization reactions -calculations of Ka, pKa, Kb, and pKb

Equations: 7) m = DV

1) n = m 8) PV = nRT

M

9) Eº = Eºcathode – Eºanode

2) C = n

V 10) Keq = [C]c[D]d Ka, Kb and Ksp

[A]a[B]b

3) c = f

11) C1V1 = C2V2

4) E = hf

12) Kw = [H3O+] [OH–] = 1.0 x 10–14 M2

5) KE = hf –

13) pH = – log [H3O+] pOH = – log [OH–]

6) P1 V1 = P2V2

T1 T2 14) [H3O+]= antilog – pH [OH–]= antilog – pOH

Problems:

1. a) What is the frequency of green light at 552 nm? b) How much energy is there in a photon of

this light?

2. What is the binding energy of unobtainium if a photon with 9.68 x 10−21 J is emitted when

struck with infrared light at 4.95 x 1013 Hz

3. Write the a) electron configuration b) orbital notation and c) noble gas notation for zinc.

4. Magnesium has 3 stable isotopes. 24Mg has a mass of 23.985042 amu and an abundance of

78.99%. 26Mg has a mass of 25.982593 amu and an abundance of 11.01% What is the mass of

the third isotope if the average mass of magnesium is 24.305 amu?

5. Draw the Lewis structure for nitrite. Indicate its electron group geometry, molecular geometry,

bond angle, hybridization, whether it is polar or non polar, and whether it exhibits resonance.

6. An oxygen tank is filled at sea level (T = 25.00C) to a pressure of 8.00 atm. The tank is

carried to the summit of Everest (T = −32.00C). What is the pressure at this new temperature?

7. A 35.0 L propane tank is filled to 18.6 atm with 1.12 kg of propane. At what temperature did

this take place?

8. Draw: a) 3-ethyl-2,2,5-trimethylhexane b) 3-methylcyclopentene

c) 6-methyl-hept-3-yne d) pentan-2-one e) ethyl propanoate

9. Balance the following redox reaction both in acidic and basic conditions:

MnO4− + S2− → Mn2+ + S

10. Calculate the electrode potential for the above redox reaction (in acidic conditions so that

your reduction potentials are accurate) and state whether it is spontaneous.

11. If the equilibrium concentration of all the following reactants and products are 0.25 M,

calculate the equilibrium constant:

Cu(s) + 4H+(aq) + 2NO3

−(aq) 2NO2(g) + Cu2+

(aq) + 2H2O(l)

12. Use a solubility table to write the balanced, overall ionic, and net ionic equations for the

reaction between aqueous solutions of calcium chloride and potassium phosphate.

13. Calculate the Ksp of CaF2 if its solubility is 1.6 mg/100.0 g at 20C.

14. What is the molar solubility of Al(OH)3 if its Ksp is 3.7 x 10−15 at 25C

15. Find the ion product and predict whether a precipitate will form if 30.0 mL of 0.032 M

lead(II) nitrate is combined with 25.0 mL of 0.066 M sodium sulfate at 25C. The Ksp for PbSO4

is 1.1 x 10−8.

16. Write the balanced equation for the neutralization of strontium hydroxide with perchloric

acid. Show the overall ionic equation and the net ionic equation. Assume the solutions are too

dilute to form precipitates.

17. Find the pH of a 0.124 M Cr(OH)3 solution, assuming 100% dissociation.

18. If 247 mL of 0.048M Ba(OH)2 is needed to titrate 375 mL of chloric acid, what is the

concentration of the acid?

19. What is the molar concentration of a sulfuric acid solution with a pH of 1.79?

20. Calculate the Ka of HClO2 and use it to determine whether chlorous acid is a strong acid.

Given: A 0.250 M solution will ionize to produce a hydronium ion concentration of

8.42 x 10−8 M at equilibrium.

Solutions

1. a) f = 5.43 x 1014 Hz b) E = 3.60 x 10−19 J

2. E = 2.31 x 10−20 J

3. a) 1s2 2s2 2p6 3s2 3p6 4s2 3d10

b)

1s 2s 2p 3s 3p 4s 3d

c) [Ar] 4s2 3d10

4. mass = 25.0 amu

5. E.G.G = trigonal planar bond angle = <120 polarity = polar

M.G. = bent hybridization = sp2 resonance = yes

6. P = 6.47 atm

7. T = 312 K

8. a) b) c) d)

e)

9. acidic: 16H+ + 2MnO4− + 5S2− → 2Mn2+ + 5S + 8H2O

basic: 8H2O + 2MnO4− + 5S2− → 2Mn2+ + 5S + 16 OH−

10. Eº = 1.98 V

11. Keq = 64

12. 3CaCl2(aq) + 2K3PO4(aq) → Ca3(PO4)2(s) + 6KCl(aq) balanced

3Ca2+(aq) + 6Cl−

(aq) + 6K+(aq) + 2PO4

3−(aq) → Ca3(PO4)2(s) + 6K+

(aq) + 6Cl−(aq) overall ionic

3Ca2+(aq) + 2PO4

3−(aq) → Ca3(PO4)2(s) net ionic

13. Ksp = 3.4 x 10−11

14. 1.1 x 10−4 M

15. ion product = 5.2 x 10−4 , yes ppt

16. Sr(OH)2(aq) + 2HClO4(aq) → Sr(ClO4)2(aq) + 2H2O(l) balanced

Sr2+(aq) + 2OH−

(aq) + 2H3O+

(aq) + 2ClO4−

(aq) → Sr2+(aq) + 2ClO4

−(aq) + 4H2O(l) overall ionic

OH−(aq) + H3O

+(aq) → 2H2O(l) net ionic

17. pH = 13.571

18. HClO3 = 0.063 M

19. H2SO4 = 8.1 x 10−3 M

20. Ka = 2.84 x 10−14, weak acid

Documents

1 s-block 18 H The Periodic Table of the Elements 1