A DIAGNOSTIC TEST - SAT Subject Test Chemistry

SAT Subject Test Chemistry



A Diagnostic Test

The following test of 85 questions is a sample of the actual test you will take to measure your chemistry achievement. It has basically the same distribution of topics, directions, and number and types of questions. Before taking this test, read the advice given in the section entitled “Final Preparation—The Day Before the Test.” Use the answer sheet provided, and limit the test time to 1 hour.

A Periodic Table of the Elements has been included for your use on problems requiring this source of information. Use this table also with the practice tests at the end of the book. The Diagnostic Test questions contain hyperlinks to their answers and explanations. Simply click on the question numbers to move back and forth.

Turn now to the test.

** Since this is an eBook, please record all of your answers separately. Directions in the Diagnostic Test do reflect those similar to what appears on the exam. All Answer Sheets should be used for reference only.

Answer Sheet


Determine the correct answer for each question. Then, using a No. 2 pencil, blacken completely the oval containing the letter of your choice. All Answer Sheets are for reference only. Please record your answers separately.


Note: For all questions involving solutions, you should assume that the solvent is water unless otherwise noted. Reminder: You may not use a calculator on this test!

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

mol = mole(s)

mm = millimeter(s)

T = temperature

V = volt(s)

V = volume

atm = atmosphere

Part A

Directions: Every set of the given lettered choices below refers to the numbered statements or formulas that immediately follow it. Choose the one lettered choice that best fits each statement or formula; 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 elements:

(A) Fluorine

(B) Chlorine

(C) Bromine

(D) Iodine

(E) Astatine

1. The element that is most active chemically

2. The element with the smallest ionic radius

3. The element with the lowest first ionization potential

4. The element that first shows some visible metallic properties at room temperature

Questions 5–7 refer to the following sublevels:

(A) 1s

(B) 2s

(C) 3s

(D) 3p

(E) 3d

5. Contains up to 10 electrons.

6. Contains one pair of electrons in the ground-state electron configuration of the lithium atom.

7. Is exactly one-half filled in the ground-state electron configuration of the phosphorus atom.

Questions 8–12 refer to the following:

(A) Avogadro”s number

(B) P1V1 = P2V2

(C) V1T2 = V2T1

(D) Dalton”s Theory

(E) Gay-Lussac”s Law

8. Proposes basic postulates concerning elements and atoms

9. Proposes a relationship between the combining volumes of gases with respect to the reactants and gaseous products

10. Proposes a temperature-volume relationship of gases

11. Proposes a concept regarding the number of particles in a mole

12. Proposes a volume-pressure relationship of gases

Questions 13–16 refer to the following structures:

(A) R–OH

(B) R–O–R*




(* Alkyl group that is not necessarily the same as R)

13. The organic structure designation that includes the functional group of an aldehyde

14. The organic structure designation that includes the functional group of an acid

15. The organic structure designation that includes the functional group of an ester

16. The organic structure designation that includes the functional group of an ether

Questions 17–21 refer to the following:

(A) H2(g)

(B) CO2(g)

(C) 2N2O(g)

(D) 2NaCl (aq)

(E) H2SO4(dilute aq)

17. The expression that can be used to designate a linear nonpolar molecule that contains polar bonds

18. The expression that can be used to designate 2 moles of atoms

19. The expression that can be used to designate 3 moles of atoms

20. The expression that can be used to designate a maximum of 3 moles of ions

21. The expression that can be used to designate 6 moles of atoms

Questions 22–25 refer to the following pairs of substances:

(A) NH3 and N2H4

(B) 16O and 17O

(C) NH4Cl and NH4NO3

(D) CH3OCH3 and CH3CH2OH

(E) O2 and O3

22. Are isotopes

23. Have both ionic and covalent bonds

24. All allotropes

25. Are strong electrolytes in aqueous solutions

Part B


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 whether 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.




A catalyst can accelerate a chemical reaction


a catalyst can decrease the activation energy required for the reaction to occur.


Molten sodium chloride is a good electrical conductor


sodium chloride in the molten state allows ions to move freely.


Ice is less dense than liquid water


water molecules are nonpolar.


Two isotopes of the same element have the same mass number


isotopes have the same number of protons.


A 1.0 g sample of calcium citrate, Ca3(C6H5O7)2 (molar mass 498 g/mol), contains more Ca than a 1.0 g sample of calcium carbonate, CaCO3 (molar mass 100 g/mol)


there are more Ca atoms in 1.0 mol of calcium carbonate than in 1.0 mol of calcium citrate.


Two liters of CO2 can be produced by 1 gram of carbon burning completely


the amount of gas evolved in a chemical reaction can be determined by using the mole relationship of the coefficients in the balanced equation.


A reaction is at equilibrium when it reaches completion


the concentrations of the reactants in a state of equilibrium equal the concentrations of the products.


The anions in an electrolytic cell migrate to the cathode


positively charged ions are attracted to the negatively charged cathode in an electrolytic cell


A solution with pH = 5 has a higher concentration of hydronium ions than a solution with a pH = 3


pH is defined as −log [H+].


An endothermic reaction can be spontaneous


both the enthalpy and the entropy changes affect the Gibbs free-energy change of the reaction.


Weak acids have small values for the equilibrium constant, Ka,


the concentration of the hydronium ion is in the numerator of the Ka expression.


One mole of NaCl contains 2 moles of ions


NaCl is a stable salt at room temperature.


A pi bond is formed between the lobes of adjacent p orbitals in the same plane of two atoms that contain only one electron each


each of the two lobes of a single p orbital can hold two electrons of opposite spin.


H2S and H2O have a significant difference in their boiling points


hydrogen sulfide has a higher degree of hydrogen bonding than water.

Part C

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

26. Two immiscible liquids, when shaken together vigorously, may form

(A) a solution

(B) a tincture

(C) a sediment

(D) a hydrated solution

(E) a colloidal dispersion

27. A thermometer is used to record the cooling of a confined pure substance over a period of time. During which interval on the cooling graph above is the system undergoing a change of state from a liquid to a solid?

28. If a principal energy level of an atom in the ground state contains 18 electrons, they will be arranged in orbitals according to the pattern

(A) s6p6d6

(B) s2p6d10

(C) s2d 6f10

(D) s2p6f10

(E) s2p2f14

29. Which of the following molecules is a saturated hydrocarbon?

(A) C3H8

(B) C2H4

(C) C4H6



30. A liter of hydrogen is at 5.0°C temperature and under 640. torr pressure. If the temperature were raised to 60.0°C and the pressure decreased to 320. torr, how would the liter volume be modified?

31. Of the following statements about the number of subatomic particles in an ion of , which is (are) true?

I. 16 protons

II. 14 neutrons

III. 18 electrons

(A) II only

(B) III only

(C) I and II only

(D) I and III only

(E) I, II, and III

32. The most active metallic elements are found in

(A) the upper right corner of the periodic chart

(B) the lower right corner of the periodic chart

(C) the upper left corner of the periodic chart

(D) the lower left corner of the periodic chart

(E) the middle of the periodic chart, just beyond the transition elements

33. If 1 mole of each of the following substances was dissolved in 1,000 grams of water, which solution would have the highest boiling point?

(A) NaCl

(B) KCl

(C) CaCl2

(D) C6H10O5

(E) C12H22O11

34. A tetrahedral molecule, XY4, would be formed if X were using the orbital hybridization

(A) p2

(B) s 2

(C) sp

(D) sp 2

(E) sp 3

35. In the following reaction, how many liters of SO2 at STP will result from the complete burning of pure sulfur in 8 liters of oxygen?

  S(s) + O2 (g)→SO2 (g)

(A) 1

(B) 4

(C) 8

(D) 16

(E) 32

36. In the above laboratory setup to measure the pressure of the confined gas, what will be true concerning the calculated pressure on the gas?

(A) The gas pressure will be the same as the atmospheric pressure.

(B) The gas pressure will be less than the atmospheric pressure.

(C) The gas pressure will be greater than the atmospheric pressure.

(D) The difference in the height (h) of mercury levels is equal to the pressure of the gas.

(E) The height (h) of mercury has no effect on the pressure calculation since the column of mercury is only used to enclose the gas volume.

37. Which of the following changes in the experiment shown in question 36 would cause the pressure in the glass container to vary from that shown?

(A) Use a U-tube of a greater diameter and maintain the height of mercury.

(B) Increase the temperature of gas in the tube.

(C) Increase the length of the upper portion of the right side of tubing.

(D) Use a U-tube of a smaller diameter and maintain the height of mercury.

(E) Replace the flask with one that has the same volume but has a flat bottom.

38. Which of the following can be classified as amphoteric?

(A) Na3PO4

(B) HCl

(C) NaOH

(D) HSO4

(E) C2O42−

39. Standard conditions (STP) are

(A) 0°C and 2 atm

(B) 32°F and 76 torr

(C) 273 K and 760 mm Hg

(D) 4°C and 7.6 cm Hg

(E) 0 K and 760 mm Hg

40. Laboratory results showed the composition of a compound to be 58.81% barium, 13.73% sulfur, and 27.46% oxygen. What is the empirical formula of the compound?

(A) BaSO4

(B) BaS2O

(C) Ba2SO3

(D) BaS2O4

(E) Ba2SO4

41. What is the percentage composition of calcium in calcium hydroxide, Ca(OH)2? (1 mol = 74 g)

(A) 40%

(B) 43%

(C) 54%

(D) 69%

(E) 74%

42. How many grams of hydrogen gas can be produced from the following reaction if 65 grams of zinc and 65 grams of HCl are present in the reaction?

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

(A) 1.0

(B) 1.8

(C) 3.6

(D) 7.0

(E) 58

43. The following statements were recorded while preparing carbon dioxide gas in the laboratory. Which one involves an interpretation of the data rather than an observation?

(A) No liquid was transfered from the reaction bottle to the beaker.

(B) The quantity of solid minerals decreased.

(C) The cloudiness in the last bottle of limewater was caused by the product of the reaction of the colorless gas and the limewater.

(D) The bubbles of gas rising from the mineral remained colorless throughout the experiment.

(E) There was a 4°C rise in temperature in the reaction vessel during the experiment.

44. The above laboratory setup can be used to prepare which of the following?

I. CO2(g)

II. H2(g)

III. O2(g)

(A) I only

(B) III only

(C) I and III only

(D) II and III only

(E) I, II, and III

45. The missing product in the nuclear reaction represented above is

46. Which of the following is (are) true regarding the aqueous dissociation of HCN, Ka = 4.9 × 10−10, at 25°C?

I. At equilibrium, [H+] = [CN].

II. At equilibrium, [H+] > [CN].

III. HCN(aq) is a strong acid.

(A) I only

(B) II only

(C) I and II only

(D) I and III only

(E) I, II, and III

47. This question pertains to the reaction represented by the following equation:

2NO(g)+ O2 (g) 2NO2 (g)+ 150 kJ

Suppose that 0.8 mole of NO is converted to NO2 in the above reaction. What amount of heat will be evolved?

(A) 30 kJ

(B) 60 kJ

(C) 80 kJ

(D) 130 kJ

(E) 150 kJ

48. How does a Brønsted-Lowry acid differ from its conjugate base?

(A) The acid has one more proton.

(B) The acid has one less proton.

(C) The acid has one more electron.

(D) The acid has one less electron.

(E) The acid has more than one additional proton.

49. Two containers having 1 mole of hydrogen gas and 1 mole of oxygen gas, respectively, are opened. What will be the ratio of the rate of effusion of the hydrogen to that of the oxygen?


(B) 4 : 1

(C) 8 : 1

(D) 16 : 1


50. A molecule in which the electron configuration is a resonance hybrid is

(A) SO2

(B) C2H6

(C) Cl2

(D) HBr

(E) NaCl

51. What is the pH of a solution in which the [OH] is 1.0 × 10−4?

(A) −4

(B) +4

(C) +7

(D) −10

(E) +10

52. If 0.365 gram of hydrogen chloride is dissolved to make 1 liter of solution (Cl = 35.5 and H = 1.00), the pH of the solution is

(A) 0.001

(B) 0.01

(C) 1

(D) 2

(E) 12

53. In the laboratory, a sample of hydrated salt was heated at 110°C for 30 minutes until all the water was driven off. The data were as follows:

Mass of the hydrate before heating = 250 grams

Mass of the hydrate after heating = 160 grams

From these data, what was the percent of water by mass in the original sample?

(A) 26.5

(B) 36

(C) 47

(D) 56

(E) 90

54. Which of the following oxides dissolves in water to form an acidic solution?

(A) Na2O

(B) CaO

(C) Al2O3

(D) ZnO

(E) SO3

55. In the laboratory, 20.0 milliliters of an aqueous solution of calcium hydroxide, Ca(OH)2, was used in a titration. A drop of phenolphthalein was added to it to indicate the end point. The solution turned colorless after 20.0 milliliters of a standard solution of 0.050 M HCl solution was added. What was the molarity of the Ca(OH)2?

(A) 0.010 M

(B) 0.025 M

(C) 0.50 M

(D) 0.75 M

(E) 1.0 M

56. Which of the following reactions will NOT spontaneously go to completion?

(A) Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

(B) CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(aq) + CO2(g)

(C) Ag+(aq) + HCl(aq) → AgCl(s) + H+(aq)

(D) Cu(s) + 2H+(aq) → Cu2+(aq) + H2(g); E0 = −0.34 V

(E) H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(aq)

57. For a laboratory experiment, a student placed sodium hydroxide crystals on a watch glass, assembled the titration equipment, and prepared a solution of 0.10 M sulfuric acid. Then he weighed 4 grams of sodium hydroxide and added it to enough water to make 1 liter of solution. What might be a source of error in the results of the titration?

(A) Some sulfuric acid evaporated.

(B) The sulfuric acid became more concentrated.

(C) The NaOH solution gained weight, thus increasing its molarity.

(D) The NaOH crystals gained H2O weight, thus making the solution less than 0.1 M.

(E) The evaporation of sulfuric acid solution countered the absorption of H2O by the NaOH solution.

58. If 60. grams of NO is reacted with sufficient O2 to form NO2 that is removed during the reaction, how many grams of NO2 can be produced? (Molar masses: NO = 30. g/mol, NO2 = 46. g/mol)

(A) 46.

(B) 60.

(C) 92.

(D) 120

(E) 180

59. Based on the information shown, each of the following equations represents a reaction in which the change in entropy, S, is positive EXCEPT

(A) CaCO3(s) → CaO(s) + CO2(g)

(B) Zn(s) + 2H+(aq) → H2(g) + Zn2+(aq)

(C) 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(g)

(D) NaCl(s) → Na+(aq) + Cl(aq)

(E) N2(g) + 3H2(g) → 2NH3(g)

60. Cl2(g) + 2Br (excess) → ?

When 1 mole of chlorine gas reacts completely with excess KBr solution, as shown above, the products obtained are

(A) 1 mol of Cl ions and 1 mol of Br

(B) 1 mol of Cl ions and 2 mol of Br

(C) 1 mol of Cl ions and 1 mol of Br2

(D) 2 mol of Cl ions and 1 mol of Br2

(E) 2 mol of Cl ions and 2 mol of Br2

61. Two water solutions are made in the laboratory, one of glucose (molar mass = 180 g/mol), the other of sucrose (molar mass = 342 g/mol). If the glucose solution had 180 grams in 1,000 grams of water and the sucrose had 342 grams in 1,000 grams of water, which statement about the freezing points of the solutions is the most accurate?

(A) The glucose solution would have the lower freezing point.

(B) The sucrose solution would have the lower freezing point.

(C) The freezing point of the sucrose solution would be lowered twice as much as that of the glucose solution.

(D) Both solutions would have the same freezing point.

(E) The freezing points of the solutions would not be affected, because both solutes are nonpolar.

62. Which Ka value indicates the strongest acid?

(A) 1.3 × 10−2

(B) 6.7 × 10−5

(C) 5.7 × 10−10

(D) 4.4 × 10−7

(E) 1.8 × 10−16

63. What mass of CaCO3 is needed to produce 11.2 liters of CO2 at STP when the calcium carbonate is reacted with an excess amount of hydrochloric acid? (Molar masses: CaCO3 = 100. g/mol, HCl = 36.5 g/mol, CO2 = 44.0 g/mol)

(A) 25.0 g

(B) 44.0 g

(C) 50.0 g

(D) 100. g

(E) None of the above

64. By experimentation it is found that a saturated solution of BaSO4 at 25°C contains 3.9 × 10−5 mole/liter of Ba2+ ions. What is the Ksp of the BaSO4?

(A) 1.5 × 10−4

(B) 1.5 × 10−9

(C) 1.5 × 10−10

(D) 3.9 × 10−10

(E) 39 × 10−9

65. What is the H° value for the decomposition of sodium chlorate, given the following information?

NaClO3 (s) → NaCl(s)+ O2 (g)

(H°f values: NaClO3(s) = −358 J/mol, NaCl(s) = −410 J/mol, O2(g) = 0 kcal/mol)

(A) 52.0 J

(B) −52.0 J

(C) 768 J

(D) −768 J

(E) (768 J)

66. To the equilibrium reaction shown below:

AgCl(s) Ag+ + Cl

a beaker of concentrated HCl (12 M) is slowly added. Which is the best description of what will occur?

(A) More salt will go into solution, and the Ksp will remain the same.

(B) More salt will go into solution, and the Ksp will increase.

(C) Salt will come out of the solution, and the Ksp will remain the same.

(D) Salt will come out of the solution, and the Ksp will decrease.

(E) No change in concentration will occur, and the Ksp will increase.

67. When the following redox equation is balanced and all coefficients are reduced to lowest whole-number terms, what is the coefficient of H2O?

  HCl + KMnO4 → H2O + KCl + MnCl2 + Cl2

(A) 1

(B) 2

(C) 5

(D) 8

(E) 16

68. Each of the following systems is at equilibrium in a closed container. A decrease in the total volume of each container will increase the number of moles of product(s) for which system?

(A) 2NH3(g) N2(g) + 3H2(g)

(B) H2(g) + Cl2(g) 2HCl(g)

(C) 2NO(g) + O2(g) 2NO2(g)

(D) CO(g) + H2O(g) CO2(g) + H2(g)

(E) Fe3O4(s) + 4H2(g) 3Fe(s) + 4H2O(g)

69. Which of the following is the correct and complete Lewis electron-dot diagram for PF3?

70. Hydrogen gas is collected in a eudiometer tube over water as shown above. The water level inside the tube is 40.8 millimeters higher than that outside. The barometric pressure is 730. millimeters Hg. The water vapor pressure at the room temperature of 29°C is found in a handbook to be 30.0 millimeters Hg. What is the pressure of the dry hydrogen?

(A) 659.2 mm Hg

(B) 689.2 mm Hg

(C) 697.0 mm Hg

(D) 740.8 mm Hg

(E) 800.8 mm Hg

71. How many moles of electrons are required to reduce 2.93 grams of nickel ions from melted NiCl2? (Molar mass of Ni = 58.7 g/mol)

(A) 0.050

(B) 0.10

(C) 1.0

(D) 1.5

(E) 2.0

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

Answer Key


1. A

14. D

102. T, T, CE

2. A

15. E

103. T, F

3. E

16. B

104. F, T

4. D

17. B

105. F, F

5. E

18. A

106. T, T, CE

6. A

19. B

107. F, F

7. D

20. E

108. F, T

8. D

21. C

109. F, T

9. E

22. B

110. T, T, CE

10. C

23. C

111. T, T, CE

11. A

24. E

112. T, T

12. B

25. C

113. T, F

13. C

101. T, T, CE

114. T, F

26. E

41. C

56. D

27. D

42. B

57. D

28. B

43. C

58. C

29. A

44. E

59. E

30. D

45. C

60. D

31. D

46. A

61. D

32. D

47. B

62. A

33. C

48. A

63. C

34. E

49. B

64. B

35. C

50. A

65. B

36. C

51. E

66. C

37. B

52. D

67. D

38. D

53. B

68. C

39. C

54. E

69. E

40. A

55. B

70. C

71. B


1. (A) In the halogen family, the most active nonmetal would be the top element, fluorine, because it has the highest electronegativity.

2. (A) As you proceed down a group, the ionic radius increases as additional energy levels are filled farther from the nucleus. Therefore fluorine, the top element, has the smallest ionic radius.

3. (E) Since astatine has the largest atomic radius and its outer electrons are shielded from the protons by a large number of interior electrons, it has the lowest first ionization potential.

4. (D) Some physical characteristics of metal are found in iodine, the fourth halogen down in the group.

5. (E) The 3d subshell has 5 orbitals, or 5 possible values of m (−2, −1, 0, +1, and +2). The Pauli exclusion principle allows each orbital a maximum of 2 electrons with opposite spins. So the 3d subshell can contain up to 10 electrons.

6. (A) The ground state is the lowest-energy, or most stable state, of an atom. The lithium atom has 3 electrons. In its stable state, the 1s subshell, or lowest-energy orbital, contains a pair of electrons. The Pauli exclusion principle allows a maximum of 2 electrons with opposite spins, so the third electron resides in the 2s orbital.

7. (D) Phosphorus has 15 electrons. In its most stable state, the lowest-energy subshells will be occupied. One electron fills the orbital of a subshell before coupling in electron spin pairs begins. For phosphorus, 12 electrons fill the 1s2, 2s2, 2p6, and 3s2 orbitals. The remaining 3 electrons fill exactly one-half of the 3p orbitals, with all of the spins in the same direction, or 3p3.

8. (D) John Dalton is credited with the basic postulates of atomic theory.

9. (E) Gay-Lussac is credited with the statement that, when gases combine, they do so in ratios of small whole numbers that are in relationship to the volumes of the reactants and the volumes of the products under the same conditions.

10. (C) Charles is credited with this temperature-volume relationship of gases:

or V1T2 = V2T1 as shown in the question.

11. (A) Avogadro is credited with the concept regarding the number of particles in a mole, 6.02 × 1023, and this number bears his name.

12. (B) Boyle is credited with the P1V1 = P2V2 relationship of gases.

13. (C) This includes the functional group of an aldehyde.

14. (D) This includes the functional group of an organic acid.

15. (E) An ester is the equivalent of an organic salt since it is usually formed from an organic alcohol,

R–OH, plus an organic acid, .

The bonding is , which gives

. (R* indicates that this hydrogen branch need not be the same as R.)

16. (B) This includes the functional group of an ether. It can be formed by the dehydration of two alcohol molecules. The reaction is

17. (B) The molecular structure of carbon dioxide is O=C=O, where the oxygens are 180° apart and, although the bonding is polar to the carbon, counteract each other to constitute a nonpolar molecule.

18. (A) A hydrogen gas molecule is diatomic; it has 2 mol of atoms in each mole of molecules, represented by H2.

19. (B) Each CO2 has three atoms per mol ecule; hence the expression can represent 3 mol of atoms in 1 mol of molecules.

20. (E) With complete ionization , or 3 mol of ions per mole of H2SO4.

21. (C) The expression 2N2O represents two triatomic molecules or 6 mol of atoms.

22. (B) The 16O and 17O are isotopes of the same element because they do not have the same number of neutrons. This difference in the number of neutrons shows up in the mass number (superscript to the left of the chemical symbol). The mass of an element is the sum of the protons and neutrons in the nucleus.

23. (C) In NH4Cl, the covalently bonded NH4+ ion is ionically bonded to the Cl ion. In NH4NO3, the covalently bonded NH4+ ion is ionically bonded to the covalently bonded NO3 ion. In each of these cations and anions are covalent bonds within the structure where the electrons are shared equally.

24. (E) The O2 and O3 are allotropes because they are different molecular forms of the same element in the same state. They have different structures due to different bond arrangements.

25. (C) Strong electrolytes will almost completely dissociate in aqueous solutions. Both NH4Cl and NH4NO3 are ionic compounds that dissociate to release the ammonium cation and either the Cl or the NO3 anion.

101. (T, T, CE) A catalyst can accelerate a chemical reaction by lowering the activation energy required for the reaction to occur.

102. (T, T, CE) Sodium chloride is an ionic substance and when molten is a good electrical conductor, and the reason is that in the molten state the ions are free to migrate to the anode and cathode.

103. (T, F) Ice is less dense than liquid water. Water molecules, however, are polar, not nonpolar, and water expands as these molecules arrange themselves into a crystal lattice.

104. (F, T) The statement that isotopes of the same element have the same mass number is false. Isotopes of the same element have the same number of protons but vary in the number of neutrons in the nucleus. Therefore, these isotopes have the same atomic number but different mass numbers.

105. (F, F) There are fewer Ca atoms in 1 mole of calcium carbonate (CaCO3) than in 1 mole of calcium citrate (Ca3(C6H5O7)2) because each mole of calcium citrate contains 3 times more Ca than each mole of calcium carbonate.

106. (T, T, CE) The reaction is:

 C + O2 → CO2

and shows 1 mol or 12.0 g of carbon produces 1 mol or 22.4 L of CO2. Then

So the statement is correct, and the assertion explains it correctly.

107. (F, F) A reaction at equilibrium has reached a point where the forward and reverse reactions are occurring at equal rates. The concentrations of the reactants and products, however, are not necessarily equal and are described by the Keq value at the temperature of the reaction.

108. (F, T) In an electrolyte reaction, anions migrate to the anode. It is true that positively charged ions (cations) are attracted to the negatively charged cathode.

109. (F, T) The value pH = 5 can be expressed as

and pH = 3 as

Thus pH = 3 represents a larger concentration of hydronium ions, H3O+.

110. (T, T, CE) The first statement is true. The change in Gibbs free energy, G, depends on the enthalpy change, H, and the entropy change, S, from the equation G = H - T S. Thus, statement II is also true and explains the first statement.

111. (T, T, CE) In the expression for the equilibrium constant of an acid, [H3O+] is in the numerator:

As [H3O+] decreases in the weak acids, the numerator becomes smaller as the denominator gets larger, therefore giving smaller Ka values.

112. (T, T) Both statements are true, but they are not related.

113. (T, F) A pi bond is formed between p lobes of adjacent atoms and in the same plane:

However, each p orbital consisting of two lobes can hold a total of two electrons, so the reason is false.

114. (T, F) The assertion that the boiling points of H2S and H2O are significantly different is true, but the reason is false. Water has the higher degree of hydrogen bonding.

26. (E) Two immiscible liquids, when shaken together vigorously, may each form small, colloidal-size particles, dispersed in the other liquid.

27. (D) In the graph, the first plateau must represent the condensation from gas to liquid because there is a second, lower plateau, which would represent the second change of state, from liquid to solid.

28. (B) If the electrons have the same principal energy level, they will fill the s 2, p6, then the d10 level. This progression is from the lowest energy sublevel to the highest, to accommodate 18 electrons.

29. (A) When the chain or ring carries the full complement of hydrogen atoms, the hydrocarbon is said to be saturated. The general formula for the saturated alkanes is CnH2n+2. The chain is CH3−CH2−CH3.

30. (D) Considering the concepts behind Charles”s Law and Boyle”s Law, you can arrive at the fraction to be used in kelvins and torrs. The volume must increase with an increase in temperature and also increase with a decrease in

pressure. Therefore the fractions would be and

Using the combined gas law equation,

and solving for V2, you get

which is the same answer.

31. (D) The mass number, 32, is the total number of neutrons and protons. Since the atomic number, 16, gives the number of protons, 32 − 16 = 16, or the number of neutrons, statement II is false. Since this ion has a charge of 2, it has two more electrons than protons, or 18 electrons. Statements I and III are true.

32. (D) The most active metallic elements are found in the lower left corner of the Periodic Table.

33. (C) The rise in the boiling point depends on the number of particles in solution. One mole of CaCl2 gives 3 mol of ions, more than any other substance listed:

 CaCl2 → Ca2++2Cl.

The number of moles of ions given by the other substances are as follows:

   (A) = 2, (B) = 2, (D) = 1, (E) = 1.

34. (E) The sp3 hybrid has the tetrahedron configuration. The sp2 (D) is trigonal planar. The sp (C) is linear. The s (B) and p (A) are the usual orbital structures.

35. (C) The reaction is

The given (8 L) and the unknown (x L) are shown above. Since the equation, according to Gay-Lussac”s Law, shows that 1 volume of oxygen yields 1 volume of sulfur dioxide, then

36. (C) The fact that the mercury level in the U-tube is higher in the right side of the tube indicates that the pressure in the flask is higher than the atmospheric pressure exerted on the open end of the tube on the right side. If the pressure inside the flask were the same as the atmospheric pressure, the height of the mercury would be the same in both sides of the U-tube.

37. (B) The only change listed that would change the pressure of the gas inside the flask is to increase the temperature of the gas. This would cause the pressure to rise.

38. (D) An amphoteric substance must be able to be a proton, (H+), donor, and a proton receiver. The bisulfate ion, HSO4, is the only choice that can either accept a proton and become H2SO4, or lose a proton and become the sulfate ion, SO42−.

39. (C) Standard conditions are 273 K and 760 mm Hg.

40. (A) Dividing the percentage of each element by the atomic mass of that element gives the basic ratio of atoms, but not necessarily in whole numbers. Thus, (Ba) 58.8 ÷ 137 = 0.43, (S) 13.7 ÷ 32 = 0.43, (O) 27.5 ÷ 16 = 1.72.

Because atoms occur in whole numbers, you now must manipulate these numbers mathematically to get whole numbers. Usually dividing each number by the smallest one helps to accomplish this: (Ba) 0.43 ÷ 0.43 = 1, (S) 0.43 ÷ 0.43 = 1, (O) 1.72 ÷ 0.43 = 4. The empirical formula is BaSO4.

41. (C) The percentage composition can be found by dividing each total atomic mass in the formula by the molar mass of the compound.

 CA = 40 40 ÷ 74 × 100% = 54% Ca
2O = 32 32 ÷ 74 × 100% = 43% O
2H =  2 ÷ 74 × 100% = 2.7% H

42. (B) The solution setup is:

Note that the equation mass is calculated under the substances that have mass units above them. According to the calculated equation masses, 73 g of HCl would be needed to react with 65 g of zinc. Since there is only 65 g of HCl, we use this and disregard the 65 g of Zn.

= 1.78 or 1.8 g of H2

43. (C) All the other statements represent observations because they merely record what was seen.

44. (E) The setup is appropriate for the collection of a basically nonsoluble gas by the displacement of water. All three gases fit this description.

45. (C) With the emission of a neutron, the total atomic mass decreases by 1. However, the number of protons is 2. The product is helium,

46. (A) The small Ka indicates that this is a weak acid, so statement III is false. When HCN ionizes, it can be shown that HCN H+ + CN. This is a molar ratio of 1 : 1, so statement I must be true.

47. (B) Because the equation shows that 2 mol of NO react to release 150 kJ, the solution is

This problem would be solved in the same manner if the heat had been expresssed in kilocalories. To convert one unit to the other, use 4.18 × 103 J = 1 kcal.

48. (A) In the Brønsted-Lowry acid-base theory:

49. (B) According to Graham”s Law, the rates of effusion of two gases are inversely proportional to the square roots of their molar masses. Since 1 mol O2 = 32 g and 1 mol H2 = 2 g

Therefore, the effusion rate of hydrogen is four times faster than that of oxygen.

50. (A) The structure of SO2 is a resonance hybrid, shown as these resonance structures:

51. (E) The Kw of water is

K w = [H+][OH] = 10−14

If [OH] = 1.0 × 10−4 , then


    pH = −log [H+] = 10

52. (D)

0.01 mol HCl → 0.01 mol H+ + 0.01 mol Cl
If [H+] = 0.01 = 1 × 10−2 mol/L, then pH = 2.

53. (B) Hydrate mass before heating = 250. g

- Hydrate mass after heating = 160. g
    Water loss = 90.0 g

54. (E) Because metallic oxides are basic anhydrides, the only nonmetallic oxide is SO3. The reaction is as follows

SO3 + H2O → H2SO4 (sulfuric acid)

55. (B) In the titration, the reaction is:

2HCl + Ca(OH)2 → CaCl2 + 2H2O

The acid to base ratio is 2 : 1, or (moles acid used) = 2(moles base used), so MaVa = 2MbVb, where M is the molarity and V is the volume expressed in liters, then

56. (D) All the reactions will go to completion except (D), which will not occur spontaneously at all. If E0 had been positive, however, this redox reaction would occur.

57. (D) Sodium hydroxide is hygroscopic and will attract water to its surface area. This water will influence its mass; consequently there will be less sodium hydroxide in the mass used.

58. (C)

This is a mass stoichiometry problem. The equation masses are placed beneath the substances that have quantities above them.

Using the equation relationship gives

Using the mole method gives

Using the equation coefficients gives


59. (E) Since in this reaction 4 volumes of gases are forming 2 volumes of product, the randomness of the system is decreasing. Therefore, entropy is decreasing and S is negative. In all the other reactions randomness is increasing.

60. (D) The reaction is

   Cl2 (g)+ 2Br(excess) → 2Cl + Br2(g)

Here 2 mol of chloride ions and 1 mol of Br2 molecules are produced.

61. (D) Since both solutions were 1 molal and neither compound ionized into more particles, the freezing points would be the same.

62. (A) HY H+ + Y

Since [H+] is in the numerator of Ka:

The stronger the acid, the greater are [H+] and the Ka value. Of the choices given, 1.3 × 10−2 is the largest.

63. (C)

This is a mass-volume problem with the mass and volume indicated below the equation.
Using the factor-label method gives

Using the mole method gives

64. (B) Ksp =[Ba2+][SO42−]

If [Ba2+] = 3.9 × 10−5, then [SO42−] must also equal the same amount, so

Ksp = [3.9 × 10−5][3.9 × 10−5]

 = 15.2 × 10−10 or 1.5 × 10−9.

65. (B) H0reaction = H0f (products) − H0f (reactants)

H0reaction = −410. J − (−358. J)

H0reaction = −52. J

66. (C) Because the HCl solution will add to the chloride ion concentration, according to Le Châtelier”s Principle the equilibrium will shift in the direction to reduce this disturbance, so the Ksp will remain the same but salt will come out of solution. This process:

AgCl Ag+ + Cl

will continue until the Ksp is reestablished. This phenomenon is called the “common ion effect.”

67. (D) HCl + KMnO4 → H2O + KCl + MnCl2 + Cl2
oxidation: 2Cl → Cl2 + 2e
reduction: 8H+ + MnO4 + 5e → Mn2+ + 4H2O
5 (oxidation reaction) and 2 (reduction reaction) will balance the e gain and loss, giving
10Cl → 5Cl2 + e
16H+ + 2MnO4 + e → 2Mn2+ + 8H2O

16HCl + 2KMnO4 → 8H2O + 2KCl + 2MnCl2 + 5Cl2

68. (C) A decrease in volume will cause the equilibrium to shift in the direction that has less volume(s) of gas(es). In every case except (C) this is the reverse reaction, which decreases the product. The coefficients give the volume relationships.

69. (E) The central P has 5 valence electrons. Of these, 2 are paired. The remaining 3 valence electrons each covalently bond with one of the 3 F to fill the outer energy level. The Lewis electron-dot diagram is

70. (C) Change the water height to the equivalent Hg height:

= 3 mm Hg

Adjust for the difference in height to get the gas pressure. Pressure on the gas is

730. mm Hg − 3 mm Hg
= 727 mm Hg.

Vapor pressure of H2O at 29°C accounts for 30 mm Hg pressure. Therefore

727 mm Hg − 30 mm Hg
= 697 mm Hg

71. (B) The reaction is

Ni2+ + 2e → Ni (s)

Since 2 mol of electrons are required to form 1 mol of nickel, and


= 0.1 mol of electrons


Your score on the diagnostic test 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 for which you have indicated 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.

Since the practice tests in this book have no large population of scores with which they can be scaled, scaled scores can only be approximated.

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 have taken the SAT Subject Test in Chemistry after using this book have 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!


This section will help you to diagnose your need to review the various categories tested by the SAT Subject Test in Chemistry.

After taking the diagnostic test, 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.

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


The percentages give you an idea of how you have done on the various major areas of the test. Because of the limited number of questions on some parts, these percentages may not be as reliable as the percentages for parts with larger numbers of questions. However, you should now have at least a rough idea of the areas in which you have done well and those in which you need more study. (There are four more practice tests in the back of this book, which may be used in a diagnostic manner as well.)

Start your study with the areas in which you are the weakest. The corresponding chapters are indicated on the next page.

After you have spent some time reviewing your weaker areas, plan a schedule of work that spans the 6 weeks before the test. Unless you set up a regular study pattern and goals, you probably will not prepare sufficiently.

The following schedule provides such a plan. Note that weekends are left free, and the time spans are held to 1- or 2-hour blocks. This will be time well spent!


The day before the test, review one of the practice tests you have already taken. Study again the directions for each type of question. Long hours of study at this point will probably only heighten your anxiety, so just look over the answer section of the practice test and refer to any chapter in the book if you need more information. This type of limited, relaxed review will probably make you feel more comfortable and better prepared.

Get together the materials that you will need. They are:

• Your admission ticket. Check the time your admission ticket specifies for arrival.

• Your identification. (You will not be admitted without some type of positive identification such as a student I.D. card with picture or a driver”s license.)

• Two No. 2 pencils with erasers.

• Watch (without an audible alarm)

Cell phone use is prohibited in both the test center and the testing room! If your cell phone is on, your scores will be canceled!

Note that calculator use is not allowed during the SAT Subject Test in Chemistry.
You should also go over this checklist:

A. Plan your activities so that you will have time for a good night”s sleep.

B. Lay out comfortable clothes for the next day. You may want to bring a snack.

C. Review the following helpful tips about taking the test:

1. Read the directions carefully.

2. In each group of questions, answer first those that you know. Temporarily skip difficult questions, but mark them in the margin so you can go back if you have time. Keep in mind that an easy question answered correctly counts as much as a difficult one.

3. Avoid haphazard guessing since this will probably lower your score. Instead, guess smart! If you can eliminate one or more of the choices to a question, it will generally be to your advantage to guess which of the remaining answers is correct. Your score will be based on the number right minus a fraction of the number answered incorrectly.

4. Omit questions when you have no idea of how to answer them. You neither gain nor lose credit for questions you do not answer.

5. Keep in mind that you have 1 hour to complete the test, and pace yourself accordingly. If you finish early, go back to questions you skipped.

6. Mark the answer grid clearly and correctly. Be sure each answer is placed in the proper space and within the oval. Erase all stray marks completely.

7. Write as much as you like in the test booklet. Use it as a scratch pad. Only the answers on the answer sheet are scored for credit.

D. Set your alarm clock so as to allow plenty of time to dress, eat your usual (or even a better) breakfast, and reach the test center without haste or anxiety.


After several weeks, most scores will be reported online at www.collegeboard.org. A full report will be available to you online a few days later. You can request a paper report. Your score will also be mailed to your high school and to the colleges, universities, and programs that you indicated. The report includes your scores, percentiles, and interpretive information. You can also get your scores—for a fee—by telephone. Call customer service at 866-756-7346 in the United States. From outside the United States, call 212-713-7789.

If your scores are not reported by eight weeks after the test date, definitely contact customer service by telephone or e-mail (sat@info.collegeboard.org). The mailing address for comments or questions about the tests is:

The College Board SAT Program
P.O. Box 025505
Miami, FL 33102