Practice Test 2

  1. Part A

  2. Directions

    Each set of lettered choices below refers to the numbered formulas or statements immediately following it. Select the one lettered choice that best fits each formula or statement. In each set, a choice may be used once, more than once, or not at all.

    Note

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

  3. Questions 1–5

    1. Ca+ and K
    2. H+ and He
    3. Cl− and F
    4. O– and S+
    5. Na+ and O−

    1. Difference of 6 electrons

    2. Same number of electrons

    3. Difference of 9 electrons

    4. Difference of 2 electrons

    5. Difference of 1 electron

  6. Questions 6–9

    1. Solute
    2. Solvent
    3. Solubility
    4. Aqueous solution
    5. Solvation

    1. Is present in a lesser amount in a solution

    2. Solvent is water

    3. Is present in greater quantity in a solution

    4. Interaction between the solute and solvent molecules

  9. Questions 10–12

    1. Sublimation
    2. Condensation
    3. Evaporation
    4. Deposition
    5. Melting

    1. Gas → Solid

    2. Gas → Liquid

    3. Solid → Gas

  12. Questions 13–16

    1. Proton
    2. Neutron
    3. Electron
    4. Isotope
    5. Ion

    1. Neutral charge; 1 amu

    2. Positive charge

    3. Negligible weight

    4. Negative charge; pairs with opposite spin

  15. Questions 17–19

    1. K
    2. As
    3. Be
    4. Se
    5. Ir

    1. Metalloid

    2. Nonmetal

    3. Class IA metal

  18. Questions 20–23

    1. O2
    2. CO2
    3. N2
    4. H2SO4
    5. He

    1. Inert gas, not very soluble in water

    2. Very soluble in water, forming a very acidic solution

    3. When ignited, burns with a blue flame and is not very soluble in water

    4. Will not burn, is not very soluble in water, and makes a weakly acidic solution

  1. Part B

  2. Directions

    Each question below consists of two statements, statement I in the left-hand column and statement II in the right-hand column. For each question, determine whether statement I is true (T) or false (F) and whether statement II is true (T) or false (F). Then, determine if statement II is a correct explanation (CE) of statement I.

  3. I II
    The most important factor in determining the chemical properties of an element is the number of electrons in the outermost shell BECAUSE the number of electrons in the outer shell determines the bonding characteristics of the element.
  4. I II
    NH3 is a Lewis base BECAUSE it can accept a proton.
  5. I II
    An element (X) with an atomic number of 16 has 14 electrons in X2− BECAUSE two protons bind the two outermost electrons.
  6. I II
    Bromine is a stronger oxidizing agent than chlorine BECAUSE it has a large atomic radius.
  7. I II
    An exothermic reaction has a positive ΔH BECAUSE heat must be added to the system for the reaction to occur.
  8. I II
    Iron is an element BECAUSE it cannot be broken into smaller units and retain its physical and chemical properties.
  9. I II
    CCl4 is a nonpolar compound BECAUSE the dipole moments are canceled out.
  10. I II
    A basic solution has more hydrogen ions than an acidic solution BECAUSE pH is defined as the −log [H+].
  11. I II
    Atomic radii decrease down a group BECAUSE the higher the atomic number within a group, the smaller the atom.
  12. I II
    When an ideal gas is cooled, its volume will increase BECAUSE temperature and volume are proportional.
  13. I II
    A 0.2 M solution of carbonic acid is a weaker conductor of electricity than a 0.2 M solution of HBr BECAUSE in solutions with the same concentration of solute molecules, H2CO3 is less dissociated than HBr.
  14. I II
    An equation where two gas molecules combine to form one gas molecule in equilibrium will increase the yield of the product when the pressure is increased BECAUSE increased pressure always favors products.
  15. I II
    Water makes a good buffer BECAUSE a good buffer will resist changes in pH.
  16. I II
    In the kinetic theory of gases, collisions between gas particles and the walls of the container are considered elastic BECAUSE gas molecules are considered pointlike, volumeless particles with no intermolecular forces and in constant, random motion.
  17. I II
    The oxidation state of Cr in Al2(Cr2O7)3 is +3 BECAUSE as a neutral compound, the sum of oxidation numbers of all the atoms must equal zero.
  18. I II
    The entropy of a solid increases when it is dissolved in solvent BECAUSE it becomes less ordered.

Return to the section of your answer sheet you started for chemistry and answer questions 24–69.

  1. Part C

  2. Directions

    Each of the incomplete statements or questions below is followed by five suggested completions or answers. Select the one that is best for each case.

  3. Which of the following is a Lewis base?

    1. NH4+
    2. CH4
    3. PH3
    4. CH3CH3
    5. HCl

  4. When one mole of sulfur burns to form SO2, 1,300 calories are released. When one mole of sulfur burns to form SO3, 3,600 calories are released. What is the ΔH when one mole of SO2 is burned to form SO3?

    1. 3,900 cal
    2. −1,950 cal
    3. 1,000 cal
    4. −500 cal
    5. −2,300 cal

  5. How many atoms are in a 36.5 g sample of SF6 gas?

    1. 1.51 × 1022
    2. 1.06 × 1022
    3. 1.51 × 1023
    4. 1.06 × 1024
    5. 1.51 × 10−24

  6. Which of the following statements about molecular and empirical formulas is (are) false?

    1. A given compound can have the same molecular and empirical formula.
    2. The molecular formula is a whole-number multiple of the empirical formula.
    3. H2O2 represents the empirical formula of hydrogen peroxide.
    1. III only
    2. I and II only
    3. II and III only
    4. I, II, and III
    5. I only

  7. The product formed when oxygen and hydrogen are mixed in a test tube at room temperature is

    1. hydrogen peroxide
    2. water
    3. a base
    4. a zwitterion
    5. no reaction takes place

  8. All of the following statements are consistent with Bohr’s model of the atom EXCEPT

    1. an electron may assume an infinite number of velocities
    2. an atom is most stable when its electronic configuration is that of the ground state
    3. the electron shell numbers represent the principal energy levels
    4. electrons in orbitals closest to the nucleus have the lowest energy
    5. they are all consistent

  9. What is the amount of heat given off by 100 g of O2 when it is used to burn an excess of sulfur according to the following reaction?

    1. 925,000 J
    2. 29,000 J
    3. 1,850 J
    4. 296 J
    5. 100 J

  10. If electricity costs 10 cents/coulomb, which of the following would have the highest cost/mole?

    1. Copper from CuSO4
    2. Sodium from NaCl
    3. Chlorine from KCl
    4. Hydrogen from H2O
    5. Iron from FeCl3

  11. What volume of water would be needed to dilute 50 mL of a 3 M HCl solution to 1 M?

    1. 25 mL
    2. 50 mL
    3. 75 mL
    4. 100 mL
    5. 150 mL

  12. If the pressure of a gas sample is doubled at constant temperature, the volume will be

    1. 4 times the original
    2. 2 times the original
    3. 1/2 of the original
    4. 1/4 of the original
    5. 1/8 of the original

  13. Which of the following molecules has a trigonal pyramidal geometry?

    1. BH3
    2. H2O
    3. CH4
    4. NH3
    5. AlCl3

  14. The field of organic chemistry focuses mainly on

    1. oxygen-containing compounds
    2. carbon-containing compounds
    3. proteins and enzymes
    4. carbohydrates
    5. nitrogen-containing compounds

  15. Which of the following is the most electronegative element?

    1. He
    2. I
    3. N
    4. O
    5. C

  16. Which of the following has the greatest affinity for electrons?

    1. Na
    2. Cl
    3. Br
    4. K
    5. C

  17. What would be the approximate weight of 1.204 × 1024 bromine atoms?

    1. 80 grams
    2. 120 grams
    3. 160 grams
    4. 180 grams
    5. 200 grams

  18. What is the oxidation state of Mn in KMnO4?

    1. −7
    2. −3
    3. 0
    4. +3
    5. +7

  19. What is the daughter element produced by technetium-99m (an isomer of 99Tc with atomic number 43 and mass number 99) after γ decay?

  20. What volume of water would be needed to dilute 50 mL of 3 M H2SO4 to 0.75 M?

    1. 50 mL
    2. 100 mL
    3. 150 mL
    4. 200 mL
    5. 250 mL

  21. All members of Group IA have similar reactivity because

    1. they have the same number of protons
    2. they have the same number of electrons
    3. they have similar outer-shell electron configurations
    4. they have valence electrons with the same quantum numbers
    5. they have the same number of neutrons

  22. If one mole of a gas originally at STP is placed in a container where the pressure is doubled and the temperature in K is tripled, what is the new volume in L?

    1. 2.2
    2. 5.6
    3. 7.5
    4. 11.2
    5. 33.6

  23. Chlorophyll, a green pigment involved in the light reactions of photosynthesis, consists of 2.4312 percent Mg. If you are given a 100 g sample of chlorophyll, how many atoms of Mg will it contain?

    1. 6.02 × 1022 atoms
    2. 6.02 × 1023 atoms
    3. 6.02 × 1024 atoms
    4. 6.02 × 1025 atoms
    5. None of the above

  24. The reverse reaction for the one shown above is exothermic. If the temperature is lowered, what will occur?

    1. The pH will decrease.
    2. The equilibrium will shift to the right.
    3. The concentration of H3O+ ions will increase.
    4. The equilibrium will shift to the left.
    5. Temperature does not affect equilibrium.

  25. Which of the following reactions shows a decrease in entropy?

    1. C (s) + 2H2 (g) → CH4 (g)
    2. H2O (g) → H2 (g) + 1/2O2 (g)
    3. 2NI3 (s) → N2 (g) +3I2 (g)
    4. 2O3 (g) → 3O2 (g)
    5. None of the above

  26. Which of the following is an incorrect association?

    1. Mendelev–periodic table
    2. Faraday–electrolytic cells
    3. Millikan–charge of electrons
    4. Rutherford–photoelectric effect
    5. They are all correct.

  27. Which of the following is true of an electrolytic cell?

    1. An electric current causes an otherwise nonspontanteous chemical reaction to occur.
    2. Reduction occurs at the anode.
    3. A spontaneous electrochemical reaction produces an electric current.
    4. The electrode to which electrons flow is where oxidation occurs.
    5. None of the above
  28. 1st Ionization Energy (eV) 2nd Ionization Energy (eV)
    U 5.6 5.6
    V 1.7 2.9
    X 1.1 13.6
    Y 12.4 2.8
    Z 2.9 1.7

    From the information given in the table above, which of the following is most probably a Group IA metal?

    1. U
    2. V
    3. X
    4. Y
    5. Z

  29. Three canisters, A, B, and C, are all at the same temperature, with volumes of 2.0, 4.0, and 6.0 liters, respectively. Canister A contains 0.976 g of argon gas at a pressure of 120 torr, canister B contains 1.37 g of nitrogen gas at a pressure of 210 torr, and canister C is completely empty at the start. Assuming ideality, what would the pressure become in canister C if the contents of A and B are completely transferred to C?

    1. 180 torr
    2. 330 torr
    3. 675 torr
    4. 0.25 atm
    5. None of the above

  30. H2O has a higher boiling point than HF because

    1. H2O is more polar than HF
    2. H2O can form more hydrogen bonds
    3. H2O has a higher molecular weight
    4. H2O has more atoms
    5. H2O does not have a higher boiling point than HF

  31. Questions 52–54 refer to the following graph.

    1. In what pressure range will the compound sublime?

      1. Pressures less than P(b)
      2. Pressures between P(b) and P(c)
      3. Pressures between P(d) and P(c)
      4. Pressures between P(b) and P(d)
      5. The compound will not sublime under any pressure.

    2. The phase change represented by crossing curve bd is

      1. evaporation
      2. sublimation
      3. condensation
      4. melting
      5. boiling

    3. The phase change represented by crossing curve bc is

      1. freezing
      2. melting
      3. deposition
      4. sublimation
      5. evaporation

  34. If you mix 3 L of 0.5 M NaCl with 9 L of 0.2777 M NaCl, what will the concentration of the final solution be, assuming that volumes are additive?

    1. 0.33 M
    2. 0.39 M
    3. 0.5777 M
    4. 0.7777 M
    5. None of the above

  35. Which of the following statements about boiling-point elevation is NOT true?

    1. Addition of a nonvolatile solute raises the vapor pressure of a solution. Therefore, a higher temperature will be needed to produce boiling.
    2. The molal concentration of solute particles dissolved in the solvent is an important factor in determining the molal boiling-point elevation constant.
    3. The identity of solute particles dissolved in the solvent is an important factor in determining the molar boiling-point elevation constant.
    1. I only
    2. II only
    3. I and III only
    4. II and III only
    5. I, II, and III

  36. What is the sum of the coefficients of the products for the following reaction, after balancing?

    …K2Cr2O7 + …HCl → …KCl + …CrCl3 + …H2O + …Cl2
    1. 10
    2. 12
    3. 13
    4. 14
    5. 15

  37. Questions 58–59 refer to the following equation:

    1. What is X?

    2. This nuclear reaction is an example of

      1. α decay
      2. β− decay
      3. β+ decay
      4. fusion
      5. γ decay

  40. Water is formed by the addition of 4.0 g of H2 (g) to an excess of O2 (g). If 27 g of H2O is recovered, what is the percent yield for the reaction?

    1. 25%
    2. 50%
    3. 75%
    4. 100%
    5. Cannot be determined

  41. For which of the following is there an increase in entropy?

    1. Na (s) + H2O (l) → NaOH (aq) + H2 (g)
    2. I2 (g) → I2 (s)
    3. H2SO4 (aq) + Ba(OH)2 (aq) → BaSO4 (s) + H2O (l)
    4. H2 (g) + 1/2 O2 (g) → H2O (l)
    5. None of the above

  42. Which of the following compounds contains the greatest percentage of oxygen by weight?

    1. C3H6O5Cl
    2. C3H6O2
    3. C5H10O5
    4. C4H8O3
    5. They are all equal.

  43. In the reaction 2SO2 + O2 → 2SO3 , 0.25 mole of sulfur dioxide is mixed with 0.25 mole of oxygen and allowed to react. What is the maximum number of moles of SO3 that can be produced?

    1. 0.0625 mole
    2. 0.125 mole
    3. 0.25 mole
    4. 0.5 mole
    5. 1.0 mole

  44. What happens to the pH of a buffer system if one halves the concentration of both the acid and the salt?

    1. Nothing
    2. pH goes up because there is less total acid in the solution.
    3. pH goes down because there is less conjugate base to mask the presence of the acid.
    4. It depends upon the original concentration of acid and salt.
    5. It is impossible to predict.

  45. What is the mass of nitrogen in a 50.0 g sample of sodium nitrite (NaNO2)?

    1. 20.2 g
    2. 16.4 g
    3. 10.1 g
    4. 8.23 g
    5. 23.4 g

  46. Which of the following conditions guarantee a spontaneous reaction?

    1. Positive ∆H, positive ∆S
    2. Positive ∆H, negative ∆S
    3. Negative ∆H, negative ∆S
    4. Negative ∆H, positive ∆S
    5. None of the above

  47. Arrange the following elements in order of decreasing nonmetallic character:

    Ge, Sn, Pb, Si

    1. Pb, Sn, Ge, Si
    2. Ge, Sn, Pb, Si
    3. Si, Ge, Sn, Pb
    4. They all have equal nonmetallic character since they are all in the same column of the periodic table.
    5. None of the above

  48. Which of the following is NOT a property of Group IA elements?

    1. Low ionization energies
    2. Low electronegativities
    3. High melting points
    4. Metallic bonding
    5. Electrical conductivity

  49. Electron affinity is defined as

    1. the change in energy when a gaseous atom in its ground state gains an electron
    2. the pull an atom has on the electrons in a chemical bond
    3. the energy required to remove a valence electron from a neutral gaseous atom in its ground state
    4. the energy difference between an electron in its ground and excited states
    5. None of the above