SAT Subject Test Chemistry

PART 3

PRACTICE TESTS

Practice Test 3

Note: For all questions involving solutions and/or chemical equations, assume that the system is in water unless otherwise stated.

Reminder: You may not use a calculator on these tests.

The following symbols have the meanings listed unless otherwise noted.

H   =   enthalpy

g   =   gram(s)

M   =   molar

J   =   joules(s)

n   =   number of moles

kJ   =   kilojoules

P   =   pressure

L   =   liter(s)

R   =   molar gas constant

mL   =   milliliter(s)

S   =   entropy

mm   =   millimeter(s)

T   =   temperature

mol   =   mole(s)

V   =   volume

V   =   volt(s)

atm   =   atmosphere

 

Part A

Directions: Every set of the given lettered choices below refers to the numbered statements or formulas immediately following it. Choose the one lettered choice that best fits each statement or formula and then fill in the corresponding oval on the answer sheet. Each choice may be used once, more than once, or not at all in each set.

Questions 1–4 refer to the following diagram:

1. The activation energy of the forward reaction is shown by

2. The activation energy of the reverse reaction is shown by

3. The heat of the reaction for the forward reaction is shown by

4. The potential energy of the reactants is shown by

Questions 5–7 refer to the following diagram:

5. To plate silver on the spoon, the position to which the wire from the spoon must be connected

6. The position of the anode

7. The position from which silver that is plated out emerges

Questions 8–11 match the following equations to the appropriate descriptions:

8. This equation shows the volume decreasing as the pressure is increased when the temperature is held constant. It is an example of Boyle’s Law.

9. This equation shows the pressure increasing as the temperature is increased when the volume is held constant. It is an example of Gay-Lussac’s Law.

10. This equation shows the volume increasing as the temperature is increased when the pressure is held constant. It is an example of Charles’s Law.

11. This equation shows that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.

Questions 12–14

(A) 1

(B) 2

(C) 3

(D) 4

(E) 5

12. When the following equation: Cu(s) + HNO3(aq) → Cu(NO3)2(aq) + H2O() + NO(g) is balanced, what will be the coefficient, in the lowest whole number, of Cu?

13. If 6 moles of Cu reacted according to the above balanced equation, what will be the number of moles of NO that would be formed?

14. If Cu(NO3)2 goes into solution as ions, what will be the number of ions into which it will dissociate?

Questions 15–18

(A) Ionic substance

(B) Polar covalent substance

(C) Nonpolar covalent substance

(D) Amorphous substance

(E) Metallic network

15. MgCl2(s)

16. HCl(g)

17. CH3–CH3(g)

18. Cu(s)

Questions 19–23

(A) Brownian movement

(B) Litmus paper reaction

(C) Phenolphthalein reaction

(D) Dehydration

(E) Deliquescent

19. The reason why a blue crystal of CuSO4 · 5H2O turns white when heated

20. The zigzag path of colloidal particles in light

21. The pink color in a basic solution

22. The pink color in an acid solution

23. The adsorbtion of water to the surface of a crystal

Part B

ON THE ACTUAL CHEMISTRY TEST, THE FOLLOWING TYPE OF QUESTION MUST BE ANSWERED ON A SPECIAL SECTION (LABELED “CHEMISTRY”) AT THE LOWER LEFT-HAND CORNER OF YOUR ANSWER SHEET. THESE QUESTIONS WILL BE NUMBERED BEGINNING WITH 101 AND MUST BE ANSWERED ACCORDING TO THE FOLLOWING DIRECTIONS.

Directions: Every question below contains two statements, I in the left-hand column and II in the right-hand column. For each question, decide if statement I is true or false and if statement II is true or false and fill in the corresponding T or F ovals on your answer sheet. *Fill in oval CE only if statement II is a correct explanation of statement I.

Sample Answer Grid:

CHEMISTRY * Fill in oval CE only if II is a correct explanation of I.

I

 

II

101. Elements in the upper/left corner of the Periodic Table are active metals

BECAUSE

metals have larger ionic radii than their atomic radii.

102. A synthesis reaction that is nonspontaneous and has a negative value for its heat of reaction will not occur until some heat is added

BECAUSE

nonspontaneous exothermic re actions need enough activation energy to get them started.

103. Transition elements in a particular period may have the same oxidation number

BECAUSE

they have a complete outer energy level.

104. When a crystal is added to a supersaturated solution of itself, the crystal does not appear to change

BECAUSE

the supersaturated solution is holding more solute than its normal solubility.

105. Equilibrium is a static condition

BECAUSE

at equilibrium, the forward reaction rate equals the reverse reaction rate.

106. The ionic bond is the strongest bond

BECAUSE

ionic bonds have electrostatic attraction due to the loss and gain of electron(s).

107. In the equilibrium reaction N2(g) + 3H2(g) ↔ 2NH3(g) + heat when the pressure in the reaction chamber is increased, the reaction shifts to the right

BECAUSE

the increase in pressure causes the reaction to shift to the right to decrease the pressure since 4 volumes on the left become 2 volumes on the right.

108. If the forward reaction of an equi librium is exothermic, adding heat to the system favors the reverse reaction

BECAUSE

additional heat causes a stress on the system, and the system moves in the direction that releases the stress.

109. An element that has an electron configuration of 1s2 2s2 2p6 3s2 3p6 3d3 4s2 is a transition element

BECAUSE

the transition elements from scandium to zinc are filling the 3d orbitals.

110. The most electronegative elements in the periodic chart are found among nonmetals

BECAUSE

electronegativity is a measure of the ability of an atom to draw valence electrons to itself.

111. Basic anhydrides react in water to form bases

BECAUSE

metallic oxides react with water to form solutions that have an excess of hydroxide ions.

112. There are 3 moles of atoms in 18 grams of water

BECAUSE

there are 6 × 1023 atoms in 1 mole.

113. Benzene is a good electrolyte

BECAUSE

a good electrolyte has charged ions that carry the electric current.

114. Normal butyl alcohol and 2-butanol are isomers

BECAUSE

isomers vary in the number of neutrons in the nucleus of the atom.

115. The reaction of CaCO3 and HCl goes to completion

BECAUSE

reactions that form a precipitate tend to go to completion.

116. A large number of alpha particles were deflected in the Rutherford experiment

BECAUSE

alpha particles that came close to the nucleus of the gold atoms were deflected.

Part C

Directions: Every question or incomplete statement below is followed by five suggested answers or completions. Choose the one that is best in each case and then fill in the corresponding oval on the answer sheet.

24. What are the simplest whole-number coefficients that balance this equation?

. . . C4H10 + . . . O2 → . . . CO2 + . . . H2O

(A) 1, 6, 4, 2

(B) 2, 13, 8, 10

(C) 1, 6, 1, 5

(D) 3, 10, 16, 20

(E) 4, 26, 16, 20

25. How many atoms are present in the formula KAl(SO4)2?

(A) 7

(B) 9

(C) 11

(D) 12

(E) 13

26. All of the following are compounds EXCEPT

(A) copper sulfate

(B) carbon dioxide

(C) washing soda

(D) air

(E) lime

27. What volume of gas, in liters, would 2 moles of hydrogen occupy at STP?

(A) 11.2

(B) 22.4

(C) 33.6

(D) 44.8

(E) 67.2

28. What is the maximum number of electrons held in the d orbitals?

(A) 2

(B) 6

(C) 8

(D) 10

(E) 14

29. If an element has an atomic number of 11, it will combine most readily with an element that has an electron configuration of

(A) 1s2 2s2 2p6 3s2 3p1

(B) 1s2 2s2 2p6 3s2 3p2

(C) 1s2 2s2 2p6 3s2 3p3

(D) 1s2 2s2 2p6 3s2 3p4

(E) 1s2 2s2 2p6 3s2 3p5

30. An example of a physical property is

(A) rusting

(B) decay

(C) souring

(D) low melting point

(E) high heat of formation

31. A gas at STP that contains 6.02 × 1023 atoms and forms diatomic molecules will occupy

(A) 11.2 L

(B) 22.4 L

(C) 33.6 L

(D) 67.2 L

(E) 1.06 qt

32. When excited electrons cascade to lower energy levels in an atom,

(A) visible light is always emitted

(B) the potential energy of the atom increases

(C) the electrons always fall back to the first energy level

(D) the electrons fall indiscriminately to all levels

(E) the electrons fall back to a lower unfilled energy level

33. Mass spectroscopy uses the concept that

(A) charged particles are evenly deflected in a magnetic field

(B) charged particles are deflected in a magnetic field inversely to the mass of the particles

(C) particles of heavier mass are deflected in a magnetic field to a greater degree than lighter particles

(D) particles are evenly deflected in a magnetic field

34. The bond that includes an upper and a lower sharing of electron orbitals is called

(A) a pi bond

(B) a sigma bond

(C) a hydrogen bond

(D) a covalent bond

(E) an ionic bond

35. What is the boiling point of water at the top of Pikes Peak?

(A) It is 100°C.

(B) It is >100°C since the pressure is less than at ground level.

(C) It is <100°C since the pressure is less than at ground level.

(D) It is >100°C since the pressure is greater than at ground level.

(E) It is <100°C since the pressure is greater than at ground level.

36. The atomic structure of the alkane series contains the hybrid orbitals designated as

(A) sp

(B) sp2

(C) sp3

(D) sp3d2

(E) sp4d3

37. Which of the following is (are) true for this reaction?

       Cu + 4HNO3 → Cu(NO3)2 + 2H2O + 2NO2(g)

  I. It is an oxidation-reduction reaction.

 II. Copper is oxidized.

III. The oxidation number of nitrogen goes from +5 to +4.

(A) I only

(B) III only

(C) I and II only

(D) II and III only

(E) I, II, and III

38. Which of the following properties can be attributed to water?

  I. It has a permanent dipole moment attributed to its molecular structure.

 II. It is a very good conductor of electricity.

III. It has its polar covalent bonds with hydrogen on opposite sides of the oxygen atom, so that the molecule is linear.

(A) I only

(B) III only

(C) I and II only

(D) II and III only

(E) I, II, and III

39. All of the following statements are true for this reaction EXCEPT

HCl(g) + H2O(l) → H3O+(aq) + Cl(aq)

(A) H3O+ is the conjugate acid of H2O.

(B) Cl is the conjugate base of HCl.

(C) H2O is behaving as a Brønsted-Lowry base.

(D) HCl is a weaker Brønsted-Lowry acid than H2O.

(E) The reaction essentially goes to completion.

40. A nuclear reactor must include which of the following parts?

I. Electric generator

II. Fissionable fuel elements

III. Moderator

(A) I only

(B) III only

(C) I and II only

(D) II and III only

(E) I, II, and III

41. Which of the following salts will hydrolyze in water to form basic solutions?

  I. NaCl

 II. CuSO4

III. K3PO4

(A) I only

(B) III only

(C) I and II only

(D) II and III only

(E) I, II, and III

42. When 1 mole of NaCl is dissolved in 1,000 grams of water, the boiling point of the water is changed to

(A) 100.51°C

(B) 101.02°C

(C) 101.53°C

(D) 101.86°C

(E) 103.62°C

43. What is the structure associated with the BF3 molecule?

(A) Linear

(B) Trigonal planar

(C) Tetrahedron

(D) Trigonal pyramidal

(E) Bent or V-shaped

Questions 44 and 45 refer to the following setup:

44. What letter designates an error in this laboratory setup?

(A) A (upper part of tube)

(B) B (lower part of tube)

(C) C

(D) D

(E) E

45. If the reaction in question 44 created a gas, where would the contents of the flask be expelled under these conditions?

(A) A

(B) B

(C) C

(D) D

(E) E


46. The most active nonmetal has

(A) a high electronegativity

(B) a low electronegativity

(C) a medium electronegativity

(D) large atomic radii

(E) a deliquescent property

47. In the reaction Fe + S → FeS, which is true?

(A) Fe + 2e → Fe2+

(B) Fe → Fe2+ + 2e

(C) Fe2+ → Fe + 2e

(D) S → S2– + 2e

(E) S2– + 2e → S

48. What is the pH of a solution with a hydroxide ion concentration of 0.00001 mole/liter?

(A) –5

(B) –1

(C) 5

(D) 9

(E) 14

49. Electrolysis of a dilute solution of sodium chloride results in the cathode product

(A) sodium

(B) hydrogen

(C) chlorine

(D) oxygen

(E) peroxide

50. .....C2H4(g) + .....O2(g) → .....CO2(g) + .....H2O(l)
If the equation for the above reaction is balanced with whole-number coefficients, what is the coefficient for oxygen gas?

(A) 1

(B) 2

(C) 3

(D) 4

(E) 5

51. 5.00 liters of gas at STP have a mass of 12.5 grams. What is the molar mass of the gas?

(A) 12.5 g/mol

(B) 25.0 g/mol

(C) 47.5 g/mol

(D) 56.0 g/mol

(E) 125 g/mol

52. A compound whose molecular mass is 90.0 grams contains 40.0% carbon, 6.67% hydrogen, and 53.33% oxygen. What is the true formula of the compound?

(A) C2H2O4

(B) CH2O4

(C) C3H6O

(D) C3HO3

(E) C3H6O3

53. How many moles of CaO are needed to react with an excess of water to form 370 grams of calcium hydroxide?

(A) 1.0

(B) 2.0

(C) 3.0

(D) 4.0

(E) 5.0

54. To what volume, in milliliters, must 50.0 milliliters of 3.50 M H2SO4 be diluted in order to make 2.00 M H2SO4?

(A) 25.0

(B) 60.1

(C) 87.5

(D) 93.2

(E) 101

55. A small value of Keq indicates that equilibrium occurs

(A) at a low product concentration

(B) at a high product concentration

(C) after considerable time

(D) with the help of a catalyst

(E) with no forward reaction

56. A student measured 10.0 milliliters of an HCl solution into a beaker and titrated it with a standard NaOH solution that was 0.09 M. The initial NaOH burette reading was 34.7 milliliters while the final reading showed 49.2 milliliters.

What is the molarity of the HCl solution?

(A) 0.13

(B) 0.47

(C) 0.52

(D) 1.57

(E) 2.43

57. A student made the following observations in the laboratory:

(a) Sodium metal reacted vigorously with water while a strip of magnesium did not seem to react at all.

(b) The magnesium strip reacted with dilute hydrochloric acid faster than an iron strip.

(c) A copper rivet suspended in silver nitrate solution was covered with silver-colored stalactites in several days, and the resulting solution had a blue color.

(d) Iron filings dropped into the blue solution were coated with an orange color.

The order of decreasing strength as reducing agents is:

(A) Na, Mg, Fe, Ag, Cu

(B) Mg, Na, Fe, Cu, Ag

(C) Ag, Cu, Fe, Mg, Na

(D) Na, Fe, Mg, Cu, Ag

(E) Na, Mg, Fe, Cu, Ag

58. A student placed water, sodium chloride, potassium dichromate, sand, chalk, and hydrogen sulfide into a distilling flask and proceeded to distill. What ingredient besides water would be found in the distillate?

(A) Sodium chloride

(B) Chalk

(C) Sand

(D) Hydrogen sulfide

(E) Chrome sulfate

59. Which of these statements is NOT correct?

(A) In an exothermic reaction, ΔH is negative and the enthalpy decreases.

(B) In an endothermic reaction, ΔH is positive and the enthalpy increases.

(C) In a reaction where ΔG is negative, the forward reaction is spontaneous.

(D) In a reaction where ΔG is positive, ΔS may also be positive.

(E) In a reaction where ΔH is positive and ΔS is negative, the forward reaction is spontaneous.

60. A student filled a steam-jacketed eudiometer with 32. milliliters of oxygen and 4.0 milliliters of hydrogen over mercury. How much of which gas would be left uncombined after the mixture was sparked?

(A) None of either

(B) 3.0 mL H2

(C) 24 mL O2

(D) 28 mL O2

(E) 30. mL O2

61. What would be the total volume, in milliliters, of gases in question 60 after sparking?

(A) 16

(B) 24

(C) 34

(D) 36

(E) 40

62. How can the addition of a catalyst affect an exothermic reaction?

  I. Speed up the reaction.

 II. Slow down the reaction.

III. Increase the amount of product formed.

(A) I only

(B) II only

(C) I and II only

(D) II and III only

(E) I, II, and III

63. In which period of the periodic table is the most electronegative element found?

(A) 1

(B) 2

(C) 3

(D) 4

(E) 5

64. What could be the equilibrium constant for this reaction: aA + b cC + dD, if A and D are solids?

(A) 

(B) 

(C) 

(D) 

(E) [A]a[B]b[C]c[D]d

65. Which of the following does NOT react with a dilute solution of sulfuric acid?

(A) NaNO3

(B) Na2S

(C) Na3PO4

(D) Na2CO3

(E) NaOH

66. Which of these statements is the best explanation for the sp3 hybridization of carbon’s electrons in methane, CH4?

(A) The new orbitals are one s orbital and three p orbitals.

(B) The s electron is promoted to the p orbitals.

(C) The s orbital is deformed into a p orbital.

(D) Four new and equivalent orbitals are formed.

(E) The s orbital electron loses energy to fall back into a partially filled p orbital.

67. The intermolecular force that is most significant in explaining the variation of the boiling point of water from the boiling points of similarly structured molecules is

(A) hydrogen bonding

(B) van der Waals forces

(C) covalent bonding

(D) ionic bonding

(E) coordinate covalent bonding

68. If K for the reaction H2 + I2  2HI is equal to 45.9 at 450°C, and 1 mole of H2 and 1 mole of I2 are introduced into a 1-liter box at that temperature, what will be the expression for K at equilibrium?

(A) 

(B) 

(C) 

(D) 

(E) 

69. What is the molar mass of a nonionizing solid if 10. grams of this solid, dissolved in 200. grams of water, formed a solution that froze at –3.72°C?

(A) 25. g/mol

(B) 50. g/mol

(C) 100. g/mol

(D) 150. g/mol

(E) 1,000. g/mol

If you finish before one hour is up, you may go back to check your work or complete unanswered questions.

Answer Key

P  R  A  C  T  I  C  E   T  E  S  T   3

1. B

13. D

101. T, F

2. D

14. C

102. T, T, CE

3. C

15. A

103. T, F

4. A

16. B

104. F, T

5. B

17. C

105. F, T

6. D

18. E

106. T, T, CE

7. E

19. D

107. T, T, CE

8. C

20. A

108. T, T, CE

9. B

21. C

109. T, T, CE

10. A

22. B

110. T, T, CE

11. D

23. E

111. T, T, CE

12. C

 

112. T, T

   

113. F, T

   

114. T, F

   

115. T, T

   

116. F, T

24. B

40. D

56. A

25. D

41. B

57. E

26. D

42. B

58. D

27. D

43. B

59. E

28. D

44. C

60. E

29. E

45. A

61. C

30. D

46. A

62. A

31. A

47. B

63. B

32. E

48. D

64. C

33. B

49. B

65. A

34. A

50. C

66. D

35. C

51. D

67. A

36. C

52. E

68. B

37. E

53. E

69. A

38. A

54. C

 

39. D

55. A

 

ANSWERS EXPLAINED

1. (B) The activation energy of the forward reaction is the energy needed to begin the reaction.

2. (D) For the reverse reaction to occur, activation energy equal to the sum of (B) + (C) is needed. This is shown by (D).

3. (C) The heat of the reaction is the heat liberated between the level of potential energy of the reactants and that of the products. This is quantity (C) on the diagram.

4. (A) The potential energy of the reactants is the total of the original potential energies of the reactants shown by (A).

5. (B) The spoon must be made the cathode to attract the Ag+ ions.

6. (D) The silver plate is the anode.

7. (E) The solution of Ag+ provides the silver for plating.

8. (C) This equation shows the volume decreasing as the pressure is increased when the temperature is held constant. It is an example of Boyle’s Law.

9. (B) This equation shows the pressure increasing as the temperature is increased when the volume is held constant. It is an example of Gay-Lussac’s Law.

10. (A) This equation shows the volume increasing as the temperature is increased when the pressure is held constant. It is an example of Charles’s Law.

11. (D) This equation shows that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.

12–14. The balanced equations with half-reactions are as follows:

12. (C)

13. (D) The coefficients show that 3 mol of Cu produces 2 mol of NO, so 6 mol of Cu produces 4 mol of NO.

14. (C) The expression Cu(NO3)2 → Cu2+ + 2NO3 shows that the dissociation yields 3 ions.

15. (A) MgCl2 is ionic because it is the product of an active metal (Mg) combining with a very active nonmetal (Cl).

16. (B) The electronegativity difference be tween H and Cl is between 0.5 and 1.7. This indicates an unequal sharing of electrons, which results in a polar covalent bond.

17. (C) Ethane, CH3–CH3(g), has symmetrically arranged C–H polar bonds so that the ethane molecule is nonpolar covalent.

18. (E) Cu(s) is a metal.

19. (D) When hydrated copper sulfate is heated, the crystal crumples as the water is forced out of the structure, and a white powder is the result.

20. (A) Brownian movement is due to molecular collisions with colloidal particles, which knock the particles about in a zigzag path noted by the reflected light from these particles.

21. (C) The indicator phenolphthalein turns pink in a basic solution.

22. (B) Litmus paper turns pink in an acid solution.

23. (E) A substance that is deliquescent, such as a crystal, draws water to its surface. At times it can draw enough water to form a water solution.

101. (T, F) The assertion is true that the most active metals are found in the upper left corner of the Periodic Table because of their ability to lose their outer electron(s). These metals actually have smaller ionic radii than their atomic radii.

102. (T, T, CE) The assertion is explained by the reason. The graphic display of this is:

103. (T, F) The assertion is true but the reasoning is false. Transition elements have incomplete inner energy levels that are being filled with the additional electrons, thus leaving the outer energy level the same in most cases. As a result, these elements have common oxidation numbers.

104. (F, T) The assertion is false and the reason is true. A supersaturated solution is holding more than its normal solubility, and the addition of a crystal causes crystallization to occur.

105. (F, T) Equilibrium is a dynamic condition because of the reason stated. The assertion is false; the reason, true.

106. (T, T, CE) Both statements are true, and the reason explains why ionic bonding is the strongest.

107. (T, T, CE) An equilibrium system must have a gaseous reactant or product for pressure to affect the equilibrium. Then increased pressure will cause the reaction to go in the direction that reduces the concentration of gaseous substances. Since four volumes of gases on the left became two volumes on the right, the reaction shifts to the right.

108. (T, T, CE) The assertion is explained by the reason; both are true.

109. (T, T, CE) Both the assertion and the reason are true; they explain that the element’s orbital designation places it in the first transition series of filling the 3d orbitals.

110. (T, T, CE) The assertion is true; the reason is true and explains why nonmetals have the highest electronegativity.

111. (T, T, CE) The assertion is explained by the reason; both are true.

112. (T, T) There are 3 moles of atoms in 18 g of water because 18 g is 1 mol of water molecules and each molecule has three atoms. The reason does not explain the assertion but is also true.

113. (F, T) Benzene is a nonionizing substance and therefore a nonelectrolyte. The reason is a true statement.

114. (T, F) Isomers have the same empirical formula but vary in their structural formulas.

115. (T, T) The reaction does go to completion, but a gaseous product is formed, not a precipitate, so II does not explain I.

116. (F, T) In the Rutherford experiment relatively few alpha particles were deflected, indicating a great deal of empty space in the atom. The reason is a true statement.

24. (B) The correct coefficients are 2, 13, 8, and 10.

25. (D) 1K + 1Al + 2S + 8O = 12 total

26. (D) Air is a mixture; all others are compounds. Washing soda (C) is sodium carbonate, and lime (E) is calcium oxide.

27. (C) One mole of a gas at STP occupies 22.4 L. So two moles of a gas at STP occupy 2.0 mol × 22.4 L = 44.8 L.

28. (D) The maximum number of electrons in each kind of orbital is:

s = 2 in one orbital

p = 6 in three orbitals

d = 10 in five orbitals

f = 14 in seven orbitals

29. (E) The element with atomic number 11 is sodium with 1 electron in the 3s orbital. It would readily combine with the element that has 3p5 as the outer orbital since it needs only 1 more electron to fill it.

30. (D) The only physical property named in the list is low melting point.

31. (A) If the gas is diatomic, then 6.02 × 1023 atoms will form 6.02 × 1023/2 molecules. At STP, 6.02 × 1023 molecules occupy 22.4 L, so half that number will occupy 11.2 L.

32. (E) Cascading excited electrons can fall only to lower energy levels that are unfilled.

33. (B) Mass spectroscopy uses a magnetic field to separate isotopes by bending their path. The lighter ones are bent farther than the heavier ones.

34. (A) The pi bond is a bond between two p orbitals, like this:

35. (C) At Pikes Peak (alt. approx. 14,000 ft) the pressure is lower than at ground level; therefore the vapor pressure at a lower temperature will equal the out side pressure and boiling will occur.

36. (C) The alkanes contain the sp3 hybrid orbitals.

37. (E) I, II, and III are correct.

38. (A) Only I is correct.

39. (D) HCl is actually a stronger Brønsted-Lowry acid than H2O, which is why the reaction occurs as shown.

40. (D) I is not necessary for the reactor, but often nuclear energy is used to oper ate an electric generator. The others, II and III, are necessary for fuel and neutron-speed control, respectively.

41. (B) III is a salt from a strong base and a weak acid, which hydrolyzes to form a basic solution with water.

42. (B) Since the boiling point is increased by 0.51°C for each mole of particles, 1 mol of NaCl → Na+ + Cl gives 2 mol of particles. Therefore the boiling point will be 1.02° higher or 101.02°C.

43. (B) The VSEPR model shows BF3 is trigonal planar and so is related to the triangle shape on one plane.

44. (C) The delivery tube is below the fluid level in the flask and will cause liquid to be forced up the thistle tube when gas is evolved in the reaction.

45. (A) At first the fluid will be expelled up the thistle tube by the gas generated and exerting pressure in the reaction flask. When the level of the fluid falls below the end of the thistle tube, the gas will then be released through the thistle tube.

46. (A) The most active nonmetal has a high attraction for another electron— thus high electronegativity.

47. (B) Fe loses electrons to form the FE2+ ion.

48. (D) The Kw of water = [H+][OH] = 10–14. If [OH] = 10–5 mol/L, then

                              [H+]= 10–14/10–5 = 10–9
                               pH = –log[H+](by definition)
                               pH = –[–9]
                               pH = 9

49. (B) When dilute NaCl solution is electrolyzed, hydrogen is given off at the cathode, chlorine is given off at the anode, and sodium hydroxide is left in the container.

50. (C) The correctly balanced equation is

C2H4(g)+3O2(g) → 2CO2(g) + 2H2O(l)

51. (D) One mole of a gas at STP occupies 22.4 L.

52. (E) To find the simple or empirical formula, divide each % by the element’s atomic mass.

Carbon
40 ÷ 12 = 3.333

Hydrogen
6.67 ÷ 1 = 6.67

Oxygen
53.33 ÷ 16 = 3.33

Next, divide each quotient by the smallest quotient in an attempt to get small whole numbers.

3.33 ÷ 3.33 = 1 C
6.67 ÷ 3.33 = 2 H
3.33 ÷ 3.33 = 1 O

The simplest formula is CH2O, which has a molecular mass of 30. The true molecular mass is given as 90.0, which is three times the simplest. Therefore the true formula is C3H6O3.

53. (E) The reaction is:

x mol                  370 g
Cao + H2O → Ca(OH)2
Ca(OH)2 molecular mass = 74

370 g ÷ 74 = 5.0 mol of Ca(OH)2 is wanted. The reaction equation shows 1 mol of CaO produces 1 mol of Ca(OH)2, so the answer is 5.0 mol.

54. (C) In dilution problems, this formula can be used:

Mbefore × Vbefore = Mafter × Vafter

Substituting gives:

3.50 × 50.0 = 2 × (?x)

x = 87.5 mL, new volume after dilution

55. (A) For Keq to be small, the numerator, which is made up of the concentration( s) of the product(s) at equilibrium, must be smaller than the denominator.

56. (A) The amount of NaOH used is
49.2 – 34.7 = 14.5 mL
Using M1 × V1 = M2 × V2 gives
0.09 M × 14.5 mL = M2 × 10 mL
M2 = 0.13 M

57. (E) The reactions recorded indicated that the ease of losing electrons is greater in sodium than magnesium, greater in magnesium than iron, greater in iron than copper, and finally greater in copper than silver.

58. (D) Distillation removes only dissolved solids from the distillate. The volatile gases, such as H2S, will be carried into the distillate.

59. (E) All the first four statements are correct.

The Gibbs free-energy equation is:

ΔG = ΔH – TΔS

In choice (E), if ΔH is positive and ΔS is negative, then ΔG will definitely be positive, which means that the forward reaction will not occur spontaneously.

60. (E) H2 to O2 ratio by volume is 2 : 1 in the formation of water. Therefore, 4.0 mL H2 will react with 2.0 mL of O2 to make 4.0 mL of steam.

2H2(g) + O2(g) → 2H2O(g)

This leaves 30. mL of O2 uncombined.

61. (C) There will be 30. mL of O2 + 4.0 mL of steam = 34 mL total.

62. (A) By definition, a catalyst can be used to speed up a reaction without itself being consumed, so I is correct.

63. (B) The most electronegative element is fluorine (F), found in period 2.

64. (C) Solids are incorporated into the K value and therefore do not appear on the right side of the equation.

65. (A) Only NaNO3 will not react because it requires heat to react.

66. (D) When hybridization forms the sp3 orbitals in methane, CH4, four entirely new orbitals, different from but equivalent to the former s and p orbitals, result.

67. (A) Hydrogen bonding between water molecules causes the boiling point to be higher than would be expected.

68. (B) At the beginning of the reaction

[H2] = 1 mol/L
 [I2] = 1 mol/L
[HI] = 0

At equilibrium

(H2 + I2  2HI)
(Let x = moles/liter of
H2 and I2 in HI form)

    [H2] = (1–x) mol/L
     [I2] = (1–x) mol/L
    [HI] = 2x mol/L

Then, substituting the above values into the equation, you get:

69. (A) 10. g/200. g of water = 50. g/1,000. g of water (5 times as much)

The freezing point depression, 3.72°, is divided by 1.86°, which is the depression caused by 1 mol in 1,000. g of water, to find how many moles are dissolved.

3.72° ÷ 1.86° = 2 mol

If 50. g caused this depression and is equal to 2 mol, then 1 mol would be  of 50. g, or 25 g. So, the molar mass of the solid is 25 g/mol.

CALCULATING YOUR SCORE

Your score on Practice Test 3 can now be computed manually. The actual test will be scored by machine, but the same method is used to arrive at the raw score. You get one point for each correct answer. For each wrong answer, you lose one-fourth of a point. Questions that you omit or that have more than one answer are not counted. On your answer sheet mark all correct answers with a “C” and all incorrect answers with an “X”.

Determining Your Raw Test Score

Total the number of correct answers you have recorded on your answer sheet. It should be the same as the total of all the numbers you place in the block in the lower left corner of each area of the Subject Area summary in the next section.

  A. Enter the total number of correct answers here: ________ Now count the number of wrong answers you recorded on your answer sheet.

  B. Enter the total number of wrong answers here: ________ Multiply the number of wrong answers in B by 0.25.

  C. Enter that product here: ________ Subtract the result in C from the total number of right answers in A.

  D. Enter the result of your subtraction here: ________

  E. Round the result in D to the nearest whole number: ________.
This is your raw test score.

Conversion of Raw Scores to Scaled Scores

Your raw score is converted by the College Board into a scaled score. The College Board scores range from 200 to 800. This conversion is done to ensure that a score earned on any edition of a particular SAT Subject Test in Chemistry is comparable to the same scaled score earned on any other edition of the same test. Because some editions of the tests may be slightly easier or more difficult than others, scaled scores are adjusted so that they indicate the same level of performance regardless of the edition of the test taken and the ability of the group that takes it. Consequently, a specific raw score on one edition of a particular test will not necessarily translate to the same scaled score on another edition of the same test.

Because the practice tests in this book have no large population of scores with which they can be scaled, scaled scores cannot be determined.

Results from previous SAT Chemistry tests appear to indicate that the conversion of raw scores to scaled scores GENERALLY follows this pattern:

Note that this scale provides only a general idea of what a raw score may translate into on a scaled score range of 800–200. Scaling on every test is usually slightly different. Some students who had taken the SAT Subject Test in Chemistry after using this book had reported that they have scored slightly higher on the SAT test than on the practice tests in this book. They all reported that preparing well for the test paid off in a better score!

DIAGNOSING YOUR NEEDS

After taking Practice Test 3, check your answers against the correct ones. Then fill in the chart below.

In the space under each question number, place a check if you answered that question correctly.

EXAMPLE:

If your answer to question 5 was correct, place a check in the appropriate box.

Next, total the check marks for each section and insert the number in the designated block. Now do the arithmetic indicated and insert your percent for each area.

*The subject areas have been expanded to identify specific areas in the text.

* The subject areas have been expanded to identify specific areas in the text.

*The subject areas have been expanded to identify specific areas in the text.

* The subject areas have been expanded to identify specific areas in the text.

Answer Sheet

P  R  A  C  T  I  C  E   T  E  S  T   4

Determine the correct answer for each question. Then, using a No. 2 pencil, blacken completely the oval containing the letter of your choice.