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Chemistry 1
Final Exam Unit Reviews
You may use The Chemistry Reference Tables, your notes and/or textbook to answer the following review
questions. Write the name of the unit and place all work and answers to that section on your own paper.
Unit 1: Matter and Change
1. Give one example of each of the following.
a) element
b) compound
c) heterogeneous mixture
d) homogeneous mixture
2. What are the differences between an element and a compound?
3. How do chemical properties differ from physical properties?
4. What is the law of conservation of mass?
5. What are the differences between a homogeneous mixture and a heterogeneous mixture?
6. List the three main states of matter and describe their basic properties.
7. Classify each of the following as either an element, compound, homogeneous mixture (solution) or
heterogeneous mixture.
a) copper (II) sulfate
b) Kool Aid
c) wood
d) plastic
e) lined paper
f) sulfur
8. Label each process as a physical or chemical change.
a) perfume evaporating on your skin
c) wood rotting
e) autumn leaves changing color
g) melting copper metal
i) mixing sugar in water
b) butter melting
d) charcoal heating a grill
f) a hot glass cracking when placed in cold water
h) burning sugar
j) digesting food
9. Use The Chemistry Reference Tables to solve the following density related problems.
a) Calculate the mass of sucrose in 15.0 ml. (1 mL = 1 cm3).
b) Calculate the volume of nitrogen that has a mass of 23.5 g.
c) A 22.5 g sample of a metal has a volume of 1.274 cm3. Calculate the density of the metal.
d) What is the mass of a sample of sodium chloride if its volume is 134 cm3?
e) Calculate the volume of 145 g of ethanol.
10. Label each process as a physical or chemical change.
a) Moth balls gradually vaporize in a closet.
b) Hydrofluoric acid attacks glass (used to etch glassware).
c) A chef making a sauce with brandy is able to burn off the alcohol from the brandy, leaving just the
brandy flavoring.
d) Chlorine gas liquefies at -35 °C under normal pressure.
e) Hydrogen burns in chlorine gas.
11. Describe the process of chromatography.
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Unit 2: Atoms, Moles and Nuclear Chemistry
1. How many protons, neutrons, and electrons are there in 40K?
2. How many liters are in 64.0 grams of CO2?
3. How many grams of CH4 are there in 1.23 x 1024 molecules?
4. What are the molar masses of the following molecules?
a) PbBr2
b) AgNO3
c) copper (II) sulfate
d) ammonium carbonate
5. How many grams are there in 18.6 moles of nitrogen dioxide (NO2)?
6. What did Thomson discover during his cathode ray experiment?
7. How many protons, neutrons and electrons does an atom of 261Rf have?
8. How many atoms are there in 3.09 x 10-4 g of Fe?
9. How many moles are in 35 g of CuSO4•5H2O?
10. Write the symbols for an alpha particle, a beta particle, a gamma ray and a positron.
11. Strontium-90 is present in nuclear fallout. Because it is in the same family as calcium, it can be found in
milk and later, bones. If you start with 1.94 x 1017 atoms of strontium-90, how many atoms will remain after
140.5 years? The half-life of strontium-90 is 28.1 years.
12. The half-life of uranium-227 is 1.3 minutes. How long will it take for a sample of uranium-227 to decay to
1/32 its original mass?
13. Complete the table below.
Atomic
symbol
B
Atomic
number
Protons
Neutrons
6
Electrons
11
Atomic
mass
24
31
37
39
29
89
35
43
100
207
Pb
102
89
Mo
70
225
53
81
100
206
159
No
Yb
261
172
106
159
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14. Using your knowledge of nuclear chemistry, write the equations for the following processes.
a) The alpha decay of iridium-174
b) The beta decay of platinum-199
c) Positron emission from sulfur-31
d) Krypton-76 undergoes electron capture
15. If the half-life for the radioactive decay of zirconium-84 is 26 minutes and you start with a 175 gram
sample, how much will be left over after 104 minutes?
16. Why is it difficult to make a fusion reaction occur?
17. State the experimental findings of the following scientists.
a) Democritus
b) Ernest Rutherford
d) Robert Millikan
e) James Chadwick
c) Niels Bohr
f) John Dalton
18. Define isotopes.
19. How many liters are in 3.20 x 1024 molecules of CO?
20. The atomic mass of an element is the weighted average mass of the isotopes of an element. What 3 things
are used to calculate the (average) atomic mass of an element?
21. Distinguish between the penetrating ability of alpha, beta and gamma radiation.
Unit 3: Electrons and the Periodic Table
1. What is the electron configuration of cadmium (Cd)?
2. What is the noble gas (shorthand) electron configuration of copper (Cu)?
3. What group on the periodic table comprises the alkali metals?
4. What do elements in the same group have in common with one another?
5. Why does electronegativity decrease as you move down a column in the periodic table?
6. Answer the following True or False. If false, change the statement to be correct.
a) Electronegativity indicates the relative ability of an atom to attract electrons in a chemical bond.
b) Sodium sulfate will react with magnesium metal to form magnesium sulfate and sodium.
c) A rubidium atom is larger than a strontium atom.
d) The first ionization energy of rubidium is larger than the first ionization energy of strontium.
e) The octet rule is a useful rule of thumb but is not always valid.
7. How many valence electrons do the following elements have?
a) Rn
b) Ra
c) Xe
d) K
e) Hg
f) As
8. Name an element for each of the following.
a) a halogen
b) an actinide metal
c) a member of Group 15
d) a lanthanide metal
e) an element that has 3 valence electrons
f) 3 transition metals
g) a noble gas
h) 2 metalloids
i) an alkali metal
j) a chalcogen (Group 16)
k) an alkaline earth metal
l) an inner transition metal
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9. Draw the electron (Lewis) dot structures for the following elements.
a) Cs
b) I
c) Sn
d) Po
e) Ca
f) Zn
10. What is the electron sea theory (model)?
11. Name an element for each of the following:
a) the element in Group 2 and Period 5
c) the element in Group 18 and Period 4
e) an element with a -3 oxidation number
g) an element with a +2 oxidation number
i) an element that has a valence of s2p1
k) an element that has the valence of s2p4
b) the element whose e- configuration ends with 5p1
d) the element whose e- configuration ends with 7s2
f) the elements whose e- configuration ends with 4d4
h) an element with a +1 oxidation number
j) the element whose e- configuration ends with 3d1
l) the element whose e- configuration ends with 5d8
12. Write down the periodic and group trends for:
a) ionization energy
b) atomic radius
d) activity (or reactivity)
e) ionic radius
c) electronegativity
13. Within each of the following groups, state which is larger in size (radius).
a) Li and Be
b) F and Ne
c) Rb and Rb+1
-2
e) O and O
f) Br and I and Cl
d) Au and Ag
14. Within each of the following groups, state which has the largest ionization energy.
a) Li and Be
b) F and Ne
c) C and O
d) B and Al
e) K and Na
f) N and Na
15. Within each of the following groups, state which has the greatest electronegativity.
a) Li and Be
b) F and Fr
c) F and Ne
d) C and O
e) Hg and Zr
f) Xe and I
g) Cs and K
16. Identify the element based on following electron configuration.
a) [Kr] 5s2 4d3
b) [Ne] 3s2 3p5
c) [Xe] 6s1
17. Write noble gas configurations for:
a) Mn
b) Te
c) Ra
18. Rank the following from least to most active (within each set).
a) Mg, Al, Ba
b) Br, Ar, I, Ne
c) Si, P, S, O
19. What is the relationship in size between neutral metallic atoms and the ions they form AND neutral nonmetallic atoms and the ions they form?
20. Are frequency and energy inversely or directly proportional?
21. If each orbital can hold a maximum of two electrons, how many electrons can each of the
following sublevels hold?
a) 2s
b) 5p
c) 4f
d) 3d
e) 4d
22. What is the shape of an s orbital?
23. How many s orbitals can there be in an energy level?
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24. How many electrons can occupy an f orbital?
25. What is the shape of a p orbital?
26. How many p orbitals can there be in an energy level?
27. Which is the lowest energy level that can have an s orbital?
28. Which is the lowest energy level that can have a d orbital?
29. Are frequency and wavelength inversely or directly proportional?
30. Why does it take more energy to remove an electron from Al+ than from Al?
31. Use the Bohr Model for the Hydrogen Atom in the Chemistry Reference Tables to answer the following.
a) An electron falls from energy level 3 to energy level 2. What color of visible light is emitted?
b) An electron falls from energy level 6 to energy level 3. What is the wavelength of the light emitted?
32. As an electron falls from a higher energy level to a lower one, energy is (absorbed, released).
33. Define the following terms.
a) orbital
b) atomic number
c) excited state
d) mass number
34. Wave–particle duality postulates that all matter, including an electron, exhibits both wave and particle
properties. The ______ effect supports the particle theory.
Unit 4: Chemical Bonding
1. For the molecule diboron tetrahydride (B2H4), draw the Lewis structure and determine the molecular shape.
2. What does it mean when we refer to a molecule as being “polar”?
3. Name the intermolecular force (ionic, hydrogen bonding, dipole-dipole, London dispersion) most important
for each of the following compounds.
a) NH3
b) BCl3
c) HF
d) H2O
e) CH4
f) HCl
g) C2H5OH
h) O2
4. Draw a Lewis structure and determine the molecular shape for
a) the cyanide ion (CN-)
b) nitrogen trichloride (NCl3)
5. Identify the differences between the properties of ionic and covalent compounds.
6. Identify the type of bond (ionic, polar covalent, nonpolar covalent or metallic) in each of the following.
a) Li2O
b) CCl4
c) MgF2
d) CsF
e) N2
f) SO2
g) Br2
h) NI3
i) Cu-Cu
7. Draw the dash-dot diagram for each of the following covalently bonded molecules.
a) NF3
b) SiI4
c) P2
d) CH2I2
e) HCN
g) OCl2
h) HAsI2
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f) NBr3
8. Determine the shape AND polarity of the molecules in question 7.
9. Why is the bond angle between the hydrogen atoms in water smaller than the bond angle between the
hydrogen atoms in ammonia (NH3)?
10. Determine the polarity of each of the following substances. Remember, you have to identify the bond type
first in order to determine the substance’s polarity. Also determine the intermolecular force that might exist
between each molecule.
a) LiF
b) K2O
c) F2
d) CH4
e) NCl3
f) Cr-Cr
11. Which compounds in question 10
a) will conduct electricity as a solid?
c) will dissolve in water?
b) will conduct electricity in an aqueous solution?
d) will have a high melting point?
12. How many valence electrons must an atom have in its outer energy level in order to be considered stable?
13. Provide 3 characteristics/properties of metals.
14. A cation has a ______ charge, while an anion has a _______ charge.
15. Two elements which have chemically bonded have an electronegativity difference of 1.32. Is this bond
ionic, polar covalent or nonpolar covalent?
16. What is the relationship between bond energy and length of single, double, and triple bonds?
17. Which 2 atoms can form multiple bonds in order to form network covalent solids?
Unit 5: Chemical Formulas and Composition Stoichiometry
1. Write the formulas for the following compounds.
a) zinc chloride
b) manganese (IV) sulfate
d) chlorine gas
e) hydrochloric acid
c) sulfur hexafluoride
f) chloric acid
2. What is the percent composition by mass of Cl in CH2Cl2?
3. Name the following chemical compounds.
a) PbBr2
b) NH3
e) H2SO4
f) V(CO3)2
i) AgCN
j) N2O4
c) F2
g) P2O5
k) CH4
d) CaS
h) BeF2
l) HBr
4. A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its
empirical formula?
5. Write the chemical formulas for the following compounds.
a) ammonium nitrate
b) fluorine gas
d) iron (III) phosphate
e) nitric acid
g) dinitrogen tetrachloride
h) tin (IV) sulfide
c) boron trichloride
f) potassium carbonate
i) silver perchlorate
6. The molecular formula of the antifreeze ethylene glycol is C2H6O2. What is the empirical formula?
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7. A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What
is its empirical formula?
8. A compound is known to have an empirical formula of CH and a molar mass of 78.11 g/mol. What is its
molecular formula?
9. A well-known reagent in analytical chemistry, dimethylglyoxime, has the empirical formula C2H4NO. If its
molar mass is 116.1 g/mol, what is the molecular formula of the compound?
10. Determine the percent chromium in potassium dichromate (K2Cr2O7).
11. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. One of them is a
blue solid that comes apart, reversibly, in the gas phase. It contains 36.84% N. What is the empirical formula
of this oxide?
12. Write the chemical formulas for the following compounds.
a) ammonium sulfide
b) nickel (II) iodide
d) mercury (I) oxide
e) copper (II) bromide
g) aluminum sulfate
h) hydroiodic acid
j) iron (II) hydrogen carbonate
k) dichlorine monoxide
13. Name the following chemical compounds.
a) LiH
b) Be(OH)2
c) SF6
f) Sr(OH)2
g) KCN
h) P4S10
c) sodium nitrate
f) lead (II) chlorite
i) potassium nitrate
l) magnesium nitrate
d) Sr(HCO3)2
i) K2O
e) SnI4
j) H3PO4
Unit 6: Chemical Reactions and Stoichiometry
1. How many grams of carbon dioxide will be made when 100 grams of CH4 burn in an excess of oxygen gas?
CH4 + 2O2  CO2 + 2H2O
2. Write complete, balanced equations for each of the following.
a) When the butane (C4H10) in a gas camp stove is burned in oxygen, the products of the reaction are carbon
dioxide gas and water vapor.
b) When chlorine gas is combined with ethylene gas (C2H4) at high temperature and pressure,
tetrachloroethylene gas (C2Cl4) and hydrogen gas are formed.
3. Balance the following chemical equations.
a) ___NaN3(s)  ___Na(s) + ___N2(g)
b) ___H2(g) + ___N2(g)  ___NH3(g)
4. In the reaction 2CO + O2  2CO2, how many moles of carbon monoxide react with how many moles of
oxygen to make how many moles of carbon dioxide?
5. What mass of water is produced when 150. g of tungsten are also produced? WO3 + 3H2  W + 3H2O
6. How many moles of CrCl2 are produced from reacting 18 g of Zn in the following reaction?
Zn + 2CrCl3  2CrCl2 + ZnCl2
7. Identify the differences between synthesis and decomposition reactions.
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8. If 8.20 moles of sulfuric acid react, how many grams of sulfur dioxide are produced?
Cu + 2H2SO4  CuSO4 + SO2 + 2H2O
9. If 3.2 moles of hydrogen react, how many moles of oxygen are required to make water? 2H2 + O2  2H2O
10. Identify the differences between single and double displacement reactions.
11. What does each symbol mean?
a) 
b) (l)
c) 
d)


12. Balance the following equations.
a) __Mg + __HCl  __MgCl2 + __H2
c) __RbCl + __O2  __RbClO4
e) __Ca(OH)2  __CaO + __H2O
g) __HCl + __Ba(OH)2  __BaCl2 + __H2O
i) __Li + __H2O  __LiOH + __H2
e) (s)
b)
d)
f)
h)
__Sb +
__Sn +
__C4H10
__Fe +
f) (g)
g) (aq)
__Cl2  __SbCl3
__KOH  ___K2SnO2 + __H2
+ __O2  __CO2 + __H2O
__Cl2  __FeCl3
13. For each of the following, tell the type of reaction AND predict the products. Do not balance the equation.
a) Cu(s) + AgNO3(aq) 
b) Br2(g) + NaF 
c) H2SO4(aq) + KOH(aq) 
d) Ni + MgCl2 
e) CoO 
f) RbI + AgNO3 
g) Cs + F2 
h) Ti(OH)4 
i) Ca(NO3)2 + Al(OH)3 
j) C6H12 + O2 
14. How many moles of H2 are needed to completely react with two moles of nitrogen?
N2 + 3H2  2NH3
15. How does a burning splint react in each of the following: oxygen, hydrogen and carbon dioxide?
16. What are 4 signs that a chemical reaction has occurred?
Unit 7: Kinetic Theory and the Gas Laws
1. A gas mixture at STP includes nitrogen (0.781 atm), carbon dioxide (0.001 atm), argon (0.009 atm) and
oxygen. According to Dalton’s Law, what is the partial pressure of oxygen in atm if the total pressure is at STP.
2. A mixture of a gas contains 50.0 kPa of chlorine, 22.3 kPa of He and 43.7 kPa of bromine. What is the total
pressure of this deadly mixture? Is this at STP?
3. The pressure of 3.5 L of nitrous oxide anesthetic gas is changed from 760 mm Hg to 364 mm Hg. Assuming
the temperature remains constant, what will the resulting volume be?
4. If a sample of He gas occupies 12.1 L at 332 C, what will be its new volume at 47C, if the pressure
remains constant?
5. If a sample of carbon dioxide occupies 5.2 L at 80C and at 200 kPa, what will be its volume at STP?
6. Calculate the number of liters occupied at STP by 6.8 moles of Kr.
7. How many moles of fluorine gas occupy 8.2 L at a temperature of 350 K with a pressure of 1.5 atm?
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8. What pressure will be exerted by 1.45 moles of hydrogen gas at 25C if the volume is 2.5 mL?
9. What gas can travel faster – carbon dioxide or fluorine? Why?
10. Which of the following behaves most like an ideal gas: He, N2, CO2 or NH3?
11. What is the pressure of 15 L of gas that was originally 75C and 250 kPa and was changed to 50C and
2.1 L?
12. Convert 745 mm Hg to atmospheres.
13. Convert 3.45 atm to kilopascals.
14. What is the partial pressure of carbon dioxide in a container that holds 5.00 moles of carbon dioxide,
3.00 moles of nitrogen, and 1.00 mole of hydrogen and has a total pressure of 1.05 atm?
15. The kinetic molecular theory of ideal gases states that:
a) All matter is composed of tiny, discrete particles (molecules or atoms).
b) Ideal gases consist of small particles (molecules or atoms) that are far apart in comparison to their own
size. The molecules of a gas are very __________ compared to the distances between them.
c) These particles are considered to be dimensionless points which occupy zero volume. The volume of real
gas molecules is assumed to be __________ for most purposes. This above statement is NOT TRUE.
Real gas molecules do occupy volume and it does have an impact on the behavior of the gas. This
impact WILL BE IGNORED when discussing ideal gases.
d) These particles are in rapid, _________, constant straight line motion. This motion can be described by
well-defined and established laws of motion.
e) There are no _________ forces between gas molecules or between molecules and the sides of the
container with which they collide. In a real gas, there actually is attraction between the molecules of a
gas. Once again, this attraction WILL BE IGNORED when discussing ideal gases.
f) Molecules collide with one another and the sides of the container. Energy can be transferred in collisions
among molecules. Energy is ________ in these collisions, although one molecule may gain energy at
the expense of the other.
16) 5.00 L of a gas is known to contain 0.965 mol. If the amount of gas is increased to 1.80 mol, what new
volume will result (at an unchanged temperature and pressure)?
17) 10.0 L of a gas is found to exert 97.0 kPa at 25.0°C. What would be the required temperature (in Celsius) to
change the pressure to standard pressure?
18. A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases
are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container?
19. A container with two gases, helium and argon, is 30.0% by volume helium. Calculate the partial pressure of
helium and argon if the total pressure inside the container is 4.00 atm.
20. If 60.0 L of nitrogen is collected over water at 40.0°C when the atmospheric pressure is 760.0 mm Hg, what
is the partial pressure of the nitrogen? The vapor pressure of water at 40.0°C is 55.3 mm Hg.
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21. Vapor pressure (decreases, increases) as temperature increases.
Unit 8: States of Matter, Phase Changes, Solutions and Solubility
1. Define the following.
a) soluble
b) insoluble
e) solvent
f) solute
i) immiscible
j) unsaturated
c) electrolyte
g) saturated
k) supersaturated
2. Which of the following are electrolytes?
a) LiCl
b) CH3OH
c) CH4
f) CaSO4
g) CO2
h) AlPO4
d) SiO2
i) Na2O
d) nonelectrolyte
h) miscible
e) Ba(NO3)2
j) KMnO4
3. For each substance in #2 that dissociates into electrolytes, write the number of ions formed.
4. What is the molar concentration of a solution where 4.00 moles of KCl are dissolved in 8000. mL of water?
5. For each of the questions (a) – (f), refer to the phase diagram for mysterious compound X.
a) What is the critical temperature of compound
X?
b) If you were to have a bottle containing
compound X in your closet, what phase
would it most likely be in?
c) At what temperature and pressure will all
three phases coexist?
d) If I have a bottle of compound X at a pressure
of 45 atm and temperature of 100oC, what
will happen if I raise the temperature to
400oC?
e) Why can’t compound X be boiled at a
temperature of 200oC?
f) If you wanted to, could you drink compound
X?
6. Calculate the molar concentration when 350. g of NaOH are dissolved in 2.5 L of water.
7. How many grams of MgCl2 were dissolved in 500. mL of water to make a 2.1 M solution?
8. You are given 10.0 mL of a 4.00 M solution. You want to dilute it to 1.50 M. What will the total volume of
the new solution be? How much water must be added to dilute the original solution?
9. You are given a stock 18 M H2SO4 solution. How many mL of the stock acid must be diluted to 2.0 M in
order to have a total volume of 200. mL?
10. How many grams of BaCO3(s) are formed when 20.0 mL of 0.200 M K2CO3 are added to BaCl2?
K2CO3 + BaCl2  2 KCl + BaCO3
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11. What is the maximum amount of KNO3 that can be
dissolved in 100 g of water at 55 C?
12. At what temperature do KCl and KClO3 have the same
saturation point?
13. What are the three gases in the solubility curve?
14. How many grams of NaCl can be dissolved in 200 g of
water at 60C?
15. 70 g of NH4Cl are dissolved in 100 g of water at 70C.
What is the best description of this solution: saturated,
unsaturated or supersaturated?
16. Indicate the solvent that will be best at dissolving the given solute in each of the following problems.
a) Solute: lithium acetate.
Solvents: carbon disulfide, CH2Cl2
b) Solute: boron trichloride.
Solvents: carbon tetrachloride, water.
c) Solute: phosphorus triiodide.
Solvents: ammonia, water.
17. What does “like dissolve like” mean?
18. Give an example of surface tension you’re familiar with and an example of a surfactant around your house.
19. Identify the precipitate formed when barium chloride (BaCl2) and sodium sulfate (Na2SO4) combine. Write
the net ionic equation for this reaction.
Unit 9: Acids and Bases
1. How many milliliters of 0.120 M HCl are required to completely neutralize 50.0 mL of 0.220 M KOH?
HCl + KOH  KCl + H2O
2. Identify the differences between the endpoint and equivalence point in titrations.
3. Describe and give 2 examples each of Arrhenius Acids and Arrhenius Bases.
4. How would you describe a Bronsted-Lowry Acid? A Bronsted-Lowry Base? Give 2 chemical formulas of
each of these.
5. Write the reaction when sulfuric acid (H2SO4) is mixed with water. Label the acid and base.
6. An Arrhenius acid reacts with an Arrhenius base to accomplish a “neutralization” reaction always yielding a
salt and water. (Double replacement) Complete the following neutralization reactions. Write in the products
you would expect to be formed and then balance the equations.
a) HCl + NaOH 
b) HNO3 + KOH 
c) Ca(OH)2 + H2SO4 
d) H3PO4 + Mg(OH)2 
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7. Give 2 examples of monoprotic acids.
8. Give 2 examples of diprotic and 1 triprotic acids. (These are polyprotic acids.)
9. If the hydrogen ion concentration [H+] is 4.7 x 10-5 mol/L, what is the solution’s pH?
10. Show the self-ionization reaction of water where 2 water molecules react together forming hydronium ion
and the hydroxide ion. Label the acid and base.
11. If the hydronium ion concentration [H3O+] is 3.4 x 10-9 M, what is the solution’s pH?
12. If the hydroxide concentration is 0.0000083 M, what is the pH?
13. If the hydroxide concentration of a solution is 3.5 x 10-2 M, what is the hydrogen ion concentration?
14. If the [H+] of a solution is 9.2 x 10-7 M, what is the pH?
15. What is the pH of a 0.0059 M NaOH solution?
16. What is the [H3O+] of a solution whose pH is 5.4?
17. What is the [OH-] of a solution that has a pH of 10.7?
18. In a titration, 17.4 mL of a 0.154 M solution of Ba(OH)2 is added to 20.0 mL of HCl solution in order to
neutralize the mixture. What is the molarity of the acid? Ba(OH)2 + 2HCl  BaCl2 + 2H2O
19. A 15.5 mL sample of 2.15 M KOH solution required 21.2 mL of acetic acid solution in a titration
experiment. Calculate the molarity of the acetic acid.
KOH + HC2H3O2  KC2H3O2 + H2O
20. By titration, 17.6 mL of aqueous sulfuric acid just neutralized 27.4 mL of 1.65 M LiOH solution. What
was the molarity of the acid? 2LiOH + H2SO4  Li2SO4 + 2H2O
21. Identify the following as acidic, basic or neutral in terms of the pH.
a) [H3O+] = 1 x 10-7 M
b) [H3O+] = 6.6 x 10-10 M
c) [H3O+] = [OH-]
-8
-4
d) [OH ] = 2.3 x 10 M
e) [OH ] = 7.9 x 10 M
f) pOH = 3.7
g) pH = 13.4
22. What volume of 0.35 M potassium hydroxide would be necessary to neutralize 55 mL of a 0.15 M solution
of nitric acid? KOH + HNO3  KNO3 + H2O
23. What volume of 0.11 M sulfurous acid would be necessary to neutralize 17.4 mL of a 0.15 M solution of
cesium hydroxide? H2SO3 + 2CsOH  Cs2SO3 + 2H2O
24. What determines whether or not an acid or base is considered strong?
Unit 10: Reaction Energy and Kinetics
1. Identify each phase change as being an endothermic change or an exothermic change.
a) melting
b) vaporizing
c) condensation
d) freezing
e) deposition
f) sublimation
g) fusion
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2. Determine the total amount of energy required for a 25.0 g sample of ice at – 20.0oC be heated to a
temperature of 20.0oC.
3. Define the following terms.
a) specific heat capacity
b) temperature
c) kinetic energy
d) heat
4. During an endothermic process, heat is (released, absorbed).
5. A substance has the following properties:
a. Hvaporization = 20 kJ/mol
b. Hfusion = 5 kJ/mol
d. normal mp = -15oC
e. Cp of solid = 3.0 J/goC
g. Cp of gas = 1.0 J/goC
a) Sketch a heating or warming curve for the substance starting at -50oC.
b) Sketch a phase diagram for the same substance.
c. normal bp = 75oC
f. Cp of liquid = 2.0 J/goC
6. What happens to the intermolecular distance of molecules during a phase change as heat is added?
7. What happens to the kinetic energy as the temperature increases?
8. When 12.29 g of finely divided brass (60% Cu, 40% Zn) at 95.0oC is quickly stirred into 40.00 g of water at
22.0oC in a calorimeter, the water temperature rises to 24.0oC. Find the specific heat of brass.
a) State what phases of matter are found in
the 5 labeled regions.
b) In which regions do phase changes occur?
c) During the heating process, which of the
five parts of the heating curve experience
a change in temperature?
9.
10. Calculate the heat necessary to raise 27.0 g of water from 10.0°C to 90.0°C.
11. How much energy is required to completely boil away 70.0 g of water at 100.0°C?
12. Sketch an exothermic reaction pathway graph.
a) What is the activation energy of the
un-catalyzed reaction?
b) What is the activation energy of the catalyzed
reaction?
c) What is the heat of reaction (ΔH)?
d) Is the reaction endothermic or exothermic?
13.
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14. Which state of matter contains the greatest entropy?
15. How much energy is required to melt 20.0 g of water at 0 °C?
16. Calculate the heat necessary to raise 2.50 g of aluminum from 120.0°C to 290.0°C.
17. How is the rate of a reaction dependent on each of the following?
a) surface area
b) temperature
c) concentration of reactants
d) gas pressure
18. A catalyst _____ the activation energy by providing an alternate ______ for the reaction to occur.
19. Determine whether the entropy increases, decreases or remains the same as the reaction proceeds.
a) CH3OH (l) → CH3OH (g)
b) N2O4 (g) → 2NO2 (g)
c) CO (g) + H2O (g) → CO2 (g) + H2 (g)
d) 2KClO3 (s) → 2KCl (s) + 3O2 (g)
e) 2NH3 (g) + H2SO4 (aq) → (NH4)2SO4 (aq)
20. Which member of the following pairs has the greater predicted amount of entropy?
a) CO2 (g) or CO2 (s)
b) PbS (s) or PbF2(s)
c) H2 (g) in a 1 L vessel or H2 (g) in a 2 L vessel
21. Which member of the following pairs has the lesser predicted amount of entropy?
a) Fe(s) at 25°C or Fe(s) at 100°C
b) H2O (l) or H2O (s)
c) FeCl2 (s) or FeCl3(s)
22. The entropy of the universe is (decreasing, increasing).
23. Consider the collision theory. Molecules must ______ in order to react, and they must collide in the correct
or appropriate _______ and with sufficient energy to equal or exceed the _______ energy.
24. For the reaction below, which change would cause the equilibrium to shift to the right?
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g)
a) Decrease the concentration of hydrosulfuric acid.
b) Increase the pressure on the system.
c) Increase the temperature of the system. The reaction is exothermic.
d) Increase the concentration of carbon disulfide.
e) Decrease the concentration of methane (CH4).
25. What would happen to the position of the equilibrium when the following changes are made to the
equilibrium system below?
2SO3(g) ↔ 2SO2(g) + O2(g)
a) Sulfur dioxide is added to the system.
b) Sulfur trioxide is removed from the system.
c) Oxygen is added to the system.
d) Decrease the temperature of the system if energy is absorbed.
26. Write the equilibrium expression for each of the following reactions.
a) 2HgO(s) ↔ Hg(l) + O2(g)
b) 4HCl(g) + O2(g) ↔ 2H2O(g) + 2Cl2(g)
c) 2H2O(g) + N2(g) ↔ 2H2(g) + 2NO(g)
d) SiO2(s) + 4HF(g) ↔ SiF4(g) + 2H2O(g)
e) CO(g) + H2(g) ↔ C(s) + H2O(g)
f) H2(g) + Cl2(g) ↔ 2HCl(g) + 49.7 kJ
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