Name:__[KEY]____________________Date:_____________ Period:____

Review: Volusia County Pre-AP Chemistry EOC 2018

Unit 1: Matter and Measurement

Unit 2: Atomic Theory and Structure

Unit 3: Electrons and Modern Atomic Theory

Unit 4: The Periodic Table

Unit 5: Ionic Bonding and Nomenclature

Unit 6: Covalent Bonding and Nomenclature

Unit 7: Chemical Reactions

Unit 8: Chemical Composition

Unit 9: Stoichiometry

Unit 10: Energy and States of Matter

Unit 11: Gas Laws

Unit 12: Solutions

Unit 13: Acids and Bases

Unit 14: Reaction Rates and Equilibrium

Unit 15: Nuclear Chemistry

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Unit 1: Matter and Measurement

1. These values were recorded as the mass of products when the same chemical reaction was carried out three separate times: 3.20 g; 2.87 g; 3.89 g.

The correct mass of products from that reaction was actually 3.45 g.

Are the measured masses accurate? _N__

Are the measured masses precise? _N__

Explain: _ The measurements were not consistently close to each other (not precise), and the average (3.32) was not close to the true value (not accurate)_

2. What is the quantity 65.5 m expressed in cm? 6550 cm

3. Convert 2.3 mL into L. 0.023 L

4. The mass of a soft lump of metal is 214 g, and the volume is 11.3 cm3.What is the density of the metal in g/cm3 ?

Using the table of metal densities,

identify the metal: Gold

The soft metal lump is then smushed together with another lump of the exact same soft metal.

The density of the new larger lump is greater than / equal to / less than the original lump.

5. Differentiate between a controlled variable, independent variable, and dependent variable.

A controlled variable is not changed but held constant.An independent variable is manipulated or tested by changing it intentionally.A dependent variable is measured or observed and depends on the independent variable.__________________________________________________________________________________

6. How many significant digits are in each of the following measurements?

0.0100 kg 2000 m 30.50 g 0.0003 mL__3__ __1__ __4__ __1__

2

Density Table

Metal Density(in g/cm3)

Aluminum 2.7Iron 7.9Gold 19.3Platinum 21.4

d = (214) = 18.9 g/cm3

(11.3)

7. Round each of the following measurements to the stated number of significant digits.

20.09 to three sig digs. ___20.1____

0.00045 to one sig dig. __0.0005___

26.98 to three sig digs. ___27.0____

8. Complete the operation and round your answer to the correct number of significant digits.

a) 3.28 + 1.2 = __4.5___ (keep fewest decimal places)

b) 84.4 ÷ 4.0 = __21_____ (keep fewest sig digs)

c) 6000 ÷ 8.0 = __800___ (keep fewest sig digs)

d) 21 – 2.0 x 2.00 = ___17____

first… (2.0 x 2.00) = 4.0 (keep fewest sig digs)…

then… 21 – 4.0 = 17 (keep fewest decimal places)

Unit 2: Atomic Theory and Structure

9. Which of the following is FALSE about subatomic particles?

A. Electrons are negatively charged and are the lightest subatomic particle.B. Protons are positively charged and have nearly the same mass as neutrons.C. Neutrons have no charge and have no mass.D. The mass of a neutron nearly equals the mass of a proton.

10. All atoms of the same element have the same ____.

A. number of neutronsB. number of protonsC. mass numbersD. mass

11. What are atoms of the same element with different numbers of neutrons?

A. ionsB. atomsC. numbers of electronsD. isotopes

12. Explain why isotopes of the same element are not considered different elements.

Isotopes of the same element have the same number of protons (same atomic number) so they must be atoms of the same element.

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X3575

13. Element X has an atomic number of 35 and a mass number of 75.How many of each subatomic particle are in a neutral atom of the element?

A. 35 protons, 35 neutrons, and 70 electronsB. 35 protons, 75 neutrons, and 35 electronsC. 75 protons, 35 neutrons, and 40 electronsD. 35 protons, 40 neutrons, and 35 electrons

14. The element silver has two naturally occurring isotopes:silver-107 has an isotopic mass of 106.905 amu and a relative abundance of 51.84%silver-109 has an isotopic mass of 108.905 amu and a relative abundance of 48.16%Calculate the average atomic mass of the element silver. (SHOW ALL WORK)

(106.905 amu)(0.5184) + (108.905 amu)(0.4816) = 107.87 amu

15. The fictitious element Z has two naturally occurring isotopes, Z-310 and Z-313.The average atomic mass of Z is found to be 311.02 amu. Which of the two isotopes of the fictitious element Z is more abundant?

Z-310 or Z-313 or equal abundance (circle one)

How could you tell? the weighted average atomic mass (311.02 amu) is closest to the isotopic mass of Z-310

16. Define Mole: the number of items in a mole of items is 6.022 x 10 23

17. 15 moles of sodium (Na) and 15 moles of carbon (C) have the same number of _________.

A. atomsB. gramsC. A and BD. NONE of the above

18. What is the molar mass of Fe3(PO4)2 ? 3(Fe) + 2(P) + 8(O) =

3(55.85) + 2(30.97) + 8(16.00) = 357.49 g/mol

19. Convert 4.50 moles of Fe to atoms of Fe.

4.50 mol Fe x 6.022 x 10 23 atoms Fe = 2.71 x 1024 atoms Fe 1 mol Fe

20. Convert 4.03 x 1022 molecules of water to moles of water.

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4.03 x 1022 molecules H2O x 1 mol H2O = 0.0669 mol H2O 6.02 x 1023 molecules H2O

21. What conclusions were drawn from Rutherford’s Gold Foil Experiment about atomic structure?

Most of the alpha particles traveled through the gold atoms showing atoms are mostly empty space. Very few of the positively charged alpha particles deflected which revealed a tiny, dense, positive region in atoms.

Unit 3: Electrons and Modern Atomic Theory

22. Examine the Bohr Model in Figure 1. Place a “G” in the box that points to the

“Ground State” electron orbit.

Place an “E” in the box that points to the“Excited State” electron orbit.

How is the “excited state” different from the“ground state?” an electron in the “excited” state has a higher energy than the “ground” state

23. Study Figures A and B below:

In which figure is energy absorbed? A or BHow do you know? _energy is absorbed to separate an electron (move further) from the nucleus

In which figure is energy emitted? A or BHow do you know? _energy is emitted when an electron moves closer to the nucleus

Which figure shows a photon of the greatest energy? A or B

How do you know? _greater change in energy levels (3 levels from 5 to 2)______________

5

Nucleus

level 1 e-

Einitial

photon

level 3

Efinal

E)level 2

e-

Efinal

photon

level 5

E initial

E)e-

e-

Figure A:electron move from

high to low orbit

Figure B:electron move from

low to high orbit

Einitial

Figure 1:Bohr Model

EG

24. Write the number of valence electrons for the elements in each group in the boxes.

25. Complete the table.

Element Symbol# of

VALENCE Electrons

Electron Configuration Lewis Dot Diagram Bohr Model

argon Ar 8 1s2 2s2 2p6 3s2 3p6

nitrogen N 5 1s2 2s2 2p3

bromine Br 7 1s22s22p63s23p64s23d104p5

manganese Mn 2 [Ar] 4s2 3d5

For #26-28, consider the following equations and constants:

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1 2 3 4 5 6 7 8

Ar

N

Br

Mn

26. What is the wavelength (in meters) of infrared light with a frequency of 3.50 x 109 Hz?

27. What is the energy (in joules) of blue light with a frequency of 6.30 x 1014 Hz?

28. What is the energy (in joules) of an x-ray with a wavelength of 3.30 x 10–10 m?

Unit 4: The Periodic Table

29. A chemist needs calcium to perform an experiment in the lab and discovers that she does not have any calcium. List one element she could use for this experiment to best replace calcium.

Mg because b/c they have the same number of valence electrons (2) so they will react similarly. (Sr is also acceptable)

30. For the elements listed above, which element has the largest atomic radius? Li Which element has the smallest atomic radius? Ne

These atoms change size from left to right across a period. Why do atomic sizes exhibit this trend?

Atoms get smaller left to right across a period as more protons are added pulling electrons closer without adding more energy levels.

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c = c = 3.00 x 108 m/sE = h h = 6.63 x 10–34 J∙s

c = λ

(3.00 x 108) = λ (3.50 x 109)

E = hE = 4.18 x 10–19 JE = (6.63 x 10–34) (6.30 x 1014)

(3.00 x 108) = λ (3.50 x 109)λ = 0.0857 m

31. Select the element with the highest ionization energy from above. O

Select the element with the lowest ionization energy from above. P

__________________________________________________________________________________

Unit 5: Ionic Bonding and Nomenclature

32. In an ionic bond, electrons are shared / transferred / connected between atoms. (circle one)

33. Ionic bonds always form between metals and nonmetals. (metals, nonmetals, metalloids)

34. Differentiate between anions and cations.Cations are positive charged ions formed by losing negative electrons (usually metals)Anions are negative charged ions formed by gaining negative electrons (usually nonmetals)

35. For the pairs of elements listed below, circle pairs that would likely form ionic bonds.

C and H Na and F Hg and Ag Mg and S N and C K and O

36. How did you know which elements in the question above would form ionic bonds?

Ionic bonds always form between __metals__ and __nonmetals_.

37. Which property corresponds to ionic compounds with ionic bonds? (circle one from each row)

High melting point or Low melting point Crystal Solid or Liquids and Gases Conducts Electricity or Non Conductor Soluble in Water or Non Soluble in Water

38. What is the charge of the unknown element X in the compound MgX2?

A. 1– B. 2– C. 1+ D. 2+

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(conducts as ions only when melted or dissolved)

39. Name the following compounds. (Some compounds may need a Roman numeral.)

Fe2O3 iron(III) oxide

NaBr sodium bromide

CaCl2 calcium chloride

SnCl4 tin(IV) chloride

40. Write a formula for the following chemical compounds.

copper(II) oxide CuO

calcium sulfide CaS

magnesium iodide MgI2

nickel(II) bromide NiBr2

Unit 6: Covalent Bonding and Nomenclature

41. What is the difference between polar and nonpolar covalent bonds?

Polar covalent bonds share electrons unequally due to a difference in electronegativity of the atoms. Nonpolar covalent bonds share electrons equally.

42. Which is TRUE of a nonpolar covalent bond?

A. electrons are shared unequally between atoms .B. a cation is bonded to an anionC. electrons are transferred between atomsD. electrons are shared equally between atoms

43. List the 7 diatomic elements.

H2 N2 O2 F2 Cl2 Br2 I2

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For #44-46, write the NAME or FORMULA for the following molecular compounds:

44. CCI4 carbon tetrachloride

45. S2O3 disulfur trioxide

46. tetraphosphorus pentoxide P4O5

For #47-49 ,

-Draw the Lewis dot structures for the following compounds.

-Count the number of electron domains around the central atom which repel each other.

-Identify the molecule as polar or nonpolar by circling the correct label.

-Name the molecular geometry and list the bond angle.

47. NI3 Domains: 4 polar or nonpolar geometry: trigonal pyramidalbond angle: 109.5o

48. SO2 Domains: 3 polar or nonpolar geometry: bentbond angle: 120o

49. CH2O Domains: 3 polar or nonpolar geometry: trigonal planarbond angle: 120o

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(originally tetrahedral due to 4 electron domains)

I – N – II

(originally trigonal planar due to 3 electron domains)O S – O

Unit 7: Chemical Reactions

50. Label each of the following : product, reactants, subscript, coefficient, yields

H2 + Cl2 2 HCl

51. After a chemical reaction, the mass of products is _____ equal to the original mass of reactants.

A. neverB. sometimesC. always

52. Circle ALL of the following that are TRUE about what happens in ALL chemical reactions.

A. Atoms are rearranged.B. More energy is released than absorbedC. Energy is absorbed to break the bonds of the reactants.D. Energy is released when the bonds of the products are formed.

For Questions 53-55: BALANCE the reaction using coefficients when necessary. CLASSIFY the reaction as one of the five types by writing in the blank.

decomposition synthesis combustion single-replacement double-replacement

53. single replacement __ Zn + 2 HNO3 __ Zn(NO3)2 + __ H2

54. double replacement __ CaCO3 + 2 HCl __ CaCl2 + __ H2CO3

55. decomposition 2 K2SO3 2 K2S + 3 O2

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reactants yields product

subscript coefficient

56. Define oxidation: losing electrons

57. Define reduction: gaining electrons

58. Identify which element is being oxidized and which element is being reduced.

Mg + S MgS Mg is oxidized, S is reduced

59. Consider the double-replacement reaction below.

Ba(HCO3)2 + FeSO4 ______ + ______

Which of the following would you expect to be one of the products for this reaction?

A. BaFe B. SO4Ba C. FeBa D. Fe(HCO3)2

Unit 8: Chemical Composition

60. What is the percent composition of carbon in Na2CO3 ?

12.01 g C . x 100 = 11.33 % C105.99 g total

61. What is the empirical formula of a compound that is 74.8% C and 25.2% H ?

74.8% C 74.8 g C x 1 mol C = 6.23 mol C ÷ 6.23 = 1 C 12.01 g C

25.2% H 25.2 g H x 1 mol H = 24.9 mol H ÷ 6.23 = 4 H 1.01 g H

62. What is the empirical formula of a compound that is 36.84% N and 63.16% O by mass?

36.84% N 36.84 g N x 1 mol N = 2.630 mol N ÷ 2.630 = 1 N x 2 = 2 N 14.01 g N

63.16% O 63.16 g O x 1 mol O = 3.948 mol O ÷ 2.630 = 1.5 O x 2 = 3 O 16.00 g O

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CH4

N2O3

63. A compound has an empirical formula of CH2 and a molecular weight of about 56 g/mol. What is the molecular formula of the compound?

molecular mass = multiple 56 = 4 4(CH2) = C4H8

empirical mass 14.03

64. A certain compound has an empirical formula of C2H4O and a molecular weight of about 44 g/mol. What is the molecular formula of this molecule?

molecular mass = multiple 44 = 1 1(C2H4O) = C2H4O empirical mass 44.06

65. What is the empirical formula for the molecule C8H18O2 ?

C4H9O

Unit 9: Stoichiometry

Ca(OH)2 + FeCl2 CaCl2 + Fe(OH)2

74 g 127 g 111 g ? g66. According to the reaction above, how many grams of Fe(OH)2 should be formed if both reactants

are reacted completely? 90 g

67. 4 Na + O2 2 Na2Oa) How many moles of sodium will react completely with 3.82 moles of oxygen (O2)?

3.82 mol O2 x (4 mol Na) = 15.3 mol Na (1 mol O2)

b) How many moles of Na2O can be produced from 13.5 mol Na?

13.5 mol Na x (2 mol Na2O) = 6.75 mol Na2O (4 mol Na)

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68. C2H4 + 3 O2 2 CO2 + 2 H2Oa) How many grams of C2H4 (28.06 g/mol) are needed to produce 66.7 grams of water?

66.7 g H2O x (1 mol H 2O) x (1 mol C 2H4) x (28.06 g C2H4) = 51.9 g C2H4

(18.02 g H2O) (2 mol H2O) (1 mol C2H4)

b) How many grams of H2O can form from 2.56 g C2H4 (28.06 g/mol)?

2.56 g C2H4 x (1 mol C 2H4) x (2 mol H 2O ) x (18.02 g H2O) = 3.29 g H2O (28.06 g C2H4) (1 mol C2H4) (1 mol H2O)

c) If 3.0 g of H2O actually forms what is the percent yield?

3.0 g H2O actual yield x 100 = 91 % 3.29 g H2O theoretical yield

2 C(s) + O2(g) 2 CO(g)

69. Solid carbon reacts with oxygen gas to produce carbon monoxide gas as shown above. Identify the limiting reactant when 3.25 mol O2 is added to 48.5 g C .

48.5 g C x (1 mol C) x (1 mol O2) = 2.02 mol O2 needed(12.01 g C) (2 mol C)

(3.25 mol O2) available > needed (2.02 mol O2)

O2 is excess

C is limiting

P4(s) + 6 H2(g) 4 PH3(g) 70. Solid phosphorus reacts with hydrogen gas to form phosphine gas (PH3) as shown above. What is

the theoretical yield of PH3 (34.00 g/mol) that can be formed when 6.20 g P4 reacts with 4.00 g H2?

6.20 g P4 x (1 mol P4 ) x (6 mol H2 ) x (2.02 g H2 ) = 0.607 g H2 needed (123.88 g P4) (1 mol P4) (1 mol H2)

(4.00 g H2) available > needed (0.607 g H2)

H2 is excess

P4 is limiting

6.20 g P4 x (1 mol P 4) x (4 mol PH 3) x (34.00 g PH3) = 6.81 g PH3 theoretically produced (123.88 g P4) (1 mol P4) (1 mol PH3)

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UNIT 10: Energy and States of Matter

71. Label each of the following processes as endothermic or exothermic:

endothermic solid ice melting into liquid

endothermic liquid water evaporating into gas

exothermic water vapor condensing into liquid

exothermic liquid water freezing into solid ice

72. The energy in the system at the beginning of an endo thermic reaction is less than the

energy in the system at end of the reaction because the system gains / loses energy.

73. Sketch a potential energy diagram for an exothermic reaction.

Label the reactants and products.

Circle the correct statement .

In an exothermic reaction, heat is absorbed from / released to the surroundings,

and the surroundings cool down / warm up.

Touching a beaker containing this reaction would feel cool / warm .

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potential energy

reaction progress

R

P

Eact

E

74. Sketch a potential energy diagram for an endothermic reaction.Label the reactants and products.

Circle the correct statement .

In an endothermic reaction, heat is absorbed from / released to the surroundings,

and the surroundings cool down / warm up.

Touching a beaker containing this reaction would feel cool / warm .

75. There are strong attractions between polar water molecules which cause water to have all of the following properties EXCEPT ____.

A. surface tensionB. liquid of greater density than solid (ice)C. attraction to nonpolar moleculesD. higher boiling point

76. Hydrogen sulfide (H2S) boils at –60oC. Even though water is a smaller molecule that should become a gas easier than H2S, water doesn’t boil until it reaches 100oC.

Why do water molecules require a much higher temperature to become a gas?

The H-bonds between water molecules are stronger intermolecular attractions than the dipole-dipole forces of H2S. The stronger attractions require more energy to be overcome.

16

potential energy

reaction progress

R

PEact

E

Use the diagram below to answer questions 77 & 78.

77. What state(s) of matter are present from:

A-B solid

B-C solid and liquid

C-D liquid

D-E liquid and gas

E-F gas

78. Identify sections on the diagram in which the following occur:

Kinetic energy increasing: A-B C-D E-F

Potential energy increasing: B-C D-E

79. What is the boiling point and freezing point of water in Celsius and in Kelvin? 100oC , 373 K

For #80, you may use the relationship:

80. A cup of water contains 55 g of water at a temperature of 21.4oC. How much heat must be removed from the water to lower its temperature to 2.5oC? (the specific heat of water is 4.18 J/goC)

q = mcΔTq = (55) (4.18) (2.5 – 21.4)q = –4300 J

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q = mcT

For #81-86, refer to the phase diagram below for water.

81. The phase change from A to C is called sublimation and from C to B is condensation.

82. The boiling point of the substance is shown at Point _2_ which is the point at which

liquid and gas phases coexist in equilibrium.

83. Point 4 represents the triple point, which is the point at which…

solid, liquid, and gas phases coexist in equilibrium

84. The critical point is shown at Point _5_ which represents the temperature above which a

liquid could not exist and the pressure above which a gas could not exist.

85. A sample of the substance is held constant at a temperature of 300 K while the

pressure is decreased from 10 atm to 0.01 atm. The phase change that occurs is vaporization.

86. A sample of the substance is held constant at a pressure of 1 atm while the temperature

is increased from 150 K atm to 250 K. The phase change that occurs is melting.

__________________________________________________________________________________Unit 11: Gas Laws

87. List 3 variables and how you would change them to increase the pressure of a gas.

a) increase moles of gas in container

b) increase temperature of gas

c) decrease volume of gas

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2

3

4

5

Label each section on the

diagram

(A, B, C)

with the correct phase

(s, l, g)

A

B

C

1

solid

liquid

gas

(collisions with walls of container)

88. What happens to gas pressure if its volume is decreased? increase or decrease

89. What happens to the volume of a gas if the pressure is increased? increase or decrease

90. What happens to the volume of a gas if the temperature is increased? increase or decrease

91. What happens to the temperature of a gas if the volume is increased? increase or decrease

For #92-93, you may refer to the following relationships:

92. 2.00 L of a gas at 2.50 atm is compressed to a volume of 0.50 L. What is the pressure if the temperature is constant?

V is halved.V and P are inversely proportional.

Therefore, P is doubled from 2.50 atm to 5.00 atm.

93. 5.00 L of a gas at 273C and 760 mmHg is stored in a flexible container. What is the volume at STP?

K = 273 + 273 = 546 K STP: T = 273 K P = 1 atm (760 mmHg)

T is halved. P is constant.T and V are directly proportional.

Therefore, V is halved from 5.00 L to 2.50 L.

94. A 3.50 mol sample of a gas at 305 K and a pressure of 800 mmHg. What is the volume of the gas?

800 mmHg x 1 atm = 1.05 atm 760 mmHg

PV = nRT

(1.05 atm) V = (3.50 mol)(0.08206)(305 K) V = (3.50)(0.08206)(305) V = 83.4 L (1.05)

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K = oC + 273

1 atm = 760 mmHg = 101.3 kPaPV = nRT R = 0.08206 L∙atm mol∙K

STP = 273 K & 1.0 atm Ideal gas at STP = 22.4 L∙mol–1

For #95-96, you may refer to the following relationships:

95. Consider the reaction: N2(g) + 3 H2(g) 2 NH3(g)

What is the total number of liters of NH3 produced when 11.2 liters of H2 reacts completely at STP?

11.2 L H2 x 1 mol H2 x 2 mol NH3 x 22.4 L NH3 = 7.47 L NH3

22.4 L H2 3 mol H2 1 mol NH3

96. Consider the reaction: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

When 48.6 grams of Mg(s) reacts completely, what is the volume of H2(g) produced if the reaction occurs at 22.0°C and 0.910 atm?

48.6 g Mg x 1 mol Mg x 1 mol H2 = 2.00 mol H2

24.31 g Mg 1 mol Mg

PV = nRT(0.910 atm) V = (2.00 mol)(0.08206)(295 K) V = (2.00)(0.08206)(295) V = 53.2 L H2

(0.910)

UNIT 12: Solutions

97. Complete the phrase that describes the types of substances that will dissolve in each other:

solvent dissolves solute.

Therefore, polar solvents (like water) can dissolve polar solutes (like alcohols, sugars, ionic compounds, etc.).

But nonpolar solutes (like fats, oils, hydrocarbons, etc.) will dissolve in nonpolar solvents.

20

K = oC + 273

1 atm = 760 mmHg = 101.3 kPaPV = nRT R = 0.08206 L∙atm mol∙K

STP = 273 K & 1.0 atm Ideal gas at STP = 22.4 L∙mol–1

98. Label each solution as saturated, unsaturated, or supersaturated based on the effect caused by adding more particles of solute:

unsaturated supersaturated saturated

For # 99 - 103, consider the graph of solubility curves below.

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99. Which substance has the lowest solubility at 20oC?

KClO3

100. If 40 g of KCl is dissolved in 100 g of water

at 80oC, then the solution is unsaturated.

101. If 80 g of KNO3 is dissolved in 100 g of water

at 40oC, then the solution is supersaturated.

102. If 80 g of KNO3 is dissolved in 100 g of water

at 50oC, then the solution is saturated.

103. How many grams of NH4Cl can dissolve in 100 g

of water at 70oC? 60 g

For #104-107, you may use the following relationships:

104. What is the molarity of a solution containing 1.98 moles NaCl of solute in 775 mL of solution?

1.98 mol = 2.55 M 0.775 L

105. How many mL of a 0.150 M NaBr solution are needed to make 0.100 L of 0.0500 M NaBr?

M1V1 = M2V2

(0.150 M)(V1) = (0.0500 M)(100 mL)

V1 = 33.3 mL

Cu(s) + 2 AgNO3(aq) Cu(NO3)2(aq) + 2 Ag(s)

106. Solid copper is added to an aqueous solution of silver nitrate shown in the reaction above. What volume of 1.65 M AgNO3(aq) solution is needed to react completely with 0.250 moles of Cu(s)?

0.250 mol Cu x 2 mol AgNO3 x 1 L AgNO3 = 0.303 L AgNO3

1 mol Cu 1.65 mol AgNO3

2 Al(s) + 3 ZnCl2(aq) 2 AlCl3(aq) + 3 Zn(s)

107. Solid aluminum is added to an aqueous solution of zinc chloride shown in the reaction above. What volume of 1.20 M ZnCl2(aq) solution is needed to react completely with 9.44 grams of Al(s)?

9.44 g Al x 1 mol Al x 3 mol ZnCl2 x 1 L ZnCl2 = 0.437 L ZnCl2 26.98 g Al 2 mol Al 1.20 mol ZnCl2

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M1V1 = M2V2 mol of solute . Liter of solutionM =

UNIT 13: Acids and Bases

108. List 3 properties of acids: List 3 properties of bases:

taste: sour taste: bitterlitmus color: red litmus color: bluereacts with metals to form H2 gas feels slippery

109. What property to acids and bases both have in common? electrolytes(conduct electricity in solution due to ions)

110. A conjugate acid-base pair is related by the transfer of a proton (H + ion) between them.

111. Circle the conjugate base in the following reaction? HNO3 + H2O NO3 + H3O+

112. Which of the following is a conjugate acid-base pair? (circle one) (A) HSO4

– and PO43–

(B) H2SO4 and HSO4–(C) H2O and HCl(D) CO and CO2

For #113 - 114, you may use the following relationships:

113. Determine the pH and Label each of the following as acidic (A), basic (B), or neutral (N).

pH acidic ( A ), basic ( B ), neutral ( N )

a) hydrogen ion concentration of 1 x10–3 M 3 A

b) [H+] = 1 x10–9 M 9 B

c) [OH–] = 1 x10–8 M 6 A

b) [H3O+] = 1.0 x10–7 M 7.0 N

c) 0.150 M hydronium ion 0.824 A

114. Calculate the pOH for a solution of pH = 1.80.

pH + pOH = 14.00

1.80 + pOH = 14.00

pOH = 12.20

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pH = –log[H+] pH + pOH = 14 Kw = [H+][OH–] = 1 x 10–14

For #115 - 118, you may use the following relationships:

115. Calculate the pH for a solution of 1 x 10–9 OH.

pOH = –log(1 x 10–9) pH + pOH = 14

pOH = 9 pH + 9 = 14

pH = 5

116. Calculate the [H+] for a solution of 9.16 x 10–8 M OH.

[H+]∙[OH–] = 1.0 x 10–14

[H+]∙(9.16 x 10–8) = 1.0 x 10–14

[H+] = 1.09 x 10–7 M

117. 10.0 mL of NaOH of unknown concentration is titrated by adding exactly 15.8 mL of 0.150 M HCl to completely neutralize the base.

a. Write the balanced equation for the neutralization of this reaction.

NaOH + HCl H2O + NaCl

b. What was the concentration of NaOH?

0.0158 L HCl x 0.150 mol HCl x 1 mol NaOH = 0.00237 mol NaOH 1 L HCl 1 mol HCl

0.00237 mol NaOH = 0.237 M NaOH 0.0100 L NaOH

118. 22.5 mL of Sr(OH)2 of unknown concentration is titrated by adding exactly 25.0 mL of 0.0100 M HCl to completely neutralize the base. What was the concentration of Sr(OH)2?

a. Write the balanced equation for the neutralization of this reaction.

Sr(OH)2 + 2 HCl 2 H2O + SrCl2

b. What was the concentration of Sr(OH)2?

0.0250 L HCl x 0.100 mol HCl x 1 mol Sr(OH)2 = 0.00125 mol Sr(OH)2 1 L HCl 2 mol HCl

0.00125 mol Sr(OH) 2 = 0.0556 M Sr(OH)2

0.0225 L Sr(OH)2

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pH = –log[H+] pH + pOH = 14 Kw = [H+][OH–] = 1 x 10–14

UNIT 14: Reaction Rates and Equilibrium

119. Give one reason why increasing the concentration of reactants increases the reaction rate:

more collisions between reactant particles

120. List two reasons why a reaction rate increases with an increase in temperature.

more frequent collisions between reactant particlesmore frequent collisions of greater energy (so more particles have the activation energy).

121. The smaller the particle size of a reactant, the greater surface area available to react.

This results in more frequent collisions between reactant particles and a higher reaction rate.

122.Using the diagram above, label A, B, C, and D.

[product(s), uncatalyzed reaction, catalyzed reaction, reactant(s)]

A: reactants

B: uncatalyzed reaction

C: catalyzed reaction

D: products

Explain the difference between reaction B and reaction C. Which reaction will be faster? Why?

The different heights of reactions B and C represent the different activation energies of the reaction.

Reaction _ C _ is a catalyzed reaction, while Reaction _ B _ is an uncatalyzed reaction.

Reaction B occurs at a slower rate because it has a higher activation energy.

123.When a reaction has reached equilibrium, which of the following is TRUE?

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reaction progress

A. the reaction stopsB. the forward reaction continuesC. the reverse reaction continuesD. BOTH the forward and reverse reactions continue

124. A chemical reaction in is dynamic equilibrium when the relative forward and reverse rates of the reaction are equal, and the total amounts of reactant and product are constant.

R(g) ⇄ P(g)

125. Two gases were placed in a container and allowed to reach equilibrium shown in the equation above. The diagram below shows how the concentrations of gases R and P in this system changed over time.

At time A, the concentrations of R and P are:constant / equal / zero .

At time B, the forward rate of reaction is:greater than / equal to / less thanthe reverse rate of the reaction.

At time C, concentrations of R and P are: constant / equal / zero .

__________________________________________________________________________________UNIT 15: Nuclear Chemistry

126. Identify each as a chemical or a nuclear reaction.

a) Involves electrons. chemical

b) Mass of reactants = mass of products. chemical

c) One element changes into another. nuclear

d) Some matter is converted into energy (E = mc2). nuclear

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time (s)

conc

entr

atio

n (M

) [R]

A B C

[P]