Name_____________________________
Honors Chemistry
FINAL EXAM REVIEW
FINAL EXAM DATES – THURSDAY, JUNE 5, 2014 and TUESDAY, JUNE 10, 2014.
The final exam is worth 100 points and is multiple choice format. Part 1 will assess Chapters 1 – 6;
Part 2 will assess Chapters 7 - 13.
FINAL REVIEW PRACTICE
DIRECTIONS: Complete on a separate sheet of paper. Show all work.
CHAPTER 1 AND 2 – INTRODUCTION TO CHEMISTRY AND MEASUREMENT
1. Is the following set of data accurate, precise, both, or neither if the actual value is 23.4 mL?
20.4 mL
20.3 mL
20.1 mL
20.5 mL
The measurements are precise but not accurate.
2. Indicate the number of significant figures in the following:
a. 1040 3
b. 600 1
c. 12.40
4
d. 0.00230
3
e. 0.82 2
3. Solve each expression using correct number of significant figures and units.
a. 24.4g + 110g
130 g
b. 145.30 mL – 12.8 mL
c. 9.8 g / .123 L
80. g/L
d. 1.4500 cm x 2.55 cm x 1.5 cm 5.5 cm3
132.5 mL
4. Solve the following using dimensional analysis (factor label method).
a. Convert 23.5 yd3 to ft3. 635 ft3
b. A pound of coffee beans yields 50 cups (this number is obtained by counting) of coffee.
How many milliliters of coffee can be obtained from 1.00 g of coffee beans? (1 cup = 240mL)
26.4 mL/g
c. Convert 23.50 m/s2 to km/min2. 84.60 km/min2
5. Convert the following to the indicated temperatures
a. -34 ̊C to ̊F -29 F
b. 213 K to ̊C -60. C
6. Explain how a Bunsen burner works. Gas and air are controlled by knobs on the burner. The gas
and air mix resulting in a smokeless flame that produces very high temperatures.
7. What is the most practical and accurate tool used to measure volume in a laboratory?
Graduated Cylinder
8. Explain how you would set up a laboratory apparatus for
a. heating a crucible of magnesium – ring stand with an iron ring supporting a clay triangle
b. heating a beaker of copper II oxide and water – ring stand with an iron ring supporting a
wire gauze. The beaker will be placed on the gauze and supported by a second iron ring.
9. A student takes an object with an accepted mass of 200.00 grams and masses it on his own
balance. He records the mass of the object as 196.5 g. What is his percent error? 1.75%
10. What value does each of the following metric unit represent?
a. centi .01 or 1 x 10-2
b. nano .000000001 or 1 x 10-9 c. kilo 1000 or 1 x 10 3
d. milli .001 or 1 x 10-3
e. micro .000001 or 1 x 10-6
CHAPTER 3 – MATTER AND ENERGY
1. Solve the following density problems using significant figures.
a. What is the density (g/mL) of a block of metal 6cm x 10cm x 100cm that weighs 6820 dg?
What would be the volume (in L) of 2.8 kg of the metal? 24.5 L = 20 L
b. A block of metal has a mass of 42.3 g. It is placed in a graduated cylinder containing 56.0
cm3 of water. If the final volume is 75.1 mL, what is the density of the metal block? 2.21 g/mL
2. How would you separate…..
a. iron filings and pepper magnet
evaporate the water
b. salt and sand dissolve the salt and filter the sand then
c. methanol and water distillation
3. What is the specific heat of water? 4.184 J/g C
4.
Which requires more heat to warm from 22.0 ̊C to 85.0̊C if they have exactly the same mass, a
sample of aluminum (c = .902 J/ g ̊C) or a sample of water. Explain. It takes more energy to heat
one gram of water than 1 gram of aluminum because water has a higher specific heat.
5. A .0234 kg sample of copper cools from 53 ̊C to - 9.4 ̊C. What is the change in its energy if the
specific heat of copper is .387 J/g ̊C? - 570 J
6. Distinguish between a physical and chemical change. A physical change is a change that results
in no change to the components of the substance . A chemical change is a change that results in a
change to the components of the substance. The identity of the substance will change due to a
physical change.
7. Classify the following as either a physical or chemical change:
a. spoiling of milk
C
c. burning a piece of paper C
b. softening glass to bend it into a new shape P
d. rusting of a nail C
8. Identify the following as a pure substance, heterogeneous mixture, or homogeneous mixture.
a. copper
Pure substance - element
b. sweetened tea Homogeneous mixture - solution
c. calcium carbonate Pure substance - compound
d. sandy water – Heterogeneous mixture - suspension
9. What is the difference between an exothermic reaction and an endothermic reaction? Relate each
reaction type to enthalpy. Endothermic reactions absorb energy and the products have more energy
than the reactants. The enthalpy change for an endothermic reaction is positive. Exothermic
reactions release energy and the products will have less energy than the reactants. Exothermic
reactions have a negative enthalpy change.
10. Using the potential energy diagram above, answer the following questions.
a. How much energy do the reactants have? 200 kJ
b. How much energy is needed to get the reaction started? 250 kJ
c. What type of reaction is this, endothermic or exotherimic? endothermic
CHAPTER 4 – CHEMICAL FOUNDATIONS
1. What is the atomic mass of silver? What is silver’s mass number? 107.87 amu ; 108
2. Determine the symbol for the atom, ion, and or isotope that fits the description below.
a. 16 p, 18 e, 16 n S-2
b. 26 p, 26 e, 30 n Fe
c. 13 p, 10 e, 11 n Al+3
d. 35 p, 35 e, 45 n Br
3. Determine the identity of an element with a mass number of 75, that contains 42 neutrons. As
4. Determine the number of protons, neutrons, and electrons in the following:
a. Ar – 40
18p; 22n; 18 e
d. Co 27p; 32n; 27e
b. F-1 9p; 10n; 10e
e.
56 Fe
c.
29
Al
13p; 16n; 13e
26p; 30n; 26e
26
5. Distinguish between an alpha and beta particle.
Alpha decay is the emission of a helium
nucleus (2 protons and 2 neutrons = alpha particle). Beta decay is the splitting of a neutron into a
proton and an electron = beta particle. The proton is added to the nucleus of the atom, changing the
identity of the atom.
6. If 36 grams of water are produced and 4 grams of hydrogen are used in its production, how many
grams of oxygen is needed? What law is represented here? 32.0 grams of O2 = Law of the
Conversation mass
7. What is the difference between the mass number for Carbon 14 and carbon’s average atomic
mass of 12.011 amu? Mass number gives the total number of protons and neutrons for one isotope
of carbon. The average atomic mass is the calculated sum of all of the isotopes of carbon and their
% abundance in nature.
8. Calculate the atomic mass of lithium if one isotope has a mass of 6.0151 amu and a percent
abundance of 7.59% and a second isotope has a mass of 7.0600 amu and a percent abundance of
92.41%. 6.98 amu
9. Describe the debate between Democritus and Aristotle in terms of the atom. Democritus – small
pieces cannot be broken into smaller pieces forever – sooner or later, a single particle so small that
cannot be broken down into anything further will remain. He calls this end particle atomos. He states
that all matter consists of a collection of atoms and if there was space between the atoms, it contains
nothing. Aristotle – he believes that it makes more sense to believe that everything can be broken
into smaller and smaller pieces forever. He questions that if matter is composed of particles and
empty space, what holds the particles together? Due to Aristotle’s fame and popularity, he wins the
debate.
10. List the five postulates to Dalton’s theory. Explain two reasons why it is no longer accepted.
1. All matter consists of tiny particles.
2. Atoms of an element are unchangeable and indestructible.
3. Elements are characterized by the mass of their elements.
– All atoms of the same element have the same mass.
– Atoms of different elements have different masses.
4. When elements react, their atoms combine in simple, whole number ratios.
– Different atoms combine in simple whole number ratios to form compounds.
5. When elements react, their atoms sometimes combine in more than one simple, whole number
ratio.
11. Compare and contrast Rutherford and Thomson’s experiments and atomic models.
Thomson – experiments using cathode ray tubes and discovers the electron. His model is referred to
as the Plum Pudding Model where atoms are sphere shaped and composed of a space of positive
charge. The negative electrons are found within this area of positive charge.
Rutherford – experiments using sheets of gold foil and alpha particles and discovers that the alpha
particles are deflected back differently than what the plum pudding model would predict. His model is
referred to as the nuclear model where the electrons are found in the empty space and the positive
charge or protons are found in the center of the atom in what he refers to as the nucleus.
12. What significance does each subatomic particle have in regards to an element’s properties?
Protons – used to identify an individual atom
Neutrons – effect the average atomic mass of an atom
Electrons – determine an atom’s chemical behavior
13. Radioactive strontium (Sr - 90) has a half life of 29 years. If you had a 10.0 gram sample, how
many years will it take you to get 2.5 grams of sample?
10.0 grams  5.0 grams = 29 years; 5.0 grams  2.5 grams = 29 years TOTAL = 58 years
14. What element undergoes alpha decay to form lead-208? Explain.
When a radioactive element undergoes alpha decay, it emits an alpha particle that is 2 protons and 2
neutrons (Helium nucleus). The element is Po – 212.
CHAPTER 5 – MODERN ATOMIC THEORY/QUANTUM MECHANICS
1. What is the wavelength (in m) of a photon of light containing 3.19 x 10 -13 J of energy?
6.23 x 10 -13 m
2. How much energy in Joules will a photon of light with a wavelength of 345.0 nm possess?
5.762 x 10-19 J
3. What is the maximum number of electrons that can occupy the following?
a. 3d
10
b. 3rd energy level 18
c. 4s 2
d. one orbital 2
e. 4th energy level 32
4. Draw the orbital diagrams for the following elements:
a. argon
b. copper
c. nitrogen
5. Write the electron configuration for the following elements/ions:
a. calcium
1s22s22p63s23p64s2
c. molybdenum
b. arsenic
1s22s22p63s23p64s23d104p3
1s22s22p63s23p64s23d104p6 5s14d5
d. magnesium ion 1s22s22p6
e. sulfur ion
1s22s22p63s23p6
6. Write the noble gas configuration for the following elements.
a. Tin
[Kr] 5s24d105p2
b. Gold [Xe] 6s1 4f14 5d10
c. Chromium [Ar] 4s13d5
7. Answer the following questions about the electron configuration of aluminum:
a. What does the 3 in 3s2 mean? Third energy level
b. What does the 2 in 3s2 mean? Two electrons in sublevel
c. How many valence electrons does this atom have? 3
d. How many electrons will an atom of aluminum (lose or gain) to form an ion? Lose 3
e. What orbitals will aluminum lose its electrons from? 3s and 3p
8. What is needed for an electron to move from one energy level to a higher energy level?
A quanta of energy must be absorbed
9. What can be determined from an atom’s atomic emission spectra? The
wavelength/frequency/energy of light emitted by an atom when electrons return to ground state.
10. What happens when an electron falls back to its ground state? It emits light of a specific
wavelength/frequency/energy.
CHAPTER 6 – THE PERIODIC TABLE
1. Write the symbol of the element that fits each description:
a. alkali metal in period 3
Sodium
b. alkaline earth metal with 5 energy levels
Strontium
c. the only metalloid chalcogen (oxygen group) Tellurium (Polonium is acceptable too)
d. halogen with four energy levels Bromine
2. How many valence electrons does each of the following possess? What ion does each form?
a. calcium 2; Ca+2
b. silver 1; Ag+1
c. phosphorus 5; P-3
d. lead 2/4; Pb+2/Pb+4
e. carbon 4; C-4
3. Which is smaller?
a. Ca or Ca+2
b. P or O
c. N or F
d. Sulfur ion or Aluminum ion
4. Does each of the following trends increase or decrease down a family or across (left to right) a
period?
Down a Family
Across a Period
Electronegativity
Decreases
Increases
Ionization Energy
Decreases
increases
Atomic Radius
Increases
decreases
5. Define electron affinity. Which elements have a negative value for electron affinity? Why?
Electron affinity is the energy change that occurs when a neutral atom acquires an electron. The
more negative the value, the greater the amount of energy that is released. Elements that readily
gain electrons like the halogens will have the most negative values because little energy is required to
force an atom to acquire the electron.
Metals such as the alkaline earth metals will have positive values. Metals wish to lose electrons, not
gain them. In order to add electrons to a metal, energy must be added forcing the addition of the
electron.
6. Draw the Lewis dot diagrams for the following elements.
a.
b.
c.
7. Why are the alkali metals one of the most reactive families of elements? They have low ionization
energy.
8. Which groups make up the representative or main group elements? 1, 2, 13, 14, 15, 16, 17, and
18
9. Give three physical properties that differ between metals and nonmetals. Metals will lose
electrons; are good conductors of electricity and heat; have luster; exist as solids; high boiling points
and melting points. Nonmetals will gain electrons; are poor conductors of electricity and heat, lack
luster, are normally found as gases or solid; are brittle; low boiling points and melting points.
10. What is a metalloid? List each of them. Elements that behave as both a metal and a nonmetal.
Boron, Silicon, Germanium, Arsenic, Antimony,Tellurium, Polonium, and Astatine
CHAPTER 7 – CHEMICAL BONDING
1. What is the difference between an ionic and a covalent bond?
An ionic bond is a bond commonly formed between a metal and a nonmetal. The electronegativity
difference is normally greater than 1.8 causing a transfer of electrons from one atom to the other. A
covalent bond is a bond commonly formed between two nonmetals. The electronegativity difference
is less than 1.8 causing the electrons to be shared (and not transferred) from one atom to the other.
2. Would the following pairs most likely form an ionic or molecular compound?
a. Mg and Cl - ionic
b.
b. I and F - molecular
c. P and Cl - molecular
d. Sn and O - ionic e. Ag and S - ionic
3. Draw a Lewis dot diagram for each of the following:
a.
b.
c.
d.
e.
4. Using electronegativity values, write the ranges for a polar covalent, non polar covalent, and ionic
bonds. What is the significance of these electronegativity differences?
1.8 and greater – ionic; .4 – 1.7 polar covalent; 0 - .3 nonpolar covalent – the electronegativity
difference determines whether or not the electrons are able to be transferred or shared.
5. Describe the three types of intermolecular forces. Give one example of each.
Intermolecular forces are weak attractive forces between MOLECULES in a molecular compound.

Dispersion force – weak force between nonpolar molecules. Any diatomic molecule like
F2
 Dipole – dipole force – force between oppositely charged ends of two polar molecules.
The more polar the molecule, the stronger the force. SCl2
 Hydrogen bond – especially strong attraction between the hydrogen end and a fluorine,
nitrogen, or oxygen atom on the other dipole. Water
6. Draw the Lewis structures for each of the following compounds:
i. Identify the total number of valence electrons for each.
ii. What is the molecular shape (linear, bent, trigonal pyramidal, trigonal planar,
tetrahedral)?
iii. What is the bond angle for each shape?
iv. Polar or nonpolar covalent bonds?
v. Polar or nonpolar molecule?
a. N2 - 10; linear; 180; nonpolar; nonpolar
b. O3 – 18; Bent; less than 120; nonpolar, nonpolar
RESONANCE STRUCTURES
c. H2S – 8; Bent; less than 109.5; polar, polar
d. NCl3 – 26; trigonal pyramidal; less than 109.5; slightly polar, polar
e. SiF4 – 32; tetrahedral; 109.5; polar, nonpolar
CHAPTER 8 – NOMENCLATURE
1. Name each of the following:
a. Ba(CN)2
m. HNO3 (aq)
b.
CoBr2
n. Al(OH)3
c.
P2O5 (STOCK)
o. HBrO4 (aq)
d.
NaH
p. KMnO4
e.
MgI2
q. Hg3(PO3)2 (CLASSIC)
f.
Ca3(AsO4)
r. Na2O2ME
g.
HCl
s. LiC2H3O2
h.
Sn(SO4)2
t. FeCl3
i.
SO2
u. CCl4
j.
BaSO3
v. AgNO3
k.
Zn(ClO4)2
w. N2O3 (STOCK)
l.
Be(BrO)2
x. Pb3N2 (STOCK)
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
(CLASSIC)
Barium Cyanide
Cobalt (II) Bromide
Phosphorus (V) Oxide
Sodium Hydride
Magnesium Iodide
Calcium Arsenate
Hydrogen Chloride
Tin (IV) Sulfate
Sulfur Dioxide
Barium Sulfite
Zinc Perchlorate
Beryllium Hypobromite
m.
n.
o.
p.
q.
r.
s.
t.
u.
v.
w.
x.
∙ 6 H2O (CLASSIC)
Nitric Acid
Aluminum Hydroxide
Perbromic Acid
Potassium Permanganate
Mercuric Phosphite
Sodium Peroxide
Lithium Acetate
Ferric Chloride Hexahydrate
Carbon Tetrachloride
Silver Nitrate
Nitrogen (III) Oxide
Lead (II) Nitride
Write the formula for each of the following:
a.
Diboron Hexahydride
m. Mercuric Sulfate
b.
Ammonium Chloride
n.
Auric Bromide
c.
Iron (III) Nitrate
o.
Cadmium Oxide
d.
Phosphoric Acid
p.
Calcium Hypochlorite
e.
Titanium (IV) Bromide
q.
Lead (II) Phosphate
f.
Cuprous Phosphide
r.
Magnesium Sulfate Heptahydrate
g.
Selenium (VI) Fluoride
s.
Silicon Dioxide
h.
Sodium Cyanide
t.
Sulfuric Acid
i.
Hydroselenic Acid
u.
Aluminum Chlorate
j.
Phosphorus Triiodide
v.
Magnesium Sulfide
k.
Barium Chloride Dihydrate
w.
Plumbous Acetate
l.
Potassium Hydroxide
x.
Manganese (III) Sulfite
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
B2H6
NH4Cl
Fe(NO3)3
H3PO4
TiBr4
Cu3P
SeF6
NaCN
H2Se
PI3
BaCl2 · 2H2O
KOH
m.
n.
o.
p.
q.
r.
s.
t.
u.
v.
w.
x.
HgSO4
AuBr3
CdO
Ca(ClO)2
Pb3(PO4)2
MgSO4 · 7H2O
SiO2
H2SO4
Al(ClO3)3
MgS
Pb(C2H3O2)2
Mn2(SO3)3
CHAPTER 9 – THE MOLE
1. What is the molar mass of each of the following?
a. Al(C2H3O2)3?
204.13 g/mol
b. Ammonium chloride
53.50 g/mol
c. Potassium bromate 167.00 g/mol
2. I know that 8.401 x 1024 atoms of a certain element has a mass of 391.8 g.
What element is this?
Silicon
3. How many Sulfide ions are present in 12.24 mg of potassium sulfide? 6.683 x 1019 ions
4. How many carbon atoms are in 1.00 mole of dicarbon hexahydride? 1.20 x 1024 carbon atoms
5. How many grams are 2.34 x 1023 molecules of C6H10OS2? 63.1 grams
6. How many molecules can be found in a sample of .847 grams of carbon dioxide?
1.16x 1022 molecules
7. Determine the percent composition of all of the elements found in the following compounds:
a. Sodium chlorate NaClO3 45.10% O; 33.30% Cl; 21.60% Na
b. Cuprous sulfide Cu2S 20.15% S; 79.15% Cu
8. What is the difference between an empirical and molecular formula? The empirical formula shows
the most simplified ratio of the elements in a compound. The molecular formula shows the actual
ratio of the elements in the compound.
9. Determine the empirical formula for a compound using the following information:
a. 58.00% Rb; 9.50% N; 32.50% O
RbNO3
b. 49.62% C; 10.82% H; 39.56% O
C5H13O
10. What is the molecular formula of a compound with an empirical formula of CHOCl (chlorine) and
a molecular weight of 129 grams? C2H2O2Cl2
11. What are the 7 diatomic elements? Br2, I2, N2, Cl2, H2, O2, F2
12. Determine the molecular formula for a compound containing 7.70% carbon, 1.30% hydrogen,
and 81.90% iodine. Its molar mass is 281 g/mol. C2H4I2
13. If you take 20.8 g of hydrated aluminum oxide and heat it to remove all of the water, what is the
formula of the original hydrate if the mass of the remaining solid after heating is 15.4 g? Al2O3 · 2H2O
14. 1.00 mol of Al2S3 consists of how many total mol of ions? 5.00 mol (2.00 mol Al and 3.00 mol S)
CHAPTER 10 – CHEMICAL REACTIONS
1. For each of the following:
a. Indicate the type of the reaction
b. Predict the products by name
c. Write the equation for the complete reaction. If the reaction does not occur, write NR.
d. For DD reactions include the states
e. Balance the reaction only if it occurs
a. silver nitrate + cupric chloride  DD (precipitate) – silver chloride and cupric nitrate
2AgNO3(aq) + CuCl2(aq)  2AgCl(s) + Cu(NO3)2(aq)
b. copper (II) + aluminum sulfate  NO REACTION
c. sodium nitrite + ammonium sulfate  NO REACTION
d. iodic acid  D4 - diiodine pentoxide and water
2HIO3  I2O5 + H2O
e. hydrochloric acid +barium hydroxide DD (neutralization)-water and barium chloride
2HCl(aq) + Ba(OH)2(aq)  BaCl2(aq) + H2O(l)
f. sodium + chromium (II) chlorate  SD – sodium chlorate and chromium
2Na + Cr(ClO3)2  2 NaClO3
+ Cr
g. cupric oxide  D2 – copper (II) and oxygen
2CuO  2Cu + O2
h. C4H8 + oxygen (complete) C – carbon dioxide and water
C4H8 + 6O2  4CO2 + 4H2O
i.
dibromine pentoxide + water  S4 – bromic acid
Br2O5 + H2O  2HBrO3
j.
iron (II) oxide + carbon dioxide  S5 – iron (II) carbonate
FeO + CO2  FeCO3
k. chlorous acid + magnesium hydroxide 
DD (neutralization) – water and magnesium chlorite
2HClO2(aq) + Mg(OH)2(s)  Mg(ClO)2(aq) + 2H2O(l)
l.
aluminum carbonate  D5 – aluminum oxide and carbon dioxide
Al2(CO3)3  Al2O3
+
3CO2
m. Sodium sulfide + hydrochloric acid 
DD (gaseous) – sodium chloride and hydrogen sulfide gas
Na2S + 2HCl(aq)  2NaCl (aq) + H2S(g)
3
n. hydrobromic acid + calcium carbonate 
DD (gaseous) -calcium bromide, carbon dioxide, and water (carbonic acid breaks down)
HBr(aq) + CaCO3(s)
 CaBr2(aq) + CO2(g) + H2O(l)
2. What is the difference between oxidation and reduction? Oxidation is a loss of electrons by one
reactant. Reduction is a gain of electrons by one reactant.
3. Write the net ionic equations for the reaction of
a. hydrochloric acid + barium hydroxide H+(aq) + OH-(aq)  H2O(l)
b. chlorous acid + sodium hydroxide HClO2(aq) + OH-(aq)  H2O(l) + ClO2 - (aq)
c. strontium sulfide + hydrochloric acid  2H+(aq) + S -(aq)  H2S(g)
d. silver nitrate + cupric chloride  Ag+(aq) + Cl -(aq)  AgCl (s)
4. What are the four main signs that a chemical reaction has occurred? Evolution of a gas;
temperature change; precipitate; color change
5. Balance the following redox reaction in a basic environment:
MnO4- + SO2  SO4 -2 + Mn+2
4OH - + 5SO2 + 2MnO4 - 2Mn+2 + 5SO4-2 + 2H2O
CHAPTER 11 – STOICHIOMETRY
1. If you react 25.4 g of bromine with an excess of sodium iodide and you produce 2320 mg of
iodine, what is the percent yield of iodine? 5.75%
2. How many grams of potassium sulfate are produced if 10.0 g of sulfuric acid and 7000.0 mg of
potassium hydroxide are mixed together? What is the limiting reactant? LR = KOH; 10.872 g
3. How many moles of chloric acid are required to react completely with 2.05 g of sodium carbonate?
.0387 mol
4. The reusable booster rockets of the U.S. space shuttle employ a mixture of aluminum and
ammonium perchlorate for fuel. The balanced equation is
3Al(s) + 3NH4ClO4  Al2O3(s) + 3NO(g) + 6H2O(l) NH3(g)
What mass of ammonium perchlorate should be used in a fuel mixture for every 1.00 kg of
aluminum used? 4.36 kg
5. How many grams of magnesium chloride are produced if 420 cm 3 of chlorine gas reacts with 14.8
g of Magnesium at STP? 1.8 g MgCl2
6. If the following reaction produces 6840 J of energy, how many grams of CrO 3 is used? 127 g
CrO3 + H2O  H2CrO4 + 5.4 kJ
7. How much heat is absorbed/released when 193 g of ammonium bromide react according to the
following equation?
NH3 + HBr  NH4Br
ΔH = 188.32 kJ
371 kJ released (enthalpy is reversed because reaction is reversed
CHAPTER 12 – SOLUTIONS
1. If 960 g of sodium hydroxide are used to prepare 16000 mL solution. What is the molarity of the
solution? 1.5 M
2. If we add .15 L of a 6.02M HCl solution and some water to a beaker, creating a 2.42 M HCl
solution, what is the volume of the new solution? .37L
3. A 2.34 M solution contains 10. g of acetic acid. What is the volume of the solution in mL? 71 mL
4. What mass of sodium oxide is consumed if it reacts with 12.3 mL of a 3.45 M HCl solution? 1.32g
5. How many grams of .67 M KClO3 are needed to prepare 1.00 Liters of solution? 82 g
6. Describe how 100 ml of a .10 M solution of sodium hydroxide would be made. Fill a volumetric
flask approximate 2/3 full of water. Add .4 grams of sodium hydroxide to the flask. Swirl the flask.
Add remaining 1/3 of water to the solution until meniscus is at the line on the flask. Mix.
7. Using the pH scale, define an acid and a base. An acid is a substance with a pH value of 0 – 6.99
and a base is a substance with a pH value of 7.01 - 14
8. Calculate the pH when [H+] = 3.9 x 10 -5.
4.41 Calculate the [H+] when the pH value is 9.57.
2.7 x 10-10
9. Which solution is a poor conductor of electricity? 1M NaCl; 1MHC2H3O2 (weak acids do
not/partially ionize) ; 1M NaOH; 1M HCl
10. What volume of 3.1M H2SO4 will neutralize 250 mL of .30 M Al(OH)3? 360 mL
CHAPTER 13 – GASES
1. Describe the conditions that allow a liquid to boil. The vapor pressure is equal to the atmospheric
pressure allowing the liquid to evaporate.
2. Compare the particles within a solid, liquid, and a gas. The particles of a solid are close together
and fixed in place, liquid the particles can slide past each other, and a gas the particles are spread
out to move freely and randomly.
3. Relate the intermolecular forces within a solid, liquid, and gas. The intermolecular forces within
solid water are much stronger than liquid water. The intermolecular forces within water vapor are
minimal.
4. Identify and define the different phase changes of matter.
a. Sublimation – solid to gas
b. Deposition – gas to solid
c. Condensation – gas to liquid
d. Evaporation – liquid to gas
e. Freezing – liquid to solid
f. Melting – solid to liquid
5. What is vapor pressure? The force exerted by a vapor that has been released by a liquid.
6. A balloon at 22 °C holds 2.00L of oxygen. If the gas in the balloon is heated to 90° C, what
volume will the balloon be? 2.5 L
7. What is the temperature of .70 moles of a gas that occupies 470 mL at a pressure of 150 kPa?
12 K
8. 500.0 mL of O2 gas is collected over water. Atmospheric pressure is 101.0 kPa. The temperature
of the water is 22 °C. What is the pressure of the gas? You will need a chart from your notes to
complete this problem! 737.8 mmHg
9. A gas is at 1.33 atm of pressure and a volume of 682 mL. What will the pressure be if the volume
is reduced to 0.419 L? 2.16 atm
10. Nitrogen gas is being held in a 14.3 m3 tank at a temperature of 62oC. What will the volume be
when the temperature drops to 24oC? 12.6 m3
11. A gas storage tank is a 1.72 atm and 35oC. What temperature is the gas at if the pressure
increases to 2.00 atm? 358 K
12. A gas with a volume of 1.00 L is at 135oC and 844 mm Hg. What is the volume if the conditions
change to 14o C and 748 mm Hg? .794 L
13. Calculate the mass of 162 L of chlorine gas, measured at STP.
512 g
14. Find the pressure in mm Hg produced by 2.35 g of carbon dioxide in a 5.00 L flask at 18 oC.
194 mmHg
15. How many grams of carbon monoxide must be placed into a 40.0 L tank to develop a pressure of
965 mm Hg at 23oC? 58.5 g
16. At what Celsius temperature will argon have a density of 10.3 g/L and a pressure of 6.43 atm?
30.6 °C