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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


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)


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+


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)


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


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)


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


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


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


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


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


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


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 ___________________________________________________________________


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


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


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.


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


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


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.


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)



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)


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)


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 =


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–


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

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