Upload
trinhdieu
View
216
Download
0
Embed Size (px)
Citation preview
D:\bkup_170430\320\320_99_Tests & Quizzes\Mid-Year.Final\320.Final Exam\320_Final_Review_KEY.docx
KEY
CHEMISTRY
320 FINAL EXAM REVIEW
THE FINE PRINT This review is a ‘best attempt’ to help you review the year’s material for the upcoming final
examination in chemistry. We compiled this review to help you, not trick you. However, we make no
guarantee that it is entirely complete. Some material may have been inadvertently omitted; some
material may be included that, although covered during the year, is not on the final exam. You are
responsible for reviewing the entire year’s material for the test just as you are responsible for doing
your best on it.
Review. Ask questions. Study hard. Good Luck.
Chemistry 320 Final Exam Review p.1
CHAPTER 1: INTRODUCTION (1 - 6)
1. Define chemistry __________________________________________________________________
study of the composition, structure, and properties of matter and the changes it undergoes
2. Properties: chemical versus physical. Give an example of each of the following:
A. chemical properties: _____ involves change identity of the substance (e.g., wood burns; iron rusts)
B. physical properties: _____ does not involve change in identity of the substance (e.g., ice to water)
3. Change: chemical versus physical. Give an example of each of the following:
A. chemical change: ______________________________________________________ (q.v. above)
B. physical change: _______________________________________________________ (q.v. above)
4. Matter:
A. define & given an example: ___________ occupies space and has mass (e.g., rock paper, scissors)
5. Elements & symbols: Give the name & symbol for the elements with old symbols - i.e., names and
symbols not easily related. (See Table 1-2).
antimony (Sb) potassium (K)
copper (Cu) Silver (Ag)
Gold (Au) Sodium (Na
Iron (Fe) Tin (Sn
Lead (Pb) Tungsten (W)
Mercury (Hg)
6. Parts of a chemical equation: For the following equation Mg + 2HCl MgCl2 + H2:
A. identify the reactant(s): _______ Mg & HCl identify the product(s): _________ MgCl2 & H2
B. what is the ‘2’ in ‘2HCl’ called? coefficient what does it mean? ______number of molecules
C. what is the ‘2’ in ‘MgCl2’ called? _ subscript what does it mean? number of preceding atoms
Chemistry 320 Final Exam Review p.2
CHAPTER 2 – MEASUREMENTS (7-12)
7. Scientific method. Name the parts of the scientific method and briefly describe each.
observation “Discovery consists of seeing what everybody has seen, and thinking what
nobody has thought. ” 1
hypothesis testable statement (not a guess)
testing hypothesis e.g., experiment
theory broad generalization explaining body of facts or phenomena; model is an
explanation of phenomena or related events (e..g, Bohr model of the atom)
publishing results The job’s not done until all of the paperwork is completed and the lab
report handed in.
8. Observations. Define each of the following and give an example:
A. Quantitative: ___________________ observation that includes numbers (e.g., 96oC, 3.0 x10
8 m/s)
B. Qualitative: ____________________ observations that do not include numbers (e.g., tall, hot, fast)
9. SI Units & prefixes. Fill in the following tables for (A) quantities, units and abbreviations, (B)
prefixes.
A. Quantity Unit Abbreviation
length meter m
mass gram g
volume liter L
density gram/milliters g/mL
temperture Kelvin K
energy Joule J
B. Prefix Abbrev. Factor Meters (m) Liters (L) Grams (g)
kilo- k- 1000 1000 m = 1 km 1000 L = 1 kL 1000 g = 1 kg
centi- c- 0.01 1 m = 100 cm 1L = 100 cL 1 g = 100 cm
milli- m- 0.001 1 m = 1000 mm 1L = 1000 mL 1 g = 1000 mg
micro- - 10–6
1 m = 10–6
m 1 L = 106 L 1 g = 10
6 g
1 Albert Szent-Györgyi (Nobel prize for physiology or medicine)
Chemistry 320 Final Exam Review p.3
10. Units. Be able to identify the most appropriate unit for measuring the following (circle correct answer):
A. the mass of water in a bathtub: kilogram gram centigram milliliter
B. the volume of a teaspoon: kiloliter liter milliliter microliter
C. the length of the classroom: kilometer meter centimeter milliliter
11. Conversion. Fill in the following table.
A. 39.8 mL = 0.0398 L B. 76.8 kg
=
76 800 g
C. 54 m = 5 400 cm D. 25oC = 298 K
12. Scientific notation.
A. Express 4.93 (mm) in meters (m) using scientific notation: ___________________ 4.93 x 10–3
m
B. Write 6.02 x 1023 in conventional format: _________________ 602 000 000 000 000 000 000 000
CHAPTER 3 – ATOMS (13-20)
13. Subatomic particles. Fill in the following table.
Particle symbol location charge relative mass
electron e– electron shell -1 1/2000
proton p+ nucleus +1 1
neutron no nucleus 0 1
14. Isotopes.
A. Define isotope: ________________________________________________________________
atoms of the same element but having different atomic masses (different number of neutrons)
B. Fill in the following table:
Element
Isotope
Symbol
Atomic
Number
Mass
Number
Number of Net
Charge Protons Neutrons Electrons
oxygen 16O
2– 8 16 8 8 10 2+
lithium 74Li 3 7 3 4 3 0
iron 5426Fe 26 54 26 28 23 3+
Chemistry 320 Final Exam Review p.4
15. How many protons are present in:
A. one atom of 3H (hydrogen-3)? __________________________________________________ one
B. one atom of 3H+ (hydrogen ion)? _______________________________________________ one
16. Molecular formulas. Answer the following questions about C5H5Fe(CO)2CH3:
A. How many carbon atoms are present? ______________________________________ (5+2+1) = 8
B. How many total atoms are present? ______________________________ (5+5+1+2+2+1+3) = 19
C. What is its formula mass? _____ (C: 8*12.01)+(H: 8*1.01)+(Fe: 1*55.85)+(O: 2*16.00) = 191.93
17. Convert mass to number of atoms and vice versa. Calculate the number of total atoms in 1.00 gram of
platinum?
1.00 g 1 mol Pt 6.02E+23 atoms Pt 3.09E+21 atoms Pt
195.08 g Pt 1 mol Pt
18. Identify atomic numbers on the periodic table. Complete the following table.
Name Symbol Atomic Number Atomic Mass
A. Lithium Li 3 6.94
B. Rubidium Rb 37 85.47
19. Noble gases:
A. What is the one identifying characteristic of noble gases? _ unreactive (full valence electron shell)
B. Name the noble gases: _________________________ helium, neon, argon, krypton, xenon, radon
20. Identifying particles.
A. What is the smallest particle of an element that retains the properties of that element? _____ atom
B. The number of ___protons___ inside an atom determines the identity of the atom.
CHAPTER 4 – ELECTRONS IN ATOMS (21-27)
21. Waves. For the following parts of a wave, provide the name, abbreviation, and a brief description.
A. wavelength (; lambda), distance from
peak to peak on adjacent wave
B. frequency (, nu), number of wave peaks
hitting a given point per unit time (s–1)
Chemistry 320 Final Exam Review p.5
22. Bohr model of the atom.
A. Describe how the Bohr model of the atom explains the line-emission spectrum of the hydrogen
atom. Use diagrams to accompany your textual description.
A.
B.
C.
D.
Figure 1 (study guide 04). Bohr model of the atom used to explain the line-emission spectrum of
hydrogen. A. In this model, the hydrogen is shown initially at the ground state (electron is at n=1).
B. Energy (e.g., photon of light) hits the electron with sufficient energy to excite it to the third
excited state (n=4). C. However, the attraction between the negatively-charged electron and the
positive nucleus causes the electron to move back towards the nucleus. In this case, it only returns to
the first excited state (n=2) whereby it emits energy in order to do this. D. Eventually, the electron
does return to the ground state (n=1) by further emitting energy.
23. Briefly describe the importance of each to our current quantum mechanical model of the atom.
A. Schrödinger's Wave Equation: _____________________________________________________
describes mathematically the wave properties of electrons and other very small particles.
B. Aufbau principle: _______________________________________________________________
an electron occupies the lowest-energy orbital that can receive it
C. Hund’s rule: ___________________________________________________________________
electron enters the lowest-energy, unoccupied orbital available; then fills the half-occupied
orbitals before filling the next available orbitals (e.g., explains ↑↓ ↑ ↑ ↑ 1s 2s
D. Heisenberg’s uncertainty principle:
one cannot know simultaneously the location and momentum of an electron.
24. What is the difference between an orbit and an orbital? ____________________________________
orbit is a regular and predictable circling around a central object (wha the earth does around the sun)
orbital is a 3-dimensional region around a nucleus that indicates the probable location of an electron
Chemistry 320 Final Exam Review p.6
25. Sketch the general electron cloud for the s- and p- orbitals:
s-orbital
p-orbitals
26. Electron configuration. Write the electron configuration for the following atoms:
oxygen [He] 2s2 2p4
tin [Kr] 4d10 5s2 5p2
27. Orbital notation. Write the orbital notation for the following atoms and identify the atom associated
with the given orbital notation:
bromine
manganese:
CHAPTER 5 – PERIODIC TABLE (28-34)
28. Identify the contribution made by each person:
A. Democritus: _____________________ idea that all matter is made of indivisible particle - atomos
B. Dalton: ________________________________________ (5) postulates of modern atomic theory
C. Mendeleev: ____________________________________________ created modern periodic table
D. Thomson: ________________________ discovered electrons; ‘Plum-pudding’ model of the atom
E. Rutherford: ____________________________________________________________________
discovered proton, nucleus, nucleus is small compared to entire atom; ‘Solar system’ model
F. Moseley: ________ arranged elements on periodic table according to atomic number (# of protons)
Chemistry 320 Final Exam Review p.7
29. On the below table (or another) periodic table, identify the following:
A. groups ↔ B. families ↕ C. periods (q.v., families)
D. alkali metals E. alkaline earth metals F. transition metals
G. noble gases H. lanthanide series I. actinide series
J. inner transition metals K. metals L. nonmetals
M. halogens
30. Identify each as an element, atom, ion, molecule and/or ionic compound?
A. Li B. Li+ C. NaCl D. 218O
element / atom ion ionic compound ion
31. Arrangement of elements in periodic table. Which elements are most related in each of the
following? Why?
A. sodium, argon, cesium B. fluorine, iodine, oxygen C. strontium, beryllium, barium
{same group/family}
32. What does the staircase line represent on the periodic table (e.g., between aluminum and silicon)
________________ separation between metals and nonmetals (elements on line are often metalloids)
33. Valence electrons.
A. What is the pattern for the number of valence electrons of an atom and its group or family?
_______________ elements in the same group/family have the same number of valence electrons
34. Valence electrons. How many valence electrons does each of the following have? _______________
A. Li one B. Li+ zero C. Pb 4 D. Br seven
Chemistry 320 Final Exam Review p.8
35. Draw arrows on the below periodic table indicating the trends in atomic radius as one goes (1) from
left to right across a row and (2) from top to bottom down a column. Explain the two trends.
Periods: As one moves left to right across a period, one adds not only electrons but also protons. The
increased pull by the more highly charged nucleus on the electrons is stronger than the repulsion
between the electrons. Thus, the atomic radii across a period decrease.
Families: one adds electron shells as one goes down a group, increasing atomic radius (or size)
CHAPTER 6 – BONDING (35-41)
36. Chemical bonds – General.
A. Name three types of chemical bonds: ______________ ionic, nonpolar covalent, & polar covalent
B. What two particles are attracted together to form a chemical bond? _________ protons & electrons
37. Types of compounds. Identify each of the following as an ionic compound or molecule. (Table 5-20
and information from Figure 6-2 will be supplied to you.)
A. HCl - ionic (3.0 – 2.1 = 0.9) B. Br2 – molecule (2.8 – 2.8)
C. CH4 - molecule (2.5 – 2.1 = 0.4) D. H2S – molecule (2.5 – 2.1 = 0.4)
38. Draw the Lewis dot structures for each of the following molecules:
A. H2O
B. CCl4
C. C2H4
Chemistry 320 Final Exam Review p.9
39. Bond properties. Fill in the following table:
A. Property Ionic bond Covalent bond Metallic bond
B. relative melting point high low high
C. relative boiling point high low high
D. relative hardness high low high
E. location of electrons unequal distribution
around specific
atom(s)
polar or nonpolar
distribution around
specific atom(s)
delocalized (‘sea of
electrons’)
40. Molecular shape. Predict the shape of the following molecules. (Table 6-5 will be supplied to you
(but without the “Formula example” column.)
A. water – AB2E2 = bent B. chlorine gas – AB2 = linear
C. NH3 – AB3E1 = trigonal-pyramidal D. CCl4 – AB4 = tetrahedral
41. Bond strength. Rank the following intermolecular bonds from strongest (1) to weakest (3):
A. dipole-dipole interactions 2
B. hydrogen bonds 1
C. London dispersion forces 3
CHAPTER 7 – FORMULAS & COMPOUNDS (42-56)
42. Naming, ionic compound: name formula. What is the formula for each of the following?
A. calcium hydroxide Ca(OH)2
B. lithium sulfate Li2SO4
43. Naming, ionic compound: formula name. What is the name for each of the following?
A. CuO copper(II) oxide
B. Cu2O copper(I) oxide
C. FeSO4 iron(II) sulfate
44. Naming, ionic compound: formula name.
A. NH4OH ammonium hydroxide
B. potassium permanganate KMnO4
45. Molar mass. Calculate the molar mass for each of the following.
A. NH4NO3 ammonium nitrate
B. FeSO4 iron(II) sulfate
Chemistry 320 Final Exam Review p.10
46. Naming, ionic compound: formula name?
A. PbCl2 lead(II) chloride
B. strontium fluoride SrF2
47. Oxidation number. What is the oxidation number (charge) of calcium in CaO? _______________ 2+
48. Predicting chemical formulas. Predict the compound formula formed from the following.
A. sodium & sulfur: Na2S
B. ammonium & nitrate: NH4NO3
49. Formula mass. What is the formula mass of KMnO4? ______________________________________
K: 1 39.10 39.10
Mn: 1 54.94 54.94
O: 4 16.00 64.00
SUM: 158.04
50. Percent composition. What is the percentage composition (by mass) of nitrogen in NH4NO3?
N: 2 14.01 28.02 / 80.06 =
H: 4 1.01 4.04
O: 3 16.00 48.00
SUM: 80.06
51. Atomic composition. How many nitrogen atoms are there in 10 mol of NH4NO3?
10 mol NH4NO3 2 mol N 6.02E+23 atoms N 1.20E+25 atoms N
1 mol NH4NO3 1 mol N
52. Empirical formula. What is the empirical formula of N2O4? ______________________________ NO2
53. Determining molecular formula from empirical formula.
A compound's empirical formula is NO. If the molar mass is 184.3 g/mol, what is the molecular
formula?
N: 14.01 184.3 = 6 N6O6
O: 16.00 30.01
sum: 30.01
54. Chemical formula.
A. How many nitrogen atoms are in one formula unit of (NH4)3PO4?_____________________ three
B. How many hydrogen atoms are in one formula unit of (NH4)3PO4? ___________________ twelve
55. Naming, molecules: formula name.
A. What is the formula for carbon tetrachloride? _____________________________________ CCl4
B. What is the formula for sulfur trioxide?___________________________________________ SO3
Chemistry 320 Final Exam Review p.11
56. Ionic compounds, nomenclature.
A. Which of the following is a binary compound? magnesium oxide or magnesium hydroxide? MgO
B. Which of the following is a binary compound? sodium nitrate or silver chloride? _________ AgCl
Chemistry 320 Final Exam Review p.12
CHAPTER 8 – EQUATIONS & REACTIONS (57-71)
57. Describe each of the following:
A. precipitate: _________ solid that is produced in a chemical reaction & separates from the solution
B. product: _______________________________________ substance formed by a chemical change
C. reactant: _____________________________________ substance that reacts in a chemical change
D. exothermic: ________________________________________________________ gives off heat
E. endothermic: ________________________________________________________ absorbs heat
F. coefficient: _ small whole number that appears in front of a chemical formula in a chem. equation
G. subscript: _______ small number in a chemical formula indicating the number of preceding atoms
58. Diatomic molecules.
A. What is a diatomic molecule? ____________ molecule containing two atoms of the same element
B. Which of the diatomic molecules are gases at STP? _________________________ HNOClBrIF
59. Types of chemical reactions. Identify the type of reaction for each equation:
A. H2(g) + Cl2(g) 2HCl(g):__________________________________________________ synthesis
B. H2O 2H2(g) + O2(g): ________________________________________________ decomposition
C. Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s): ________________ single (cationic) displacement
D. NaCl(aq) + AgNO3(aq) NaNO3(aq) + AgCl(s): ________________________ double replacement
E. C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(g): _________________________________ combustion
60. Chemical formulas. Predict the formula for the product(s) produced, and balance the following
equation:
2 HCl + 1 Ca(OH)2 1 CaCl2 + 2 H2O
61. Balancing chemical equation. Balance the following chemical equation:
_2_Al2O3 __4_Al + __3_O2
62. Types of chemical reactions. For each of the following, provide the general format for the equation.
A. synthesis reaction ______________________________________________________ A + B C
B. decomposition reaction ________________________________________________ AB A + B
C. single-replacement reaction ________________________________________ AB + C AC + B
D. double-replacement reaction ____________________________________ AB + CD AD + CB
E. combustion reaction ________________________________________ CxHy + O2 CH4 + H2O
Chemistry 320 Final Exam Review p.13
63. Properties of metals. What is the product of an active metal and a strong acid? __________________
General properties of metals:
conduct heat & electricity luster
ductile malleable high tensile strength
Product of an active metal + strong acid: hydrogen gas
64. Activity series. Predict whether or not a reaction will occur in the following. If a reaction will occur,
predict the product(s). (Table 8-3 will be provided.)
A. Cr(s) + H2O(l) __________________ CrxOy + H2() (possibilites: CrO2, CrO3, Cr2O3, Cr3O4)
B. Pt(s) + O2(g) ______________________________________________________ (no reaction)
C. Cd(s) + 2HBr(aq) __________________________________________________ CdO + H2()
D. Mg(s) + steam _____________________________________________________ MgO + H2()
65. Name, formulas, and uses of common acids. Fill in the below table.
Name Formula Use/Occurrence
A. hydrochloric acid HCl food digestion / stomach acid
B. nitric acid HNO3 fertilizer, pharmaceuticals, explosives
C. sulfuric acid H2SO4 car batteries
D. phosphoric acid H3PO4 food ingredient (e.g., Coke, Pepsi)
66. Balancing chemical equations. Predict the product(s) and balance the following chemical equations.
A. __1__Mg + ___1__H2O _____________________________________ MgO + H2()
B. ___1__Ca + __2__H2O _______________________________ Ca(OH)2(aq) + H2(g)
C. ___2_AgNO3 + ___1__Mg _____________________________ Mg(NO3)2 + __2__Ag
D. __3___Ca(OH)2 + ___2__H3PO4 _____________________________1 Ca3(PO4)2 + 6 H2O
67. Groups. Relative reactivity.
A. Which is more reactive: alkali metals or alkaline earth metals? _________________ alkali metals
B. Which is more reactive: halogens or noble gases? ______________________________ halogens
C. Which is more reactive: transition elements or alkaline earth metals? ______ alkaline earth metals
68. What are three general indicators that a chemical reaction has occurred?
A. heat or light given off ___________________________________________________________
B. gas or precipitate formed ________________________________________________________
C. color change ___________________________________________________________________
Chemistry 320 Final Exam Review p.14
69. Reaction prediction. Predict the product(s) in the following reaction.
A. 2 Ca + O2 _____________________________________________________________ 2 CaO
B. 2 Li + H2O _______________________________________________________ Li2O + H2(g)
70. Balancing equations. Once the products of a chemical reaction are determined, can you balance a
chemical equation by changing their subscripts? Yes or no. Explain
___ No. That would change the chemical identity. Changing the coefficient(s) balances the equation.
71. Balancing equations. Predict the coefficients: __2_Mg + __1_O2 __2_MgO.
CHAPTER 9 – STOICHIOMETRY (72-86)
72. Mole ratio. State the mole ratios for the following chemicals in the below chemical equation –
balance the equation first.
__1_CH4(g) + __2_O2(g) _1_CO2 + _2_H2O(g)
A. CH4 : O2 1 : 2 B. CH4 : CO2 1 : 1
C. CH4 : H2O 1 : 2 D. O2 : CO2 2 : 1
E. O2 : H2O 1 : 1 F. CO2 : H2O 1 : 2
73. Mole ratio. For the reaction in problem #72, how many moles of oxygen are required to completely
burn 12 mol of methane (CH4)? ________________________
12 mol CH4 2 mol O2 24 mol O2
1 mol CH4
74. Excess reactant. In the following reaction:
1N2(g) + 3H2(g) 2NH3(g)
A. If 3 moles of nitrogen and 6 moles of hydrogen are used,
1) What is the limiting reactant? 3 mol N2 3 mol H2 9 mol H2 N2 is limiting reactant
1 mol N2
2) What is the excess reactant? 6 mol H2 1 mol N2 2 mol N2 2 is excess reactant
3 mol H2
=
=
75. Stoichiometry, mass-mass. How many grams of oxygen gas are produced when 48.32 g of water
decompose? 2H2O(l) 2H2(g) + O2(g)
48.32 g water 1 mol water 1 mol O2 32.00 g O2 = 42.90 g O2
18.02 g water 2 mol water 1 mol O2
Chemistry 320 Final Exam Review p.15
76. Stoichiometry, moles-number of particles. How many hydrogen molecules are produced in question
#75?
48.32 g water 1 mol water 2 mol H2 6.02E+23 molecules H2 = 1.61E+24 molecules H2
18.02 g water 2 mol water 1 mol H2
77. Stoichiometry, mass-volume. How many liters of water are produced from the combustion of 67.9 g
of gasoline at STP? 2C8H18 + 25O2 18H2O + 16CO2
67.9 g gas 1 mol gas 18 mol H2O 22.4 L H2O = 1.20E+02 L H2O
114.26 g gas 2 mol gas 1 mol H2O
78. Stoichiometry, mass-moles. For the reaction U(s) + 3F2(g) UF6(g), how many moles of uranium
hexafluoride are produced from 33.06 g of uranium?
33.06 g U 1 mol U 1 mol UF6 = 1.39E-01 mol UF6
238.03 g U 1 mol U
79. Percent error. A student measures the mass and volume of a substance and calculates its density to be
1.36 g/mL. The actual value is 1.42 g/mL. What is the student’s percent error? % error: theory - experimental
theory
% error: 1.42 g/mL - 1.36 g/mL
1.42 g/mL4.23%
* 100% =
* 100% =
80. Limiting reactant. What is the limiting reactant in the formation of ammonia (N2 + 3H2 2NH3)
given 100 g of N2 and 40 g of H2.
reactant
100 g N2 1 mol N2 3 mol H2 2.02 g H2 21.6 g H2 2 is limiting reactant
28.02 g N2 1 mol N2 1 mol H2 (have enough H2)
40 g H2 1 mol H2 1 mol N2 28.02 g N2 185 g N2 2 is excess reactant
2.02 g H2 3 mol H2 1 mol N2 (don't have enough N2)
=
=
required
Chemistry 320 Final Exam Review p.16
81. Stoichiometry, mass-mass. How many grams of oxygen are in 0.18 g of water?
0.18 g water 1 mole water 1 mol O 16.00 g O 0.16 g O
18.02 g water 1 mol water 1 mol O=
82. Stoichiometry, general. What is the purpose of balancing a chemical equation?
represent accurately the reaction with the chemical equation that follows the law of conservation of mass
83. Percent yield. Chlorobenzene, an important chemical in industry, is made with the following
reaction: C6H6(l) + Cl2(g) C6H5Cl(s) + HCl(g). When 18.4 g of C6H6 react with an excess of Cl2,
the actual yield is C6H5Cl is 19.4 g. What is the percent yield of chlorobenzene?
18.4 g C6H6 1 mol C6H6 1 mol C6H5Cl 112.46 g C6H5Cl = 26.49 g C6H5Cl
78.11 g C6H6 1 mol C6H6 1 mol C6H5Cl
% yield: 19.4 g 73.23 %
26.49 g=
theory
expt'l* 100%
84. Percent yield. If the expected yield of NO is 73.6% from the reaction 4NH3(g) + 5O2(g) 4NO(g) +
6H2O(g), how many grams of NO would be recovered from 170 g of NH3 and an unlimited amount of
oxygen gas.
170 g NH3 1 mol NH3 4 mol NO 30.01 g NO = 300 g NO
17.04 g NH3 4 mol NH3 1 mol NO
% yield: % yield theory
100 %
73.6 300 = 221 g NO
100 %=
=expt'l
expt'l
theory
expt'l* 100%
85. Stoichiometry, mol-mol. For the reaction HCl + NaOH NaCl + H2O, how many moles of water
can be produced from 16.0 mol of hydrochloric acid and an unlimited amount of sodium hydroxide?
16.0 mol HCl 1 mol H2O = 16.0 mol H2O
1 mol HCl
86. Stoichiometry. Define stoichiometry:
______ deals with the mass(amount) relationships between reactants and products in a chemical reaction.
Chemistry 320 Final Exam Review p.17
CHAPTERS 10/11: GASES (87-102)
87. Dalton’s law of partial pressures. In a sample of dry air, the partial pressure of O2 is 0.285 torr and N2
is 593.525 torr. What is the partial pressure of CO2, the third gas? (assume STP; Table A-8 will be
provided.)
PT = PO2 + PN2 + PCO2
PCO2 = PT PO2 PN2
PCO2 = 760.00 0.285 593.525
PCO2 = 166.190 mm Hg
ALL UNITS ARE IN mmHg
88. Ideal gas law. What is the pressure exerted by 3.0 mol of gas occupying 0.75 L at 22oC?
=
=
= 3.0 mol 0.0821 (L-atm)/(K-mol) 295 K = 96.878 atm
0.75 L
V
P
R T
P n R T
P V n
89. Measurements. Identify what is measured by each of the following instruments.
A. barometer ____________________________________________________ atmospheric pressure
B. spectrophotometer ______________________________________________ wavelengths of light
C. sphygmomanometer __________________________________________________ blood pressure
D. thermometer _________________________________________________________ temperature
90. Relationship. Identify each as a directly related or inversely related. PV = nRT
A. temperature and volume direct B. volume and pressure inverse
C. pressure and
temperature direct
D. temperature and number of moles
inverse
91. Unit conversion. Convert the following: atm mm Hg torr psi
A. 1.2 912 912 17.6
B. 0.971 737.6 737.6 14.3
C. 2.00 1520 1520 29.4
Chemistry 320 Final Exam Review p.18
92. Dalton’s law of partial pressure. Hydrogen gas, generated by the reaction of magnesium and water, is
collected over water at 25.0oC. If the barometric pressure is 745.7 mm Hg, what is the partial pressure
of the hydrogen gas?
PT = PH2 + PH2O
PH2 = PT PH2O
PH2 = 745.7 23.8
= 721.9 mmHg
ALL UNITS ARE IN mmHg
( from SG 10'11, appendix)
93. Combined gas law. If a 3.0 L sample of gas has a pressure of 0.89 atm at 20oC, what is its volume at
1.5 atm and 15.oC?
P1 V1 P2 V2
T1 T2
V2 = P1 V1 T2
T1 P2
V2 = 0.89 atm 3.0 L 15oC
20oC 1.5 atm
V2 = 1.3 L
=
94. Stoichiometry, gas. What is the volume of oxygen gas produced from the decomposition of 3L of
water? 2H2O(l) 2H2(g) + 1O2(g)
3 L H2O 1000 mL H2O 1 g H2O 1 mol H2O 1 mol O2 22.4 L O2 = 1865 L O2
1 L H2O 1 mL H2O 18.02 g H2O 2 mol H2O 1 mol O2
95. STP. How many moles of gas occupy an empty 2-L soda bottle at STP? ______________________
P =
n =
P = 1.0 atm 2 L = 0.089 mol
0.0821 (L-atm)/(K-mol) 273 K
T
P V
V n
R T
Chemistry 320 Final Exam Review p.19
96. Ideal gas law, molar mass. A 10 L sample of gas has a mass of 18 g at STP. What is the molar mass
of the gas? M =
M = 18 g 0.082 (L-atm)/(K-mol) 273 K = 40 g/mol
1 atm 10 L
P V
Tm R
97. Gas laws, rearrangement of equation. Rearrange the combined gas law to solve for V2. ___________
P1 V1 P2 V2
V2 = P1 V1 T2
T1 P2
=T1 T2
98. Ideal gas law. How many moles of gas are in a sample having a volume of 32.9 L at 23oC and
exerting a pressure of 767 mm Hg.
P = 767 mm Hg 1 atm = 1.01 atm
760 mm Hg
V = 32.9 L
n = ?
T = 23 oC + 273 = 296 K
P =
n =
P = 1.01 atm 32.9 L = 1.37 mol
0.0821 (L-atm)/(K-mol) 296 K
R T
T
P V
V n
99. Kinetic-molecular theory of matter. Which has more kinetic energy: a VW Jetta, having a mass of
approximately 2200 lbs and traveling at 65 mph, or a H2, having a mass of approximately 4500 lb
traveling at 30 mph? Explain.
½
Jetta: = ½ 2200 lb 4225 m2/hr
2= 4,600,000 lb*m
2/hr
2
H2: = ½ 4500 lb 900 m2/hr
2= 2,000,000 lb*m
2/hr
2
Kinetic Energy = m v2
The kinetic energy is related to the velocity squared but the mass to only the first
order. Thus, velocity has more of an impact on kinetic energy than mass.
Chemistry 320 Final Exam Review p.20
100. Combined gas law. The temperature of a 4.0 mL balloon is raised from 4oC to 35oC. What is the
new volume of the balloon? P1 V1 P2 V2 T1 = 4 oC = 277 K
T2 = 35 oC = 308 K
V2 = P1 4.0 mL 308 K
277 K P2
V2 = 4.45 mL
=T1 T2
101. Terms. Describe each of the following.
A. fluidity: ________________________________________________ ability to flow (or be poured)
B. compressibility: ______________________________________ volume can be greatly decreased
C. diffusion & effusion: _____ diffusion: spontaneous mixing of particles of two or more substances
caused by their random motion
effusion: gas particles, under pressure, pass through a tiny opening (e.g., leak on a bicycle tire)
102. Density of gases. Rank in order of decreasing density (assume the same pressure & temperature):
Rank Molar Mass
A. H2 4 2.02 g
B. HCl 1 36.46 g
C. O2 2 32.00 g
D. HF 3 20.01 g
all gases occupy the same volume at the
same pressure & temperature (Avogadro's
Law) so density is directly proportional to
molar mass
Explanation
CHAPTER 12/13: SOLUTIONS (103-111)
103. Terms. Define & give an example of:
A. heterogeneous mixture: __________________________________________________________
a blend of two or more kinds of matter, each retains its own identity and properties
B. solution: ________________ a homogeneous mixture of two or more substances in a single phase
C. suspension:
___ a mixture in which the particles are large enough to settle out (unless the mixture is agitated)
D. colloid suspension: ______________________________________________________________
a mixture consisting of particles that are intermediate in size between those in solution and
supsensions (produces Tyndall effect)
Chemistry 320 Final Exam Review p.21
104. Terms. Define & give an example of:
A. melting solid liquid
(ice water)
B. freezing liquid solid
(water ice)
C. vaporization liquid gas (water steam)
D. condensation gas liquid (dew)
E. sublimation solid gas (dry ice)
F. deposition gas solid (frost)
105. Terms. Define & give an example of:
A. supersaturated solution: __ solution that conatins more dissolved solute than a saturated solution
contains under the same conditions (honey; hot packs in lab [sodium acetate])
B. saturated solution: ____________ solution that contains the maximum amount of dissolved solute
(coffee that contains just enough sugar so adding any more will cause it to fall out of solution)
C. unsaturated solution: __ solution that contains less solute than a saturated solution (coffee w/ 1 tsp
sugar)
106. Terms. Define & give an example of:
A. electrolyte: ____ substance that dissolves in water to give a solution that conducts electric current
(NaCl)
B. nonelectrolyte: _ substance that dissolves in water to give a solution that does not conduct electric
current (glucose)
C. soluble: _____________________________________ capable of being dissolved (sugar in water)
D. immiscible: ____________________ two liquids that do not dissolve in each other (oil and water)
107. Terms. Explain the Tyndall effect.
_______________________ property of colloid suspensions: ray of light is visible (headlights in fog)
108. Solubility. What are three ways to increase the rate at which a solid dissolves in a solvent:
A. heat B. particle size (SA-V ratio) C. agitation (stirring)
109. Molarity. What is the molarity of a solution that contains 25.3 g of NaCl in 500 mL of solution?
25.3 g NaCl 1 mol NaCl 1000 mL 0.866 M
500 mL 58.4 g NaCl 1 L
110. Kinetic-molecular theory.
A. Which has the most kinetic energy: particles in a gas, particles in a liquid, or particles in a solid?
B. Which diffuses faster, HCl or O2? Why? ____________ lower molar mass faster rate of motion
22
2
1
2
1 :Law sGraham' BBAA vmvm =
Chemistry 320 Final Exam Review p.22
111. Terms. Define & give an example of the following:
A. precipitate: _ something that ‘falls out of solution’ (solid produced in a liquid solution; rain, snow)
B. surface tension: force that tends to pull together part’s of a liquid’s surface, thereby decreasing the
surface area to the smallest possible size.
CHAPTER 15/16: ACIDS & BASES (112-120)
112. Properties, acids & bases. List five properties of each.
ACIDS BASES
taste sour taste bitter
change color of acid-base indicators
(e.g., turn litmus red)
change color of acid-base indicators (e.g., turn
litmus blue)
some react with active metals to produce
hydrogen gas
dilute solutions feel slippery
react with bases to produce water & a
salt
react with acids to produce water & a salt
conduct electric current conduct electric current
113. Conjugate acids & bases. Complete the following equation and identify the conjugate acid-base
pairs. HCl + H2O _____ + _____
114. Common acids. Fill in the following table for name, formula and occurrence of common acids:
Name Formula Use/Occurrence
A. hydrochloric acid HCl digest food (stomach)
B. nitric acid HNO3 fertilizer, pharmaceuticals, explosives
C. acetic acid HC2H3O2 vinegar (is 3% acetic acid + flavors)
D. sulfuric acid H2SO4 car batteries
E. phosphoric acid H3PO4 food ingredient (e.g., Coke, Pepsi)
F. lactic acid (n/a) milk
G. citric acid C6H8O7 fruit
H. malic acid C4H6O5 apples
115. Conjugate acids & bases. Complete the following equation and identify the conjugate acid-base
pairs. NH3 + H2O __NH4+___ + __OH–___
116. Compare Arrhenius’s (traditional) and Brønsted-Lowry definitions of acids and bases:
Acid Base
Arrhenius’s Definition donates H+ donates OH–
Chemistry 320 Final Exam Review p.23
Brønsted-Lowry donates H+ accepts H
+
117. Identify each of the following as a strong or weak acid (table will be provided):
A. HCl strong acid B. HNO3 strong acid
C. HC2H3O2 weak acid D. H2SO4 strong acid
118. Acid-base neutralization reactions. Predict the product(s) formed from the following reactions:
A. HC2H3O2 + NaOH NaC2H3O2 + H2O ___________________________________________
B. 2H3PO4 + 3Ca(OH)2 Ca3(PO4)2 + 6H2O _________________________________________
119. Acid-base indicators. What is the color of litmus in:
A. acids: red D
red
aci
B. base: blue
lue
aseB
120. pH. Determine the pH or pOH of a given solution:
A. What is the pH of a 0.05 M Mg(OH)2 solution? _______________________________________
pOH = -log([OH])
Mg(OH)2 Mg2+
+ 2OH
molarity: 0.05 0.05 0.10
pOH = -log(0.10) = 1.0
14 = pH + pOH
pH = 14 - pOH
pH = 14 - 1.0
pH = 13
B. What is the [OH–] of a 1.02 x 10–12 M H3PO4 solution? _________________________________
1 H3PO4 3 H+
+ 1 PO43
molarity: 1.02E-12 3.06E-12 1.02E-12
pH = -log([H3O+]) = 11.5
14 = pH + pOH
pOH = 14 - pH
pOH = 14 - 11.5
pOH = 2.49 1