SAT Chemistry Subject Test - Practice Test 2

SAT Chemistry Subject Test - Practice Test 2

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
M = molar
n = number of moles
P = pressure
R = molar gas constant
S = entropy
T = temperature
V = volume
atm = atmosphere
g = gram(s)
J = joules(s)
kj = kilojoules
L = liter(s)
mL = milliliter(s)
mm = millimeter(s)
mol = mole(s)
V = volt(s)
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 terms:
(A) Boiling point
(B) Melting point
(C) Critical point
(D) Freezing point
(E) Triple point

1. The temperature and pressure at which three states of a substance may coexist

2. The temperature at which a solid becomes a liquid

3. The temperature of 373 K for H20 at standard pressure

4. The temperature at which the vapor pressure of a liquid equals the atmospheric pressure

Questions 5-7:
refer to the following diagram:

5. The ΔH of the reaction to form CO from C + O2

6. The ΔH of the reaction to form CO2 from CO + O2

7. The ΔH of the reaction to form CO2 from C + O2

Questions 8-11:
(A) Hydrogen bond
(B) Ionic bond
(C) Polar covalent bond
(D) Nonpolar covalent bond
(E) Metallic bond

8. The type of bond between atoms of potassium and chloride when they form a crystal of potassium chloride

9. The type of bond between the atoms in a nitrogen molecule

10. The type of bond between the atoms in a molecule of C02 (electronegativity difference = 1)

11. The type of bond between the atoms of calcium in a crystal of calcium

Questions 12-14:
refer to the following graphs:
12. The graph of volume vs. pressure for a gas at constant temperature

13. The graph of pressure vs. temperature for a gas at constant volume

14. The graph of volume vs. temperature for a gas at constant pressure

Questions 15-18:
(A) Least-reactive family of elements
(B) Alkali metals
(C) Halogen family of elements
(D) Noble gases
(E) Family whose oxides form acids in water

15. The elements that most actively react with water to release hydrogen

16. The elements least likely to become involved in chemical reactions

17. Family that contains elements in the colored gaseous state, in the liquid state, and with metallic properties

18. Group of nonmetallic elements containing N and P

Questions 19-23:
(A) 1s
(B) 2s
(C) 3s
(D) 3 p
(E) 3d

19. Electron energy sublevel filled by the first period of transition metals

20. The lowest energy orbital of those shown

21. Of the electron energy sublevels shown, the one that holds a maximum of 6 electrons

22. Largest of the orbitals with a spherical probability distribution

23. Orbital that describes the probability distribution for sodium’s outermost electron in the ground state



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 if statement II is true or false and fill in the corresponding T or F ovals on your answer sheet. *Fill in oval CE oniv 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.


Directions: Each of the questions or incomplete statements below is followed by five suggested answers or completions. Select the one that is best in each case and then fill in the corresponding oval on the answer sheet.

Question 24:
All of the following involve a chemical change EXCEPT
(A) the formation of HCl from H2 and Cl2
(B) the color change when NO is exposed to air
(C) the formation of steam from burning H2 and O2
(D) the solidification of vegetable oil at low temperatures
(E) the odor of NH3 when NH4Cl is rubbed together with Ca(OH)2 powder

Question 25:
When most fuels bum, the products include carbon dioxide and
(A) hydrocarbons
(B) hydrogen
(C) water
(D) hydroxide
(E) hydrogen peroxide

Question 26:
In the metric system, the prefix kilo- means
(A) 10°
(B) 10-1
(C) 1o-2
(D) 102
(E) 103

Question 27:
How many atoms are in 1 mole of water?

Question 28:
Which of the following elements normally exist as monoatomic molecules?
(A) Cl
(B) H
(C) O
(D) N
(E) He

Question 29:
The shape of a PCl3 molecule is described as
(A) bent
(B) trigonal planar
(C) linear
(D) trigonal pyramidal
(E) tetrahedral

Question 30:
The complete loss of an electron of one atom to another atom with the consequent formation of electrostatic charges is referred to as
(A) a covalent bond
(B) a polar covalent bond
(C) an ionic bond
(D) a coordinate covalent bond
(E) a pi bond between p orbitals

Question 31:
In the decomposition of water with electricity (electrolysis), the following reaction occurs.
2H2O() → 2 H2(g) + O2(g)
The hydrogen is
(A) oxidized from +1 to 0
(B) oxidized from 0 to +1
(C) reduced from 0 to +1
(D) reduced from +1 to 0
(E) not changing oxidation states

Question 32:
Which of the following radiation emissions has no mass?
(A) Alpha particle
(B) Beta particle
(C) Proton
(D) Neutron
(E) Gamma ray

Question 33:
If a radioactive element with a half-life of 100 years is found to have transmutated so that only 25% of the original sample remains, what is the age, in years, of the sample?
(A) 25
(B) 50
(C) 100
(D) 200
(E) 400

Question 34:
What is the pH of an acetic acid solution if the[H3O+] = 1 × 10-4 mole/liter?
(A) 1
(B) 2
(C) 3
(D) 4
(E) 5

Question 35:
The polarity of water is useful in explaining which of the following?
I. The solution process
The ionization process
III. The high conductivity of distilled water
(A) I only
(B) II only
(C) I and II only
(D) II and III only
(E) I, II, and III

Question 36:
When sulfur dioxide is bubbled through water, the solution will contain
(A) sulfurous acid
(B) sulfuric acid
(C) hyposulfuric acid
(D) persulfuric acid
(E) anhydrous sulfuric acid

Question 37:
Four grams of hydrogen gas at STP contain
(A) 6.02 × 1023 atoms
(B) 12.04 × 1023 atoms
(C) 12.04 × 1046 atoms
(D) 1.2 × 1023 molecules
(E) 12.04 × 1023 molecules

Question 38:
Analysis of a gas gave: C = 85.7% and H = 14.3%. If the formula mass of this gas is 42 atomic mass units, what are the empirical formula and the true formula?
(A) CH; C4H4
(B) CH2; C3H6
(C) CH3; C3H9
(D) C2H2; C3H6
(E) C2H4; C3H6

Question 39:
Which fraction would be used to correct a given volume of gas at 303 K to its new volume when it is heated to 333 K and the pressure is kept constant?
 Question 40:
Which of the substances listed decreases the freezing point of benzene (C6H6) more than the others if a lab tech tries to dissolve 5.00 grams in 500.0 g of benzene?
(A) paradichlorobenzene, C6H4Cl2
(B) sodium chloride, NaCl
(C) aluminum chloride, AlCl3
(D) ethanol, C2H5OH
(E) sucrose, C12H22O11

Question 41:
What is the approximate pH of a 0.005 M ' solution of H2SO4?
(A) 1
(B) 2
(C) 5
(D) 9
(E) 13

Question 42:
How many grams of NaOH are needed to make 100 grams of a 5% solution?
(A) 2
(B) 5
(C) 20
(D) 40
(E) 95

Question 43:
For the Haber process: N2 + 3H2 ⇌ 2NH3 + heat (at equilibrium), which of the following statements concerning the reaction rate is/are true?
I. The reaction to the right will increase when pressure is increased.
II. The reaction to the right will decrease when the temperature is increased.
III. The reaction to the right will decrease when NH3 is removed from the chamber.
(A) I only
(B) II only
(C) I and II only
(D) II and III only
(C) I, II, and III

Question 44:
If you titrate 1.0 M H2SO4 solution against 50. milliliters of 1.0 M NaOH solution, what volume of H2SO4, in milliliters, will be needed for neutralization?
(A) 10.
(B) 25.
(C) 40.
(D) 50.
(E) 100.

Question 45:
How many grams of CO2 can be prepared from 150 grams of calcium carbonate reacting with an excess of hydrochloric acid solution?
(A) 11
(B) 22
(C) 33
(D) 44
(E) 66

Question 46:
refers to the following diagram:

46. The diagram represents a setup that may be used to prepare and collect
(A) NH3
(B) NO
(C) h2
(D) S03
(E) C02

Question 47:
The lab setup shown above was used for the gravimetric analysis of the empirical formula of MgO. In synthesizing MgO from a Mg strip in the crucible, which of the following is NOT true?
(A) The initial strip of Mg should be cleaned.
(B) The lid of the crucible should fit tightly to exclude oxygen.
(C) The heating of the covered crucible should continue until the Mg is fully reacted.
(D) The crucible, lid, and the contents should be cooled to room temperature before measuring their mass.
(E) When the Mg appears to be fully reacted, the crucible lid should be partially removed and heating continued.

Questions 48-50
refer to the following experimen­tal scenario and data:
A sample of the hydrated form of barium chloride is placed into an evaporating dish and heated strongly for 20 minutes on a hot plate. After cooling in a desiccator, the mass of the dish and contents is determined. The whole system is then heated strongly again for 5 minutes and cooled. Then the mass is measured again. The heating, cooling, and measuring of the mass is repeated a third time.
Mass of evaporating dish                     = 14.30 g
Mass of dish and hydrated BaCl2       = 26.50 g
Mass of dish and anhydrous BaCl2    = 24.77 g (after 1st heating)
Mass of dish and anhydrous BaCl2     = 24.70 g (after 2nd heating)
Mass of dish and anhydrous BaCl2    = 24.70 g (after 3rd heating)

48. The purpose of cooling the anhydrous BaCl2 in the desiccator is to
(A) protect the experimenter from the heat of the dish
(B) cool the dish slowly
(C) keep the contents of the dish from rehydrating
(D) cool the dish quickly
(E) rehydrate the contents of the dish

49. If the person running the experiment failed to notice a small amount of the solid substance jump out of the evaporating dish during heating, the resulting data would suggest the amount of water in the hydrate is
(A) less than what it should be
(B) the same as if nothing jumped out of the dish
(C) more than what it should be
(D) equal to the mass of the BaCl2 that also left the dish
(E) more than the mass of the BaCl2 that also left the dish

50. The number of moles of water released from the hydrated barium chloride after three rounds of heating is
(A) 0.0200 mol
(B) 0.0400 mol
(C) 0.0800 mol
(D) 0.100 mol
(E) 1.20 mol

Questions 51:
What is the mass, in grams, of 1 mole of KAl(SO4)2 12H2O?
(A) 132
(B) 180
(C) 394
(D) 474
(E) 516

Questions 52:
What mass of aluminum will be completely oxidized by 2 moles of oxygen at STP?
(A) 18 g
(B) 8 g
(C) 4 g
(D) 0 g
(E) 8 g

Questions 53:
In general, when metal oxides react with water, they form solutions that are
(A) acidic
(B) basic
(C) neutral
(D) unstable
(E) colored

Questions 54-56
refer to the following spontane­ous reactions:
Zn(s) + SnCl2(aq) → Sn(s) + ZnCl2(aq)
Sn(s) + CuCl2(aq) → Cu(s) + SnCl2(aq)

54. The significant driving force associated with both of these reactions is
(A) the transfer of electrons
(B) the transfer of protons
(C) the transfer of mass
(D) the transfer of ions
(E) the transfer of atoms

55. The most active of the metals shown is
(A) Sn
(B) Cu
(C) Zn
(D) Cl
(E) ZnCl2

56. Substances acting as oxidizing agents include
(A) tin(II) chloride
(B) tin
(C) zinc chloride
(D) both tin and copper(II) chloride
(E) both zinc chloride and copper(II) chloride

Questions 57:
How many liters of oxygen (STP) can be prepared from the decomposition of 212 grams of sodium chlorate (1 mol = 106 g)?
(A) 2
(B) 4
(C) 8
(D) 2
(E) 4

Questions 58:
In this equation: Al(OH)3 + H2SO4 → Al2(SO4)3 + H2O, the whole-number coefficients of the balanced equation are
(A) 1, 3, 1, 2
(B) 2, 3, 2, 6
(C) 2, 3, 1, 6
(D) 2, 6, 1, 3
(E) 1, 3, 1, 6

Questions 59:
(A) -183.9 kcal
(B) -91.9 kcal
(C) +45.3 kcal
(D) +22.5 kcal
(E) -12.5 kcal

Questions 60:
Isotopes of an element are related because which of the following is (are) the same in these isotopes?
I. Atomic mass
II. Atomic number
III. Arrangement of orbital electrons
(A) I only
(B) II only
(C) I and II only
(D) II and III only
(E) I, II, and III

Questions 61:
In the reaction of zinc with dilute HCl to form H2, which of the following will increase the reaction rate?
I. Increasing the temperature
II. Increasing the exposed surface of zinc
III. Using a more concentrated solution of HCl
(A) I only
(B) II only
(C) I and III only
(D) II and III only
(E) I, II, and III

Questions 62:
The laboratory setup shown above can be used to prepare a
(A) gas less dense than air and soluble in water
(B) gas more dense than air and soluble in water
(C) gas soluble in water that reacts with water
(D) gas insoluble in water
(E) gas that reacts with water

Questions 63:
In this reaction: CaCO3 + 2HCl → CaCl2 + H2O + CO2. If 4.0 moles of HCl are available to the reaction with an unlimited supply of CaCO3, how many moles of CO2 can be produced at STP?
(A) 0
(B) 5
(C) 0
(D) 5
(E) 0

Questions 64:
A saturated solution of PbS at 25°C contains 5.0 × 10-14 mole/liter of Pb2+ What is the Ksp value of this salt?
(A) 5.0 × 10-14
(B) 5.0× 10-15
(c)5.0× 10-14
(D) 5.0 × 10-28
(E) 5.0× 10-27

Questions 65:
If 0.1 mole of K2S was added to the solution in question 64, what would happen to the Pb2+ concentration?
(A) It would increase.
(B) It would decrease.
(C) It would remain the same.
(D) It would first increase, then decrease.
(E) It would first decrease, then increase.

Questions 66:
Which of the following will definitely cause the volume of a gas to increase?
I. Decreasing the pressure with the temperature held constant.
II. Increasing the pressure with a temperature decrease.
III. Increasing the temperature with a pressure increase.
(A) I only
(B) II only
(C) I and III only
(D) II and III only
(E) I, II, and III

Questions 67:
The number of oxygen atoms in 0.50 mole of Al2(CO3)3 is
(A) 4.5 × 1023
(B) 9.0 × 1023
(C) 3.6 × 1024
(D) 2.7 × 1024
(E) 5.4 × 1024

Question 68:
refers to a solution of 1 M acid, HA, with Ka = 1 × 10-6.

68. What is the H3O+ concentration? (Assume [HA] = 1, [H3O+] = x, [A-] = x.)
(A) 1 × 10-5
(B) 1 × 10-4
(C) 1 × 10-2
(D) 1 × 10-3
(E) 0.9 × 10-3

Question 69:
What is the percent dissociation of acetic acid in a 0.1 M solution if the [H3O+] is 1 × 10-3 mole /liter?
(A) 0.01%
(B) 0.1%
(C) 1.0%
(D) 1.5%
(E) 2.0%

Practice Test 2


Answer 1:
(E) A phase diagram shows that all three states can exist at the triple point.

Answer 2:
(B) When a solid turns into a liquid, the process is called melting. This happens at a particular temperature when the pressure on the solid is constant. It’s called the melting point.

Answer 3:
(A) Water boils at 100.0°C under standard pressure (1.00 atmosphere). On the absolute temperature scale, that’s 373 K.

Answer 4:
(A) When the vapor pressure of a liquid is the same as the atmospheric pressure, bubbles can be created in the body of the liquid and boiling can take place.

Answer 5:
(B) The first step is the ΔHfor C + 1/2 O2 -> CO. It releases -110.5kJ of heat. This is written as -110.5 kj because it is exothermic.

Answer 6:
(C) This is the second step on the diagram. It releases -283.0 kj of heat.

Answer 7:
(A) To arrive at the AH,take the total drop (-393.5 kj) or add these reactions:

Answer 8:
(B) Potassium and chlorine have a large enough difference in their electronegativities to form ionic bonds. The respective positions of these two elements in the periodic chart also are indicative of the large difference in their electronegativity values.

Answer 9:
(D) TWo atoms of an element that forms a diatomic molecule always have a nonpolar covalent bond between them since the electron attraction or electronegativity of the two atoms is the same.

Answer 10:
(C) Electronegativity differences between 0.5 and 1.7 are usually indicative of polar covalent bonding. CO2 is an interesting example of a nonpolar molecule with polar covalent bonds since the bonds are symmetrical in the molecule.

Answer 11:
(E) Calcium is a metal and forms a metallic bond between atoms.

Answer 12:
(E) This graph shows the volume decreasing as the pressure is increased and the tem­perature is held constant. It is an example of Boyle’s Law (PV = k).

Answer 13:
(A) This graph shows the pressure increasing as the temperature is increased and the volume is held constant. It is an example of Gay-Lussac’s Law (P/T = k).

Answer 14:
(C) This graph shows the volume increasing as the temperature is increased and the pressure is held constant. It is an example of Charles’s Law (V/T= k).

Answer 15:
(B) The alkali metals react with water to form hydroxides and release hydrogen. A typical reaction is:
2Na(s) + 2H2O(ℓ) → 2NaOH(aq) + H2(g)

Answer 16:
(D) The noble gases are the least reactive because of their completed outer orbital.

Answer 17:
(C) The halogen family contains the colored gases fluorine and chlorine at room tem¬peratures, the reddish liquid bromine, and metallic-like purple iodine.

Answer 18:
These nonmetals, when they are oxides, react as acidic anhydrides with water to form acid solutions.

Answer 19:
(E) The first transition metal on the Periodic Table is scandium. It is in the third period. Scandium’s electron notation shows that this element is the first to have the 3d sublevel filled. The nine elements that follow scandium fill up the 3d sublevel.

Answer 20:
The lowest-energy orbital in any atom is the Is. The Is orbital represents the prob¬ability distribution for the electron when it is closest to the nucleus.

Answer 21:
The type of sublevel that can hold 6 electrons is called p type. A p sublevel contains 3 degenerate orbitals, each with a capacity of 2 electrons.

Answer 22:
(C) An s-type orbital has a spherical probability distribution for the electrons in that energy state. Energy level 3 represents a higher energy than energy levels 1 or 2; there¬fore, the 3s orbital is larger than the Is or 2s orbitals. The higher the energy of the elec¬trons, the farther away the electrons will likely be found from the nucleus.

Answer 23:
The electron notation for sodium is 1s2 2s2 2p6 3s1. Hence, the outermost electron in the sodium atom is in the 3s orbital.

Answer 101:
(T, F)
Sulfur trioxide is shown by three structural formulas because each bond is "hybrid” of a single and double bond. Resonance in chemistry does not mean that the bonds resonate between the structures shown in the structural drawing.

Answer 102:
(T, F)
When AG is negative in the Gibbs equation, the reaction is spontaneous. However, the total equation determines this, not just the AH. The Gibbs equation is:
ΔG = ΔH - TΔS.

Answer 103:
(T, T, CE)
One mole of each of the substances contains 6.02 × 1023 molecules, but their molecular masses are different. CO2 is found by adding one C = 12 and two O = 32, or a total of 44 amu. The H2O, however, adds up to two H = 2 plus one O = 16, or a total of 18 amu. Thus, it is true that 1 mol of CO2 at 44 g/mol is heavier than 1 mol of H2O at 18 g/mol.

Answer 104:
(T, T, CE)
Hydrosulfuric acid is a weak acid but is used in qualitative tests because of the distinctly colored precipitates of sulfides that it forms with many metallic ions.

Answer 105:
(T, T, CE) Sodium chloride is an ionic crystal, not a molecule, and its ions are hydrated by the polar water molecules.

Answer 106:
(F, F) The addition of more H3PO4 causes the equilibrium to shift to the right and increase the concentration of H3O+ ions until equilibrium is restored. Therefore, the first statement is false. The second is also false since the equilibrium constant remains the same at a given temperature.

Answer 107:
(T, T, CE) The statement is true, and the reason is also true and explains the statement.

Answer 108:
(T, F) The statement is true, but not the reason. In an equilibrium reaction, concentrations can be shown to progress like this:
until equilibrium is reached. Then the concentrations stabilize.

Answer 109:
(F, T) The forward and reverse reactions are occurring at equal rates when equilibrium is reached. The reactions do not stop. The concentrations remain the same at this point.

Answer 110:
(T, T, CE) Since acetylene is known to be a linear molecule with a triple bond between the two carbons, the sp orbitals along the central axis with the hydrogens bonded on either end fit the experimental evidence.

Answer 111:
(F, F)
The weakest bonds between molecules are van der Waals forces, not coordinate covalent bonds.

Answer 112:
(T, T, CE)
The terms dilute and concentrated merely indicate a relatively small amount of solvent and a large amount of solvent, respectively. You can have a dilute saturated solution if the solute is only slightly soluble.

Answer 113:
(T, T)
Lithium replaces any other metal in a single replacement reaction; therefore, it is the most active metal. It’s also true that lithium has only one electron in its outer energy level. Although lithium’s outer configuration of only one electron aids in its high activity, it is not the only reason for its high activity. Other metals (like sodium and potassium) have only one electron in their outer energy levels and are fairly active, but they aren’t as active as lithium.

Answer 114:
(F, T)
The oxidation state of carbon is often +4 but not always. Carbon does have 4 valence electrons that it often relinquishes responsibility for in terms of oxidation numbers (which is why carbon can have an oxidation state of +4); however, it doesn’t have to do so, and in many compounds it exhibits a different oxidation state.

Answer 115:
(T, T, CE) There are as many electrons as there are protons in a neutral atom, and the atomic number represents the number of protons.

Answer 116:
(F, T)
The first two principal energy levels fill up at 2 and 8 electrons, respectively. That leaves 7 electrons to fill the 3s and 3p orbitals like this: 3s2, 3p5. With only one electron missing in the 3p orbitals, the most likely oxidation number is -1.

Answer 24:
(D) The solidification of vegetable oil is merely a physical change, like the formation of ice from liquid water at lower temperatures. All the other choices involve actual recombinations of atoms and thus are chemical changes.

Answer 25:
Water is formed because most common fuels contain hydrogen in their structures.

Answer 26:
The other choices, in order, represent 1, 1/10 or deci-,1/100 or centi-, and 100 or hecto-.

Answer 27:
One mole of any substance contains 6.02 × 1023 molecules. Since each water mol¬ecule is triatomic, there would be 3(6.02 × 1023) atoms present.

Answer 28:
The noble gases are all monatomic because of their complete outer energy levels. A rule to help you remember diatomic gases is:- Gases ending in -gen or -ine usually form diatomic molecules.

Answer 29:
By both the VSEPR (Valence Shell Electron Pair Repulsion) method and the orbital structure method, the PCl3 molecule is trigonal pyramidal:

Answer 30:
The complete loss and gain of electrons is an ionic bond. All other bonds indicated are “sharing of electrons” type bonds or some form of covalent bonding.

Answer 31:
The cathode reaction releases only H2 gas. This half-reaction is as given in (D).

Answer 32:
The beta particle is a high-speed electron and has the smallest mass of the first four choices. However, gamma rays are electromagnetic waves. They have no mass.

Answer 33:
(D) If 25% of the sample now remains, then 100 years ago 50% would be present. If you go back another 100 years, the sample would contain 100% of the radioactive element. Therefore, the sample is 100 + 100 = 200 years old.

Answer 34:
pH = -log[H30+] = —log[l x 10-4] = -(-4) = 4.

Answer 35:
Only I and II are true. Distilled water does not significantly conduct an electric current. The polarity of the water molecule is helpful in ionization and in causing sub¬stances to go into solution.

Answer 36:
SO2 is the acid anhydride of H2SO3 or sulfurous acid. H2O + SO2 -> H2SO3.

Answer 37:
(E) Four grams of hydrogen gas at STP represent 2 mol of hydrogen since 2 g is the gram-molecular mass of hydrogen. Each mole of a gas contains 6.02 × 1023 molecules, so 2 mol contains 2 × 6.02 × 1023 or 12.04 × 1023 molecules.

Answer 38:
To solve percent composition problems, first, using estimation, divide the % given by the atomic mass:
The empirical formula is CH2. Since the molecular mass is 42 and the empirical formula has a molecular mass of 14, the true formula must be 3 times the empirical formula, or C3H6.

Answer 39:
Because the temperature (in kelvins) increases from 303 K to 333 K, the volume of
the gas should increase with pressure held constant. The correct fraction is 333/303.

Answer 40:
Paradichlorobenzene, a nonpolar solid, is the only substance that will appreciably dissolve in benzene. Therefore, it will be the only substance capable of depressing the freezing point of benzene.

Answer 41:
The pH is -log[H+]. A 0.005 molar solution of H2SO4 ionizes in a dilute solution to release two H+ ions per molecule of H2SO4. Therefore, the molar concentration of H+ ion is 2 × 0.005 mol/L or 0.010 mol/L. Substituting this in the formula, you have:
pH = -log [0.01] = -log [1 × 10-2]
The log of a number is the exponent to which the base 10 is raised to express that number in exponential form:
              -log [1 × 10-2] = -[-2] = 2

Answer 42:
If the solution is to be 5% sodium hydroxide, then 5% of 100 g is 5 g. Percent is always by mass unless otherwise specified.

Answer 43:
Because this equation is exothermic, higher temperatures will decrease the reac¬tion to the right and increase the reaction to the left, so II is true.
Also, I is true because with an increase of pressure the reaction will try to relieve that pressure by going in the direction that has the least volume: in this reaction, to the right. Statement III is false because removing product in this reaction would increase the forward reaction. Statements I and II are true.

Answer 44:
The reaction is:
Answer 45:

Answer 46:
(E) The other choices are wrong because:
(A) is less dense than air and will escape through hole in stopper
(B) reacts with air
(C) is less dense than air and will escape through hole in stopper
(D) needs heat to be evolved

Answer 47:
The Mg needs oxygen to form MgO so the lid cannot be tightly sealed. Oxygen is needed for the Mg to oxidize to MgO. All other choices are true.

Answer 48:
A desiccator is a device used to keep things dry. After the hydrated BaCl2 has had the water removed, the desiccator keeps the water from returning to the substance.

Answer 49:
If material jumped out of the evaporating dish while heating, barium chloride and water would both be removed. The purpose of the heating is to remove only the water. Therefore, it would seem like more water had been removed than was actually present, and the mass lost from the barium chloride would inaccurately be counted as mass lost from water.

Answer 50:
As a result of the complete heating, 1.80 grams of water is removed from the hydrate. The mass of water released from the barium chloride can be converted into moles by using dimensional analysis.
Answer 51:
1K = 39, 1A1 = 27, 2(SO4) = 2(32 + 16 × 4) = 192, and 12H2O = 12(2 + 16) = 216. This totals 474 g.

Answer 52:
Answer 53:
Metal oxides are generally basic anhydrides.

Answer 54:
Both of the reactions shown are single replacement reactions. Single replacement reactions are also redox processes in which electrons are transferred from one sub¬stance to another. They can be viewed as a driving force for reactions to occur.

Answer 55:
(C) Tin replaces copper in reaction #2 and is therefore more active than copper. Zinc replaces tin in reaction #1, however, and is therefore more active than both tin and copper. This is reflected in the activity series where zinc is at a higher level.

Answer 56:
Zinc is being oxidized (from 0 to +2) by the tin(II) chloride in the first reaction. Therefore, tin(II) chloride is the oxidizing agent in that reaction. Copper(II) chloride is the oxidizing agent in the second reaction, but is not matched up with another oxidizing agent in any of the other answers.

Answer 57:

Answer 58:
(C) The balanced equation has the coefficients 2, 3, 1, and 6: 2Al(0H)3 + 3H2SO4→ Al2(SO4)3 + 6H2O.

Answer 59:

Answer 60:
II and III are identical; isotopes differ only in the number of neutrons in the nucleus and this affects the atomic mass only.

Answer 61:
I, II, and III will increase the rate of this reaction. Each of them causes the rate of this reaction to increase.

Answer 62:
This setup depends on water displacement of an insoluble gas.

Answer 63:
The coefficients give the molar relations, so 2.0 mol of HCl give off 1.0 mol of CO2. Given 4.0 mol of HCl, you have

Answer 64:
(E) Ksp = [Pb2+] [S2-]. Since [Pb2+] is given at 5.0 × 10-14 and because the formula is a 1:1 ratio, the [S2-] must also be that value, the Ksp can be calculated by
Ksp = [5.0 × 10-14] [5.0 × 1014] = 25 × 10-28
Putting the answer in scientific notation, i.e., one digit to the left of the decimal point in the first factor, the answer is 2.5 × 10-27.

Answer 65:
The introduction of the “common ion” S2- at 0.1 molar forces the equilibrium to shift to the left and reduce the Pb2- concentration.

Answer 66:
According to the gas laws, only I will definitely cause an increase in the volume of a confined gas.

Answer 67:
In 1 mol of Al2(CO3)3, nine oxygens (three carbonates with three oxygen atoms each) are in each formula unit, or 9 mol of O atoms are in 1 mol of Al2(CO3)3. Because only 0.50 mol is given, there are 1/2 (9) or 4.5 mol of O atoms. In 4.5 mol of oxygen, there are
Answer 68:
(D) When HA ionizes, it forms equal amounts of H+ and A- ions, but these amounts are very small because the Ka is very small. Ka can be expressed as [H+] [A-]/[HA]. Because you are told to assume [HA] = 1, you have:

Answer 69:


Your score on Practice Test 2 can now be computed manually. The actual test is 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!

After taking Practice Test 2, 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.



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