Министерство образования и науки Российской Федерации
Федеральное агентство по образованию
Южно-Уральский государственный университет
Химический факультет
Ш143.21-9
Д182
Е.И.Данилина
ХИМИЯ НА АНГЛИЙСКОМ ЯЗЫКЕ
Модуль 2
ОБЩАЯ И НЕОРГАНИЧЕСКАЯ ХИМИЯ
Учебное пособие
Челябинск
Издательский центр ЮУрГУ
2009
ББК Ш143.21-923
УДК 54(075.8)
Одобрено
учебно-методической комиссией химического факультета
Рецензенты:
Балыкин В.П., д-р хим. наук, профессор кафедры аналитической и физической
химии Челябинского государственного университета,
Толчев А.В., д-р хим. наук, профессор, зав. кафедрой общетехнических
дисциплин Челябинского государственного педагогического университета
Д182
Данилина, Е.И.
Химия на английском языке. Модуль 2. Общая и неорганическая
химия: учебное пособие. – Челябинск: Издательский центр ЮУрГУ,
2009. – 48 с.
Учебное пособие составлено на английском языке по материалам
британских, канадских и американских учебников по химии для
колледжей и университетов и предназначено для практических занятий
и самостоятельной работы студентов. В учебном пособии предложены
качественные вопросы и расчетные задачи, охватывающие основные
темы курса общей и неорганической химии. В приложениях приведены
необходимые справочные материалы для численного решения задач и
их устного чтения: Периодическая таблица, таблицы констант
ионизации, произведений растворимости и стандартных электродных
потенциалов, а также транскрипция названий элементов и соединений
по правилам ИЮПАК.
Пособие предназначено для студентов 2 курса химического
факультета.
УДК 54(075.8)
© Издательский центр ЮУрГУ
2
CONTENTS
1. The Structure of the Atom................................................................................ 4
2. The Periodic Table and Periodic Law............................................................... 8
3. Energy and Chemical Change...........................................................................12
4. Rates of Chemical Reactions.............................................................................18
5. Dynamic Chemical Equilibrium........................................................................21
6. Ionic Equilibria..................................................................................................28
7. Electrochemistry................................................................................................35
Appendix 1. Periodic Table of Chemical Elements...............................................38
Appendix 2. Elements and Electronegative Components......................................40
Appendix 3. Acids and Anions..............................................................................43
Appendix 4. Ionization Constants for Acids..........................................................44
Appendix 5. Ionization Constants for Nitrogen Bases...........................................45
Appendix 6. Solubility Product Constants in Water at 25 oC.................................45
Appendix 7. Standard Reduction Potentials...........................................................46
References..............................................................................................................48
3
1. THE STRUCTURE OF THE ATOM
1.1. How many protons and electrons are in each of the following atoms?
a) boron;
c) platinum;
b) radon;
d) magnesium.
1.2. An atom of an element contains 66 electrons. What element is it?
1.3. An atom of an element contains 14 protons. What element is it?
1.4. How many protons and electrons are contained in an atom of element 44?
1.5. For each of the following chemical symbols, determine the element name and the
number of protons and electrons an atom contains.
a) V;
e) Zn;
i) Mo;
b) Ir;
f) Al;
j) Sc;
c) Mn;
g) Cs;
k) Bi;
d) S;
h) Br;
l) Cu.
1.6. A carbon atom has a mass number of 12 and an atomic number of 6. How many
neutrons does it have?
1.7. An isotope of mercury has 80 protons and 120 neutrons. What is the mass number
of this isotope?
1.8. An isotope of xenon has an atomic number of 54 and contains 77 neutrons. What is
the xenon isotope’s mass number?
1.9. How many electrons, protons, and neutrons are contained in each of the following
59
163
70
atoms: 132
55 Cs; 27 Co; 69Tm; 30 Zn.
1.10. How many electrons, protons, and neutrons are contained in each of the following
atoms: a) gallium-64; b) titanium-48; c) fluorine-23; d) helium-8.
1.11. Determine the number of protons, electrons, and neutrons for isotopes in the table.
Name each isotope, and write its symbol.
Element
Neon
Calcium
Oxygen
Iron
Mercury
Atomic number
10
20
8
26
80
4
Mass number
22
46
17
57
204
1.12. Boron has two naturally occurring isotopes, namely: boron-10 (abundance =
19.8%, mass = 10.013 amu), boron-11 (abundance = 80.2%, mass = 11.009 amu).
Calculate the atomic mass of boron.
1.13. Helium has two naturally occurring isotopes, helium-3 and helium-4. The atomic
mass of helium is 4.003 amu. Which isotope is more abundant in nature?
1.14. Chlorine, which has an atomic mass of 35.453 amu, has two naturally occurring
isotopes, Cl-35 and Cl-37. Which isotope occurs in greater abundance?
1.15. Calculate the atomic mass of magnesium. The three magnesium isotopes have
atomic masses and relative abundances of 23.985 amu (78.99%), 24.986 amu (10.00%),
and 25.982 amu (11.01%).
1.16. Silver has two isotopes, 107
49 Ag has a mass of 106.905 amu (52.00%), and
has a mass of 108.905 amu (48.00%). What is the atomic mass of silver?
109
49 Ag
1.17. Calculate the atomic mass of titanium. The five titanium isotopes have atomic
masses and relative abundances of 45.953 amu (8.00%), 46.952 amu (7.30%), 47.948
amu (73.80%), 48.948 amu (5.50%), and 49.945 amu (5.40%).
1.18. Complete the table below.
Isotope
Composition of Various Isotopes
Atomic
Mass
Number of Number of Number of
number
number
protons
neutrons
electrons
32
16
24
20
Zn-64
9
11
10
23
1.19. Write ground-state electron configurations for the following elements.
a) bromine (Br);
d) rhenium (Re);
b) strontium (Sr);
e) terbium (Tb);
c) antimony (Sb);
f) titanium (Ti).
1.20. How many electrons are in orbitals related to the third energy level of a sulfur
atom?
1.21. How many electrons occupy p orbitals in a chlorine atom?
1.22. Sketch the electron arrangements for the first 20 elements in the periodic table.
5
1.23. What element has the following ground-state electron configuration:
[Kr]5s24d105p1?
1.24. What element has the following ground-state electron configuration: [Xe]6s2?
1.25. Write out the electron configuration and draw the orbital diagram for each of the
following elements.
a) silicon;
c) calcium;
b) fluorine;
d) krypton.
1.26. What is a valence electron? Draw the electron-dot structures (Lewis structures) for
the elements in the previous problem (1.25).
1.27. Identify the number of valence electrons in the outer energy levels of the
following elements:
a) chlorine;
f) lead;
b) helium;
g) antimony;
c) indium;
h) selenium;
d) strontium;
i) arsenic;
e) rubidium;
j) xenon.
1.28. Draw electron-dot structures (Lewis structures) for the following elements:
a) magnesium;
d) rubidium;
b) sulfur;
e) thallium;
c) bromine;
f) argon.
1.29. Draw the first 20 elements in the periodic table using Lewis structures (electrondot structures).
1.30. Use the periodic table to draw Lewis (electron-dot) structures for the following
elements: barium (Ba), gallium (Ga), tin (Sn), bismuth (Bi), iodine i), cesium (Cs),
selenium (Se), neon (Ne).
1.31. Which element has the following orbital diagram?
↑↓
1s
↑↓
2s
↑
2p
1.32. Write orbital notations and complete electron configurations for atoms of the
following elements, compare them to electron-dot structures.
a) beryllium;
c) nitrogen;
b) aluminum;
d) sodium.
6
1.33. Use noble-gas notation to describe the electron configurations of the elements
represented by the following symbols.
a) Mn;
d) Zn;
g) Pb;
b) Kr;
e) Zr;
h) Ra;
c) P;
f) W;
i) Sm.
1.34. What elements are represented by each of the following electron configurations?
a) 1s22s22p5;
d) [Kr]5s24d105p4;
b) [Ar]4s2;
e) [Rn]7s25f13;
c) [Xe]6s24f4;
f) 1s22s22p63s23p64s23d104p5.
1.35. Draw electron configurations and electron-dot structures (Lewis structures) for
atoms of each of the following elements:
a) carbon;
d) potassium;
b) arsenic;
e) barium;
c) polonium;
f) indium.
1.36. Answer the following questions for the atom of antimony:
a) How many electron-containing orbitals has the atom?
b) How many of the orbitals are completely filled?
c) How many of the orbitals are associated with the atom’s n = 5 principal energy
level?
1.37. Shown below are the Lewis structures (electron-dot structures) for five elements:
sulfur (S), chlorine (Cl), argon (Ar), potassium k), and calcium (Ca). Answer the
questions below about these structures.
a) Which of the above Lewis structures (electron-dot structures) is the same as the
Lewis structure for the ion S2–? Explain your answer describing the process of electron
gain/loss.
b) Which of the above Lewis structures (electron-dot structures) is the same as the
Lewis structure for the ion Cl–? Explain your answer describing the process of electron
gain/loss.
c) Which of the above Lewis structures (electron-dot structures) is the same as the
Lewis structure for the ion K+? Explain your answer describing the process of electron
gain/loss.
d) Name an ion of calcium that has electron-dot structure similar to that of argon.
Explain your answer describing the process of electron gain/loss.
7
2. THE PERIODIC TABLE AND PERIODIC LAW
2.1. Identify the name and symbol
periodic table:
a) Group IVA, Period 2;
b) Group IB, Period 4;
c) Group VIIIA, Period 6;
d) Group IA, Period 1;
of the elements in the following locations of the
e) Group IIB, Period 5;
f) Group 2 IIA, Period 4;
g) Group VIIA, Period 5;
h) Group IIIA, Period 3.
2.2. Identify the element that is described by the following information. Refer to a
periodic table as necessary.
a) It is a Group IVA element in the third period.
b) It is a Group VA element in the fifth period.
c) It is the other element in Group VA, with smaller total number of electrons.
d) It is a halogen that exists in the liquid state at room temperature.
2.3. Develop four more element descriptions like those in problem 2.2. Exchange them
with a classmate and identify each other’s elements.
2.4. Without using the periodic table, determine the group, period, and block of an atom
with the following electron configurations.
a) [Ne]3s2;
b) [He]2s2;
c) [Kr]5s24d105p5.
2.5. Determine the group, period, and block in which each of the following elements is
located on the periodic table.
a) [Kr]5s24d1;
c) [He]2s22p6;
b) [Ar]4s23d104p3;
d) [Ne]3s23p1.
2.6. Identify the elements with the following valence electron configurations:
a) 5s1;
c) 3s2;
b) 4s23d2;
d) 4s24p3.
2.7. Write the electron configuration of the element fitting each of the following
descriptions.
a) The group 2A element in the fourth period.
b) The noble gas in the fifth period.
c) The group 2B element in the fourth period.
d) The group 6A element in the second period.
2.8. What are the symbols for the elements with the following valence electron
configurations?
a) s2d1;
b) s2p3;
c) s2p6.
2.9. Consider the following elements: H, Li, N, F, Co, Ag, Kr, I, Hg.
8
a) Sketch an outline of the periodic table, with these elements properly placed.
b) State the group number and period number each element belongs to.
c) Identify each element as a metal, metalloid, or non-metal.
d) Identify the state of each element at room temperature.
e) Draw the Lewis structure for each of these elements.
2.10. How many valence electrons are there in an atom of each of these elements?
a) strontium;
f) tin;
b) silicon;
g) chlorine;
c) bromine;
h) magnesium;
d) sulfur;
i) helium;
e) neon;
j) sodium.
Which of them have the same Lewis structures and which are different?
2.11. Draw Lewis structures for the elements: lithium, barium, boron, carbon, nitrogen.
2.12. Find the group and period of each of the following elements in the periodic table:
a) europium;
f) mercury;
b) neodymium;
g) ytterbium;
c) carbon;
h) bromine;
d) nitrogen;
i) chromium;
e) silicon;
j) krypton.
2.13. Which has the largest radius: helium (He), xenon (Xe), or argon (Ar)? Which has
the smallest?
2.14. Which has the largest radius: magnesium (Mg), silicon (Si), sulfur (S), or sodium
(Na)? The smallest?
2.15. Which has the largest atomic radius: nitrogen (N), antimony (Sb), or arsenic (As)?
The smallest?
2.16. Using only their location in the periodic table, rank the atoms in each set by
decreasing atomic size. Explain your answers.
a) Mg, Be, Ba;
f) Se, Br, Cl;
b) Ca, Se, Ga;
g) Mg, Ca, Li;
c) Br, Rb, Kr;
h) Sr, Te, Se;
d) Se, Br, Ca;
i) In, Br, I;
e) Ba, Sr, Cs;
j) S, Se, O.
2.17. Using only a periodic table, rank the atoms in each set in order of decreasing size.
Explain your ranking.
a) Na, K, H;
b) Mg, S, Si;
c) Cl, K, Ar.
9
2.18. Using only their location in a periodic table, rank each of the following sets of
elements in order of increasing atomic size. Explain your answer in each case.
a) Mg, S, Cl;
d) Rb, Xe, Te;
b) Al, B, In;
e) P, Na, F;
c) Ne, Ar, Xe;
f) O, S, N.
2.19. Can you determine which of two unknown elements has the larger radius if the
only known information is that the atomic number of one of the elements is 20 greater
than the other?
2.20. Using only their location in a periodic table, rank each of the following sets of
elements in order of decreasing ionization energy. Explain your answer in each case.
a) Cl, Br, I;
d) Na, Li, Cs;
b) Ga, Ge, Se;
e) S, Cl, Br;
c) K, Ca, Kr;
f) Cl, Ar, K.
2.21. Using only a periodic table, rank the elements in each set in order of increasing
ionization energy. Explain your ranking.
a) B, N, F;
b) F, Cl, Br;
c) Na, Cs, K.
2.22. Using only a periodic table, rank the elements in each set by increasing ionization
energy. Explain your answers.
a) Xe, He, Ar;
d) Kr, Br, K;
b) Sn, In, Sb;
e) K, Ca, Rb;
c) Sr, Ca, Ba;
f) Kr, Br, Rb.
2.23. Using only a periodic table, identify the atom in each of the following pairs with
the lower first ionization energy.
a) B, O;
d) F, N;
b) B, In;
e) Ca, K;
c) I, F;
f) B, Tl.
2.24. Using only a periodic table, rank the elements in each set in order of increasing
electron affinity. Explain your ranking.
a) Be, Ca, Mg;
b) Kr, Se, Br;
c) Na, Cs, K.
2.25. Which element in each of the following pairs will have the lower electron affinity?
Explain your answer in each case.
a) K or Ca;
c) S or Se;
b) O or Li;
d) Cs or F.
2.26. Based only on their position in the periodic table, arrange the elements in each set
in order of increasing attraction for electrons in a bond.
a) Li, Br, Zn, La, Si;
b) P, Ga, Cl, Y, Cs.
10
2.27. Which element in each pair is more electronegative?
a) K, As;
b) N, Sb;
c. Sr, Be.
2.28. For each of the following properties, indicate whether fluorine or bromine has a
larger value.
a) electronegativity;
c) atomic radius;
b) ionic radius;
d. ionization energy.
2.29. Give the correct responses in the following questions. Explain why according to
the location of the elements in the periodic table.
a) Which has the lower ionization energy? Li or K.
b) Which would be more polar? HF or HBr.
c) Which is more nonmetallic? F or I.
d) Which is more electronegative? K or Rb.
e) Which has more outer shell electrons? Ca or C.
f) Which would you expect to be more ionic? LiF or HCl.
2.30. Element A, with three electrons in its outer energy level, is in Period 4 of the
periodic table. How does the number of its valence electrons compare with that of
Element B, which is in Group IIIA and Period 6? Use Lewis structures to help you
express your answer.
2.31. Predict the composition and ionic or molecular character for binary compounds of
the following elements with hydrogen:
a) Li, Sn, Se;
b) K, C, Se;
c) Na, B, Sb.
2.32. Write formulas for the products of reactions of the following elements with the
excess of oxygen. Are they of acidic or basic character?
a) Na, Mg, Al, Si, P, S, Cl;
b) Rb, Sr, In, Sn, Sb, Te, I.
2.33. Describe an element, using only the periodic table. The following answers must be
given fully.
a) What is the name and the symbol of the element?
b) What is its location in the periodic table (group and period)?
c) How many protons and neutrons are in the nucleus of its atom?
d) How many electrons are in its atom?
e) What is the electron configuration of the element?
f) How many valence electrons does it have (according to Lewis structure)?
g) What element block does it belong to?
h) Is it a metal, semi-metal, nonmetal?
i) What is its compound with hydrogen, is it of ionic or molecular character?
j) What is the formula of its oxide (in the higher oxidation state), is it of acidic or
basic character?
11
3. ENERGY AND CHEMICAL CHANGE
3.1. An exothermic reaction releases 86.5 kJ. How many kilocalories of energy are
released?
3.2. If an endothermic process absorbs 256 J, how many kilocalories are absorbed?
3.3. What is the equivalent in joules of 126 calories?
3.4. Convert 455 kilojoules to kilocalories.
3.5. In the construction of bridges and skyscrapers, gaps must be left between adjoining
steel beams to allow for the expansion and contraction of the metal due to heating and
cooling. The temperature of a sample of iron with a mass of 10.0 g changed from
50.4°C to 25.0°C with the release of 114 J heat. What is the specific heat of iron?
3.6. If the temperature of 34.4 g of ethanol increases from 25.0°C to 78.8°C, how much
heat has been absorbed by the ethanol?
3.7. A nugget of pure gold with the mass of 4.50 g absorbed 276 J of heat. What was the
final temperature of the gold if the initial temperature was 25.0°C? The specific heat of
gold is in the table (problem 3.8.).
3.8. A 155 g sample of an unknown substance was heated from 25.0°C to 40.0°C. In the
process, the substance absorbed 5696 J of energy. What is the specific heat of the
substance? Identify the substance among those listed in the table:
Specific Heats of Common Substances at 298 K (25 oC)
Substance
Specific heat, J/g oC
Substance
Specific heat, J/g oC
Water (l)
4.184
Aluminum (s)
0.897
Water (s) = ice
2.03
Iron (s)
0.449
Water (g) = steam
2.01
Lead (s)
0.129
Ethanol (l)
2.44
Silver (s)
0.235
Granite (s)
0.803
Gold (s)
0.129
3.9. What is the specific heat of an unknown substance if a 2.50-g sample releases 12.0
cal as its temperature changes from 25.0°C to 20.0°C?
3.10. The temperature of 55.6 grams of a material decreases by 14.8°C when it loses
3080 J of heat. What is its specific heat?
3.11. What is the specific heat of a metal if the temperature of a 12.5 g sample increases
from 19.5°C to 33.6°C when it absorbs 37.7 J of heat?
12
3.12. A sample of ethylene glycol, used in car radiators, has a mass of 34.8 g. The
sample liberates 783 J of heat. The initial temperature of the sample is 22.1°C. What is
the final temperature?
3.13. A sample of ethanol, C2H5OH, absorbs 23.4 kJ of energy. The temperature of the
sample increases from 5.6°C to 19.8°C. What is the mass of the ethanol sample? The
specific heat capacity of ethanol is 2.46 J/(g ⋅ °C).
3.14. A child’s swimming pool contains 1000 L of water. When the water is warmed by
solar energy, its temperature increases from 15.3°C to 21.8°C. How much heat does the
water absorb?
3.15. What temperature change results from the loss of 255 kJ from a 10.0 kg sample of
water?
3.16. If 335 g water at 65.5°C loses 9750 J of heat, what is the final temperature of the
water?
3.17. The temperature of a sample of water increases from 20.0°C to 46.6°C as it
absorbs 5650 J of heat. What is the mass of the sample?
3.18. Explain how you could calculate the heat released in freezing 0.250 mol water.
3.19. How much heat is required to warm 122 g of water by 23.0°C?
3.20. Which of the following processes are exothermic? Endothermic?
a) C2H5OH(l) → C2H5OH(g).
b) NH3(g) → NH3(l).
c) Br2(l) → Br2(s).
d) NaCl(s) → NaCl(l).
e) C5H12(g) + 8 O2(g) → 5 CO2(g) + 6 H2O(l).
3.21. Write the correct sign of ΔH for each of the following changes in physical state.
a) C2H5OH(s) → C2H5OH(l).
b) H2O(g) → H2O(l).
c) CH3OH(l) → CH3OH(g).
d) NH3(l) → NH3(s).
3.22. A reaction is characterized by ΔH = –500 kJ/mol. Does the reaction mixture
absorb heat from the surroundings or release heat to them?
3.23. A reaction is characterized by ΔH = +280 kJ/mol. Does the mixture of reactants
and products release heat to the surroundings or absorb heat from them?
13
3.24. For each of the following reactions, determine: (a) does the enthalpy increase or
decrease; (b) is Hreactant > Hproduct or is Hproduct > Hreactant; (c) is ΔH positive or negative?
a) Al2O3(s) → 2Al(s) + 3/2 O2(g) (endothermic);
b) Sn(s) + Cl2(g) → SnCl2(s) (exothermic).
3.25. Consider the following reaction:
N2(g) + O2(g) → 2 NO(g); ΔHo = –180.6 kJ.
a) Draw an enthalpy diagram for the reaction.
b) What is the enthalpy change for the formation of one mole of nitrogen
monoxide?
c) What is the enthalpy change for the reaction of 1.00⋅102 g of nitrogen with
sufficient oxygen?
3.26. The reaction of iron with oxygen is very familiar. You can see the resulting rust on
buildings, vehicles, and bridges. You may be surprised, however, at the large amount of
heat that is produced by this reaction.
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + 1.65⋅103 kJ.
a) What is the enthalpy change for this reaction?
b) Draw an enthalpy diagram that corresponds to the thermochemical equation.
c) What is the enthalpy change for the formation of 23.6 g of iron(III) oxide?
3.27. Tetraphosphorus decoxide, P4O10, is an acidic oxide. It reacts with water to
produce phosphoric acid, H3PO4, in an exothermic reaction.
P4O10(s) + 6 H2O(l) → 4 H3PO4(aq); ΔHo = –257.2 kJ
a) How much energy is released when 5.00 mol of P4O10 reacts with excess
water?
c) How much energy is released when 235 g of H3PO4 is formed?
3.28. Calcium oxide, CaO, reacts with carbon in the form of graphite. Calcium carbide,
CaC2, and carbon monoxide, CO, are produced in an endothermic reaction.
CaO(s) + 3 C(s) + 462.3 kJ → CaC2(s) + CO(g)
a) 246.7 kJ of energy is available to react. What mass of calcium carbide is
produced, assuming sufficient reactants?
b) What is the enthalpy change for the reaction of 46.7 g of graphite with excess
calcium oxide?
c) 1.38⋅1024 molecules of calcium oxide react with excess graphite. How much
energy is needed?
3.29. Acetylene, C2H2, undergoes complete combustion in oxygen. Carbon dioxide and
water are formed. When one mole of acetylene reacts, 1.3⋅103 kJ of energy is released.
a) Draw a diagram to represent the thermochemical equation.
b) How much energy is released when the complete combustion of acetylene
produces 1.50 g of water?
14
3.30. Ethene, C2H4, reacts with water to form ethanol, CH3CH2OH:
C2H4(g) + H2O(l) → CH3CH2OH(l).
Determine the enthalpy change of this reaction, given the following thermochemical
equations.
1. CH3CH2OH(l) + 3 O2(g) → 3 H2O(l) + 2 CO2(g); ΔHo = –1367 kJ.
2. C2H4(g) + 3 O2(g) → 2 H2O(l) + 2 CO2(g); ΔHo = –1411 kJ.
3.31. A typical automobile engine uses a lead-acid battery. During discharge, the
following chemical reaction takes place:
Pb(s) + PbO2(s) + 2 H2SO4(l) → 2 PbSO4(aq) + 2 H2O(l).
Determine the enthalpy change of this reaction, given the following equations.
1. Pb(s) + PbO2(s) + 2 SO3(g) → 2 PbSO4(s); ΔHo = –775 kJ.
2. SO3(g) + H2O(l) → H2SO4(l); ΔHo = –133 kJ.
3.32. Mixing household cleansers can result in the production of hydrogen chloride gas,
HCl(g). Not only is this gas dangerous in its own right, but it also reacts with oxygen to
form chlorine gas and water vapour.
4 HCl(g) + O2(g) → 2 Cl2(g) + 2 H2O(g).
Determine the enthalpy change of this reaction, given the following equations.
1. H2(g) + Cl2(g) → 2 HCl(g); ΔHo = –185 kJ.
2. H2(g) + 1/2 O2(g) → H2O(l); ΔHo = –285.8 kJ.
3. H2O(g) → H2O(l); ΔHo = –40.7 kJ.
3.33. Calculate the enthalpy change of the following reaction between nitrogen gas and
oxygen gas, given thermochemical equations (1), (2), and (3).
2 N2(g) + 5 O2(g) → 2 N2O5(g).
1. 2 H2(g) + O2(g) → 2 H2O(l); ΔHo = –572 kJ.
2. N2O5(g) + H2O(l) → 2 HNO3(l); ΔHo = –77 kJ.
3. 1/2 N2(g) + 3/2 O2(g) + 1/2 H2(g) → HNO3(l); ΔHo = –174 kJ.
3.34. Calculate ΔHo for the reaction 2 C(s) + 2 H2(g) → C2H4(g) given the following
thermochemical equations:
1. 2 CO2(g) + 2 H2O(l) → C2O4(g) + 3 O2(g); ΔHo = 1411 kJ.
2. C(s) + O2(g) → CO2(g); ΔHo = –393.5 kJ.
3. 2 H2(g) + O2(g) → 2 H2O(l); ΔHo = –572 kJ.
3.35. Calculate ΔHo for the reaction HCl(g) + NH3(g) → NH4Cl(s) given the following
thermochemical equations:
1. H2(g) + Cl2(g) → 2 HCl(g); ΔHo = –184 kJ.
2. N2(g) + 3 H2(g) → 2 NH3(g); ΔHo = –92 kJ.
3. N2(g) + 4 H2(g) + Cl2(g) → 2 NH4Cl(s); ΔHo = –628 kJ.
15
3.36. Calculate the enthalpy change of the following reaction, given equations (1), (2),
and (3):
2 H3BO3(aq) → B2O3(s) + 3 H2O(l).
1. H3BO3(aq) → HBO2(aq) + H2O(l); ΔHo = –0.02 kJ.
2. H2B4O7(s) + H2O(l) → 4 HBO2(aq); ΔHo = –11.3 kJ.
3. H2B4O7(s) → 2 B2O3(s) + H2O(l); ΔHo = 17.5 kJ.
3.37. Hydrogen can be added to ethene, C2H4, to obtain ethane, C2H6.
C2H4(g) + H2(g) → C2H6(g).
Show that the equations for the formation of ethene and ethane from their elements can
be algebraically combined to obtain the equation for the addition of hydrogen to ethene.
3.38. Zinc sulfide reacts with oxygen gas to produce zinc oxide and sulfur dioxide.
2 ZnS(s) + 3 O2(g) → 2 ZnO(s) + 2 SO2(g).
Write the chemical equation for the formation of the indicated number of moles of each
compound from its elements. Algebraically combine these equations to obtain the given
equation.
3.39. The standard molar enthalpy of formation of calcium carbonate is –1207.6 kJ/mol.
Calculate the enthalpy of formation of calcium oxide, given the following equation:
CaO(s) + CO2(g) → CaCO3; ΔHo = –178.1 kJ.
3.40. Small amounts of oxygen gas can be produced in a laboratory by heating
potassium chlorate, KClO3.
2 KClO3(s) → 2 KCl(s) + 3 O2(g).
Calculate the enthalpy change of this reaction, using enthalpies of formation, namely:
ΔHof (KClO3) = –397.7 kJ/mol; ΔHof (KCl) = –436.5 kJ/mol.
3.41. Use the following equation to answer the questions below.
CH3OH(l) + 1.5 O2(g) → CO2(g) + 2 H2O(g).
a) Calculate the enthalpy change of the complete combustion of one mole of
methanol, using enthalpies of formation: ΔHof (CH3OH) = –239.2 kJ/mol; ΔHof (CO2) =
–393.5 kJ/mol; ΔHof (H2O(g)) = –241.8 kJ/mol.
b) How much energy is released when 125 g of methanol undergoes complete
combustion?
3.42. In the early 1960s, Neil Bartlett, at the University of British Columbia, was the
first person to synthesize compounds of the noble gas xenon. A number of noble gas
compounds (such as XeF2, XeF4, XeF6, and XeO3) have since been synthesized.
Consider the reaction of xenon difluoride with fluorine gas to produce xenon
tetrafluoride: XeF2(g) + F2(g) → XeF4(s). Use the following standard molar
enthalpies of formation to calculate the enthalpy change for this reaction: ΔHof (XeF2) =
– 108 kJ/mol; ΔHof (XeF4) = –251 kJ/mol.
16
3.43. Hydrogen is a very appealing fuel, in part because burning it produces only nonpolluting water. One of the challenges that researchers face in making hydrogen fuel a
reality is how to produce hydrogen economically. Researchers are investigating
methods of producing hydrogen indirectly. The following series of equations represent
one such method.
3 FeCl2(s) + 4 H2O(g) → Fe3O4(s) + 6 HCl(g) + H2(g); ΔHo = 318 kJ;
Fe3O4(s) + 3/2 Cl2(g) + 6 HCl(g) → 3 FeCl3(s) + 3 H2O(g) + 1/2 O2(g); ΔHo = –249 kJ;
3 FeCl3(s) → 3 FeCl2(s) + 3/2 Cl2(g); ΔHo = 173 kJ.
a) Show that the net result of the three reactions is the decomposition of water to
produce hydrogen and oxygen.
b) Use Hess’s law and the enthalpy changes for the reactions to determine the
enthalpy change for the decomposition of one mole of water. Check your answer, using
the enthalpy of formation of water, ΔHof (H2O(g)) = –241.8 kJ/mol.
3.44. Predict the sign of entropy change ΔS for the following reaction. Explain the basis
for your prediction: 2 H2(g) + O2(g) → 2 H2O (g).
3.45. Predict the sign of entropy change ΔS for each reaction or process.
a) FeS(s) → Fe2+(aq) + S2–(aq)
b) SO2(g) + H2O(l) → H2SO3(aq)
3.46. If ΔHo = 285.4 kJ/mol and ΔSo = 137.55 J/mol K, calculate Gibbs free energy
change ΔGo at 25 oC for the supposed process. Is the reaction spontaneous? Is either or
both of the driving forces (ΔHo and ΔSo) for the reaction favorable?
3.47. Calculate ΔGo at 25°C for the reaction 2 NO2(g) → N2O4(g) given functions ΔHo =
–57.20 kJ/mol and ΔSo = –175.83 J/mol K. Is this reaction spontaneous? What is the
driving force for spontaneity?
3.48. The standard Gibbs free energy of formation has the values: –286.06 kJ/mol for
NaI(s), –261.90 kJ/mol for Na+(aq), and –51.57 kJ/mol for I–(aq) at 25°C. Calculate ΔGo
for the reaction in water: NaI(s) ⇄ Na+(aq) + I–(aq).
3.49. Calculate ΔGo at 700 K for the reduction of the oxides of iron and copper by
carbon, represented by the equations:
1. 2 Fe2O3(s) + 3 C(graphite) → 4 Fe(s) + 3 CO2(g);
2. 2 CuO(s) + C(graphite) → 2 Cu(s) + CO2(g).
Values of ΔGof at 700 K are –92 kJ/mol for CuO(s), –632 kJ/mol for Fe2O3(s),
and –395 kJ/mol for CO2(g). Which oxide can be reduced using carbon in a wood fire
(which has a temperature of about 700 K), assuming standard state conditions?
17
4. RATES OF CHEMICAL REACTIONS
4.1. The following reaction is second order in A and first order in B: 2 A + B → 3 C.
What is the rate law equation for the reaction below? (Assume that A, B, and C are the
same compounds for each reaction.)
4 A + 2 B → 6 C.
4.2. Consider the general reaction below:
a A + b B → c C + d D.
Based on this equation, is it correct to write the following rate law equation by
inspection? Explain your answer.
Rate = k [A]a [B]b.
4.3. Consider the following rate law equation.
Rate = k [A]2 [B].
a) How does the reaction rate change if [A] decreases by a factor of 2 and [B]
increases by a factor of 4?
b) How does the reaction rate change if [A] and [B] are doubled?
4.4. Consider the following rate law equation:
Rate = k [HCrO4–] [HSO3–]2 [H+].
a) What is the order with respect to each reactant?
b) What is the overall reaction order?
c) What are the units for the rate constant?
4.5. Cyclopropane, C3H6, is used in the synthesis of organic compounds and as a fastacting anesthetic. It undergoes rearrangement to form propene C3H6. If cyclopropane
disappears at a rate of 0.25 mol/s, at what rate is propene being produced?
4.6. Ammonia, NH3, reacts with oxygen to produce nitric oxide, NO, and water vapour.
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g).
At a specific time in the reaction, ammonia is disappearing at rate of 0.068 mol/(L ⋅ s).
What is the corresponding rate of production of water?
4.7. Hydrogen bromide reacts with oxygen to produce bromine and water vapour.
4 HBr(g) + O2(g) → 2 Br2(g) + 2 H2O(g).
How does the rate of decomposition of HBr (in mol/(L ⋅ s) ) compare with the rate of
formation of Br2 (also in mol/(L ⋅ s) )? Express your answer as an equation.
4.8. Magnesium metal reacts with hydrochloric acid to produce magnesium chloride and
hydrogen gas:
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g).
Over an interval of 1.00 s, the mass of Mg(s) changes by −0.011 g.
18
a) What is the corresponding rate of consumption of HCl(aq) (in mol/s)?
b) Calculate the corresponding rate of production of H2(g) (in L/s) at 20°C and
101 kPa.
4.9. In the following reaction, the rate of production of sulfate ions is calculated to be
1.25⋅10–3 mol/(L ⋅ s).
2 HCrO4– + 3 HSO3– + 5 H+ → 2 Cr3+ + 3 SO42– + 5 H2O.
a) What is the corresponding rate at which [HSO3–] decreases over the same time
interval?
b) What is the corresponding rate at which [HCrO4–] decreases over the same
time interval?
4.10. Use the data in the following table to calculate the average reaction rates.
Experimental Data for H2 + Cl2 → 2 HCl
Time (s)
[H2] (M)
[Cl2] (M)
[HCl] (M)
0.00
0.030
0.050
0.000
4.00
0.020
0.040
0.020
a) Calculate the average reaction rate expressed in moles H2 consumed per liter
per second.
b) Calculate the average reaction rate expressed in moles Cl2 consumed per liter
per second.
c) Calculate the average reaction rate expressed in moles HCl produced per liter
per second.
4.11. When heated, ethylene oxide decomposes to produce methane and carbon
monoxide: C2H4O(g) → CH4(g) + CO(g). At 415°C, the following initial rate data
were recorded. Determine the rate law equation and the rate constant at 415 oC.
Experiment
1
2
3
[C2H4O] (mol/L)
0.00285
0.00428
0.00570
Initial rate (mol/(L ⋅ s)
5.84 ⋅ 10-7
8.76 ⋅ 10-7
1.17 ⋅ 10-6
4.12. Iodine chloride reacts with hydrogen to produce iodine and hydrogen chloride:
2 ICl + H2 → I2 + 2 HCl. At temperature T, the following initial rate data were
recorded. Determine the rate law equation and the rate constant at temperature T.
Experiment
1
2
3
[ICl]0 (mol/L)
0.20
0.40
0.20
[H2]0 (mol/L)
0.050
0.050
0.200
19
Initial rate (mol/(L ⋅ s)
0.0015
0.0030
0.0060
4.13. Sulfuryl chloride (also known as chlorosulfuric acid and thionyl chloride), SO2Cl2,
is used in a variety of applications. At a certain temperature, the rate of decomposition
of sulfuryl chloride was studied: SO2Cl2(g) → SO2(g) + Cl2(g).
Experiment
1
2
3
[SO2Cl2] (mol/L)
0.150
0.300
0.450
Initial rate (mol/(L ⋅ s)
3.3 ⋅ 10–6
6.6 ⋅ 10–6
9.9 ⋅ 10–6
a) Write the rate law equation for the decomposition of sulfuryl chloride.
b) Determine the rate constant, k, for the reaction, with the appropriate units.
4.14. A first-order decomposition reaction has a rate constant of 2.34 ⋅ 10–2 year–1. What
is the half-life of the reaction? Express your answer in years and in seconds.
4.15. When cyclopropane, C3H6, undergoes rearrangement to propene at 1000°C, the
first-order rate constant for the decomposition of cyclopropane is 9.2 s−1.
a) Determine the half-life of the reaction.
b) What percent of the original concentration of cyclopropane will remain after 4
half-lives?
4.16. The following reaction is exothermic.
2 ClO(g) → Cl2(g) + O2(g).
Draw and label a potential energy diagram for the reaction. Propose a reasonable
activated complex.
4.17. Consider the following reaction.
AB + C → AC + B; ΔH = 65 kJ; Ea(rev) = 34 kJ.
a) Draw and label a potential energy diagram for this reaction.
b) Calculate and label Ea(fwd). Include a possible structure for the activated
complex.
4.18. Consider the reaction: C + D → CD; ΔH = –132 kJ; Ea(fwd) = 61 kJ.
a) Draw and label a potential energy diagram for this reaction.
b) Calculate and label Ea(rev). Include a possible structure for the activated
complex.
4.19. In the upper atmosphere, oxygen exists in forms other than O2(g). For example, it
exists as ozone, O3(g), and as single oxygen atoms, O(g). Ozone and atomic oxygen
react to form two molecules of oxygen. For this reaction, the enthalpy change is −392 kJ
and the activation energy is 19 kJ. Draw and label a potential energy diagram. Include a
value for Ea(rev). Propose a structure for the activated complex.
20
5. DYNAMIC CHEMICAL EQUILIBRIUM
5.1. Write the equilibrium expression for each homogeneous reaction.
a) The reaction between nitrogen gas and oxygen gas at high temperatures:
N2(g) + O2(g) ⇄ 2 NO(g).
b) The reaction between hydrogen gas and oxygen gas to form water vapour:
2 H2(g) + O2(g) ⇄ 2 H2O(g).
c) The reduction-oxidation equilibrium of iron and iodine ions in aqueous
solution:
2 Fe3+(aq) + 2 I–(aq) ⇄ 2 Fe2+(aq) + I2(aq).
d) The oxidation of ammonia:
4 NH3(g) + 5 O2(g) ⇄ 4 NO(g) + 6 H2O(g).
5.2. Write equilibrium constant expressions for these equilibria.
a) N2O4(g) ⇄ 2 NO2(g);
b) CO(g) + 3 H2(g) ⇄ CH4(g) + H2O(g);
c) 2 H2S(g) ⇄ 2 H2(g) + S2(g).
5.3. Write equilibrium constant expressions for these heterogeneous equilibria.
a) C10H8(s) ⇄ C10H8(g);
b) CaCO3(s) ⇄ CaO(s) + CO2(g);
c) H2O(l) ⇄ H2O(g);
d) C(s) + H2O(g) ⇄ CO(g) + H2(g);
e) FeO(s) + CO(g) ⇄ Fe(s) + CO2(g).
5.4. Three reactions, and their equilibrium constants, are given below.
1. N2(g) + O2(g) ⇄ 2 NO(g); Keq = 4.7⋅10–31.
2. 2 NO(g) + O2(g) ⇄ 2 NO2(g); Keq = 1.8⋅10–6.
3. N2O4(g) ⇄ 2 NO2(g); Keq = 0.025.
Arrange these reactions in the order of their tendency to form products.
5.5. Identify each reaction as essentially going to completion or not taking place.
a) N2(g) + 3 Cl2(g) ⇄ 2 NCl3(g); Keq = 3.0⋅1011.
b) 2 CH4(g) ⇄ C2H6(g) + H2(g); Keq = 9.5⋅10–13.
c) 2 NO(g) + 2 CO(g) ⇄ N2(g) + 2 CO2(g); Keq = 2.2⋅1059.
21
5.6. Most metal ions combine with other ions in solution. For example, in aqueous
ammonia, silver(I) ions are at equilibrium with different complex ions.
[Ag(H2O)2]+(aq) + 2 NH3(aq) ⇄ [Ag(NH3)2]+(aq) + 2 H2O(l).
At room temperature, Keq for this reaction is 1⋅107. Which of the two silver complex
ions is more stable? Explain your reasoning.
5.7. Consider the following reaction: H2(g) + Cl2(g) ⇄ 2 HCl(g). Its equilibrium
constant Keq = 2.4⋅1033 at 25 oC. HCl(g) is placed in a reaction vessel. To what extent do
you expect the equilibrium mixture to dissociate into H2(g) and Cl2(g)?
5.8. The following reaction took place in a sealed flask at 250°C:
PCl5(g) ⇄ PCl3(g) + Cl2(g).
At equilibrium, the gases in the flask had the following concentrations:
[PCl5] = 1.2⋅10–2 mol/L, [PCl3] = 1.5⋅10–2 mol/L, and [Cl2] = 1.5⋅10–2 mol/L. Calculate
the value of Keq at 250°C.
5.9. Iodine and bromine react to form iodine monobromide, IBr.
I2(g) + Br2(g) ⇄ 2 IBr(g).
At 250°C, an equilibrium mixture in a 2.0 L flask contained 0.024 mol of I2(g),
0.050 mol of Br2(g), and 0.38 mol of IBr(g). What is the value of Keq for the reaction?
5.10. At high temperatures, carbon dioxide gas decomposes into carbon monoxide and
oxygen gas. At equilibrium, the gases have the following concentrations: [CO2(g)] =
= 1.2 mol/L, [CO(g)] = 0.35 mol/L, and [O2(g)] = 0.15 mol/L. Determine Keq at the
temperature of the reaction.
5.11. Hydrogen sulfide is a pungent, poisonous gas. At 1400 K, an equilibrium mixture
was found to contain 0.013 mol/L hydrogen, 0.046 mol/L sulfur in the form of S2(g),
and 0.18 mol/L hydrogen sulfide. Calculate the equilibrium constant, at 1400 K, for the
following reaction: 2 H2S(g) ⇄ 2 H2(g) + S2(g).
5.12. Methane, ethyne, and hydrogen form the following equilibrium mixture:
2 CH4(g) ⇄ C2H2(g) + 3 H2(g).
While studying this reaction mixture, a chemist analyzed a 4.0 L sealed flask at
1700°C. The chemist found 0.46 mol of CH4(g), 0.64 mol of C2H2(g), and 0.92 mol of
H2(g). What is the value of Keq for the reaction at 1700°C?
5.13. Sulfur atoms combine to form molecules that have different numbers of atoms
depending on the temperature. At about 1050°C, the following dissociation occurs:
S8(g) ⇄ 4 S2(g).
22
The initial concentration of S8(g) in a flask is 9.2⋅10−3 mol/L, and the equilibrium
concentration of the same gas is 2.3⋅10−3 mol/L. What is the value of Keq?
5.14. Consider an equilibrium in which oxygen gas reacts with gaseous hydrogen
chloride to form gaseous water and chlorine gas. At equilibrium, the gases have the
following concentrations: [O2] = 8.6⋅10−2 mol/L, [HCl] = 2.7⋅10−2 mol/L, [H2O] =
= 7.8⋅10−3 mol/L, [Cl2] = 3.6⋅10−3 mol/L.
a) Write a balanced chemical equation for this reaction.
b) Calculate the value of the equilibrium constant.
5.15. Calculate Keq for the equilibrium:
N2O4(g) ⇄ 2 NO2(g),
using the data [N2O4] = 0.0185 mol/L and [NO2] = 0.0627 mol/L.
5.16. Calculate Keq for the equilibrium:
CO(g) + 3 H2(g) ⇄ CH4(g) + H2O(g),
using the data [CO] = 0.0613 mol/L, [H2] = 0.1839 mol/L, [CH4] = 0.0387 mol/L, and
[H2O] = 0.0387 mol/L.
5.17 Determine the value of equilibrium constant at 400 K for the decomposition of
phosphorus pentachloride, if [PCl5] = 0.135 mol/L, [PCl3] = 0.550 mol/L, and [Cl2] =
= 0.550 mol/L. The equation for the reaction is: PCl5(g) ⇄ PCl3(g) + Cl2(g).
5.18. At 25 °C, the value of Keq for the following reaction is 82.
I2(g) + Cl2(g) ⇄ 2 ICl(g).
0.83 mol of I2(g) and 0.83 mol of Cl2(g) are placed in a 10 L container at 25 °C. What
are the concentrations of the three gases at equilibrium?
5.19. At a certain temperature, Keq = 4.0 for the following reaction:
2 HF(g) ⇄ H2(g) + F2(g).
A 1.0 L reaction vessel contained 0.045 mol of F2(g) at equilibrium. What was the
initial amount of HF in the reaction vessel?
5.20. A chemist was studying the following reaction:
SO2(g) + NO2(g) ⇄ NO(g) + SO3(g).
In a 1.0 L container, the chemist added 0.17 mol of SO2 to 0.11 mol of NO2. The value
of Keq for for the reaction at a certain temperature is 4.8. What is the equilibrium
concentration of SO3 at this temperature?
23
5.21. Phosgene, COCl2, is an extremely toxic gas. It is prepared by mixing carbon
monoxide and chlorine gas:
CO(g) + Cl2(g) ⇄ COCl2(g).
0.055 mol of CO and 0.072 mol of Cl2 are placed in a 5.0 L container. At 870 K, the
equilibrium constant is 0.20. What are the equilibrium concentrations of the mixture?
5.22. Hydrogen bromide decomposes at 700 K.
2 HBr(g) ⇄ H2(g) + Br2(g); Keq = 4.2⋅10–9.
0.090 mol of HBr is placed in a 2.0 L reaction vessel and heated to 700 K. What is the
equilibrium concentration of each gas?
5.23. The following equation represents the equilibrium reaction for the dissociation of
phosgene gas: COCl2(g) ⇄ CO(g) + Cl2(g). At 100 °C, the value of Keq for this
reaction is 2.2⋅10−8. The initial concentration of COCl2 in a closed container at 100 °C is
1.5 mol/L. What are the equilibrium concentrations of CO and Cl2?
5.24. Hydrogen sulfide dissociates into hydrogen and sulfur in gaseous state at 1400 °C,
with Keq equal to 2.4⋅10−4. Suppose, 4.0 mol of H2S is placed in a 3.0 L container. What
is the equilibrium concentration of H2 at 1400°C?
5.25. At a certain temperature, the value of Keq for the following reaction is 3.3⋅10–12.
2 NCl3(g) ⇄ N2(g) + 3 Cl2(g).
A certain amount of nitrogen trichloride, NCl3, is put in a 1.0 L reaction vessel at this
temperature. At equilibrium, 4.6⋅10−4 mol of N2 is present. What amount of NCl3 was
put in the reaction vessel?
5.26. At a certain temperature, the value of Keq for the following reaction is 4.2⋅10–8.
N2(g) + O2(g) ⇄ 2 NO(g).
0.45 mol of N2 and 0.26 mol of O2 are put in a 6.0 L reaction vessel. What is the
equilibrium concentration of NO at this temperature?
5.27. At a particular temperature, Keq for the decomposition of carbon dioxide gas
equals 2.0⋅10–6.
2 CO2(g) ⇄ 2 CO(g) + O2(g).
3.0 mol of CO2 is put in a 5.0 L container. Calculate the equilibrium concentration of
each gas.
5.28. 0.50 mol of CO and 0.50 mol of H2O are placed in a 10 L container at 700 K. The
following reaction occurs: CO(g) + H2O(g) ⇄ H2(g) + CO2(g); Keq = 8.3.
What is the concentration of each gas that is present at equilibrium?
24
5.29. At a certain temperature, Keq = 10.5 for the equilibrium
CO(g) + 2 H2(g) ⇄ CH3OH(g).
Calculate these concentrations:
a) [CO] in an equilibrium mixture with 0.933 mol/L H2 and 1.32 mol/L CH3OH;
b) [H2] in an equilibrium mixture with 1.09 mol/L CO and 0.325 mol/L CH3OH;
c) [CH3OH] in an equilibrium mixture with 0.0661 mol/L H2 and 3.85 mol/L CO.
5.30. Equilibrium constant is 1.60 at 933 K for this reaction:
H2(g) + CO2(g) ⇄ H2O(g) + CO(g).
Calculate the equilibrium concentration of hydrogen when [CO2] = 0.320 mol/L,
[H2O] = 0.240 mol/L, and [CO] = 0.280 mol/L.
5.31. At 2273 K, Keq = 6.2⋅10–4 for the reaction:
N2(g) + O2(g) ⇄ 2 NO(g).
If [N2] = 0.05200 mol/L and [O2] = 0.00120 mol/L, what is the concentration of
NO at equilibrium?
5.32. How would decreasing the volume of the reaction vessel affect these equilibria?
a) 2 SO2(g) + O2(g) ⇄ 2 SO3(g).
b) H2(g) + Cl2(g) ⇄ 2 HCl(g).
c) 2 NOBr(g) ⇄ 2 NO(g) + Br2(g).
5.33. Use Le Chatelier’s principle to predict how each of these changes would affect the
ammonia equilibrium system:
N2(g) + 3 H2(g) ⇄ 2 NH3(g);
a) removing hydrogen from the system;
b) adding ammonia to the system;
c) adding hydrogen to the system.
5.34. In the following equilibrium, would you raise or lower the temperature to obtain
these results?
C2H2(g) + H2O(g) ⇄ CH3CHO(g); ΔH° = –151 kJ;
a) an increase in the amount of CH3CHO;
b) a decrease in the amount of C2H2;
c) an increase in the amount of H2O.
5.35. Predict how this equilibrium would respond to a simultaneous increase in both
temperature and pressure.
CO(g) + Cl2(g) ⇄ COCl2(g); ΔH° = –220 kJ.
25
5.36. Use Le Chatelier’s principle to predict how each of the following changes would
affect this equilibrium:
H2(g) + CO2(g) ⇄ H2O(g) + CO(g);
a) adding H2O(g) to the system;
b) removing CO(g) from the system;
c) adding H2(g) to the system;
d) adding something to the system to absorb CO2(g).
5.37. How would increasing the volume of the reaction vessel affect these equilibria?
a) NH4Cl(s) ⇄ NH3(g) + HCl(g).
b) N2(g) + O2(g) ⇄ 2 NO(g).
5.38. How would decreasing the volume of the reaction vessel affect these equilibria?
a) 2 N2H4(g) + 2 NO2(g) ⇄ 3 N2(g) + 4 H2O(g).
b) 2 H2O(g) ⇄ 2 H2(g) + O2(g).
5.39. How would these equilibria be affected by increasing the temperature?
a) 4 NH3(g) + 5 O2(g) ⇄ 4 NO(g) + 6 H2O(g) + heat.
b) heat + NaCl(s) ⇄ Na+(aq) + Cl–(aq).
5.40. Ethylene (C2H4) reacts with hydrogen to form ethane (C2H6):
C2H4(g) + H2(g) ⇄ C2H6(g) + heat.
How would you regulate the temperature of this equilibrium to do the following?
a) increase the yield of ethane;
b) decrease the concentration of ethylene;
c) increase the amount of hydrogen in the system.
5.41. How would simultaneously decreasing the temperature and volume of the system
affect these equilibria?
a) heat + CaCO3(s) ⇄ CaO(s) + CO2(g).
b) 4 NH3(g) + 5 O2(g) ⇄ 4 NO(g) + 6 H2O(g) + heat.
5.42. Consider the following reaction: H2(g) + I2(g) + 52 kJ ⇄ 2 HI(g). In which
direction does the equilibrium shift if there is an increase in temperature?
5.43. Why does changing the volume of the reaction vessel have no effect on this
equilibrium?
CO(g) + Fe3O4(s) ⇄ CO2(g) + 3 FeO(s).
26
5.44. A decrease in the pressure of each system below is caused by increasing the
volume of the reaction container. In which direction does the equilibrium shift?
a) CO2(g) + H2(g) ⇄ CO(g) + H2O(g).
b) 2 NO2(g) ⇄ N2O4(g).
c) 2 CO2(g) ⇄ 2 CO(g) + O2(g).
d) CH4(g) + 2 H2S(g) ⇄ CS2(g) + 4 H2(g).
5.45. For each reversible reaction, determine whether the forward reaction is favored by
high temperatures or low temperatures.
a) N2O4(g) ⇄ 2 NO2(g); ΔH = +59 kJ.
b) 2 ICl(g) ⇄ I2(g) + Cl2(g); ΔH = −35 kJ.
c) 2 HF(g) ⇄ H2(g) + F2(g); ΔH = −536 kJ.
5.46. The following reaction is exothermic: 2 NO(g) + 2 H2(g) ⇄ N2(g) + 2 H2O(g). In
which direction does the equilibrium shift as a result of each change?
a) removing the hydrogen gas;
b) increasing the pressure of gases in the reaction vessel by decreasing the
volume;
c) increasing the pressure of gases in the reaction vessel by pumping in argon gas
while keeping the volume of the vessel constant;
d) increasing the temperature;
e) using a catalyst.
5.47. In which direction does the equilibrium shift as a result of the change to each
homogeneous equilibrium system?
a) Adding Cl2(g): 2 Cl2(g) + O2(g) ⇄ 2 Cl2O(g).
b) Removing N2(g): 2 NO2(g) ⇄ N2(g) + 2 O2(g).
c) Using a catalyst: CH4(g) + H2O(g) ⇄ CO2(g) + H2(g).
d) Decreasing the total volume of the reaction container: 2 NO2(g) ⇄ N2O4(g).
e) Increasing the temperature:
CO(g) + 3 H2(g) ⇄ CH4(g) + H2O(g); ΔH = −230 kJ.
27
6. IONIC EQUILIBRIA
6.1. Name and write the formula of the conjugate base of each molecule or ion.
a) HCl;
b) HCO3−;
c) H2SO4;
d) N2H5+.
6.2. Name and write the formula of the conjugate acid of each molecule or ion.
a) NO3–;
b) OH–;
c) H2O;
d) HCO3–.
6.3. Identify the conjugate acid-base pairs in each reaction.
a) HS–(aq) + H2O(l) ⇄ H2S(aq) + OH–;
b) O2–(aq) + H2O(l) → 2 OH–.
6.4. Identify the conjugate acid-base pairs in each reaction.
a) H2S(aq) + NH3(aq) ⇄ NH4+(aq) + HS–(aq);
b) H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4–(aq).
6.5. Calculate the concentration of hydronium ions in each solution.
a) 4.5 mol/L HCl(aq);'
b) 30.0 mL of 4.50 mol/L HBr(aq) diluted to 100.0 mL;
c) 18.6 mL of 2.60 mol/L HClO4(aq) added to 24.8 mL of 1.92 mol/L NaOH(aq).
6. 6. Calculate the concentration of hydroxide ions in each solution.
a) 3.1 mol/L KOH(aq);
b) 21.0 mL of 3.1 mol/L KOH diluted to 75.0 mL;
c) 23.2 mL of 1.58 mol/L HCl(aq) added to 18.9 mL of 3.50 mol/L NaOH(aq).
6.7. Determine whether reacting each pair of solutions results in an acidic solution or a
basic solution. Then calculate the concentration of the ion that causes the solution to be
acidic or basic. (Assume that the volumes in part (a) are additive. Assume that the
volumes in part (b) stay the same.)
a) 31.9 mL of 2.75 mol/L HCl(aq) added to 125 mL of 0.0500 mol/L
Mg(OH)2(aq);
b) 4.87 g of NaOH(s) added to 80.0 mL of 3.50 mol/L HBr(aq).
6.8. 2.75 g of MgO(s) is added to 70.0 mL of 2.40 mol/L HNO3(aq). Is the solution that
results from the reaction acidic or basic? What is the concentration of the ion that is
responsible for the character of the solution?
6.9. Phosphoric acid, H3PO4(aq) is triprotic. It has three hydrogen ions that may be
dissociated.
a) Write an equation to show the dissociation of each proton.
b) Show that H2PO4– can act as either an acid or a base.
c) Which is the stronger acid, H3PO4(aq) or H2PO4−(aq)? Explain your answer.
28
6.10. Para-aminobenzoic acid (PABA) is a weak monoprotic acid that is used in some
sunscreen lotions. Its formula is C6H4NH2COOH. What is the formula of the conjugate
base of PABA?
6.11. Boric acid, B(OH)3(aq), is used as a mild antiseptic in eye-wash solutions. The
following reaction takes place in aqueous solution:
B(OH)3(aq) + 2 H2O(l) ⇄ B(OH)4–(aq) + H3O+(aq).
a) Identify the conjugate acid-base pairs.
b) Is boric acid strong or weak? How do you know?
6.12. Classify each compound as a strong acid, weak acid, strong base, or weak base.
a) butyric acid, CH3CH2CH2COOH (responsible for the odour of rancid butter);
b) hydroiodic acid, HI(aq) (added to some cough syrups);
c) potassium hydroxide, KOH (used in the manufacture of soft soaps);
d) red iron oxide, Fe2O3 (used as a colouring pigment in paints).
6.13. Determine [H3O+] and [OH−] in each solution:
a) 0.45 mol/L hydrochloric acid;
b) 1.1 mol/L sodium hydroxide.
6.14. Determine [H3O+] and [OH−] in each solution.
a) 0.95 mol/L hydrobromic acid;
b) 0.012 mol/L calcium hydroxide.
6.15. [OH−] is 5.6⋅10–14 mol/L in a solution of hydrochloric acid. What is the molar
concentration of the HCl(aq)?
6.16. [H3O+] is 1.7⋅10–14 mol/L in a solution of calcium hydroxide. What is the molar
concentration of the Ca(OH)2(aq)?
6.17. [H3O+] of a sample of milk is found to be 3.98⋅10–7 mol/L. Is the milk acidic,
neutral, or basic? Calculate the pH and [OH−] of the sample.
6.18. A sample of household ammonia has a pH of 11.9. What is the pOH and [OH−] of
the sample?
6.19. Phenol, C6H5OH, is used as a disinfectant. An aqueous solution of phenol was
found to have a pH of 4.72. Is phenol acidic, neutral, or basic? Calculate [H3O+], [OH−],
and pOH of the solution.
29
6.20. At normal body temperature, 37 °C, the value of Kw for water equals 2.5⋅10–14.
Calculate [H3O+] and [OH−] at this temperature. Is pure water at 37 °C acidic, neutral, or
basic?
6.21. A sample of baking soda was dissolved in water and the pOH of the solution was
found to be 5.81 at 25 °C. Is the solution acidic, basic, or neutral? Calculate the pH,
[H3O+], and [OH−] of the solution.
6.22. A chemist dissolved some Aspirin™ in water. The chemist then measured the pH
of the solution and found it to be 2.73 at 25 °C. What are [H3O+] and [OH−] of the
solution?
6.23. Calculate the pH of a sample of vinegar that contains 0.83 mol/L acetic acid. What
is the percent dissociation of the vinegar?
6.24. In low doses, barbiturates act as sedatives. The formula of barbituric acid is
C4H4N2O3. A chemist prepares a 0.10 mol/L solution of barbituric acid and finds the pH
of the solution to be 2.50. What is the acid dissociation constant for barbituric acid?
What percent of its molecules dissociate?
6.25. A solution of hydrofluoric acid has a molar concentration of 0.0100 mol/L. What
is the pH of this solution?
6.26. Hypochlorous acid, HOCl, is used as a bleach and a germ-killer. A chemist finds
that 0.027% of hypochlorous acid molecules are dissociated in a 0.40 mol/L solution of
the acid. What is the value of Ka for the acid?
6.27. The word “butter” comes from the Greek butyros. Butanoic acid (common name:
butyric acid) gives rancid butter its distinctive odour. Calculate the pH of a 1.0⋅10–2 M
solution of butanoic acid (Ka = 1.51⋅10–5).
6.28. Caproic acid, C5H11COOH, occurs naturally in coconut and palm oil. It is a weak
monoprotic acid, with Ka = 1.3⋅10–5. A certain aqueous solution of caproic acid has a pH
of 2.94. How much acid was dissolved to make 100 mL of this solution?
6.29. Carbonated beverages contain a solution of carbonic acid. Carbonic acid is also
important for forming the ions that are present in blood.
CO2(aq) + H2O(l) ⇄ H2CO3(aq);
H2CO3(aq) + H2O(l) ⇄ HCO3–(aq) + H3O+(aq);
HCO3–(aq) + H2O(l) ⇄ CO32–(aq) + H3O+(aq).
Calculate the pH of a solution of 5.0⋅10–4 mol/L carbonic acid (use Appendix 4).
What is [CO32–] in the solution?
30
6.30. Hydrosulfuric acid, H2S(aq), is a weak diprotic acid that is sometimes used in
analytical work. Calculate the pH and [HS−(aq)] of a 7.5⋅10–3 mol/L solution.
6.31. A 0.10 mol/L solution of a weak monoprotic acid was found to be 5.0%
dissociated. Calculate Ka.
6.32. Oxalic acid, HOOCCOOH, is a weak diprotic acid that occurs naturally in some
foods, including rhubarb. Calculate the pH of a solution of oxalic acid that is prepared
by dissolving 2.5 g in 1.0 L of water. What is the concentration of hydrogen oxalate,
HOOCCOO−, in the solution?
6.33. A sample of blood was taken from a patient and sent to a laboratory for testing.
Chemists found that the blood pH was 7.40. They also found that the hydrogen
carbonate ion concentration was 2.6⋅10–2 mol/L. What was the concentration of carbonic
acid in the blood?
6.34. An aqueous solution of household ammonia has a molar concentration of 0.105 M.
Calculate the pH of the solution.
6.35. Hydrazine, N2H4, has been used as a rocket fuel. The concentration of an aqueous
solution of hydrazine is 5.9⋅10–2 mol/L. Calculate the pH of the solution.
6.36. Morphine, C17H19NO3, is a naturally occurring base that is used to control pain. A
4.5⋅10–3 mol/L solution has a pH of 9.93. Calculate Kb for morphine.
6.37. Methylamine, CH3NH2, is used to manufacture several prescription drugs.
Calculate [OH−] and pOH of a 0.25 mol/L aqueous solution of methylamine.
6.38. At room temperature, trimethylamine, (CH3)3N, is a gas with a strong ammonialike odour. Calculate [OH−] and the percent of trimethylamine molecules that react with
water in a 0.22 mol/L aqueous solution.
6.39. An aqueous solution of ammonia has a pH of 10.85. What is the concentration of
the solution?
6.40. Use the table of Ka values in Appendix 4 to list the conjugate bases of the
following acids in order of increasing base strength: formic acid, HCOOH; hydrofluoric
acid, HF(aq); benzoic acid, C6H5COOH; phenol, C6H5OH.
6.41. Compare Kb for ammonia, NH3, and for trimethylamine, (CH3)3N. Which is the
stronger acid, NH4+ or (CH3)3NH+?
6.42. A buffer solution is made by mixing 250 mL of 0.200 mol/L aqueous ammonia
and 400 mL of 0.150 mol/L ammonium chloride. Calculate the pH of solution.
31
6.43. A buffer solution is made by mixing 200 mL of 0.200 mol/L aqueous ammonia
and 450 mL of 0.150 mol/L ammonium chloride. Calculate the pH of solution.
6.44. A buffer solution contains 0.200 mol/L nitrous acid, HNO2(aq), and 0.140 mol/L
potassium nitrite, KNO2(aq). What is the pH of the buffer solution?
6.45. A buffer solution is prepared by dissolving 1.80 g of benzoic acid, C6H5COOH,
and 1.95 g of sodium benzoate, NaC6H5COO, in 800 mL of water. Calculate the pH of
the buffer solution.
6.46. Predict whether an aqueous solution of each salt is neutral, acidic, or basic.
a) NaCN;
c) Mg(NO3)2;
b) LiF;
d) NH4I.
6.47. Is the solution of each salt acidic, basic, or neutral? For solutions that are not
neutral, write equations that support your predictions.
a) NH4BrO4;
c) NaOBr;
b) NaBrO4;
d) NH4Br.
6.48. Compare Ka for benzoic acid, C6H5COOH, and for phenol, C6H5OH. Which is the
stronger base, C6H5COO−(aq) or C6H5O−(aq)? Explain your answer.
6.49. Sodium hydrogen sulfite, NaHSO3, is a preservative that is used to prevent the
discolouration of dried fruit. In aqueous solution, the hydrogen sulfite ion can act as
either an acid or a base. Predict whether NaHSO3 dissolves to form an acidic solution or
a basic solution. (Refer to Appendix 4 for ionization data.)
6.50. Sodium carbonate and sodium hydrogen carbonate both dissolve to form basic
solutions. Comparing solutions with the same concentration, which of these salts forms
the more basic solution? Explain.
6.51. Potassium phosphate and potassium dihydrogen phosphate both dissolve to form
basic solutions. Comparing solutions with the same concentration, which of these salts
forms the more basic solution? Explain.
6.52. Determine whether or not each ion reacts with water. If the ion does react, write
the chemical equation for the reaction. Then predict whether the ion forms an acidic
solution or a basic solution.
a) Br–;
c) CH3NH3+;
b) ClO4–;
d) OCl–.
6.53. A chemist measures the pH of aqueous solutions of Ca(OH)2, CaF2, NH4NO3,
KNO3, and HNO3. Each solution has the same concentration. Arrange the solutions
from most basic to most acidic.
32
6.54. Write the balanced chemical equation that represents the dissociation of each
compound in water. Then write the corresponding solubility product expression.
a) copper(I) chloride;
d) calcium phosphate.
b) barium fluoride;
e) silver carbonate;
c) silver sulfate;
f) ammonium magnesium phosphate.
6.55. The maximum solubility of silver cyanide, AgCN, is 1.5⋅10–8 mol/L at 25°C.
Calculate Ksp for silver cyanide.
6.56. A saturated solution of copper(II) phosphate, Cu3(PO4)2 , has a concentration of
6.1⋅10–7 g Cu3(PO4)2 per 100 mL of solution at 25 °C. What is Ksp for Cu3(PO4)2?
6.57. A saturated solution of CaF2 contains 1.2⋅1020 molecules of calcium fluoride per
liter of solution. Calculate Ksp for CaF2.
6.58. Ksp for silver chloride, AgCl, is 1.8⋅10–10 at 25 oC.
a) Calculate the molar solubility of AgCl in a saturated solution at 25 °C.
b) How many molecules of AgCl are dissolved in 1.0 L of saturated silver
chloride solution?
c) What is the percent (m/v) of AgCl in a saturated solution at 25 °C?
6.59. Calculate the molar solubility of Fe(OH)3 at 25 °C (ref. Appendix 6).
6.60. Calculate the solubility (in mol/L and in g/L) of Zn(IO3)2 in a saturated solution.
6.61. Determine the molar solubility of AgCl:
a) in pure water; b) in 0.15 mol/L NaCl.
6.62. Determine the molar solubility of lead(II) iodide, PbI2, in 0.050 mol/L NaI.
6.63. Calculate the molar solubility of calcium sulfate, CaSO4:
a) in pure water; b) in 0.25 mol/L Na2SO4.
6.64. Calculate the molar solubility of lead(II) chloride, PbCl2:
a) in pure water; b) in 0.10 mol/L CaCl2.
6.65. The maximum solubility of barium fluoride, BaF2, at 25 °C, is 1.3 g/L.
a) Calculate Ksp for BaF2 at 25 °C.
b) Calculate the solubility of BaF2 in molecules of BaF2/L.
6.66. A solution of BaCl2 is added to a solution of Na2SO4.
a) Calculate the solubility (in mol/L and in g/L) of BaSO4 in pure water.
b) Calculate the solubility (in mol/L and in g/L) of BaSO4 in 0.085 M Na2SO4.
33
6.67. A solution contains 0.15 mol/L of NaCl and 0.0034 mol/L Pb(NO3)2. Does a
precipitate form? Include a balanced chemical equation for the formation of the possible
precipitate.
6.68. One drop (0.050 mL) of 1.5 mol/L potassium chromate, K2CrO4, is added to 250
mL of 0.10 mol/L AgNO3. Does a precipitate form?
6.69. A chemist adds 0.010 g of CaCl2 to 500 mL of 0.0015 mol/L sodium carbonate,
Na2CO3. Does a precipitate of calcium carbonate form?
6.70. 0.10 mg of magnesium chloride, MgCl2, is added to 250 mL of 0.0010 mol/L
NaOH. Does a precipitate of magnesium hydroxide form?
6.71. 100 mL of 1.0⋅10–3 mol/L Pb(NO3)2 is added to 40 mL of 0.040 mol/L NaCl.
Does a precipitate form?
6.72. 25 mL of 0.10 mol/L NaOH is added to 500 mL of 0.00010 mol/L cobalt(II)
chloride, CoCl2. Does a precipitate form?
6.73. 250 mL of 0.0011 mol/L Al2(SO4)3 is added to 50 mL of 0.022 mol/L BaCl2. Does
a precipitate form? Include a balanced chemical equation for the formation of the
possible precipitate.
6.74. A chemist adds 1.0 mg of NaI to 50 mL of a 0.010 mol/L solution of Pb(NO3)2.
Does a precipitate form?
6.75. How many milligrams of Na2SO4 will just begin to precipitate calcium sulfate,
CaSO4, from 500 mL of a 0.10 mol/L solution of CaCl2?
6.76. How many drops of 0.0010 mol/L silver nitrate solution will just begin to
precipitate AgCl from 5000 mL of a 0.90% (m/v) solution of NaCl? (Assume that one
drop equals 0.050 mL.)
6.77. Compare the values of solubility products constants for two salts with the same
anions: Ksp for CaSO4 is 2.4⋅10–5, and Ksp for SrSO4 is 3.2⋅10–7. Suppose that you have a
1.0 L solution that is 0.20 mol/L in Ca2+ ions, and 0.20 mol/L in Sr2+ ions. You slowly
begin to add solid Na2SO4.
a) Explain why SrSO4 precipitates first.
b) How many milligrams of Al2(SO4)3 will just begin to precipitate SrSO4 from
the solution?
34
7. ELECTROCHEMISTRY
7.1. Write a net ionic equation for a reaction in which:
a) Fe2+ acts as an oxidizing agent;
b) Al acts as a reducing agent;
c) Au3+ acts as an oxidizing agent;
d) Cu acts as a reducing agent;
e) Sn2+ acts as an oxidizing agent and as a reducing agent.
7.2. Write the oxidation half-reaction, the reduction half-reaction, and the overall cell
reaction for each of the following galvanic cells. Identify the anode and the cathode in
each case. (In part (b), platinum is present as an inert electrode.)
a) Sn(s) | Sn2+(aq) || Tl+(aq) | Tl(s);
b) Cd(s) | Cd2+(aq) || H+(aq) | H2(g) | Pt(s).
7.3. A galvanic cell involves the overall reaction of iodide ions with acidified
permanganate ions to form manganese(II) ions and iodine. The salt bridge contains
potassium nitrate.
a) Write the half-reactions, and the overall cell reaction.
b) Identify the oxidizing agent and the reducing agent.
c) The inert anode and cathode are both made of graphite. Solid iodine forms on
one of them. Which one?
7.4. Look up the standard reduction potentials of the following half-reactions (ref.
Appendix 7). Predict whether acidified nitrate ions will oxidize manganese(II) ions to
manganese(IV) oxide under standard conditions.
MnO2(s) + 4 H+(aq) + 2e− → Mn2+(aq) + 2 H2O(l);
NO3–(aq) + 4 H+(aq) + 3e− → NO(g) + 2 H2O(l).
7.5. Predict whether each reaction is spontaneous or non-spontaneous under standard
conditions.
a) 2 Cr(s) + 3 Cl2(g) → 2 Cr3+(aq) + 6 Cl–(aq).
b) Zn2+(aq) + Fe(s) → Zn(s) + Fe2+(aq).
c) 5 Ag(s) + MnO4–(aq) + 8 H+(aq) → 5 Ag+(aq) + Mn2+(aq) + 4 H2O(l).
7.6. Predict whether each reaction is spontaneous or non-spontaneous under standard
conditions in an acidic solution.
a) H2O2(aq) → H2(g) + O2(g).
b) 3 H2(g) + Cr2O72–(aq) + 8 H+(aq) → 2 Cr3+(aq) + 7 H2O(l).
7.7. Determine the standard cell potential for each of the following redox reactions.
a) CuSO4(aq) + Ni(s) → NiSO4(aq) + Cu(s).
b) Fe(s) + 4 HNO3(aq) → Fe(NO3)3(aq) + NO(g) + 2 H2O(l).
35
7.8. Determine if each of the following balanced redox reactions is spontaneous as
written, calculate the cell potential.
a) Sn(s) + 2 Cu+(aq) → Sn2+(aq) + 2 Cu(s).
b) Mg(s) + Pb2+(aq) → Pb(s) + Mg2+(aq).
c) 2Mn2+(aq) + 8H2O(l) + 10Hg2+(aq) → 2MnO4–(aq) + 16H+(aq) + 5Hg22+(aq).
7.9. Write the two half-reactions for the following redox reactions. Subtract the two
reduction potentials to find the standard cell potential for a galvanic cell in which this
reaction occurs.
a) Cl2(g) + 2 Br–(aq) → 2 Cl–(aq) + Br2(l).
b) 2 Cu+(aq) + 2 H+(aq) + O2(g) → 2 Cu2+(aq) + H2O2(aq).
7.10. Determine the standard cell potential for each of the following redox reactions.
a) 3 Mg(s) + 2 Al3+(aq) → 3 Mg2+(aq) + 2 Al(s).
b) 2 K(s) + F2(g) → 2 K+(aq) + 2 F–(aq).
c) Cr2O72–(aq) + 14 H+(aq) + 6 Ag(s) → 2 Cr3+(aq) + 6 Ag+(aq) + 7 H2O(l).
7.11. The cell potential for the following galvanic cell is given: Eocell = 1.750 V.
Zn | Zn2+ (1 mol/L) || Pd2+ (1 mol/L) | Pd.
Determine the standard reduction potential for the following half-reaction:
Pd2+(aq) + 2e− → Pd(s).
7.12. These equations represent overall cell reactions. Determine the standard potential
for each cell and identify the reactions as spontaneous or nonspontaneous as written.
a) 2 Al3+(aq) + 3 Cu(s) → 2 Cu2+(aq) + 2 Al(s).
b) Hg2+(aq) + 2 Cu+(aq) → 3 Cu2+(aq) + Hg(l).
c) Cd(s) + 2 NO3–(aq) + 4 H+(aq) → 2 Cd2+(aq) + 2 NO2(g) + 2 H2O(l).
7.13. Write the standard cell notation for the following cells in which the half-cell listed
is connected to the standard hydrogen electrode. An example is Na | Na+ || H+ | H2.
Determine the voltage of the cells formed.
a) Zn | Zn2+;
d) Cu | Cu2+;
b) Hg | Hg2+;
e) Al | Al3+.
c) Ga | Ga3+;
f) NO2 | NO3–.
7.14. Calculate the cell potential of voltaic cells that contain the following pairs of halfcells.
a) Chromium in a solution of Cr3+ ions; copper in a solution of Cu2+ ions.
b) Zinc in a solution of Zn2+ ions; platinum in a solution of Pt2+ ions.
c) A half-cell containing both HgCl2 and Hg2Cl2; lead in a solution of Pb2+ ions.
d) Tin in a solution of Sn2+ ions; iodine in a solution of I– ions.
7.15. Predict the products of the electrolysis of a 1 mol/L solution of sodium chloride.
36
7.16. The electrolysis of molten calcium chloride produces calcium and chlorine. Write
a) the half-reaction that occurs at the anode;
b) the half-reaction that occurs at the cathode;
c) the chemical equation for the overall cell reaction.
7.17. For the electrolysis of molten lithium bromide, write
a) the half-reaction that occurs at the negative electrode;
b) the half-reaction that occurs at the positive electrode;
c) the net ionic equation for the overall cell reaction.
7.18. Explain why calcium can be produced by the electrolysis of molten calcium
chloride, but not by the electrolysis of aqueous calcium chloride.
7.19. One half-cell of a galvanic cell has a nickel electrode in a 1 mol/L nickel(II)
chloride solution. The other half-cell has a cadmium electrode in a 1 mol/L cadmium
chloride solution.
a) Find the cell potential.
b) Identify the anode and the cathode.
c) Write the oxidation half-reaction, the reduction half-reaction, and the overall
cell reaction.
7.20. An external voltage is applied to change the galvanic cell in question 7.21 into an
electrolytic cell. Repeat parts (a) to (c) for the electrolytic cell.
7.21. Calculate the mass of zinc plated onto the cathode of an electrolytic cell by a
current of 750 mA in 3.25 h.
7.22. How many minutes does it take to plate 0.925 g of silver onto the cathode of an
electrolytic cell using a current of 1.55 A?
7.23. The nickel anode in an electrolytic cell decreases in mass by 1.20 g in 35.5 min.
The oxidation half-reaction converts nickel atoms to nickel(II) ions. What is the
constant current?
7.24. The following two half-reactions take place in an electrolytic cell with an iron
anode and a chromium cathode.
Oxidation: Fe(s) → Fe2+(aq) + 2e−.
Reduction: Cr3+(aq) + 3e− → Cr(s).
During the process, the mass of the iron anode decreases by 1.75 g.
a) Find the change in mass of the chromium cathode.
b) Explain why you do not need to know the electric current or the time to
complete part (a).
37
PERIODIC TABLE
I
II
III
IV
V
H
1
2
3
4
5
6
7
Hydrogen
1
1.0079
Li
Lithium
3
6.94
Na
Sodium
11
22.99
K
Potassium
19
39.098
Cu
Copper
29
63.54
Rb
Rubidium
37
85.47
Ag
Silver
47
107.87
Cs
Cesium
55
132.905
Au
Gold
79
196.97
Fr
Francium
87
[223]
Lanthanides
Gd
64
Gadolinium
157.2
Actinides
Cm
96
Curium
[247]
B
Be
Beryllium
4
9.012
Mg
Magnesium
12
24.305
Ca
Calcium
20
40.08
Zn
Zinc
30
65.38
Sr
Strontium
38
87.62
Cd
Cadmium
48
112.41
Ba
Barium
56
137.33
Hg
Mercury
80
200.5
Ra
Radium
88
226.03
La
57
Tb
65
Ac
89
Bk
97
C
N
Nitrogen
Carbon
14.0067
6
12.011 7
P
Si
Phosphorus
Silicon
30.974
14
28.085 15
Ti
V
Titanium
Vanadium
22
47.90 23
50.94
As
Ge
Arsenic
Germanium
74.92
32
72.59 33
Zr
Nb
Zirconium
Niobium
40
91.22 41
92.906
Sn
Sb
Tin
Antimony
50
118.69 51
121.75
Hf
Ta
Hafnium
Tantalum
La–Lu
72
178.49 73
180.95
Bi
Pb
Tl
Bismuth
Lead
Thallium
208.98
207.2 83
81
204.3 82
Rf
Db
Rutherfordium
Dubnium
Ac–(Lr)
104
[261] 105
[262]
Boron
5
10.81
Al
Aluminum
13
26.98
Sc
Scandium
21
44.956
Ga
Gallium
31
69.72
Y
Yttrium
39
88.91
In
Indium
49
114.82
Lanthanum
138.905
Terbium
158.93
Actinium
[227]
Berkelium
[247]
Ce
58
Dy
66
Th
90
Cf
98
38
Cerium
140.12
Dysprosium
162.50
Thorium
232.038
Californium
[251]
Pr Praseodymium
59
140.9077
Ho
Holmium
67
164.93
Pa
Protactinium
91
231.036
Es
Einsteinium
99
[254]
OF CHEMICAL ELEMENTS
VI
Appendix 1
VII
VIII
(H)
O
F
Fluorine
Oxygen
18.998
8
15.999 9
Cl
S
Chlorine
Sulfur
35.453
16
32.06 17
Cr
Mn
Chromium
Manganese
24
51.996 25
54.938
Br
Se
Bromine
Selenium
79.904
34
78.96 35
Mo
Tc
Molybdenum
Technetium
42
95.94 43
98.906
Te
I
Tellurium
Iodine
52
127.6 53
126.90
W
Re
Tungsten
Rhenium
74
183.8 75
186.21
At
Po
Astatine
Polonium
[210]
84
[209] 85
Sg
Bh
Seaborgium
Bohrium
106
[266] 107
[264]
Nd
60
Er
68
U
92
Fm
100
He
Fe
26
Iron
55.847
Ru
Ruthenium
44
101.07
Os
Osmium
76
190.2
Hs
Hassium
108
[277]
Neodymium Pm
Promethium
144.24 61
[145]
Erbium Tm
Thulium
167.26 69
168.93
Uranium Np
Neptunium
238.029 93
237.048
Fermium Md Mendelevium
[257] 101
[258]
39
Sm
62
Yb
70
Pu
94
No
102
Helium
2
4.0026
Ne
Neon
10
20.179
Ar
Argon
18
39.948
Co
Ni
Cobalt
Nickel
27
58.933 28
58.70
Kr
Krypton
36
83.80
Rh
Pd
Rhodium
Palladium
45
102.9 46
106.4
Xe
Xenon
54
131.3
Ir
Pt
Iridium
Platinum
77
199.2 78
195.1
Rn
Radon
86
[222]
Mt
Meitnerium
109
[268]
Samarium
150.4
Ytterbium
173.04
Plutonium
[244]
Nobelium
[255]
Eu
63
Lu
71
Am
95
Lr
103
Europium
151.96
Lutetium
174.967
Americium
[243]
Lawrencium
[256]
Appendix 2
Elements and Electronegative Components
Symbol
Ac
Al
Ag
Am
Ar
As
At
Au
B
Ba
Be
Bh
Bi
Bk
Br
C
Ca
Cd
Ce
Cf
Cl
Cm
Co
Cr
Cs
Cu
Db
Dy
Er
Es
Eu
F
Fe
Fm
Fr
Name
Transcription
actinium
aluminum
silver
americium
argon
arsenic
astatine
gold
boron
barium
beryllium
bohrium
bismuth
berkelium
bromine
carbon
calcium
cadmium
cerium
californium
chlorine
curium
cobalt
chromium
cesium
copper
dubnium
dysprosium
erbium
einsteinium
europium
fluorine
iron
fermium
francium
αk_'tin_i: _əm
ə_'lu:m_ə_nəm
'sil_vər
αm_ə_'ris_i: _əm
'a:r_gən
'a:rs_ən_ik
'αs_tə_ti:n
gould
'bo:_rən
'bαr_i:_əm
bə_'ril_i: _əm
'bo:r_i: _əm
'biz_məθ
'bə:r_kli: _əm
'brou_mi:n
'ka:r_bən
'kαl_si:_əm
'kαd_mi:_əm
'sir_i:_əm
kalə_'fo:r_ni:_əm
'klo:r_i:n
'kju:r_i:_əm
'kou_bo:lt
'krou_mi:_əm
'si:_zi:_əm
'kop_ər
'du:b_ni:_əm
dis_'prou_zi:_əm
'ə:r_bi:_əm
aın_'staın_i:_əm
yu:_'rou_pi:_əm
'flu:r_i:n
'aı_ərn
'fer_mi:_əm
'frαn_si:_əm
40
Electronegative
component
Transcription
arsenide
'a:rs_ən_aıd
boride
'bo:r_aıd
beryllide
bə_'ril_aid
bromide
carbide
'brou_maıd
'ka:r_baıd
chloride
'klo:r_aıd
fluoride
'flu:r_aıd
Appendix 2 (continued)
Symbol
Ga
Gd
Ge
H
He
Hf
Hg
Ho
Hs
I
In
Ir
K
Kr
La
Li
Lr
Lu
Md
Mg
Mn
Mo
Mt
N
Na
Nb
Nd
Ne
Ni
No
Np
O
Os
P
Pa
Pb
Pd
Name
gallium
gadolinium
germanium
hydrogen
helium
hafnium
mercury
holmium
hassium
iodine
indium
iridium
potassium
krypton
lanthanum
lithium
lawrencium
lutetium
mendelevium
magnesium
manganese
molybdenum
meitnerium
nitrogen
sodium
niobium
neodymium
neon
nickel
nobelium
neptunium
oxygen
osmium
phosphorus
protactinium
lead
palladium
Transcription
'gαl_i:_əm
gαd_əl_'in_i:_əm
jə:r_'meın_i:_əm
'haı_drə_jən
'hi:_li:_əm
'hαf_ni:_əm
'mə:r_kyə_ri:
'houl_mi:_əm
'ha:_si:_əm
'aı_ə_daın
'in_di:_əm
i_'rid_i:_əm
pə_'tαs_i:_əm
'krip_tən
'lαn_θə_nəm
'liθ_i:_əm
'lou_'ren_si:_əm
lu:_ti: _shəm
'men_də_li:_vi:_əm
mαg_'ni:_zi:_əm
'mαŋ_gə_ni:s
mə_'lib_de_nəm
maıt_'nir_i:_əm
'naı_trə_jən
'soud_i:_əm
naı_'ou_bi:_əm
ni:_ou_'dim_i:_əm
'ni:_on
'nik_əl
nou_'bel_i:_əm
nep_'tu:_ni:_əm
'ok_sə_jən
'oz_mi:_əm
'fos_fə_rəs
prout_αk_'tin_i:_əm
led
pə_'leıd_i:_əm
41
Electronegative
component
Transcription
germanide
hydride
'jə:r_mə_naıd
'haı_draıd
iodide
'aı_ə_daıd
nitride
'naı_traıd
oxide
'ok_saıd
phosphide
'fo_sfaıd
plumbide
'pləm_baıd
Appendix 2 (end)
Symbol
Name
Transcription
Po
Pm
Pr
Pt
Pu
Ra
Rb
Re
Rf
Rh
Rn
Ru
S
Sb
Sc
Se
Sg
Si
Sm
Sn
Sr
Ta
Tb
Tc
Te
Th
Ti
Tl
Tm
U
V
W
Xe
Y
Yb
Zn
Zr
polonium
promethium
praseodymium
platinum
plutonium
radium
rubidium
rhenium
rutherfordium
rhodium
radon
ruthenium
sulfur
antimony
scandium
selenium
seaborgium
silicon
samarium
tin
strontium
tantalum
terbium
technetium
tellurium
thorium
titanium
thallium
thulium
uranium
vanadium
tungsten
xenon
yttrium
ytterbium
zinc
zirconium
pə_'lou_ni:_əm
prə_'mi:_thi:_əm
preı_zi:_ou_'dim_i:_əm
'plαt_ən_əm
plu:_'tou_ni:_əm
'reı_d_i:_əm
ru:_ 'bid_i:_əm
'ri:_ni:_əm
rəð_ər_'fo:r_di:_əm
'roud_i:_əm
'reı_dən
ru:_ 'thi:_ni:_əm
'səl_fər
'αn_tə_'mou_ni:
'skαn_di:_əm
sə_'li:_ni:_əm
si:_ 'bo:rg_i:_əm
'sil_ə_kən
sə_'mαr_i:_əm
tin
'stron_ti:_əm
'tαnt_əl_əm
'tə:r_bi:_əm
tek_'ni:_shi:_əm
tə_'lu_ri:_əm
'tho:r_i:_əm
taı_'teı_ni:_əm
'θαl_i:_əm
'θu:_li:_əm
yə_'reı_ni:_əm
və_'neıd_i:_əm
'təŋ_stən
'zi:_non
'i_tri:_əm
i_'tə:r_bi:_əm
ziŋk
zə:r_'kou_ni:_əm
42
Electronegative
component
Transcription
sulfide
'səl_faıd
selenide
'sel_ə_naıd
silicide
'sil_ə_saıd
telluride
tə_'lu_raıd
Appendix 3
Acids and Anions
Formula
Acid
HCl
HClO
HClO2
HClO3
HClO4
hydrochloric
hypochlorous
chlorous
chloric
perchloric
Transcription
haı_drə_'klo:r_ik
haı_pə_'klo:r_əs
'klo:r_əs
'klo:r_ik
pə:r_'klo:r_ik
Anion
chloride
hypochlorite
chlorite
chlorate
perchlorate
Transcription
'klo:r_aıd
haı_pə_'klo:r_aıt
'klo:r_aıt
'klo:r_eıt
pə:r_'klo:r_eıt
(similar with other halogens)
HCN
HMnO4
HNO2
HNO3
HOCN
HSCN
CH3COOH
H2C2O4
H2CO3
H2Cr2O7
H2CrO4
H2S
H2SiO3
H2S2O3
H2SO3
H2SO4
H3AsO3
H3AsO4
H3BO3
H3PO3
H3PO4
hydrocyanic
permanganic
nitrous
nitric
cyanic
thiocyanic
acetic
oxalic
carbonic
dichromic
chromic
hydrosulfuric
silicic
thiosulfuric
sulfurous
sulfuric
arsenious
arsenic
boric
phosphorous
phosphoric
haı_drou_saı_'an_ik
pə:r_mαn_'gαn_ik
'naı_trəs
'naı_trik
saı_'αn_ik
θaı_ou_saı_'αn_ik
ə_'si:t_ik
ok_'sαl_ik
ka:r_'bon_ik
daı_'krou_mik
'krou_mik
haı_drə_səl_'fyur_ik
sə_'lis_ik
θaı_ə_səl_'fyur_ik
'səl_fə_rəs
'səl'fyur_ik
a:r_'si:n_i:_əs
a:r_'sen_ik
'bo:r_ik
'fos_fə_rəs
fos_'fo:r_ik
43
cyanide
permanganate
nitrite
nitrate
cyanate
thiocyanate
acetate
oxalate
carbonate
dichromate
chromate
sulfide
silicate
thiosulfate
sulfite
sulfate
arsenite
arsenate
borate
phosphite
phosphate
'saı_ə_naıd
pə:r_'mαŋ_gə_neıt
'naı_traıt
'naı_treıt
'saı_ə_neıt
θaı_ou_'saı_ə_neıt
'αs_ə_teıt
'ok_sə_leıt
'ka:r_bə_nət
daı_'krou_meıt
'krou_meıt
'səl_faıd
'sil_ə_kət
θaı_ə_'səl_feıt
'səl_faıt
'səl_feıt
'a:r_sə_naıt
'a:rs_ən_eıt
'bo: _reıt
'fos_faıt
'fos_'feıt
Appendix 4
Ionization Constants for Acids
Acid
Formula
Conjugate base
Ka
Monoprotic Acids
Acetic acid
Benzoic acid
Chlorous acid
Cyanic acid
Formic acid
Hydrobromic acid
Hydrochloric acid
Hydrocyanic acid
Hydrofluoric acid
Hydrogen oxide
Hypobromous acid
Phenol
CH3COO–
C6H5COO–
ClO2–
OCN–
COOH–
Br–
Cl–
CN–
F–
OH–
BrO–
C6H5O–
CH3COOH
C6H5COOH
HClO2
HOCN
HCOOH
HBr
HCl
HCN
HF
H2O
HOBr
C6H5OH
1.8 ⋅ 10–5
6.3 ⋅ 10–5
1.1 ⋅ 10–2
3.5 ⋅ 10–4
1.8 ⋅ 10–4
1.0 ⋅ 102
1.3 ⋅ 106
6.2 ⋅ 10–10
6.3 ⋅ 10–4
1.0 ⋅ 10–14
2.8 ⋅ 10–9
1.3 ⋅ 10–10
Polyprotic Acids
Boric acid
Carbonic acid
Citric acid
Oxalic acid
Phosphoric acid
Hydrosulfuric acid
Sulfuric acid
Sulfurous acid
Tartaric acid
H2BO3–
HBO32–
HCO3–
CO32–
H2C6H5O7–
HC6H5O72–
C6H5O73–
HC2O4–
C2O42–
H2PO4–
HPO42–
PO43–
HS–
S2–
HSO4–
SO42–
HSO3–
SO32–
HC4H4O6–
C4H4O62–
H3BO3
H2BO3–
H2CO3
HCO3–
H3C6H5O7
H2C6H5O7–
HC6H5O72–
H2C2O4
HC2O4–
H3PO4
H2PO4–
HPO42–
H2S
HS–
H2SO4
HSO4–
H2SO3
HSO3–
H2C4H4O6
HC4H4O6–
44
5.4 ⋅ 10–10
< 1.0 ⋅ 10–14
4.5 ⋅ 10–7
4.7 ⋅ 10–11
7.4 ⋅ 10–4
1.7 ⋅ 10–5
4.0 ⋅ 10–7
5.6 ⋅ 10–2
1.5 ⋅ 10–4
6.9 ⋅ 10–3
6.2 ⋅ 10–8
4.8 ⋅ 10–13
8.9 ⋅ 10–8
1.0 ⋅ 10–10
1.0 ⋅ 103
1.0 ⋅ 10–2
1.4 ⋅ 10–2
6.3 ⋅ 10–8
9.3 ⋅ 10–4
4.3 ⋅ 10–5
Appendix 5
Ionization Constants for Nitrogen Bases
Base
Formula
Conjugate acid
Kb
1.2-Diaminoethane
Dimethylamine
Ethanamine
Methanamine
Trimethylamine
Ammonia
Hydrazine
Hydroxylamine
Pyridine
Aniline
Urea
NH2CH2CH2NH2
(CH3)2NH
C2H5NH2
CH3NH2
(CH3)3N
NH3
N2H4
NH2OH
C5H5N
C6H5NH2
NH2CONH2
NH2CH2CH2NH3+
(CH3)2NH2+
C2H5NH3+
CH3NH3+
(CH3)3NH+
NH4+
N2H5+
NH3OH+
C5H5NH+
C6H5NH3+
NH2CONH3+
8.4 ⋅ 10–5
5.4 ⋅ 10–4
4.5 ⋅ 10–4
4.6 ⋅ 10–4
6.4 ⋅ 10–5
1.8 ⋅ 10–5
1.3 ⋅ 10–6
8.8 ⋅ 10–9
1.7 ⋅ 10–9
7.5 ⋅ 10–10
1.3 ⋅ 10–14
o
Appendix 6
Solubility Product Constants in Water at 25 C
Compound
Ksp
Compound
Ksp
Compound
Ksp
Ag2CO3
Ag2CrO4
AgBr
AgBrO3
AgCl
AgCN
AgI
AlPO4
Ba(IO3)2
BaCO3
BaCrO4
BaF2
BaSO4
Be(OH)2
Ca(IO3)2
Ca(OH)2
Ca3(PO4)2
8.46 ⋅ 10–12
1.12 ⋅ 10–12
5.35 ⋅ 10–13
5.38 ⋅ 10–5
1.77 ⋅ 10–10
5.97 ⋅ 10–17
8.52 ⋅ 10–17
9.84 ⋅ 10–21
4.01 ⋅ 10–9
2.58 ⋅ 10–9
1.12 ⋅ 10–10
1.84 ⋅ 10–7
1.08 ⋅ 10–10
6.92 ⋅ 10–22
6.47 ⋅ 10–6
5.02 ⋅ 10–6
2.07 ⋅ 10–33
CaCO3
CaF2
CaSO4
Cd(OH)2
CdF2
Co(OH)2
Co3(PO4)2
Cu3(PO4)2
CuBr
CuCl
CuCN
CuI
CuSCN
Eu(OH)3
Fe(OH)2
Fe(OH)3
FeF2
3.36 ⋅ 10–9
3.45 ⋅ 10–11
4.93 ⋅ 10–5
7.20 ⋅ 10–15
6.44 ⋅ 10–3
5.92 ⋅ 10–15
2.05 ⋅ 10–35
1.40 ⋅ 10–37
6.27 ⋅ 10–9
1.72 ⋅ 10–9
3.47 ⋅ 10–20
1.27 ⋅ 10–12
1.08 ⋅ 10–13
9.38 ⋅ 10–27
4.87 ⋅ 10–17
2.79 ⋅ 10–39
2.36 ⋅ 10–6
Hg2SO4
Mg(OH)2
MgCO3
Ni(OH)2
Ni3(PO4)2
Pb(OH)2
Pb(SCN)2
PbBr2
PbCl2
PbCO3
PbCrO4
PbI2
Sn(OH)2
Sr(IO3)2
TlBrO3
Y(IO3)3
Zn(OH)2
6.50 ⋅ 10–7
5.61 ⋅ 10–12
6.82 ⋅ 10–6
5.48 ⋅ 10–16
4.74 ⋅ 10–32
1.43 ⋅ 10–20
4.39 ⋅ 10–23
6.60 ⋅ 10–6
1.70 ⋅ 10–5
7.40 ⋅ 10–14
2.30 ⋅ 10–13
9.80 ⋅ 10–9
5.45 ⋅ 10–27
1.14 ⋅ 10–7
1.10 ⋅ 10–4
1.12 ⋅ 10–10
3.00 ⋅ 10–17
45
Appendix 7
Standard Reduction Potentials
Reduction half-reaction
Eo, V
F2(g) + 2e− ⇄ 2F−(aq)
2.866
Co3+(aq) + e− ⇄ Co2+(aq)
1.920
H2O2(aq) + 2H+(aq) + 2e− ⇄ 2H2O(l)
1.776
Ce4+(aq) + e− ⇄ Ce3+(aq)
1.720
PbO2(s) + 4H+(aq) + SO42−(aq) + 2e− ⇄ PbSO4(s) + H2O(l)
1.691
MnO4−(aq) + 8H+(aq) + 5e− ⇄ Mn2+(aq) + 4H2O(l)
1.507
Au3+(aq) + 3e− ⇄ Au(s)
1.498
PbO2(s) + 4H+(aq) + 2e− ⇄ Pb2+(aq) + 2H2O(l)
1.455
Cl2(g) + 2e− ⇄ 2Cl−(aq)
1.358
Cr2O72−(aq) + 14H+(aq) + 6e− ⇄ 2Cr3+(aq) + 7H2O(l)
1.232
O2(g) + 4H+(aq) + 4e− ⇄ 2H2O(l)
1.229
MnO2(s) + 4H+(aq) + 2e− ⇄ Mn2+(aq) + 2H2O(l)
1.224
IO3−(aq) + 6H+(aq) + 6e− ⇄ I−(aq) + 3H2O(l)
1.085
Br2(l) + 2e− ⇄ 2Br−(aq)
1.066
AuCl4−(aq) + 3e− ⇄ Au(s) + 4Cl−(aq)
1.002
NO3−(aq) + 4H+(aq) + 3e− ⇄ NO(g) + 2H2O(l)
0.957
2Hg2+(aq) + 2e− ⇄ Hg22+(aq)
0.920
Ag+(aq) + e− ⇄ Ag(s)
0.800
Hg22+(aq) + 2e− ⇄ 2Hg(l)
0.797
Fe3+(aq) + e− ⇄ Fe2+(aq)
0.771
O2(g) + 2H+(aq) + 2e− ⇄ H2O2(aq)
0.695
I2(s) + 2e− ⇄ 2I−(aq)
0.536
Cu+(aq) + e− ⇄ Cu(s)
0.521
O2(g) + 2H2O(l) + 4e− ⇄ 4OH−(aq)
0.401
Cu2+(aq) + 2e− ⇄ Cu(s)
0.342
AgCl(s) + e− ⇄ Ag(s) + Cl−(aq)
0.222
Cu2+(aq) + e− ⇄ Cu+(aq)
0.153
2H+(aq) + 2e− ⇄ H2(g)
0.000
46
Appendix 7 (end)
Reduction half-reaction
Eo, V
2H+(aq) + 2e− ⇄ H2(g)
0.000
Fe3+(aq) + 3e− ⇄ Fe(s)
−0.037
Pb2+(aq) + 2e− ⇄ Pb(s)
−0.126
Sn2+(aq) + 2e− ⇄ Sn(s)
−0.138
Ni2+(aq) + 2e− ⇄ Ni(s)
−0.257
Cd2+(aq) + 2e− ⇄ Cd(s)
−0.403
Cr3+(aq) + e− ⇄ Cr2+(aq)
−0.407
Fe2+(aq) + 2e− ⇄ Fe(s)
−0.447
Cr3+(aq) + 3e− ⇄ Cr(s)
−0.744
Zn2+(aq) + 2e− ⇄ Zn(s)
−0.762
2H2O(l) + 2e− ⇄ H2(g) + 2OH−(aq)
−0.828
Al3+(aq) + 3e− ⇄ Al(s)
−1.662
Mg2+(aq) + 2e− ⇄ Mg(s)
−2.372
La3+(aq) + 3e− ⇄ La(s)
−2.379
Na+(aq) + e− ⇄ Na(s)
−2.711
Ca2+(aq) + 2e− ⇄ Ca(s)
−2.868
Ba2+(aq) + 2e− ⇄ Ba(s)
−2.912
K+(aq) + e− ⇄ K(s)
−2.931
Li+(aq) + e− ⇄ Li(s)
−3.040
47
REFERENCES
1. Dingrando, Laurel. Glencoe Chemistry: Matter and Change / Laurel Dingrando,
Kathleen V. Gregg, Nicolas Hainen, Cheryl Wistrom. – Glencoe / McGraw-Hill, 2004.
– 1021 p.
2. Houk, Clifford C. Chemistry Concepts and Problems / Clifford C. Houk, Richard
Post. – John Wiley and Sons, Inc., 1996. – 313 p.
3. Lewis, Rob. Chemistry / Rob Lewis, Wynne Evans. – Palgrave Macmillan, 2006.
– 479 p.
4. Masterton, William L. Chemistry Principles and Reactions / William L.
Masterton, Cecile N. Hurley. – Brooks/Cole Cengage Learning, 2004. – 727 p.
5. Mustoe, Frank. Chemistry / Frank Mustoe, Michael Jansen. – McGraw-Hill /
Ryerson, 2002. – 623 p.
6. Zumdahl, Steven S. World of Chemistry / Steven S. Zumdahl, Susan S. Zumdahl,
Donald J. DeCoste. – McDougal Littell, 2007. – 865 p.
48