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Chemistry 11
Final Examination Review - Answers
Part A - True or False.
Indicate whether each of the following statements is true or false. Correct the false statements.
T
F
F
1. Dalton's atomic theory includes a statement that says atoms of the same element are identical.
2. Dalton's theory states that atoms have variable masses.
Dalton’s atomic theory states that atoms of the same element have
the same mass and that atoms of different elements have different
masses.
3. J.J. Thomson is credited with the discovery of protons.
J.J. Thomson is credited with the discovery of electrons.
F
4. The mass of an electron is equal to the mass of a proton.
T
T
T
F
5. The mass of a proton is approximately equal to the mass of a neutron.
The mass of an electron is less than the mass of a proton.
T
F
6. The atomic number represents the number of protons in a nucleus.
7. The proton has a mass of approximately 1 u.
8. The difference in mass of isotopes of the same element is due to the different number of
protons in the nucleus.
The difference in mass of isotopes of the same element is due to the
different number of neutrons in the nucleus.
9. The isotope carbon-12 is used as the relative mass standard for the atomic mass scale.
10. The mass of the most common isotope of each element is listed on the periodic table.
The average mass of the naturally occurring isotopes of each element
is listed on the periodic table.
Part B - Multiple Choice
D
A
A
D
B
A
B
1. A(n) __ is used to represent a compound.
a) symbol
b) equation
c) subscript
d) formula
2. ___ atoms or groups of atoms are called ions.
a) charged
b) diatomic
c) neutral
d) monatomic
3. For the formula of a compound to be correct, the algebraic addition of the charges on the
atoms or ions in the compound must add up to __.
a) zero
b) one
c) two
d) four
4. Potassium bromide is an example of a(n) __ compound.
a) molecular
b) organic
c) polyatomic
d) ionic
5. The only common polyatomic ion that has a positive charge is the __ ion.
a) phosphate
b) ammonium
c) sulfate
d) nitrate
6. In the formula H2SO4, the number 4 would be called a(n) __.
a) subscript
b) oxidation number
c) coefficient
d) charge
7. Which subatomic particle contributes the least to the mass of an atom?
a) nucleon
b) electron
c) proton
d) neutron
Chemistry 11 – Final Examination Review – Answers – Page 1 of 27
D
C
B
B
C
A
C
B
B
D
A
B
D
B
D
8. In the free, or uncombined, state the number of protons in the nucleus of an element must
equal the __.
a) mass number
c) mass number - atomic number
b) number of neutrons in the nucleus
d) number of electrons present
9. Which of the following ideas of the Bohr model is not retained in the modern theory of
atomic structure?
a) Electrons can absorb or emit energy only in whole numbers of photons.
b) Atoms have a central positively charged nucleus.
c) Electrons move around the nucleus as planets orbit the sun.
d) Most of the volume of an atom is empty space.
10. Which of the following orbitals is spherical in shape?
a) 3p
b) 2s
c) 4d
d) 5f
11. The third energy level of an atom may have __ electrons.
a) 2
b) 18
c) 8
d) 32
12. How many sublevels are possible at the fourth energy level?
a) 2
b) 3
c) 4
d) 18
13. Lustrous, malleable, ductile elements that are good conductors of electricity and heat are
classified as __.
a) metals
b) nonmetals
c) metalloids
d) noble gases
14. The electron configuration of a certain element ends with 3p5. Which of the following
describes its position in the periodic table?
a) period 5, group 13 b) period 3, group 15
c) period 3, group 17 d) period 5, group 15
15. Which of the following is an example of a metalloid?
a) iodine
b) boron
c) bromine
d) indium
16. The periodicity of the elements is basically a function of their __.
a) nuclear stability
b) atomic numbers
c) mass numbers
d) none of these
17. An element with seven electrons in the outer level would be a __.
a) metal
b) metalloid
c) noble gas
d) nonmetal
18. As the atomic number in a period increases, the degree of nonmetallic character __.
a) increases
c) increases then decreases
b) decreases
d) remains the same
19. Elements in a group have similar chemical properties because of their similar __.
a) nuclear configurations
c) mass numbers
b) outer electron configurations
d) names
20. The period number in the periodic table designates the __ for the row.
a) total nuclear charge
c) maximum number of outer electrons
b) maximum number of nucleons
d) highest energy level
21. The radii of the atoms become smaller from sodium to chlorine across period 3. This is
primarily a result of __.
a) the shielding effect
c) the increased number of electrons
b) increased nuclear charge
d) decreased metallic character
22. Which of the following would correctly characterize a nonmetal?
a) low first ionization energy, low electron affinity
b) high first ionization energy, low electron affinity
c) low first ionization energy, high electron affinity
d) high first ionization energy, high electron affinity
Chemistry 11 – Final Examination Review – Answers – Page 2 of 27
D
B
A
C
D
A
A
D
A
D
B
D
C
A
A
C
23. Compared to the stability of the original atom, the stability of its ion that resembles a noble gas
configuration would be __.
a) identical
b) sometimes less
c) less
d) greater
24. Which of the following does NOT affect the ionization energy of an electron?
a) the distance of the electron from the nucleus c) the number of energy levels in the atom
b) the number of neutrons in the nucleus
d) the charge on the nucleus
25. For each subsequent electron removed from an atom the ionization energy required __.
a) increases
b) decreases
c) remains the same
d) no pattern
26. The formation of bonds between atoms depends on __.
a) the electron configurations of the atoms involved
b) the attraction the atoms have for electrons
c) both of the preceding factors
d) neither of the preceding factors
27. The particle that results when two or more atoms form covalent bonds is a __.
a) single charged atom b) molecule
c) polyatomic ion
d) b or c
28. Compounds that have low melting points, are brittle, and do not conduct electricity are probably __.
a) covalent
c) ionic
b) metallic
d) compounds of polyatomic ions
29. The most active __ have the highest electronegativities.
a) nonmetals
b) metalloids
c) metals
d) noble gases
30. __ compounds have high melting points, conduct electricity in the molten phase, and tend to be
soluble in water.
a) hydrogen
b) metallic
c) covalent
d) ionic
31. The element in the following group that has the lowest electronegativity is __.
a) potassium
b) arsenic
c) bromine
d) chromium
32. If there are only two electron pairs in the outer energy level of an atom in a molecule, they will be
found __.
a) at 90º to one another
c) at 120º to one another
b) on the same side of the nucleus
d) on opposite side of the nucleus
33. The __ molecule has two bonding pairs and two unshared pairs of electrons.
a) CH4
b) H2O
d) HF
c) NH3
34. A certain atom contains 34 protons, 34 electrons, and 45 neutrons. This atom has a mass
number of __.
a) 34
b) 45
c) 68
d) 79
35. An example of a compound is
a) oxygen
b) mercury
c) salt
d) diamond
36. Carbon is classed as an element rather than as a compound because it
a) cannot be chemically decomposed into two or more substances
b) has been known for many centuries
c) is formed when wood is heated out of contact with air
d) combines with oxygen to form a gas
37. The positively charged particles in the nucleus of an atom are called
a) protons
b) neutrons
c) electrons
d) ions
38. The charged particles that are found outside the nucleus of an atom are called
a) protons
b) ions
c) electrons
d) mesons
Chemistry 11 – Final Examination Review – Answers – Page 3 of 27
D
A
A
B
39. A nonmetallic atom generally becomes a negative ion by
a) losing protons
b) losing electrons
c) gaining protons
d) gaining electrons
40. Which of the following subatomic particles has the smallest mass?
a) electron
b) neutron
c) proton
d) nucleus
41. The nuclear atom was first postulated by
a) Rutherford
b) Bohr
c) Dalton
d) Thompson
42. Which of the following symbols represents an atom that contains the largest number of neutrons?
235
a) 92 U
C
C
239
16
b) 92 U
239
239
c) 93 Np
d) 94 Pu
43. The nuclide symbol 8 O represents an oxygen atom with
a) a mass of 8 u
c) a mass of 16 u
b) an atomic number of 16
d) 16 neutrons
44. If Z represents the atomic number of an element and A represents the mass number, then the
number of neutrons in one atom is
a) A
b) A + Z
c) A - Z
d) Z - A
24
25
D
45. Which of the following statements about the elemental species 11 X and 12 Z is correct?
a) They are isotopes of the same element.
b) They are nonmetals.
c) They are members of the same chemical family.
d) They have the same number of neutrons per atom.
B
46. An element X has a mass number of 32 and an atomic number of 16. The most common ion
of element X is represented by
a) X+
b) X2c) Xd) X2+
Part C - Short Answer
1.
What is the maximum number of electrons that may occupy one orbital?
two
2. The Lewis electron dot diagram is used to represent only which electrons in the atom?
Valence electrons
3. What is the diagonal rule used to predict?
The order in which orbitals are filled
4. How many sublevels are possible at the third energy level?
three
5. How many orbitals are there in the f sublevel?
seven
6. What is the maximum number of electrons that can occupy a d sublevel?
ten
7. Which sublevel may contain a maximum of three pairs of electrons?
‘p’ sublevel
Chemistry 11 – Final Examination Review – Answers – Page 4 of 27
8. What must be true about the spins of two electrons occupying the same orbital?
They must have opposite spin
9. Write the electron configuration for each of the following elements:
1s22s1
a) lithium
b) radium
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2
c) sodium
1s22s22p63s1
d) mercury
1s22s22p63s23p64s23d104p65s24d105p66s24f145d10
10. Draw the energy level diagram for each of the following elements:
a) tin
6s
↑↓
5s
↑
↑
5p
↑↓
4p
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
↑↓
↑↓
↑↓
4s
↑↓
3s
↑↓
2s
↑↓
↑↓
↑↓
4d
↑↓
↑↓
↑↓
↑↓
↑↓
3d
↑↓
↑↓
↑↓
1s
Chemistry 11 – Final Examination Review – Answers – Page 5 of 27
10. Draw the energy level diagram for each of the following elements:
b) krypton
↑↓
4p
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
↑↓
↑↓
5s
↑↓
↑↓
↑↓
4s
↑↓
3s
↑↓
2s
↑↓
3d
↑↓
↑↓
↑↓
1s
c) gold
6p
↑↓
6s
↑↓
5s
↑↓
↑↓
5p
↑↓
↑↓
4p
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
↑↓
↑↓
↑↓
4s
↑↓
3s
↑↓
2s
↑↓ ↑↓ ↑↓
5d
↑↓
↑
↑↓ ↑↓ ↑↓
4d
↑↓
↑↓
↑↓ ↑↓ ↑↓
3d
↑↓
↑↓
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4f
↑↓
1s
Chemistry 11 – Final Examination Review – Answers – Page 6 of 27
10. Draw the energy level diagram for each of the following elements:
d) potassium
↑
4s
3d
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
↑↓
↑↓
3s
↑↓
2s
↑↓
1s
11. Write the energy level population for each of the following elements:
a) calcium
2, 8, 8, 2
b) sulfur
c) scandium
d) tungsten
2. 8. 6
2, 8, 9, 2
2, 8, 18, 32, 12, 2
12. Element X has the following configuration: 1s22s22p63s22p64s23d104p65s1
a) What period is this element located in?
five
b) What group is this element in?
one (alkali metals)
c) Identify the element.
Rubidium (Rb)
d) Is the element a metal, nonmetal, or metalloid?
Metal
13. List five properties of metals and five properties of nonmetals.
Metals
-
good conductors of heat and electricity
lustre
malleable and ductile
high densities
high boiling points and melting points
resist stretching and twisting
solids at room temperature
Nonmetals
-
no lustre
poor conductors
may be solid, liquid or gas
low densities
low melting and boiling points
Chemistry 11 – Final Examination Review – Answers – Page 7 of 27
14. In a group, do the radii of atoms increase or decrease as the atomic number increases? Explain
why.
In a group, the radii increases as the atomic number increases because
electrons are being added at higher energy levels which are located
further from the nucleus producing a larger atom.
15. Does the radii of atoms within a period increase or decrease as the atomic number increases?
Explain why.
The radii of atoms in a period decrease as the atomic number increases.
As you go across (left to right) a period the electrons are being added to
the same energy level while more protons are added to the nucleus. The
attractive force is greater and the electrons are pulled closer to the
nucleus producing a smaller atom.
16. In each of the following pairs of atoms, pick the one that is larger.
b) K, Ca
c) Al, B
d) Br, Cl
e) F, N
a) Mg, Na
f) Ne, Ar
17. In a group, will the ionization energy tend to increase or decrease with increasing atomic number?
Explain why.
In a group the ionization energy decreases as the atomic number
increases. As you go down a group the size of the atom increases causing
the electrons to be located further away from the nucleus. The attraction
between the nucleus (+) and electrons (-) which decreases as the
distance between them increases, which makes it easier to remove an
electron.
18. In a period, will the ionization energy tend to increase or decrease with increasing atomic number?
Explain why.
In a period the ionization energy increases as the atomic number
increases. As you go across a period (left to right) the size of the atom
decreases, the electrons are located closer to the nucleus with a greater
attraction between the electrons and the nucleus. Because the attraction
is greater, more energy is required to remove an electron.
19. Do metals or nonmetals generally have lower ionization energies?
Metals generally have lower ionization energies than nonmetals.
20. Carbon has a first ionization energy of 1086.5 kJ/mol. Predict whether the first ionization energies of the
following elements will be more or less than that of carbon.
higher ionization energy
a) helium
b) lithium
c) fluorine
d) silicon
lower ionization energy
higher ionization energy
lower ionization energy
Chemistry 11 – Final Examination Review – Answers – Page 8 of 27
21. As the distance between the nucleus and the outer electrons of an atom increases, will the
ionization energy increase or decrease?
As the distance between the outer electrons and the nucleus increases,
the attractive force decreases which causes the ionization energy to
decrease.
22. As the positive charge on an ion increases, will the ionization energy increase or decrease?
As the positive charge on the ion increases the electrons are held more
tightly and the ionization energy increases.
23. How are substances that are gases or soft solids at room temperature classified?
Substances that are gases or soft solids at room temperature are classified
as covalent compounds.
24. Would an element with two outer electrons be a metal or a nonmetal?
An element with two outer electrons would be a metal.
25. What Russian scientist designed the first periodic table?
Mendeleev designed the first periodic table.
26. Which group of elements has eight outer electrons?
The noble gases have eight outer electrons.
27. For the transition elements, as the atomic number increases, to which sublevel are the electrons
being added?
For the transition metals, the electrons are being added to the “d”
sublevel.
28. According to the octet rule, how many pairs of outer electrons do the most stable atoms have?
According to the octet rule, the most stable atoms have four pairs of outer
electrons.
29. In the lanthanide series, as the atomic number increases, to which sublevel are electrons being
added?
In the lanthanide series, electrons are being added to the “f” sublevel.
30. Which atom in each of the following pairs would have the lower first ionization energy?
a) N, O
b) Te, Se
c) C, Ge
d) Br, I
e) I, Sb
f) Al, N
Chemistry 11 – Final Examination Review – Answers – Page 9 of 27
31. In a group, will the electron affinity tend to increase or decrease with increasing atomic number?
Explain why.
In a group, electron affinity decreases as atomic number increases. As
you go down the group, the size of the atom increases, electrons are
added further from the nucleus where the attraction is less, therefore less
energy is released when an electron is added to the atom, and electron
affinity is smaller.
32. In a period, will the electron affinity tend to increase or decrease with increasing atomic number?
Explain why.
In a period, the electron affinity increases as the atomic number
increases. As you go across (left to right) the size of the atom decreases
and the electrons are added closer to the nucleus where the attraction is
greater – causing more energy to be released when the electron is added,
increasing the electron affinity. (Noble gases do not follow this trend
because of their increased stability, they do not gain electrons easily.)
33. Which atom in each of the following pairs would have the higher electron affinity?
b) At, Rn
c) Y, Os
d) Ga, In
e) Br, I
f) O, N
a) S, Cl
34. Give examples of molecules with the following shapes:
a) linear
HCl, CO2, BeCl2
BH3, BCl3, AlBr3
NH3, NH3
CH4, CCl4
H2O, H2S
b) trigonal planer
c) trigonal pyramidal
d) tetrahedral
e) bent
35. Use electron dot formulas to show the complete balanced bonding reactions between the following
elements. Indicate the type of bond expected.
a) sodium and oxygen
b) nitrogen and hydrogen
c) zirconium and sulfur
d) magnesium and chlorine
Chemistry 11 – Final Examination Review – Answers – Page 10 of 27
36. Write the correct names for the following chemical compounds.
m) SO3
a) HCl(aq)
hydrochloric acid
potassium hydroxide
HgOH
mercury(I) hydroxide
FeCl3
iron(III) chloride
Al2(SO4)3 aluminum sulfate
N2O5
dinitrogen pentoxide
HF
hydrofluoric acid
Pb(OH)2
lead(II) hydroxide
NH4NO3 ammonium nitrate
NaHCO3
sodium bicarbonate
Zn(NO2)2 zinc nitrite
H3PO4(aq) phosphoric acid
b) KOH
n) NaC2H3O2
c)
o) HFO2(aq)
d)
e)
f)
g)
h)
i)
j)
k)
l)
p) Al(BrO4)3
q) SbF3
r) Pd(CN)2
s) Ca(MnO3)2
t) Be(NO4)2
u) NiSeO4
v) H2SO3(aq)
w) Ba(OH)2
x) PbS
sulfur trioxide
sodium acetate
flourous acid
aluminum perbromate
antimony(III) fluoride
palladium(II) cyanide
calcium manganate
beryllium pernitrate
nickel(II) selenate
sulfurous acid
barium hydroxide
lead(II) sulfide
37. Write the correct chemical formula for each compound and balance the equation.
a) sodium carbonate + calcium hydroxide ¼ sodium hydroxide + calcium carbonate
Na2CO3 + Ca(OH)2 ¼ 2NaOH + CaCO3
b) carbon dioxide + water ¼ carbonic acid
CO2 + H2O ¼ H2CO3
c) phosphorus + oxygen ¼ phosphorus pentoxide
2P + 5O2 ¼ 2PO5
d) sodium + water ¼ sodium hydroxide + hydrogen
2Na + 2H2O ¼ 2NaOH + H2
e) zinc + sulfuric acid ¼ zinc sulfate + hydrogen
Zn + H2SO4 ¼ ZnSO4 + H2
f) aluminum sulfate + calcium hydroxide ¼ aluminum hydroxide + calcium sulfate
Al2(SO4)3 + 3Ca(OH)2 ¼ 2Al(OH)3 + 3CaSO4
g) calcium oxide + water ¼ calcium hydroxide
CaO + H2O ¼ Ca(OH)2
h) iron + copper(I) nitrate ¼ iron(II) nitrate + copper
Fe + 2CuNO3 ¼ Fe(NO3)2 + 2Cu
i)
iron(II) sulfide + hydrochloric acid ¼ hydrogen sulfide + iron(II) chloride
j)
potassium oxide + water ¼ potassium hydroxide
FeS + 2HCl ¼ H2S + FeCl2
K2O + H2O ¼ 2KOH
Chemistry 11 – Final Examination Review – Answers – Page 11 of 27
k.
carbon + ferric oxide ¼ iron + carbon dioxide
3 C + 2 Fe2O3 ¼ 4 Fe + 3 CO2
l.
sulfur tetrafluoride + water ¼ sulfur dioxide + hydrofluoric acid
SF4 + 2 H2O ¼ SO2 + 4 HF
m. calcium hydroxide + phosphoric acid ¼ calcium phosphate + water
3 Ca(OH)2 + 2 H3PO4 ¼ Ca3(PO4)2 + 6 H2O
n.
ethane + oxygen ¼ carbon dioxide + water
2 C2H6 +
o.
7 O2 ¼ 4 CO2 + 6 H2O
aluminum sulfate + ammonia + water ¼ aluminum hydroxide + ammonium sulfate
Al2(SO4)3 + 6 NH3 + 6 H2O ¼ 2 Al(OH)3 + 3 (NH4)2SO4
38. Identify the type of reaction. Write the correct chemical formulas of the compounds, complete and
balance the equations and name the products formed.
calcium chloride + carbonic acid
+ H2CO3
a) calcium carbonate + hydrochloric acid ¼
CaCO3 + 2HCl ¼ CaCl2
double displacement
b) iron + sodium bromide ¼
Fe + NaBr ¼ no reaction b/c iron is less active than sodium
single displacement
ammonium chloride + iron(II) acetate
¼ 2NH4Cl + Fe(C2H3O2)2
c) ammonium acetate + iron(II)chloride ¼
2NH4C2H3O2 + FeCl2
double displacement
silver sulfate + ammonium bromide
¼ Ag2SO4 + 2NH4Br
d) silver bromide + ammonium sulfate ¼
2AgBr + (NH4)2SO4
double displacement
e) zinc + sulfuric acid ¼
Zn + H2SO4 ¼ H2
single displacement
hydrogen + zinc sulfate
+ ZnSO4
f) neon + potassium ¼
Ne + K ¼ no reaction b/c noble gases are unreactive
combination
lead(II) chloride + water
+ 2H2O
g) lead(II )hydroxide + hydrochloric acid ¼
Pb(OH)2 + 2HCl ¼ PbCl2
double displacement
iron(II) sulfide
Fe + S ¼ FeS
combination
h) iron + sulfur ¼
Chemistry 11 – Final Examination Review – Answers – Page 12 of 27
i)
potassium chloride + oxygen
potassium chlorate (heated) ¼
2KClO3 ¼ 2KCl + 3O2
decomposition
j)
calcium hydroxide
CaO + H2O ¼ Ca(OH)2
combination
calcium oxide + water ¼
nitric acid
N2O5 + H2O ¼ 2HNO3
combination
12. carbon dioxide + water ¼ carbonic acid
CO2 + H2O ¼ H2CO3
combination
13. chlorine + chromium(III) bromide ¼ bromine + chromium(III) chloride
11. dinitrogen pentoxide + water ¼
3 Cl2 + 2 CrBr3 ¼ 3 Br2 + 2 CrCl3
Single displacement
14. sulfur + oxygen ¼sulfur
dioxide or sulfur trioxide
S + O2 ¼ SO2
or S + O2 ¼ SO3
combination
15. zinc + hydrochloric acid ¼
zinc chloride + hydrogen
Zn + 2 HCl ¼ ZnCl2 + H2
Single displacement
16. sodium + water ¼
sodium hydroxide + hydrogen
2 Na + 2 H2O ¼ 2 NaOH + H2
single displacement
17. magnesium + water ¼
Mg + H2O ¼no reaction – Mg is not active enough to replace H in water
Single displacement
18. copper + stannic nitrate ¼
Cu + Sn(NO3)4 ¼ no reaction – Cu is less active than Sn
Single displacement
19. aluminum + cupric sulfate ¼
copper + aluminum sulfate
2 Al + 3 CuSO4 ¼ 3 Cu + Al2(SO4)3
Double Displacement
Chemistry 11 – Final Examination Review – Answers – Page 13 of 27
Part D – Calculations
Show all work. Express your answers using significant digits and include units with your answers.
1.
Calculate the average atomic mass of the following elements:
a) Mg-24
Mg-25
Mg-26
mass = 23.985 u
mass = 24.986 u
mass = 25.983 u
78.70%
10.13%
11.17%
(0.7870)(23.985) + (0.1013)(24.986) + (0.1117)(25.983) = 24.31 u
b) Ir-191
Ir-193
mass = 191.0 u
mass = 193.0 u
37.58%
62.42%
(0.3758)(191.0) + (0.6242)(193.0) = 192.2 u
2. Calculate the percentage of each of the isotopes of silver if silver-107 has a mass of 106.905 u and
silver-109 has a mass of 108.905 u and the average atomic mass of silver is 107.869 u
107
Ag = x,
109
Ag = y
x+y=1
106.905x + 108.905y = 107.869
-106.905x - 106.905y = -106.905
106.905x + 108.905y = 107.869
2y = 0.964
¼
107
Ag = 51.8%,
109
use elimination or substitution
y = 0.482
¼
x = 0.518
Ag = 48.2%
Chemistry 11 – Final Examination Review – Answers – Page 14 of 27
*3. Naturally, occurring silicon consists of three isotopes, 28Si, 29Si, and 30Si, whose atomic masses are 27.9769,
28.9765, and 29.9738, respectively. The most abundant isotope is 28Si, which accounts for 92.23 percent
of naturally occurring silicon. Given that the observed atomic mass of silicon is 28.0855, calculate the
percentages of 29Si and 30Si in nature.
Isotope
Mass
Percentage
28
29
Si
27.9769 g
92.23%
Si
28.9765 g
x
30
Si
29.9738 g
y
Average
28.0855
(eqn 1)
(eqn 1)
(eqn 1)
(0.9223)(27.9769) + (28.9765)(x) + (29.9738)(y) = 28.0855
25.8031 + 28.9765x + 29.9738y = 28.0855
28.9765x + 29.9738y = 2.2824
(eqn 2)
(eqn 2)
0.9223 + x + y = 1
x + y = 0.0777
(eqn 1)
(-28.9756)(eqn 2)
y = 0.0310,
use substitution or elimination
28.9765x + 29.9738y = 2.2824
-28.9765x – 28.9765y = -2.2824
0.9973y = 0.03093
x = 0.0467,
∴
29
y = 0.03093
0.9973
Si = 3.10% and
30
Si = 4.67%
4. Complete the chart below:
Element
Atomic #
Mass #
Protons
Neutrons
Electrons
calcium-43
20
43
20
23
20
lead-211
82
211
82
129
82
plutonium-242
94
242
94
148
94
chromium-50
24
50
24
26
24
29
65
29
36
27
S2-
16
34
16
18
18
128 53 I
53
128
53
75
54
208
4+
82 Pb
82
208
82
126
78
65
Cu2+
34
Chemistry 11 – Final Examination Review – Answers – Page 15 of 27
5. Convert each of the following to moles.
a. 8.8 g of potassium carbonate
8.8 g K2CO3 + 138.2052 g/mol = 0.0637 mol
b. 0.257 g of arsenic pentachloride
0.257 g AsCl5 + 252.1866 g/mol = 1.02 x 10-3 mol
c. 12.5 L of carbon dioxide at STP
12.5 L CO2 + 22.4 L/mol = 0.558 mol
d. 5.00 x 1013 atoms of iron
5.00 x 1013 atoms of Fe + 6.022 x 1023 =
e. 8.63 x 10
28
0.830 mol
molecules of water
8.63 x 1028 molecules of H2O + 6.022 x 1023 =
1.43 x 105 mol
f. 450.0 mL helium at STP
0.4500 L He + 22.4 L/mol = 0.0201 mol
g. 236.0 g of ammonium phosphate
236.0 g (NH4)3PO4 + 149.086 74 g/mol = 1.58 mol
h. 15.0 g of butanoic acid
15.0 g CH3CH2CH2COOH
+
88.106 32 g/mol = 0.170 mol
6. Calculate the mass of each of the following.
a. 2.60 mol of sodium carbonate
(2.60 mol)(105.988 74 g/mol) = 276 g
b. five million atom of gold
(5 000 000 atoms) + (6.022 x 1023) = 8.30 x 10-18 mol
(8.30 x 10-18 mol)(196.9665 g/mol) = 1.64 x 10-15 g
c. 25.0 mL of carbon dioxide at STP
(0.0250 L) + (22.4 L/mol) = 0.001 12 mol
(0.001 12 mol)(44.0098 g/mol) = 0.0491 g
d. 4.50 x 1021 molecules of decanoic acid
CH3CH2CH2CH2CH2CH2CH2CH2CH2COOH or C10H202
4.50 × 10 21
= 0.00747 mol
6.022 × 10 23
g
) = 1.29
(0.00747 mol )(172.2676 mol
g
Chemistry 11 – Final Examination Review – Answers – Page 16 of 27
7. Calculate the molarity of 825 mL of solution, which contains 30.0 g of acetic acid.
30.0 g CH 3 COOH
g
mol
60.05256
= 0.4996 g
0.4996 g
= 0.606
0.825 L
mol
L
8. What volume of solution can be made from 80.0 g of sodium hydroxide if a 2.00 M solution is required?
80.0 g NaOH
g
mol
38.887 11
V =
n
C
= 2.00 mol
2.00 mol
=
= 1.00 L
mol
L
2.00
9. What mass of calcium chloride is required to produce 750.0 mL of a 0.500 M solution?
n = C×V
=
(0.500
(0.375 mol)(110.984
mol
L
)(0.7500 L)
)
= 41.6 g
g
mol
= 0.375 mol
10. What volume 0f 14.0 M nitric acid is required to produce 750.0 mL of a 0.250 M solution?
C1 V1 = C2 V2
C V
V1 = 2 2
C1
=
(0.250 )(0.7500L) = 0.0134 L
mol
L
14.0
mol
L
or
13.4 mL
11. What concentration results when 250.0 mL of a 0.125 M solution of hydrochloric acid is mixed with 125.0
mL of a 1.00 M solution?
C ab
=
C V + C b Vb
= a a
Va + Vb
=
(0.125
0.03125 mol + 0.125 mol
0.375 L
mol
L
)(0.2500 L) + (1.00
mol
L
0.2500 L + 0.1250 L
=
0.417
mol
L
Chemistry 11 – Final Examination Review – Answers – Page 17 of 27
)(0.1250 L)
12. What volume of 0.225 M sulfuric acid must be mixed with 500.0 mL of a 0.750 M solution in order to obtain
a 0.500 M solution?
C a Va + C b Vb
Va + Vb
C ab =
0.500
(0.500
0.500
(0.275
Va =
mol
L
mol
L
(0.225
=
)(V
mol
L
a
mol
L
)(0.5000 L )
Va + 0.5000L
=
(0.225
mol
L
)V
a
+ 0.375 mol
+ 0.250 mol
=
(0.225
mol
L
)V
a
+ 0.375 mol
+ 0.5000 L
)V
a
0.125 mol
0.275
)V + (0.750
)
a
mol
Va
L
mol
L
mol
L
=
0.125 mol
= 0.455 L
E. STOICHIOMETRY - Begin each problem by writing a balanced chemical equation.
1.
Carbon dioxide is produced in the reaction between calcium carbonate and hydrochloric acid. How many
grams of calcium carbonate would be needed to react completely with 15.0 g of hydrochloric acid? How
many grams of calcium chloride would be formed?
CaCO3 + 2HCl ¼ H2O + CO2 + CaCl2
⎛ 15.0 g HCl ⎞⎛ 1 mol CaCO
⎜
⎟⎜
3
⎜
⎟
⎜
g
⎜ 36.46094
⎟⎜
HCl
mol ⎠⎝ 2 mol
⎝
⎞
⎟
⎟ 100.0892
⎟
⎠
(
⎛ 15.0 g HCl ⎞⎛ 1 mol CaCl ⎞
⎜
⎟⎜
2 ⎟
⎜
⎟
⎜
⎟ 110.986
g
⎜ 36.460 94
⎟⎜ 2 mol HCl ⎟
mol ⎠⎝
⎝
⎠
(
g
mol
g
mol
) = 20.6 g
) = 22.8 g
Chemistry 11 – Final Examination Review – Answers – Page 18 of 27
CaCO 3
CaCl2
2.
Sulfur dioxide may be catalytically oxidized to sulfur trioxide. How many grams of sulfur dioxide could be
converted by this process if 100.0 g of oxygen are available for the oxidation?
2SO2 + O2 ¼ 2SO3
⎛ 100.0 g O ⎞⎛ 2 mol SO ⎞
⎜
2 ⎟⎜
2 ⎟
⎜
⎟
⎜
⎟ 64.0648
g
⎜ 31.9988
⎟⎜ 1 mol O2 ⎟
mol ⎠⎝
⎝
⎠
(
g
mol
) = 400.4 g
SO2
3. Phosphoric acid is produced in the reaction between calcium phosphate and sulfuric acid. How much of the
phosphoric acid would be produced from 55.0 g of the calcium phosphate?
Ca3(PO4)2 + 3H2SO4 ¼ 3CaSO4 + 2H3PO4
⎛ 55.0 g Ca (PO )
⎜
3
4 2
⎜
g
310.183
⎜
mol
⎝
⎞⎛ 2 mol H PO
⎟⎜
3
4
⎟⎜
⎟⎜ 1 mol Ca3 (PO 4 )2
⎠⎝
⎞
⎟
⎟ 97.995
⎟
⎠
(
g
mol
) = 34.8 g
H3PO 4
4. How much magnesium sulfate is needed to completely react with 145 g of sodium chloride?
sodium sulfate could be produced by this reaction?
MgSO4 + 2NaCl ¼ Na2SO4 + MgCl2
⎛ 145 g NaCl ⎞⎛ 1 mol MgSO ⎞
⎜
⎟⎜
4 ⎟
g
⎜
⎟
⎜
⎟ 120.3686 mol = 149 g
g
⎜ 58.443
⎟⎜
NaCl ⎟⎠
mol ⎠⎝ 2 mol
⎝
(
⎛ 145 g NaCl ⎞⎛ 1 mol Na SO
⎜
⎟⎜
2
4
⎜
⎟
⎜
g
⎜ 58.443
⎟⎜
NaCl
mol ⎠⎝ 2 mol
⎝
⎞
⎟
⎟ 142.043
⎟
⎠
(
g
mol
)
MgSO 4
) = 176 g
Na2 SO 4
How much
5. Gold will dissolve in the acid mixture known as aqua regia according to the following reaction:
Au + HNO3 + 3 HCl Î AuCl3 + NO + 2 H2O
How much gold(III)chloride will be produced in this reaction when one starts with 5.0 mg of gold?
⎛ 5.0 mg Au ⎞⎛ 1 mol AuCl ⎞
3 ⎟
⎜
⎟⎜
⎜
⎟ 303.3255
mg ⎟⎜
⎜ 196.9665 mmol ⎟⎜ 1 mol Au ⎟
⎝
⎠⎝
⎠
(
mg
mmol
) = 7.7
mg
AuCl 3
How much hydrochloric acid must be added initially to dissolve the all this gold?
⎛ 5.0 mg Au ⎞⎛ 3 mol
⎜
⎟⎜
⎜
mg ⎟⎜
⎜ 196.9665 mmol ⎟⎜ 1 mol
⎝
⎠⎝
HCl ⎞⎟
⎟ 36.461
Au ⎟⎠
(
mg
mmol
) = 2.8
Chemistry 11 – Final Examination Review – Answers – Page 19 of 27
mg
HCl
6. How many grams of sulfuric acid will react with 400.0 g of aluminum metal?
⎛ 400.0 g Al ⎞⎛ 3 mol H SO
2
4
⎜
⎟⎜
⎜
⎟
⎜
g
⎜ 26.98154 mol ⎟⎜ 2 mol Al
⎝
⎠⎝
⎞
⎟
⎟ 98.079
⎟
⎠
(
g
mol
) = 2181
g
H2 SO4
7. It is desired to prepare 50.0 g of water by synthesis. How many litres of hydrogen must be used?
⎛ 50.0 g H O ⎞⎛ 2 mol H ⎞
2
2 ⎟
⎜
⎟⎜
⎜
⎟ 22.4
g ⎟⎜
⎜ 18.01528 mol ⎟⎜ 2 mol H 2 O ⎟
⎝
⎠⎝
⎠
(
8.
L
mol
) = 62.2
L
H2
An unknown amount of potassium chlorate was heated until no more oxygen was evolved.
potassium chloride remained in the test tube.
15.824 g of
2KClO3 ¼ 2KCl + 3O2
What mass of potassium chlorate had originally been placed in the tube?
⎛ 15.824 g
⎜
⎜
⎜ 74.5513
⎝
KCl ⎞⎟⎛⎜ 2 mol KClO3 ⎞⎟
⎟⎜
⎟ 122.5495
g
⎟
⎜
⎟
mol
KCl
2
mol
⎠⎝
⎠
(
g
mol
) = 26.0
g
KClO3
What volume of oxygen gas was evolved in the process?
⎛ 15.824 g
⎜
⎜
⎜ 74.5513
⎝
9.
KCl ⎞⎟⎛⎜ 3 mol
⎟⎜
g
⎟⎜ 2 mol
mol
⎠⎝
O2 ⎞⎟
⎟ 22.4
KCl ⎟⎠
(
L
mol
) = 7.13 L
O2
How many grams of carbon can be completely burned in 15.0 L of oxygen?
⎛ 15.0 L
⎜
⎜
⎜ 22.4
⎝
O2 ⎞⎟⎛⎜ 1 mol C ⎞⎟
⎟⎜
⎟ 12.011
L
⎟⎜ 1 mol O 2 ⎟
mol
⎠⎝
⎠
(
g
mol
) = 8.04
g
C
How many litres of carbon dioxide gas are produced in this reaction?
⎛ 15.0 L
⎜
⎜
⎜ 22.4
⎝
O2 ⎞⎟⎛⎜ 1 mol CO2 ⎞⎟
⎟⎜
⎟ 22.4
L
⎟
⎜
⎟
mol
⎠⎝ 1 mol O 2 ⎠
(
L
mol
) = 15.0
Chemistry 11 – Final Examination Review – Answers – Page 20 of 27
L
CO2
10. How many grams of magnesium metal are required to liberate 250.0 mL of hydrogen gas from hydrochloric
acid?
⎛ 0.2500 L H ⎞⎛ 1 mol
2 ⎟⎜
⎜
⎜
⎟⎜
L
22.4 mol
⎜
⎟⎜ 1 mol
⎝
⎠⎝
Mg ⎞⎟
⎟ 24.305
H 2 ⎟⎠
(
g
mol
) = 0.271
g
Mg
Exactly how much acid would be used up in this reaction?
⎛ 0.2500 L H ⎞⎛ 2 mol HC lg ⎞
2 ⎟⎜
⎜
⎟
⎜
⎟⎜
⎟ 36.46094
L
22.4 mol
⎜
⎟⎜ 1 mol H 2 ⎟
⎝
⎠⎝
⎠
(
11.
g
mol
) = 0.8141
g
HCl
Will 30.0 L of fluorine gas completely react with 50.0 L of hydrogen gas? Which gas is in excess and by
how much? What volume of hydrogen fluoride is formed in this reaction?
⎛ 30.0 L
⎜
⎜
⎜ 22.4
⎝
F2 ⎞⎟⎛⎜ 2 mol
⎟⎜
L
⎟⎜ 1 mol
mol
⎠⎝
HF ⎞⎟
⎟ 22.4
F 2 ⎟⎠
⎛ 50.0 L
⎜
⎜
⎜ 22.4
⎝
H 2 ⎞⎟⎛⎜ 2 mol
⎟⎜
L
⎟⎜ 1 mol
mol
⎠⎝
HF ⎞⎟
⎟ 22.4
H 2 ⎟⎠
(
(
L
mol
) = 60.0
L
mol
L
HF
) = 100. L
HF
60.0 L of hydrogen fluoride is produced.
Fluorine is the limiting reactant and hydrogen is in excess.
⎛ 30.0 L F ⎞⎛ 1 mol H ⎞
2 ⎟⎜
2 ⎟
⎜
L
⎜
⎟⎜
⎟ 22.4 mol = 30.0 L HF needed
L
⎜ 22.4 mol ⎟⎜ 1 mol O 2 ⎟
⎝
⎠⎝
⎠
50.0 L of hydrogen available - 30.0 L needed = 20.0 L in excess
(
)
Chemistry 11 – Final Examination Review – Answers – Page 21 of 27
12. A mixture containing 100.0 g of H2 and 100.0 g of O2 is sparked so that water is formed. How much water
is formed?
⎛ 100.0 g H ⎞⎛ 2 mol H O ⎞
2 ⎟⎜
2
⎜
⎟
⎜
⎟⎜
⎟ 18.01528
g
⎜ 2.01588 mol ⎟⎜ 2 mol H 2 ⎟
⎝
⎠⎝
⎠
g
mol
) = 894
g
H2 O
⎛ 100.0 g O ⎞⎛ 2 mol H O ⎞
2 ⎟⎜
2
⎜
⎟
⎜
⎟ 18.01528
g ⎟⎜
⎜ 31.9988 mol ⎟⎜ 1 mol o 2 ⎟
⎝
⎠⎝
⎠
g
mol
) = 113
g
H2 O
(
(
Oxygen is the limiting reactant and hydrogen is in excess.
113 g of water are formed.
13. When copper is heated with sulfur, Cu2S is formed. How many grams of Cu2S could be produced if 100.0 g
of copper is heated with 50.0 g of sulfur?
⎛ 100.0 g
⎜
⎜
⎜ 63.546
⎝
⎛ 50.0 g
⎜
⎜
⎜ 32.066
⎝
Cu ⎞⎟⎛⎜ 1 mol Cu2 S ⎞⎟
⎟⎜
⎟ 159.158
g
⎟
⎜
⎟
2
mol
Cu
mol
⎠⎝
⎠
(
S ⎞⎟⎛⎜ 1 mol Cu2 S ⎞⎟
⎟ 159.158
g ⎟⎜
⎟
⎜
⎟
1
mol
S
mol
⎠⎝
⎠
(
g
mol
g
mol
) = 125.2
) = 248
g
g
Cu2 S
Cu2 S
Copper is the limiting reactant and sulfur is in excess.
125.2 g of copper(I) sulfide is formed.
14. What volume (at STP) of carbon monoxide is required to produce 100.0 g of iron according to the equation:
(not balanced)
Fe2O3 + CO Î Fe + CO2
⎛ 100.0 g
⎜
⎜
⎜ 55.847
⎝
Fe ⎞⎟⎛⎜ 3 mol
⎟⎜
g
mol ⎟⎜ 2 mol
⎠⎝
CO ⎞⎟
⎟ 22.4
F2 ⎟
⎠
(
L
mol
) = 60.2
Chemistry 11 – Final Examination Review – Answers – Page 22 of 27
L
CO
*15. Hydrozineis a nitrogen-hydrogen
hovingthe formuloNzH+. rt is on oily, colourlessliguidthot
compound
derivedfrom it
andloils ot 113.5"c.ihe principoluseof hydrozineondcertoincompounds
freezesot1.5oC
dissolved
is o rocket fuels,but it is olsousedin fuei cells,in the treotmentof woter in boilersto removed
is the
hydrozine
ond in the plosticsindustry. Onewidelyusedmethodfor the monufoctureol
oxygengos,
is
process.The Roschigprocessinvolvesthree reoctionsteps. rn the first sfep,sodiumhydroxide
Roschig
hypochlorite
sodium
The
sodiumchlorideondwoter.
reoctedwith chlorineto proJucesodiumhypochlorite,
(NHzcl)ond
in
producedin the first step is reoctedwith ommonio the secondstep to producechloromine
with
is
reocted
step
sodiumhydroxide. The chloromineond sodiumhydroxideproducedin the second
sodiumchlorideand woter. A chemicolploneusingthe
in the third step to producehydrazine,
ommonio
g8.O%
hydrozine for every 1.00 k9 of chlorine. Whot ore the
Roschigprocessobtoins O.2gg k9 of
theoreticol,octuol,ond percentyieldsof pure hydrozine?
2NaOH * Clz ) NaOCI+ NaCl * HzO
NaOCI*NHgtNH2C|+NaOH
NH2CI+ NaOH * NHg) NzHr + NaCl * HzO
1OOOg
70.906
= 452 g
Ct,lfr
c
tnol
L mot NH,ct)ltmor n,n.l6r.oorr,
-"r ruaocr]f
;r)
NH2ct
Jt1-", % Jtt mot
NrHo
=
naoct )lL mot
)
theoretical Yield
actual yield = (98.Oolo[zSSg) = zel g = dctual yield
percentageYield =
actual yield x 1OO= 64.80/oyield
theoretical yield
F - ORoANICCHEAAI5TRY
1 . Drowthe structureof eochof the followingcompounds:
c) 1,3-cyclohexodiene
o) 2,3-dinitrophenol
ot+
,\ro.
do,
LO^
v
b) 2,3,4-'trimethyloctone
C(lt
ct{. -cl+-cil+^ Ct\-CAz-Uza\rCHzctl. ctlt
d) 4,4-dimethyl-2-Pentene
Qilt
CIlrcH=cHq cilt
crL
Review- Answers- Page23 of ?7
Chemistry11- FinolExomination
l) proponone
e) phenylethene
o
Cl+= CFl.
tl
cSg-c - ct{3
m) l-proPoxybutone
2-phenylProPone
f)
.,*rd-c*=
9) 1,3-dibromonoPhtholene
C|{tcH,-CA;o-Ct{>-CilL-CLz'C[!3
n) ethyl chloroethonoote
o
,,
6v
@."
ct
o)
h) 1,2-ethonediol
iodobenzene
L
,\
ts'
fil'otl o l+
i)
| /)t
\Z
p) 3-methYl-3-Pentonol
proponol
orI
o
I
ll
CA;cHv-l-''Hr-CIls
cHr
cA r-cr|;cr+
i)
erlr_Ctl,
lftre-D-
q ) 4-methylphenol
ocid
pentonoic
O
t,
CLl;ctl;cl;il\;C-o t't
ot
^.
Y
ctfg
r) 2-omino-3-phenylproponol o
k) propylethonoote
I
-CHr{A>CL
Cllr-C-c:
c*-" cu tt
A
\?
Review- Answers- Poge24 of ?7
Chemistry11- FinolExominotion
IU,
w) nitromethone
s) 3-methyl-3-pentonol
cr,
iJ0 z- ct\3
c l{t cHr-L -c{r- ctl,
r
e
l
o|+
t)
x) 3-ethyl-5-methyl-2-naPhthol
ethylmethylether
w:L*.
CHr'tfit o - c{t
cfit
y) 1,3,5-cyclohexonetriol
u) butyl butanoote
0u
0r l
n
I
CAr-1t
; C$; Go cttgctlr-cl{; t! t
L}-
noAlotl
v) 3-ethylhexonoicocid
o
z) propyl2,3-dimethylPentonoote
-t-ou
CtJ;cU;t[r-ctt-crt]
-t-o- cL-ut>-ctt3
Ctlr-c+,fA
f
ft
-
CH,CH'
cAz
I
Ctlr
Z. Write o boloncedeguotionfor eachof the reoctionsbelow. Identify the type of reoctionond nomethe
products.
&,
o) pentone+ bromine)
Ctls- CvI;CII;Cil;Ctl. t k - - >
g^,t"r1.S"*ta"
t- VNno
f-r#*-,,
nx n
b) 2,3-dichloro-?-penlene+ woter )
cilo-9= t -cllz-Ltlt
cr cl
"*r'cLr-clr-c{>-cd'
+ tl"-c) ->
l
cHr- ?- i- ctt>-cHt
Ct
cl
?,?-d:Nwo 4 "*+"J
ad/i lsn- (xn
c) butcnoicocid + ethonol)
o
o
+ eLr4lll-cil'-t
Ctle-tl;Crl;'i-oU' ct\-CH>-oH
(sler ili ealrtm
H6.
+ hgd"+ogt'^brorv'tolt
oA
t'f
t
I
rrYr'l
-o -cLt'c4t + HrO
{Vlrf arlw,oo'fu t
Review- Answers- Poge25 of 27
Chemistry11- FinolExominotion
palor
d) benzene+ iodine)
V
Yxva
*wbg\L(r^
)
oxYsen
e) *t[f';ti
,
+ il-!
@t
+ hgd-'ro6t'*;o/i'J"
I- iodofun#
-)
tT-
to QO. + 4 H"o
w bon d,irr; oh r aa/t'r'
A
/,',Y n-r
S/-*/
+
AOvy'hv{firn
lL
O L
rKn
2-nophthol+ oxygan)
f)
((-'rHco)
z AryoP -, -zzoz- >
46P,l,.sha*-
| x r'r
+ fluorine)
9) 2,5,6-trimethyl-3-heptyne
c- 7H- Cl'l-cHt r
c1t,-c+-c=
'
it+t c{t
br+t
aold;lsaw,rF*
ZD eO>
+
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Review- Answers- Poge26 of 27
Chemistry11- FinolExominotion
3. Name the following compounds.
a)
3-methyl-5-nitro-3-hexene
b)
1,3-dichlorocyclohexane
c)
1,3-diamino-5-nitrobenzene
d)
3,6-dibromo-2-naphthol
e)
diethyl ether or ethoxyethane
f)
1,4-cyclohexadiene
g)
2,3-dimethylbutanoic acid
h)
2,3,4-pentanetriol
i)
2-methyl-3-pentanone
j)
3-ethyl-2-methylpentanal
k)
methyl propanoate
l)
4,5-dimethyl-7-propyl-1-naphthol
m)
2,4-hexadiyne
n)
1,2-butadiene
o)
pentyl 3,3-dimethylbutanoate
p)
1,3,5-cyclohexanetriol
q)
1-amino-2-bromo-4-nitrocyclopentane
r)
2,3-dimethylpentanedioic acid
s)
butyl 2-methylpropanoate
t)
phenyl-3,4-diaminopentanoate
Chemistry 11 – Final Examination Review – Answers – Page 27 of 27