Grade : 11 · _____ 2. If gas A has a higher temperature than gas B, then the particles in gas A a. have greater average kinetic energy than those in gas B. b. have less average kinetic

Chemistry Grade : 11

Term-3/Final Exam

Revision Sheet

Exam Date: Tuesday 12/6/2018

CCS:Chem.6a,6b,6c,6d,6e,6f,7a,7b,7d,7c,7e,7f,1g

Chapter(12):Solutions Sections:1,2,3 Textbook pages 378 to 408

Chapter(16):Reaction energy Sections:1,2 Textbook pages 500 to 526

Solutions

* write the letter of the term or phrase that best completes each statement :

_____ 1. Agitation prevents settling in a(n)

a. alloy.

b. homogeneous mixture.

c. suspension.

d. gaseous mixture.

_____ 2. All of the following are heterogeneous mixtures except

a. whole wheat bread.

b. granite.

c. tap water.

d. an oil-water mixture.

_____ 3. What is the concentration of a 100. mL aqueous solution that contains 1.00 g

KCl (molar mass = 74.55 g/mol)?

a. 1.34 M KCl

b. 0.134 M KCl

c. 0.0134 M KCl

d. 0.001 34 M KCl

_____ 4. To determine the molarity of an HCl solution, you need to know the number

of

a. grams of HCl in 1 106 g of solution.

b. moles of HCl dissolved in the total moles of solution.

c. moles of HCl in 1 L of solution.

d. moles of HCl dissolved in 1 kg of solvent.

_____ 5. What type of solute-solvent combination is carbon dioxide in water?

a. gas-liquid

b. liquid-gas

c. liquid-liquid

d. cannot be determined

_____ 6. What is the molarity of a solution that contains 0.202 mol KCl (molar mass =

74.55 g/mol) in 7.98 L of solution?

a. 0.0132 M KCl

b. 0.0253 M KCl

c. 0.459 M KCl

d. 1.36 M KCl

Use this figure to answer questions (7 and 8):

_____ 7. A solution containing 35 g of Li2SO4 dissolved in 100 g of water is heated

from 10°C to 90°C. According to information in the figure, this temperature

change would result in

a. an additional 5 g of Li2SO4 in solution.

b. an additional 30 g of Li2SO4 in solution.

c. 5 g of Li2SO4 precipitate.

d. no change in Li2SO4 concentration.

_____ 8. According to saturation curves shown in the figure, which of the following

solutions is supersaturated?

a. 40 g of NaCH3COO in 100 g of water at 40°C

b. 140 g of NaCH3COO in 100 g of water at 80°C

c. 80 g of NaCH3COO in 100 g of water at 40°C

d. 80 g of NaCH3COO in 200 g of water at 40°C

_____ 9. In 100 mL of cold water, 35 g of NaCl will dissolve, but 70 g will not. This

observation implies that

a. solubility depends on temperature.

b. in order to dissolve more NaCl, you must increase the pressure.

c. solubility depends on the amounts of solute and solvent present.

d. NaCl is not easily hydrated.

_____ 10. What is the molarity of a solution that contains 125 g NaCl (molar mass =

58.44 g/mol) in 4.00 L solution?

a. 0.535 M NaCl

b. 2.14 M NaCl

c. 8.56 M NaCl

d. 31.3 M NaCl

_____ 11. Which of the following is soluble in water?

a. potassium nitrate

b. silver

c. benzene

d. carbon tetrachloride

_____ 12. In a solution at equilibrium,

a. no dissolution occurs.

b. the rate of dissolution is less than the rate of crystallization.

c. the rate of dissolution is greater than the rate of crystallization.

d. the rate of dissolution and the rate of crystallization are equal.

_____ 13. A dissolved solute that does not form ions is

a. a nonelectrolyte.

b. a weak electrolyte.

c. a strong electrolyte.

d. insoluble.

_____ 14. How many moles of HCl (molar mass = 36.46 g/mol) are present in 0.70 L of

a 0.33 M HCl solution?

a. 0.23 mol

b. 0.28 mol

c. 0.38 mol

d. 0.47 mol

_____ 15. If the temperature stays the same, the solubility of gases in liquids

a. increases with increasing pressure.

b. cannot reach equilibrium.

c. decreases with increasing pressure.

d. does not depend on pressure.

_____ 16. A NaOH solution contains 1.90 mol of NaOH (molar mass = 40.00 g/mol), and

its concentration is 0.555 M. What is its volume?

a. 0.623 L

b. 0.911 L

c. 1.05 L

d. 3.42 L

_____ 17. Which solution would be least likely to carry an electric current?

a. NaCl

b. HCl

c. C6H12O6

d. CsI

_____ 18. Which type of mixture contains the smallest particles?

a. emulsions

b. solutions

c. suspensions

d. colloids

_____ 19. Which does not affect the rate at which a solid solute dissolves?

a. the vapor pressure of the solvent

b. the temperature of the solvent

c. the surface area of the solid

d. the speed at which the solution is stirred

_____ 20. Which pair of compounds is immiscible?

a. water and alcohol

b. water and toluene

c. toluene and gasoline

d. benzene and gasoline

_____ 21. Under which conditions is more CO2 dissolved in a carbonated beverage?

a. in a glass at room temperature

b. in a bottle that has been left uncapped in the refrigerator

c. in a glass with ice cubes

d. in an unopened bottle in the refrigerator

_____ 22. A solid is dissolved in some water at 25°C in a beaker. The outside of the

beaker feels cold to the touch. What does this tell you about this solution?

a. The enthalpy of solution for the solid is negative.

b. The solution has not come to equilibrium.

c. The solution must be heated to continue the dissolving process.

d. The enthalpy of solution for the solid is positive.

_____ 23. What is the molality of an aqueous NaOH solution made with 5.00 kg of

water and 3.6 mol NaOH (molar mass = 40.00 g/mol)?

a. 3.6 m NaOH

b. 1.4 m NaOH

c. 0.72 m NaOH

d. 0.090 m NaOH

_____ 24. How much methanol, CH3OH (molar mass = 32.05 g/mol), is needed to make

a 0.90 m solution in 250 g of water?

a. 0.14 g CH3OH

b. 7.2 g CH3OH

c. 100 g CH3OH

d. 220 g CH3OH

*Answer the following questions in the space provided:

1. Solid CaCl2 does not conduct electricity. Explain why it is considered to be an

electrolyte.

______________________________________________________________

______________________________________________________________

2. Explain the following statements at the molecular level:

a. Generally, a polar liquid and a nonpolar liquid are immiscible.

______________________________________________________________

______________________________________________________________

b. Carbonated soft drinks taste flat when they warm up.

______________________________________________________________

______________________________________________________________

3. An unknown compound is observed to mix with toluene, C6H5CH3, but not with

water:

a. Is the unknown compound ionic, polar covalent, or nonpolar covalent? Explain

your answer.

______________________________________________________________

______________________________________________________________

b. Suppose the unknown compound is also a liquid. Will it be able to dissolve table

salt? Explain why or why not.

______________________________________________________________

______________________________________________________________

*Problems:

4. Consider 500. mL of a 0.30 M CuSO4 solution.

__________________ a. How many moles of solute are present in this solution?

__________________ b. How many grams of solute were used to prepare this

solution?

5. a. If a solution is electrically neutral, can all of its ions have the same type of charge?

Explain your answer.

______________________________________________________________

______________________________________________________________

__________________ b. The concentration of the OH ions in pure water is known

to be 1.0 1027 M. How many OH ions are present in

each milliliter of pure water?

6. 90. g of CaBr2 are dissolved in 900. g of water.

__________________ a. What volume does the 900. g of water occupy if its density

is 1.00 g/mL?

__________________ b. What is the molality of this solution?

Reaction energy

* write the letter of the term or phrase that best completes each statement :

_____ 1. A chemical change is likely to occur when

a. energy and randomness both increase.

b. energy and randomness both decrease.

c. energy increases and randomness decreases.

d. energy decreases and randomness increases.

_____ 2. If gas A has a higher temperature than gas B, then the particles in gas A

a. have greater average kinetic energy than those in gas B.

b. have less average kinetic energy that those in gas B.

c. contain the same average kinetic energy as those in gas B.

d. may contain more, less, or the same average kinetic energy as those in gas B.

_____ 3. An example of increasing entropy is the

a. formation of crystals from a solution.

b. formation of 1 mol of gas from 1 mol of one reactant gas and 1 mol of another

reactant gas.

c. dissolution of crystals in a solution.

d. None of the above

_____ 4. The amount of energy absorbed by a system as heat during a process at constant

pressure is the change in

a. enthalpy.

b. entropy.

c. temperature.

d. Gibbs free energy.

_____ 5. To determine the amount of energy as heat associated with the change taking

place in a calorimeter, the information that is not needed is the

a. specific heat of the calorimeter.

b. air temperature outside of the calorimeter.

c. specific heat of the water.

d. change in temperature of the water.

_____ 6. How much energy is needed to raise the temperature of 40.0 g of argon from 25C

to 40C? The specific heat capacity of argon is 0.520 J/(g·K).

a. 20.8 J

b. 208 J

c. 312 J

d. 416 J

_____ 7. An increase in temperature in a system causes a(n)

a. increase in entropy.

b. decrease in entropy.

c. increase in specific heat.

d. decrease in specific heat.

_____ 8. The change in Gibbs free energy for a substance can be found by the expression

a. H TS.

b. H + TS.

c. S TH.

d. H + TH.

_____ 9. A chemical reaction occurs spontaneously when G is

a. positive.

b. negative.

c. zero.

d. constant.

_____ 10. A chemical reaction is exothermic when H is

a. positive.

b. negative.

c. zero.

d. constant.

_____ 11. The term thermodynamics refers to the study of

a. energy changes.

b. only physical changes.

c. only chemical changes.

d. None of the above

_____ 12. What is the energy change per gram of ice when an iceberg composed of pure

water, cp = 2.06 J/(g·K), is heated from 25C to 15C?

a. 0.21 J

b. 21 J

c. 210 J

d. impossible to calculate without knowing the mass of the iceberg

_____ 13. The standard enthalpy of formation of Cl2 is

a. positive.

b. negative.

c. zero.

d. impossible to determine without more information.

Use the data in the following table to answer questions (14–17):

Standard Enthalpies and Entropies

Substance Standard enthalpy of

formation (kJ/mol)

Standard

entropy (J/mol)

H2(g) 0 130.7

O2(g) 0 205.1

H2O(g) 242 188.7

CO2(g) 393 213.8

C(s) (graphite) 0 5.7

_____ 14. For the reaction represented by the

equation )(OH)(O2

1)(H 222 ggg , the values of H and S are

a. 242 kJ and 148.3 J.

b. 242 kJ and 148.3 J.

c. 242 kJ and 46.4 J.

d. 242 kJ and 44.6 J.

_____ 15. The reaction represented by the equation C(s) + O2(g) CO2(g) is

a. spontaneous.

b. not spontaneous.

c. spontaneous at 298 K, but not at 25C.

d. impossible to determine without more information.

_____ 16. For the reaction represented by the equation )(O2

1)(H)(OH 222 ggg the

value of G is

a. always positive.

b. always negative.

c. positive at low temperatures and negative at high temperatures.

d. negative at low temperatures and positive at high temperatures.

_____ 17. The reaction represented by the equation CO2(g) C(s) + O2(g) is always

a. exothermic.

b. endothermic.

c. isothermic.

d. spontaneous.

_____ 18. For an exothermic reaction,

a. G is always positive.

b. H is always negative.

c. S is always positive.

d. All of the above

_____ 19. To determine the specific heat, you must know all of the following factors except

a. mass.

b. amount of energy needed to raise the temperature.

c. volume.

d. temperature change.

_____ 20. In order for an endothermic reaction to occur spontaneously, what must happen?

a. The enthalpy must be negative.

b. The entropy must be positive.

c. The entropy must be negative.

d. The change in Gibbs free energy must be positive.

_____ 21. Energy transferred as heat always moves spontaneously from matter

a. at a lower temperature to matter at a higher temperature.

b. at a higher temperature to matter at a lower temperature.

c. at an enthalpy of zero to matter with a negative enthalpy.

d. None of the above

_____ 22. H =

a. Hreactants Hproducts

b. Hreactants + Hproducts

c. Hproducts Hreactants

d. Hproducts Hreactants

_____ 23. The change in energy represented by a thermochemical equation is always

a. equal to the number of moles of substances undergoing a change.

b. directly proportional to the number of moles of substances undergoing a

change.

c. indirectly proportional to the number of moles of substances undergoing a

change.

d. less than the number of moles of substances undergoing a change.

_____ 24. The melting of ice is always a(n)

a. exothermic reaction.

b. endothermic reaction.

c. negative entropic reaction.

d. catalyzed reaction.

_____ 25. For a reaction that has a H of 23.0 kJ and a S of 130 J/K, what is G at

25.0C?

a. 15.7 kJ

b. +15.7 kJ

c. 61.7 kJ

d. +61.7 kJ

*Answer the following questions in the space provided:

1. Describe Hess’s law.

______________________________________________________________

______________________________________________________________

2. What determines the amount of energy absorbed by a material when it is heated?

______________________________________________________________

______________________________________________________________

3. Describe what is meant by enthalpy of combustion and how a combustion calorimeter

measures this enthalpy.

______________________________________________________________

______________________________________________________________

4. The following equation represents a reaction that is strongly favored in the forward

direction:

2C7H5(NO2)3(l) + 12O2(g) 14CO2(g) + 5H2O(g) + 3N2O(g) + energy

Why would G be negative in the above reaction?

______________________________________________________________

______________________________________________________________

*Problems:

5. Consider the following equation and data: 2NO2(g) N2O4(g)

0

fH of N2O4 = +9.2 kJ/mol

0

fH of NO2 = +33.2 kJ/mol

0G = 4.7 kJ/mol N2O4

__________________ Use Hess’s law to calculate 0H for the above reaction.

6. Calculate the energy needed to raise the

temperature of 180.0 g of water from 10.0C to

40.0C. The specific heat of water is 4.18 J/(K · g).

7. a. Calculate the change in Gibbs free energy for the

following equation at 25C.

2H2O2(l) 2H2O(l) + O2(g)

Given H = 196.0 kJ/mol

S = +125.9 J/mol

b. Is this reaction spontaneous?

8. For elements in their standard state, the value of Hf0 is _____.

9. The formation and decomposition of water can be represented by the following

thermochemical equations:

_________________ a. Is energy being taken in or is it being released as liquid H2O

decomposes?

_________________ b. What is the appropriate sign for the enthalpy change in this

decomposition reaction?

10.__________________ If 200. g of water at 20°C absorbs 41 840 J of energy,

what will its final temperature be?

11. __________________ Aluminum has a specific heat of 0.900 J/(g °C). How much

energy in kJ is needed to raise the temperature of a 625 g block

of aluminum from 30.7°C to 82.1°C?

12. The products in a reaction have an enthalpy of 458 kJ/mol, and the reactants have an

enthalpy of 658 kJ/mol.

_________________ a. What is the value of H for this reaction?

_________________ b. Which is the more stable part of this system, the reactants or

the products?

13. The enthalpy of combustion of acetylene gas is 1301.1 kJ/mol of C2H2.

a. Write the balanced thermochemical equation for the complete combustion of C2H2.

______________________________________________________________

_________________ b. If 0.25 mol of C2H2 reacts according to the equation in part a,

how much energy is released?

_________________ c. How many grams of C2H2 are needed to react, according to

the equation in part a, to release

3900 kJ of energy?

14. __________________ Determine the H for the reaction between Al and Fe2O3,

according to the equation 2Al + Fe2O3 → Al2O3 + 2Fe. The

enthalpy of formation of Al2O3

is 1676 kJ/mol. For Fe2O3 it is 826 kJ/mol.

15. __________________ Use the data in Appendix Table A-14 of the text to determine

the H for the following equation.

2H2O2(l) → 2H2O(l) + O2(g)

16. For the following examples, state whether the change in entropy favors the forward or

reverse reaction:

__________________ a. )HCl( )HCl( gl

__________________ b. )(OHC )(OHC 61266126 saq

__________________ c. )(3H )(N )(2NH 223 ggg

__________________ d. )(HC )(H3C 12642 lg

17.__________________ a. Write an equation that shows the relationship between

enthalpy, H, entropy, S, and free energy, G.

b. For a reaction to occur spontaneously, the sign of G should be ______.

18. Consider the following equation:

)(OH )(NH )O(H )(NH -

423 aqaqlg energy

__________________ a. The enthalpy factor favors the forward reaction. True or

False?

b. The sign of TH0 is negative. This means the entropy factor favors the ______.

c. Given that G0 for the above reaction in the forward direction is positive, which term

is greater in magnitude and therefore predominates, TS or H?

______________________________________________________________

______________________________________________________________

19. Consider the following equation for the vaporization of water:

H2O(l) = H2O(g) H = + 40.65 kJ/mol at 100°C

__________________ a. Is the forward reaction exothermic or endothermic?

__________________ b. Does the enthalpy factor favor the forward or reverse

reaction?

__________________ c. Does the entropy factor favor the forward or reverse reaction?