Download PRACTICE EXAM for FALL 2013 FINAL EXAM (Unit 6 + review) 1

PRACTICE EXAM for FALL 2013 FINAL EXAM (Unit 6 + review)
a
b
c
d
e
1. Describe what each box represents (e.g., drawing (a) represents a pure monatomic element)
2. Draw circle pictures to illustrate each of these. Use for A and
for B
a. 4 molecules of A2
c. 2 molecules of A2 and 2 molecules of B2
b. 4 molecules of A2B
d. 2 molecules of A2B and 2 molecules of AB
3. a.
b.
c.
b.
List the 3 subatomic particles we studied, including their charge and mass in amu.
Describe the arrangement of these particles in the atom.
Compare the size of the nucleus to the size of the atom [hint: it involves football]
Describe how mass is distributed in an atom.
4. a. What are isotopes? How will the symbols of isotopes be alike? How will they differ?
b. Define atomic number and mass number. How is each shown on the atomic symbol?
5. List the number of protons, neutrons, and electrons in an atom of
a. 40Ar
b. 52Cr
c. 27Al
d. 3H
e. 63Cu
f.
6. Identify each atom and write its symbol
a. 38 protons, 52 neutrons, 38 electrons
b. 6 protons, 8 neutrons, 6 electrons
207
Pb
c. 17 protons, 20 neutrons, 17 electrons
d. 53 protons, 78 neutrons, 53 electrons
7. Identify the family to which each element belongs:
a. Li
b. Br
c. H
d. Xe
e. Sr
8. Give a symbol for
a. the element in period 4 group 4A
b. the 3rd transition metal in period 4
c. the halogen in period 2
9. Which is larger (has a greater radius)?
a. P or Ar
b. K or Rb
c. C or N
b. the alkali metal in period 6
c. a semimetal in period 5
f. the element in period 3 group 2A
d. Cl1– or Cl
e. Ba2+ or Ba
f. S or O
10. a. Define ionization energy and describe its general trend across a period and down a family.
b. Which has a higher ionization energy?
i. Si or Cl
ii. Sr or Ca
c. Which loses an electron more easily?
i. Ne or Kr
ii. Ge or Br
11. How many particles are in a mole? What is this number called?
12. Do these mole conversions. Setups with units and appropriate sigfigs required.
a. How many atoms are in 0.560 moles of Pb?
b. How many molecules are in 3.75 mol NH3?
13. List the number of valence electrons in each atom, and draw its dot structure:
a. Ca
b. Si
c. O
14. Rewrite this paragraph, using words or phrases from the table at right to fill in
the missing parts of this description of chemical bonding. Some may be used
more than once.
When forming a chemical bond, atoms will _____ valence electrons in order to
achieve the same number of electrons as ____. That number of electrons is ___,
except for small atoms like H that try to achieve just ___ electrons in their valence
level. This bonding theory is called the _____ rule. This rule predicts that in general
metals tend to _____ valence electrons and nonmetals tend to _____ or _____
valence electrons when forming chemical bonds. Thus, an atom of Al will _____ __
electrons when forming a bond, and an atom of Cl will _____ __ electrons when
forming a bond.
eight
gain
lose
lose, gain, or share
octet
share
the nearest noble gas
one
two
three
15. Based on the rule in question #14, predict the formula for the compound that forms between the
indicated atoms. For example, O and H form H2).
a. C and H
b. K and O
c. S and H
d. Cs and Cl
16. Fill in this description of ions:
A metals tends to ______ valence electrons to form a _____ ion, which is called a __________ .
( +, – )
(gain, lose)
(cation, anion)
A nonmetal tends to ______ valence electrons to form a _____ ion, which is called a __________ .
( +, – )
(gain, lose)
(cation, anion)
A cation is __________ than its parent atoms, and an anion is __________ than its parent atom.
(larger, smaller)
(larger, smaller)
17. a. Draw Lewis dot structures to show the formation of each of these ionic crystals:
i. NaBr
ii. MgO
iii. K2S
18. Draw Lewis dot structures for each of these molecules. Include the number of valence electrons.
a. CCl4
b. COS c. H2S d. NF3 e. CH2O f. ClO21– g. NO21– h. PH3
19. Indicate the number of lone pairs on the central atom in each structure in #18.
20. a. Compare/contrast the three major types of bonding: covalent, ionic, and metallic
b. Identify the type of bonding in each of these substances:
i. NH3 ii. K2O iii. Fe iv. C6H12O6 v. H2O vi. CuCl2 vii. Au
c. Match each sketch to the label that best describes the type of substance:
Atomic gas Molecular (covalent) substance Metallic substance Atomic crystal Ionic crystal
a
b
c
d
e
21. Do these mole conversions. Setups with units and appropriate sigfigs required.
a. What is the mass of 0.0058 mol NaOH?
b. How many moles are in 25.84 g CO2?
c. Calculate the number of moles in 2.41 g AgNO3
d. Calculate the mass of 0.135 moles Na2C2O4
22. For the reaction HCl (aq) + NaHCO3 (s)  NaCl (aq) + CO2 (g) + H2O (l)
identify the reactants and the products and describe the state of each substance in the equation.
23. Balance these equations
a. Na (s) + H2O (l)  NaOH (aq) + H2 (g)
b. C5H12 (l) + O2 (g)  CO2 (g) + H2O (g)
c. Ag (s) + H2S (g) + O2 (g  Ag2S (s) + H2O (l)
d. AgNO3 (aq) + Cu (s)  Cu(NO3)2 (aq) + Ag (s)
e. C2H5OH (l) + O2 (g)  CO2 (g) + H2O (g)
24. Do these stoichiometry problems:
a. How many moles of BrF3 form when 5.98 g of Br2 react?
Br2 (l) + 3 F2 (g)  2 BrF3 (g)
b. How many grams of (NH4)2Cr2O7 are consumed when 25 g of Cr2O3 form?
(NH4)2Cr2O7 (s)  4 H2O (l) + N2 (g) + Cr2O3 (s)
c. How many grams of Cu form when 1.28 g Al react?
2 Al (s) + 3 CuCl2 (aq)  2 AlCl3 (aq) + 3 Cu (s)
25. a. Define the terms “exothermic” and “endothermic.”
b. Identify each reaction as exothermic or endothermic and do the stoichiometry problem.
i. 2 H2 + O2  2 H2O + 572 kJ
How much energy (in kJ) is released when 3.6 g H2O form?
ii. 88 kJ + PCl5  PCl3 + Cl2
How many moles of PCl5 can be decomposed with 460 kJ of energy?
26. a. Explain each of these in molecular terms:
i. gas pressure
ii. gas temperature
iii. why gas pressure increases when the container size is decreased
iv. why gas pressure increases when the temperature is increased
v. why a gas always fill its entire container
b. O2 gas is collected by displacing water from a collection tube. The total pressure in the tube is 762
mm Hg; the water vapor pressure in the experiment is 24 mm Hg. What is the pressure of the O2?
c. Two identical containers are at the same temperature and pressure; one contains Ne gas and the
other, Xe gas. Compare the kinetic energies, masses, velocities, and number of particles in the
two containers.
d. If containers of the following gases (all at the same temperature) were opened at the same time,
which would reach you first, and why?
a. Cl2 (71 g/mol) b. CO2 (44 g/mol) c. CH4 (16 g/mol) d. all of them at the same time
27. Gas law problems:
a. A balloon filled with 635 mL of oxygen gas at 23 °C is placed in a freezer, where it cools to –10
°C. What is the volume of the cold balloon? The pressure and amount of gas remain constant.
b. A small gas cylinder contains 3.22 L of argon at 11.7 atm pressure. What is the volume of the gas
at 1.05 atm? Assume temperature and amount of gas remain constant.
c. 325 mL of air at room pressure (765 mm Hg) are compressed with a piston to a volume of 42 mL.
What is the pressure of the compressed air? The temperature and amount of air remain constant.
d. A He weather balloon has a volume of 18 L at the earth's surface, where the temperature is 25 °C
and the pressure is 750 mm Hg. What is the volume of the balloon when it reaches an altitude
where the temperature is –22 °C and the pressure is 375 mm Hg? The amount of He is constant.
e. What is the volume of 0.085 mol hydrogen gas at 0.97 atm and 21 °C?
f. A canister of gas at 2.8 atm and 75 °C is cooled to 0 °C. If the volume and amount of gas are
constant, what is the pressure in the cold bottle?
g. A sample of nitrogen gas in a 275 mL container at 0.82 atm pressure and 26 °C is transferred to
larger container with volume 750 mL . The temperature in the larger container is now 17 °C.
What is the pressure of the gas in the larger container? The amount of nitrogen is constant.
h. How many moles of chlorine gas are in a 345 L tank at 7800 mm Hg and 27 °C?
28. a. What does STP stand for? What are the STP conditions? Include units!
b. What is the volume of one mole of any gas at STP?
c. How many moles are in 35.73 mL of H2 gas at STP?
d. How many mL of O2 gas at STP contain 0.00500 mol O2?
29. Do these stoichiometry problems.
a. 2 C4H10 (g) + 13 O2 (g)  8 CO2 (g) + 10 H2O (l)
How many L of O2 gas are needed to react with 3.00 L of C4H10? Both gases are at STP.
b. Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g)
How many moles of HCl are needed to produce 25.0 mL H2 gas at STP?
c. C2H5OH (l) + 3 O2 (g)  2 CO2 (g) + 3 H2O (l)
How many L of CO2 gas at STP are produced when 12.85 g C2H5OH burn?
30. Ne and Xe gases condense to form liquids but not at the same temperature: Ne condenses at 24 K
and Xe at 161 K. Explain why either gas condenses at all, and why at different temperatures?
31. Draw a heating curve that shows the temperature changes as C6H5OH (phenol) is heated from 20 °C
to 200 °C. Phenol melts at 43 °C and boils at 182 °C. Label the curve to show solid, liquid, gas,
melting, freezing, and boiling. What is the freezing point of C6H5OH? When the phenol is melting or
boiling, is it absorbing energy, releasing energy, or staying at the same energy?
32. Define each of these terms [hint: if you don’t know, try the glossary in the back of your text]:
a. melting b. freezing c. condensation d. sublimation e. boiling f. evaporation
33. a. What is the relationship between molar mass and boiling point for molecules of similar structure?
b. Predict which would have a higher boiling point: Cl2 or I2? CH3OH or C2H5OH?
34. Vapor pressure is an equilibrium property. What two processes are equal in this equilibrium? Why
does vapor pressure always increase at higher temperature?
35. At right is a graph of the vapor pressure at
various temperatures for 5 substances.
a. What is the normal boiling point of each
substance? How do you determine that?
b. Compare the strength of the attractions
between (a) molecules to the strength of the
attractions between (c) molecules, and explain
your reasoning.
c. Would the boiling point of (b) be different if
you measured it at Tioga Pass (elevation 3000
m above sea level)? If yes, would it be higher
or lower? Explain.
d. We recognize that a liquid is boiling because
bubbles are forming throughout the liquid.
What is inside the bubbles in (d) when it is
boiling?
e. Which of the substances on the graph has the
strongest attractions between its molecules?
Which has the weakest attractions between its
molecules? How did you decide?
36. a.
b.
c.
d.
e.
Define the terms solute, solvent, and solution.
Define molarity.
Calculate the concentration (in mol/L) of a solution that contains 7.18 g ZnCl2 in 125 mL solution.
How many mL of 0.0496 M HCl are needed to provide 0.00100 mol HCl?
How many moles of NaCl are in 35.0 mL of 0.148 M NaCl solution?
37. Do these stoichiometry problems:
a. How many mL of 0.100 M KOH are needed to produce 0.200 g Zn(OH)2?
ZnCl2 (aq) + 2 KOH (aq)  Zn(OH)2 (s) + 2 KCl (aq)
b. How many grams of FeCl2 can react with 26.24 mL of 0.0214 M MnO41–?
5 FeCl2 (s) + MnO41– (aq) + 8 H1+ (aq)  5 Fe3+ (aq) + Mn2+ (aq) + 10 Cl1– (aq) + 4 H2O
c. 0.186 g of H2C2O4 react with 32.62 mL of NaOH solution. What is the concentration of the NaOH
solution, in mol/L?
H2C2O4 (s) + 2 NaOH (aq)  2 H2O + Na2C2O4 (aq)
ANSWERS to Fall 13 Final review
1. a. pure monatomic element
b. pure compound
d. mixture of two different elements & one compound
c. pure diatomic element
e. pure compound
2.
a. 4 molecules of A2
3. a.
b.
c.
d.
b. 4 molecules of A2B
c. 2 molecules of A2
and 2 molecules of B2
d. 2 molecules of A2B
and 2 molecules of AB
See Table 4.1, pg 106. Proton mass ≈ neutron mass = 1 amu; e– mass = 1/1840 = 0.00054 amu.
Protons & neutrons in tiny nucleus; electrons occupy rest of atom’s volume around nucleus
If atom = football stadium, nucleus = marble on 50-yard line
Essentially all the mass is in the nucleus
4. a. Isotopes = atoms w same number of protons but different numbers of neutrons; same chemical
symbol & same atomic number (lower left corner of symbol), but different mass numbers (upper
left corner). 36 Li and 37 Li are isotopes of lithium.
b. Atomic number = number of p+ (shown in lower left corner, often omitted because element symbol
is sufficient to identify atomic number). Mass number = number of particles in nucleus = p+ + n0
(shown in upper left corner).
5. a. 18 p+, 22 n0
6. a.
90
Sr
b.
b. 24 p+, 28 n0
14
C
7. a. alkali metal
8. a. Ge
9. a. P
b. V
c.
37
Cl
c. 13 p+, 14 n0
d.
b. halogen
c. F
b. Rb
e. 29 p+, 34 n0
f. 82 p+, 125 n0
I
c. no family
b (d). Cs
c. C
131
d. 1 p+, 2 n0
c (e). Sb or Te
d. Cl1–
e. Ba
d. noble gas
e. alkaline earth
f. Mg
f. S
10. a. Ionization energy = energy needed to remove outermost e– from an atom or ion. Generally
increases across a period (left to right) and decreases down a group/family.
b. i. Cl
ii. Ca
c. i. Kr
ii. Ge
11. 6.02 x 1023; Avogadro’s number
12. a. 3.37 x 1023 atoms
b. 2.26 x 1024 molecules
13. Ca 2 valence e–, Si 4 valence e–, O 6 valence e–.
Ca
Si
O
14. When forming a chemical bond, atoms will lose, gain, or share valence electrons in order to achieve
the same number of electrons as the nearest noble gas. That number of electrons is 8, except for
small atoms like H that try to achieve just 2 electrons in their valence level. This bonding theory is
called the octet rule. This rule predicts that in general metals tend to lose valence electrons and
nonmetals tend to gain or share valence electrons when forming chemical bonds. Thus, an atom of
Al will lose 3 electrons when forming a bond, and an atom of Cl will gain 1 electron when forming a
bond.
15. a. CH4
b. K2O
c. H2S
d. CsCl [look at the dot structures to see how they achieve octet)
16. A metals tends to lose valence electrons to form a + ion, which is called a cation. A nonmetal tends
to gain valence electrons to form a – ion, which is called an anion. A cation is smaller than its parent
atom, and an anion is larger than its parent atom.
17. check in class for K2S
1+
Na
Br
Na
18. a. CCl4 (32 e–)
(check in class for
1–
b. COS (16 e–)
c. H2S (8 e–)
f. ClO21– g. NO21– h. PH3 )
2–
2+
Mg
O
Mg
Br
O
d. NF3 (26 e–)
e. CH2O (12 e–)
Cl
Cl
C
H
Cl
Cl
19. a. none
S
b. none
C
c. two
O
d. one
F
S
F
F
H
e. none
N
f. two
O
H
g. one
C
H
h. one
20 a. Covalent bonding: nonmetals share valence electrons w other nonmetals so each atom is
surrounded by 8 (or 2, for H). Atoms stay w their valence e– to form molecules. Molecule has no
charge, so it does not conduct electricity in any state.
Ionic bonding: metals transfer e– to nonmetals; metals form positive ions (cations) and nonmetals
form negative ions (anions), which stick together b/c of attraction of opposite charges. Ions collect
into crystals but do not stay w own electrons to form molecules. Ions can conduct electricity if free to
move (liquid or dissolved in water).
Metallic bonding: metal atoms release valence electrons  lattice of + ions in “sea” of mobile e– that
holds metal cations together. e– free to move, so metals conduct in both solid and liquid state.
b. i. covalent
ii. ionic iii. metallic iv. covalent
v. covalent
vi. ionic
vii. metallic
c.
a
ionic crystal
21. a. 0.23 g
b
atomic gas
b. 0.5871 mol
c
molecular substance
c. 0.0142 mol
d
metallic substance
d. 18.1 g
22. a. The reactants are HCl and NaHCO3. The products are NaCl, CO2 and H2O.
e
atomic crystal
b. NaHCO3 is solid, H2O is liquid, CO2 is a gas, and HCl and NaCl are aqueous (dissolved in water).
23. a. 2, 2, 2, 1
b. 1, 8, 5, 6
24. a. 0.0748 mol BrF3
c. 4, 2, 1, 2, 2
b. 41 g (NH4)2Cr2O7
d. 2, 1, 1, 2
e. 1, 3, 2, 3
c. 4.52 g Cu
25. a. Exothermic reactions release energy; endothermic reactions absorb energy
b. i. exothermic, 57 kJ ii. endothermic, 5.2 mol
26. a. i. Gas pressure is caused by collisions of gas molecules with container wall.
ii. Gas temperature is proportional to molecular kinetic energy.
iii. Molecules have less space to bounce around, so they collide with container walls more often.
iv. Molecules are moving faster, so they strike container walls more often and with greater force.
v. Molecules are in constant random motion, so they eventually distribute throughout entire
container, no matter how large it is.
b. 738 mm Hg
c. Both have same kinetic energy (same temperature); Xe atom is heavier, Ne atom is faster; same
number of particles in both containers (Avogadro’s hypothesis).
d. All would reach you eventually (because the molecules are in constant random motion), but CH4
would reach you first because it has the smallest molar mass and therefore the fastest molecules,
given that they are all at the same temperature (same kinetic energy).
27. a. 564 mL b. 35.9 L c. 5920 mm Hg or 7.79 atm
d. 30. L
e. 2.1 L
f. 2.2 atm
g. 0.29 atm
h. 144 mol
28. a. STP is Standard Temperature and Pressure: 0 °C or 273 K and 1 atm or 760 mm Hg
b. one mole gas == 22.4 L at STP
c. 0.001595 mol
d. 112 mL
29. a. 19.5 L
b. 0.00223 mol
c. 12.50 L
30. A gas condenses because its molecules are attracted to each other. At low enough kinetic energy
(temperature), attractions overcome molecular motion and the molecules clump together to form a
liquid. This temperature is lower for Ne than for Xe, suggesting that the attractions between Ne atoms
are weaker than the attractions between Xe atoms.
temperature (°C)
31. The curve should look like the heating curve at right,
with the melting/freezing plateau (BC) at 43 °C and
the boiling plateau (DE) at 182 °C. C6H5OH freezes
at 43 °C (same temp as melting point). Melting and
boiling are endothermic (absorbing energy, even
though temperature is not changing).
F
D
E
B
C
A
32. a. change from solid to liquid
b. change from liquid to solid
energy added 
c. change from gas or vapor to liquid
d. change from solid to gas without passing through the liquid state
e. change from liquid to gas throughout the liquid
f. change from liquid to vapor at the surface of the liquid (below the boiling point)
33. a. For molecules of similar structure, boiling point increases as molar mass increases.
b. higher b.p. = I2 and C2H5OH (the heavier of each pair)
34. Evaporation & condensation. Vapor pressure increases because at higher temperature, more
molecules have enough energy to escape from the liquid.
35. a. Normal boiling point = temperature at which vapor pressure equals 760 mm Hg or 1 atm (normal
or standard atmospheric pressure). Substance (a) boils at 35 °C, (b) at 80 °C, (c) at 100 °C, (d) at
110 °C, and (e) cannot be determined from this graph.
b. Interactions between (a) molecules are weaker than those between (c) molecules; (a) molecules
evaporate more easily and boil at a lower temperature => they are less “sticky” than (c) molecules.
c. All boiling points would be lower at Tioga Pass, because atmospheric pressure is lower at high
elevation. Boiling begins when vapor pressure matches atmospheric pressure; if atmospheric
pressure is lower, boiling point is lower.
d. (d) vapor
e. (e) molecules have the strongest attractions (highest boiling point), and (a) molecules have the
weakest attractions (lowest boiling point).
36. a. Solute is what dissolves, solvent is what it dissolves in, and the solution is the mixture.
mol solute
b. Molarity is the concentration of the solution; M
L solution
c. 0.421 M d. 20.2 mL e. 0.00518 mol
37. a. 40.2 mL
b. 0.3559 g
c. 0.127 M
Extra material
4. Fill in the blanks to complete these paragraphs. In your response, write the entire paragraph, not just
the missing words.
The Greek philosopher _____ was among the first to suggest the existence of atoms, but the modern idea of the atom
began with ______’s atomic theory in the year 1808. A few years later, in ____, Amadeo Avogadro stated his hypothesis
that _____________ to explain why gases combine in simple, whole number ratios. His hypothesis later became the basis
for the mole, the counting unit of chemistry. One mole contains _____ particles.
Almost a century later, in _____, the English physicist __________ discovered the electron and showed that the atom was
made of even smaller particles. He thought the atom was like a blueberry muffin, but in the year ____, __________ did
an experiment in which he directed alpha particles at a thin sheet of gold foil. He expected that _______________, but to
his astonishment ____________! From this experiment, he proposed the nuclear model of the atom, in which
_______________.
9. Rubidium is 72.17% 85Rb (84.9117 amu) and 27.83% 87Rb (86.9178 amu). Calculate the average
atomic mass of Rb. Don’t forget units and watch your sigfigs.
d. For each container, the difference between the mercury levels is given below the container. Draw
in the mercury in the U-tube. If the atmospheric pressure is 760 mm Hg, what is the pressure of
the gas in each of these containers:
left 35 mm above right
levels same
left 40 mm below right
247 kJ + CH4 + CO2  2 CO + 2 H2
How much energy is consumed when 54 g CO are produced?
iv. N2H4 + 2 H2O2  N2 + 4 H2O + 818 kJ
How many grams of H2O2 are needed to produce 3500 kJ?
iii.
4. Stoichiometry problems
a. How many mL of CO2 gas at STP can react with 0.18 g LiOH in this reaction:
2 LiOH (s) + CO2 (g)  Li2CO3 (s) + H2O (l)
b. A car airbag inflates when sodium azide decomposes explosively to produce nitrogen gas:
2 NaN3 (s)  2 Na (l) + 3 N2 (g)
How many grams of NaN3 must decompose to produce 41 L of N2 at STP?
c. How many L of CO2 gas at STP will form when 50.0 g of propane (C3H8) burn?
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
h.
Br2 (l) + 2 KI (aq)  2 KBr (aq) + I2 (s)
How many moles of KBr form when 0.40 moles of Br2 react?
9. A student finds that 0.182 g H2C2O4 react with 35.61 mL of NaOH in this reaction:
H2C2O4 (s) + 2 NaOH (aq)  2 H2O + Na2C2O4 (aq)
What is the concentration of the NaOH solution, in mol/L?
10. How many grams of Mg3(PO4)2 would form if 25.0 mL of 0.120 M MgCl2 reacted with excess
Na3PO4 in this reaction:
3 MgCl2 (aq) + 2 Na3PO4 (aq)  Mg3(PO4)2 (s) + 6 NaCl (aq)
f.
Na2CO3 (aq) + CaCl2 (aq)  2 NaCl (aq) + CaCO3 (s)
How many grams of NaCl will form when 25.0 mL of 0.138 M Na2CO3 react?
d. How many grams of sodium hydroxide are in 23.85 mL of 0.118 M NaOH solution?
e. How many mL of 0.200 M K2CO3 solution are needed to provide 4.15 g of K2CO3?