South Pasadena · AP Chemistry



1•Matter and Measurement

PRACTICE TEST

|1. |How many significant digits are |[pic] |

| |present in the temperature read | |

| |from the thermometer illustrated| |

| |to the right? | |

a) 1 b) 2 c) 3 d) 4

2. The dimensions of a rectangular solid are 8.00 cm long, 4.00 cm wide, and 2.00 cm high. If the density of the solid is 10.0 g/cm3, what is its mass?

a) 10/64 grams d) 320 grams

b) 10.0 grams e) 640 grams

c) 64.0 grams

3. A metal sample weighing 30.9232 grams was added to a graduated cylinder containing 23.26 mL of water. The volume of water plus the sample was 24.85 mL. Which setup will result in the density of this metal?

a) 30.9232 x (24.85-23.26)

b) [pic]

c) [pic]

d) 30.9232 x [pic]

e) [pic]

4. The number of significant digits in 0.30500 is

a) 1 d) 4

b) 2 e) 5

c) 3

5. A box measures 3.50 cm x 2.915 cm. The product of these numbers = 10.2025 cm2. What is the proper way to report the area of the box?

a) 10.20 cm2 c) 10 cm2

b) 10.2 cm2 d) 10. cm2

6. The result of 2.350 x (4.0 + 6.311) is,

a) 24 c) 24.21

b) 24.2 d) 24.205

7. A student does a calculation using her calculator and the number 280.27163 is shown on the display. If there are actually three significant figures, how should she show the final answer?

a) 280 d) 2.80 x 10-2

b) 280.3 e) 2.80 x 102

c) 280.27

8. The term that refers to the reproducibility of a laboratory measurement is

a) precision c) accuracy

b) repeatability d) exactness

9. Which measurement below is NOT written with three significant digits?

a) 2.00 cm c) 0.003 L

b) 550. grams d) 12.7 mm

10. The number 6.33 x 102 equals,

a) 6.33 c) 633

b) 0.633 d) 0.0633

11. All the following are characteristic properties of phosphorus. Which one is a chemical property?

a) Both red phosphorus and white phosphorus exist in solid allotropic forms.

b) The red form melts at about 600°C and the white form melts at 44°C.

c) The white form is soluble in liquid carbon disulfide, but is insoluble in water.

d) When exposed to air, white phosphorus will burn spontaneously, but red phosphorus will not.

12. Classify each observation as a physical or a chemical property and tally them.

Observation 1: Bubbles form on a piece of metal when it is dropped into acid.

Observation 2: The color of a crystalline substance is yellow.

Observation 3: A shiny metal melts at 650°C.

Observation 4: The density of a solution is 1.84 g/cm3

a) 2 chemical properties and 2 physical properties

b) 3 chemical properties and 1 physical properties.

c) 1 chemical properties and 3 physical properties

d) 4 chemical properties

e) 4 physical properties

13. Chromatography is a good way to separate the

a) elements in a compound

b) the components in a mixture

c) the atoms in an element

d) the phases of a pure substance

14. When a pure solid substance was heated, a student obtained another solid and a gas, each of which was a pure substance. From this information which of the following statements is ALWAYS a correct conclusion?

a) The original solid is not an element.

b) Both products are elements.

c) The original solid is a compound and the gas is an element.

d) The original solid is an element and the gas is a compound.

e) Both products are compounds.

15. The prefix “milli-” corresponds to what multiplication factor?

a) 10-6 d) 103

b) 10-3 e) 106

c) 101

16. A solution of sugar water may be defined as a

a) heterogeneous mixture

b) homogeneous mixture

c) heterogeneous compound

d) homogeneous compound

e) homogeneous element

17. “Wafting” is the proper technique for

a) neutralizing a spilled acid.

b) putting out burning clothing.

c) washing chemicals from the eye.

d) smelling a chemical substance.

e) observing the color of a chemical.

18. You measure the density of a slab of lead as 11.10 g/mL. The accepted value is 11.34 g/mL. The percent error for your measurement is

a) 2.1 % c) 3.7 %

b) 2.4 % d) 5.1 %

19. Which one of the following elements is correctly matched with its symbol?

a) Ag, gold

b) Ni, nickel

c) Fl, fluorine

d) Mg, manganese

e) H, helium

20. The marks on the following target represent someone who is:

[pic]

a) accurate, but not precise.

b) precise, but not accurate.

c) both accurate and precise.

d) neither accurate nor precise.

|Answers: |

|1.C 2.E 3.B 4.E 5.B 6.B |

|7.E 8.A 9.C 10.C 11.D |

|12.C 13.B 14.A 15.B 16.B |

|17.D 18.A 19.B 20.D |

2 ( Atoms and Elements

PRACTICE TEST

1.Certain properties are characteristic of metals.

Which property means that you can pound the substance into a foil?

a) ductility c) sectility

b) conductivity d) malleability

2. Which of the following is a metalloid?

a) As b) Ag c) S d) Pb e) He

3. Which of the following is a transition metal?

a) Cl b) Ni c) P d) Ca e) C

4. Which of the following is an alkali metal?

a) Mg b) Kr c) K d) Al e) H

5. Which of the following is an lanthanide?

a) Xe b) Eu c) Cd d) P e) W

6. Which element has the highest melting point?

a) Pb b) Au c) Os d) W e) Hg

7. Cathode rays start at the

a) negative electrode c) positive electrode

b) power source d) gas inside the tube

8. In a cathode ray tube, electrons are bent toward

a) a positively charged plate.

b) a negatively charged plate.

9. Listed below are the charges and masses of four particles. Which one will be deflected the least in a mass spectrometer?

a) +2, 2 amu c) +1, 1 amu

b) +4, 4 amu d) +1, 4 amu

10. In a Millikan oil drop type experiment, the charge on four oil drops (in Coulombs) was found to be:

|3.33 |Coulombs |

|8.88 |Coulombs |

|6.66 |Coulombs |

|11.10 |Coulombs |

What is the charge on the electron according to this experiment?

a) 1.11 Coulomb c) 4.44 Coulomb

b) 2.22 Coulomb d) 11.10 Coulomb

11. Pictured below is a schematic of the Rutherford experiment. Which scattered α-particle gives the best evidence for the nuclear atom?

[pic]

a) a b) b c) c d) d e) e

12. Which of the following is an isotope of the element with 20 protons (p=20) and 22 neutrons (n=22)?

a) titanium-22 c) calcium-40

b) zirconium-40 d) titanium-48

13. The imaginary element X has the following natural abundances and isotopic masses. What is the atomic mass of X?

X 24.02 amu 40.0%

X 26.10 amu 60.0%

Show your work:

For questions 14 - 17, use the following key:

(each answer may be used once, more than once,

or not at all)

a) alpha

b) beta

c) gamma

d) alpha and beta, but not gamma

14. A high energy form of light

15. Two protons & two neutrons

16. A high speed electron

17. Used by Ernest Rutherford as a “probe”

For questions 18 - 22, use the following key:

(each answer may be used once, more than once,

or not at all.)

a) John Dalton

b) Ernest Rutherford

c) J.J. Thomson

d) Democritus

18. His model of the atom has been called the “plum pudding” Model.

19. His model of the atom has been called the “billiard ball” model.

20. He studied matter in cathode ray tubes.

21. His philosophical idea included the term “atomos”.

22. He added to the atomic theory the idea that atoms had positive and negative parts.

23. Consider the following notation: Rn

Which statement below is correct?

a) This particle contains 86 protons

b) This particle has a mass number of 86

c) This particle has an atomic number of 220

d) This particle contains 220 neutrons

24. Which elements did Mendeleev leave spaces for in his periodic table?

_____ _____ _____

25. If copper metal is a mixture two isotopes, Cu-63, mass = 62.9298 u and Cu-65, mass = 64.9278 u. The molar mass of copper is 64.546 g/mole. Calculate the % abundances of the two isotopes of copper. Show your work.

Just For Fun:

Element names finish these sentences.

• A ridiculous inmate is a ___.

• I bumped my ___ the car door.

• I am sad when all the flowers ____.

• What the police officer does to the crook. ___

• What the doctor does to the patient. ___

• What the undertaker does if the doctor doesn’t succeed. ___

• If your cattle get away, ___.

• A famous London theatre is the ___.

• Demonstrations help keep the lectures from getting ___.

• Linoleum, tile, and hardwood are three types of ___.

3 • Molecules and Compounds

PRACTICE TEST

1. What is the formula of the ionic compound formed between Mg and Br?

a) MgBr d) Mg2Br2

b) Mg2Br e) Mg2Br3

c) MgBr2

2. What is the formula of the ionic compound formed between Ca and P?

a) Ca2P3 d) Ca2P

b) CaP e) Ca3P2

c) Ca5P10

3. What is the name of the SO32– ion?

a) sulfate d) sulfur trioxide

b) nitrate e) hydrogen sulfate

c) sulfite

4. What is the correct formula and charge for the chromate ion?

a) CrO42– d) Cr2O7–

b) CrO4– e) Cr3+

c) Cr2O72–

5. Which one of the following elements forms ions with two different valences?

a) calcium c) iron

b) arsenic d) fluorine

6. The correct name for CCl4 is

a) carbon(I) chloride

b) carbon chloride

c) carbon tetrachloride

d) monocarbon chloride(IV)

e) carbochlorinate

7. The correct formula for hydrogen telluride is

a) HTe c) H3Te

b) H2Te d) HTe2

8. The correct formula for dinitrogen tetroxide is

a) NO2 d) NO3–

b) N2O4 e) (N2O)4

c) N2O5

9. The correct name for S2Cl2 is

a) sulfur dichloride

b) sulfur(I) chloride

c) sulfur(II) chloride

d) disulfur dichloride

e) sulfur chloride

10. The correct name for NO2 is

a) nitrogen dioxide

b) nitrite

c) nitrogen oxide

d) nitrogen(II) oxide

e) nitrate

11. The molar mass of (NH4)2S is closest to:

a) 50 g/mol c) 68 g/mol

b) 82 g/mol d) 100 g/mol

12. How many atoms are in 12 molecules of glucose, C6H12O6?

a) 24 c) 2160

b) 288 d) 7.22 x 1024

13. Calculate the number of atoms in 4.0 x 10-5 g of aluminum.

a) 8.9 x 1017 c) 6.5 x 1020

b) 4.6 x 1019 d) 3.8 x 1023

14. Which of the following samples contains the smallest number of atoms?

a) 1 g H2 c) 1 g O3

b) 1 g O2 d) 1 g Cl2

15. What is the mass of one molecule of octane, C8H18?

a) 114 g c) 1.10 x 10-22 g

b) 1.89 x 10-22 g d) 4.32 x 10-23 g

16. What is the percent nitrogen (by mass) in ammonium carbonate, (NH4)2CO3?

a) 14.53% c) 29.16%

b) 27.83% d) 33.34%

17. Of the following, the only empirical formula is

a) N2F2 c) H2C2

b) N2F4 d) HNF2

18. A compound consists of the following elements by weight percent:

carbon - 40.0%

oxygen - 53.3%

hydrogen - 6.7%

The ratio of carbon : oxygen : hydrogen in the empirical formula is

a) 1:2:1 c) 1:1:2

b) 1:1:1 d) 2:1:2

19. An organic compound which has the empirical formula CHO has a molar mass of 232. Its molecular formula is:

a) CHO c) C4H4O4

b) C2H2O2 d) C8H8O8

20. When CaSO4·y H2O is heated, all of the water is driven off. If 34.0 g of CaSO4 [molar mass = 136] is formed from 43.0 g of CaSO4·y H2O, what is the value of y?

a) 1 c) 3

b) 2 d) 4

Answers:

|1. |c |6. |c |11. |c |16. |c |

|2. |e |7. |b |12. |b |17. |d |

|3. |c |8. |b |13. |a |18. |c |

|4. |a |9. |d |14. |d |19. |d |

|5. |c |10. |a |15. |b |20. |b |

4 • Chemical Equations and Stoichiometry

PRACTICE TEST

1. Balance the following equation:

___NH3 + ___O2 ( ___NO2 + ___H2O

The balanced equation shows that 1.00 mole of NH3 requires ___ mole(s) of O2.

a) 0.57 c) 1.33

b) 1.25 d) 1.75

2. Write a balanced equation for the combustion of acetaldehyde, CH3CHO.

When properly balanced, the equation indicates that ___ mole(s) of O2 are required for each mole of CH3CHO.

a) 1 c) 2.5

b) 2 d) 3

3. Balance the following equation with the SMALLEST WHOLE NUMBER COEFFICIENTS possible. Select the number that is the sum of the coefficients in the balanced equation:

___KClO3 ( ___KCl + ___O2

a) 5 b) 6 c) 7 d) 8

4. Write a balanced equation for the combustion of propane, C3H8.

When properly balanced, the equation indicates that ___ moles of O2 are required for each mole of C3H8.

a) 3 b) 3.5 c) 5 d) 8

5. What is the total mass of products formed when 16 grams of CH4 is burned with excess oxygen?

a) 80 g c) 36 g

b) 44 g d) 32 g

6. Calculate the mass of hydrogen formed when 25 g of aluminum reacts with excess hydrochloric acid.

2Al + 6HCl ( 2 AlCl3 + 3 H2

a) 0.41 g c) 1.2 g

b) 0.92 g d) 2.8 g

7. How many grams of the mixed oxide, Fe3O4, are formed when 6.00 g of O2 react with Fe according to

3Fe + 2O2 ( Fe3O4

a) 43.4 c) 174

b) 86.8 d) 21.7

8. For the reaction:

2MnO2 + 4KOH + O2 + Cl2 (2KMnO4 + 2KCl + 2H2O

there is 100. g of each reactant available. Which reagent is the limiting reagent?

[Molar Masses: MnO2=86.9; KOH=56.1; O2=32.0; Cl2=70.9]

a) MnO2 c) O2

b) KOH d) Cl2

9. How many grams of nitric acid, HNO3, can be prepared from the reaction of 92.0 g of NO2 with 36.0 g H2O?

3NO2 + H2O ( 2HNO3 + NO

a) 64 c) 84

b) 76 d) 116

10. The reaction of 25.0 g benzene, C6H6, with excess HNO3 resulted in 21.4 g C6H5NO2. What is the percentage yield?

C6H6 + HNO3 ( C6H5NO2 + H2O

a) 100% c) 54.3%

b) 27.4% d) 85.6%

11. How many grams of H2O will be formed when 16.0 g H2 is allowed to react with 16.0 g O2 according to

2H2 + O2 ( 2H2O?

a) 18.0 g c) 9.00 g

b) 144 g d) 32.0 g

12. When 8.00 g of H2 reacts with 32.0 g of O2 in an explosion, 2H2 + O2 ( 2H2O, the final gas mixture will contain:

a) H2, H2O, and O2 c) O2 and H2O only

b) H2 and H2O only d) H2 and O2 only

13. 1.056 g of metal carbonate, containing an unknown metal, M, were heated to give the metal oxide and 0.376 g CO2.

MCO3(s) + heat ( MO(s) + CO2(g)

What is the identity of the metal M?

a) Mg c) Zn

b) Cu d) Ba

14. Styrene, the building block of polystyrene, is a hydrocarbon, a compound containing only C and H. A given sample is burned completely and it produces 1.481 g of CO2 and 0.303 g of H2O. Determine the empirical formula of the compound.

a) CH c) C2H3

b) CH2 d) C2H5

7&8 • Atomic Structure & Periodicity

PRACTICE TEST

A = 2.18 x 10-18 J h = 6.626 x 10-34 J·s

R = 1.097 x 107 m–1 c = 3.00 x 108 m·s–1

mass of an electron = 9.11 x 10-31 kg

1. What wavelength corresponds to a frequency of 8.22 x 109 Hz?

a) 0.307 m d) 0.110 m

b) 0.0365 m e) 27.4 m

c) 0.122 m

2. A radio station transmits at 110 MHz (110 x 106 Hz). What wavelength is this radio wave?

a) 3.65 x 10–5 m c) 3.81 x 10–5 m

b) 3.30 m d) 2.73 m

3. Which one of the following is NOT a proper unit for frequency?

a) Hz c) m·s–1

b) s–1 d)

4. Calculate the wavelength of the fourth line in the Balmer series (the visible series) of the hydrogen spectrum.

a) 0.12334 m d) 4.1029 x 10–7 m

b) 24.373 m e) 36.559 m

c) 2.7353 x 10–7 m

5. What is the relationship between the energy of a photon of light and its frequency?

a) E = ν d) E =

b) E = e) E =

c) E = hν

6. What is the energy needed to raise an electron in the hydrogen atom from the second energy level to the third energy level?

a) 1.52 x 104 J d) 4.48 x 10–19 J

b) 3.63 x 10–19 J e) 3.03 x 10–19 J

c) 2.18 x 10–19 J

7. What is the de Broglie wavelength of an electron moving at 80.0% the speed of light.

a) 3.03 x 10–12 m c) 3.30 x 1011 m

b) 2.42 x 10–12 m d) 1.59 x 10–25 m

8. What resultant is expected from the interference of the two waves shown below?

[pic]

a) [pic] c) [pic]

b) [pic] d) [pic]

9. Which quantum number determines the subshell occupied by an electron (s, p, d, f, etc.)?

a) n c) ml

b) l d) ms

10. What position on the standing wave shown below corresponds to a crest?

[pic]

a) A b) B c) C d) D e) E

11. How many orbitals make up the 4d subshell?

a) 0 b) 1 c) 3 d) 5 e) 7

12. The value of l that is related to the following orbital is:

[pic]

a) 0 b) 1 c) 2 d) 3 e) 4

13. The correct electron configuration for nitrogen is

a) 1s2 2s2 2p6 3s2 3p2

b) 1s2 2s2 2p6 2d4

c) 1s2 2s2 2p3

d) 1s2 2s2 3s2 4s1

e) 1s2 1p5

14. The electron configuration of the indicated atom in the ground state is correctly written for which atom?

a) Ga [Ar] 3d12 4s2

b) Ni [Ar] 3d10

c) Ni [Ar] 3s2 3p8

d) Cu [Ar] 3d10 4s1

15. Which of the following sets of quantum numbers is possible for a 3d electron?

a) n = 3, l = 3, ml = –2, ms = +

b) n = 2, l = 1, ml = +1, ms = –

c) n = 3, l = 1, ml = 0, ms = –

d) n = 3, l = 2, ml = –2, ms = +

e) n = 4, l = 1, ml = +1, ms = +

16. In what section of the periodic table is the 4f subshell being filled?

a) period 4

b) transition elements Y to Cd

c) noble gases

d) group IA

e) lanthanides

17. Which one of the following elements has 3 electrons in a p subshell?

a) Sb b) Na c) Sc d) V e) Nd

18. Which of the following distributions of electrons is correct for three electrons in p-subshell?

|a) |( | |( | |( |

|b) |(( | |( | | |

|c) |( | |( | |( |

|d) |( | |(( | | |

|e) |(( | |( | | |

| | | | | | |

19. Which of the following particles would be most paramagnetic?

a) P

b) Ga

c) Br

d) Cl-

e) Na+

20. Which of the following correctly represents the ionization of an atom?

a) Cl(g) + e– → Cl–(g)

b) Na(g) → Na+(g) + e–

c) Na(s) – e– → Na+(g)

d) Cl2(g) → 2 Cl(g)

21. Which of the following is likely to have the largest atomic radius?

a) H b) Mn c) Cl d) Rb e) Ag

22. Which one of the following isoelectronic species has the smallest radius?

a) Mg2+ d) F –

b) Na+ e) O2–

c) Ne

NOTE: explain your reasoning on the last page.

23. Which of the following has the greatest ionization energy?

a) K b) Ca c) Fe d) Ga e) Br

24. Which of the following has the lowest ionization energy?

a) Li b) Na c) K d) Rb e) Cs

25. The successive ionization energies for one of the period three elements is listed below. Which element is referred to?

|E1 |577.4 kJ/mol |

|E2 |1,816 kJ/mol |

|E3 |2,744 kJ/mol |

|E4 |11,580 kJ/mol |

|E5 |15,030 kJ/mol |

a) Na b) Mg c) Al d) Si e) P

NOTE: explain your reasoning on the last page.

26. Draw the orbital diagram for a neutral Ag atom:

[pic]

Write the electron configuration for silver:

27. Long form:

28. Short form:

29. Explain your answer to question 22.

Which one of the following isoelectronic species has the smallest radius?

a) Mg2+

b) Na+

c) Ne

d) F –

e) O2–

30. Explain your answer to question 25.

The successive ionization energies for one of the period three elements is listed below. Which element is referred to?

|E1 |577.4 kJ/mol |

|E2 |1,816 kJ/mol |

|E3 |2,744 kJ/mol |

|E4 |11,580 kJ/mol |

|E5 |15,030 kJ/mol |

a) Na b) Mg c) Al d) Si e) P

5 • Reactions In Aqueous Solution

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

1. On the basis of the solubility rules, which of the following is insoluble?

a) K2O d) (NH4)2SO4

b) Na2CO3 e) Ba(C2H3O2)2

c) PbS

2. In a double replacement reaction, formation of which of the following does not necessarily lead to a chemical change?

a) HC2H3O2 d) H2S

b) AgCl e) NaCl

c) CO2

3. Reaction of an acid with a carbonate (such as CaCO3) always results in the formation of

a) O2 d) O3

b) C(diamond) e) CO2

c) CH4

4. Which of the following is incorrect?

a) all salts containing NH4+ are soluble.

b) all salts containing NO3– are soluble.

c) all fluorides are soluble.

d) all sulfates (except those of Ca2+, Sr2+, Ba2+, and Pb2+) are soluble.

e) most hydroxides are insoluble, except those of Ca2+, Sr2+, Ba2+, the alkali metals and NH4+.

5. One of the gases shown below is NOT usually formed in a double replacement reaction. Which one?

a) N2 d) NH3

b) CO2 e) H2S

c) SO2

6. Write the balanced molecular equation for the reaction of washing soda, Na2CO3 and vinegar, HC2H3O2.

7. The net ionic equation for the above reaction is:

8. How many moles of H+ are associated with the acid, H2SO3, during neutralization?

a) 0 b) 1 c) 2 d) 3

9. How many moles Al2O3 are needed to neutralize 1 mole of HCl?

a) 1/3 d) 6

b) 2/3 e) 12

c) 2 f) 1/6

10. Write the net reaction that will occur when solid ammonium carbonate is added to a solution of hydrosulfuric acid.

11. When H2SO4 and Ba(OH)2 are reacted in a double replacement reaction, one of the products of the reaction is…

a) H2 d) BaH2

b) H2O e) SO2

c) BaS

12. In the double replacement reaction between the weak acid, HC2H3O2 and strong base, NaOH, which ion(s) are spectator ions?

a) Na+, C2H3O2– d) H+, C2H3O2–

b) Na+, OH– e) Na+ only

c) OH– only

13. Which of the following is a base?

a) KOH d) CH3OH

b) C2H5OH e) CO2

c) Br–

14. Which of the following is a strong acid?

a) H2CO3 d) HClO3

b) HF e) HNO3

c) H3PO4

15. Which of the following is an acid in aqueous solutions?

a) H2CO3 d) H2O

b) Al2O3 e) BaO

c) CH4

16. SO2 turns into which acid in solution?

a) HNO3 d) H2S

b) H2SO3 e) HNO2

c) H2SO4

17. What is the oxidation number of C in CO32–?

a) +6 d) +1

b) +4 e) –1

c) +2

18. What is the oxidation number of Br in KBrO4?

a) +1 b) –1 c) +5 d) +7 e) +8

19. For each change below, label the change of the underlined element as Oxidation, Reduction, or Neither

___ Cu2+ ( Cu(

___ CH4 ( CO2

___ H2O2 ( H2O

___ CO2 ( H2CO3

20. How many milliliters of 0.123 M NaOH solution contain 25.0 g of NaOH (molar mass = 40.00 g/mol)?

a) 5.08 mL d) 625 mL

b) 50.8 mL e) 5080 mL

c) 508 mL

21. If you need 1.00 L of 0.125 M H2SO4, how would you prepare this solution?

a) Add 950. mL of water to 50.0 mL of 3.00 M H2SO4.

b) Add 500. mL of water to 500. mL of 0.500 M H2SO4.

c) Add 750 mL of water to 250 mL of 0.375 M H2SO4.

d) Dilute 36.0 mL of 1.25 M H2SO4 to a volume of 1.00 L.

e) Dilute 20.8 mL of 6.00 M H2SO4 to a volume of 1.00 L.

22. What is the ion concentration in a 0.12 M solution of BaCl2?

a) [Ba2+] = 0.12 M and [Cl−] = 0.12 M.

b) [Ba2+] = 0.12 M and [Cl−] = 0.060 M.

c) [Ba2+] = 0.12 M and [Cl−] = 0.24 M.

d) [Ba2+] = 0.060 M and [Cl−] = 0.060 M.

e) [Ba+] = 0.12 M and [Cl2−] = 0.12 M.

23. What is the molarity of the solution that results when 60.0 g NaOH is added to enough water to make 500. mL solution?

a) 1.33 M d) 8.0 M

b) 12.0 M e) 1.50 M

c) 3.00 M

24. What is the molarity of the solution that results when 45.0 g HCl is dissolved in enough water to make 250. mL solution?

a) 4.94 M d) 1.80 M

b) 4.50 M e) 1.46 M

c) 3.24 M

25. What is the concentration of Cl– ion in 0.60 M AlCl3 solution?

a) 1.8 M d) 0.30 M

b) 0.60 M e) 0.10 M

c) 0.20 M

26. How many grams of Na2CO3 (molar mass = 106.0 g/mol) are required for complete reaction with 25.0 mL of 0.155 M HNO3?

Na2CO3 + 2HNO3 ( 2NaNO3 + CO2 + H2O

a) 0.122 g d) 20.5 g

b) 0.205 g e) 205 g

c) 0.410 g

27. What volume of 0.150 M NaOH is needed to react completely with 3.45 g iodine according to the equation:

3 I2 + 6 NaOH ( 5 NaI + NaIO3 + 3 H2O

a) 181 mL d) 2.04 mL

b) 45.3 mL e) 1.02 mL

c) 4.08 mL

28. What is the concentration of an NaOH solution if it takes 16.25 mL of a 0.100 M HCl solution to titrate 25.00 mL of the NaOH solution?

a) 0.0165 M d) 0.100 M

b) 0.151 M e) 0.413 M

c) 0.0650 M

29. A 4.00 M solution of H3PO4 will contain ___g of H3PO4 in 0.250 L of solution.

a) 196 g d) 24.0 g

b) 98.0 g e) 12.0 g c) 49.0 g

16 ( Chemical Equilibrium

PRACTICE TEST

1. Consider the reaction system,

CoO(s) + H2(g) [pic] Co(s) + H2O(g).

The equilibrium constant expression is

a) [pic] d) [pic]

b) [pic] e) [pic]

c) [pic]

2. Given the equilibrium,

2SO2(g) + O2(g) [pic] 2SO3(g), if this equilibrium is established by beginning with equal number of moles of SO2 and O2 in a 1.0 Liter bulb, then the following must be true at equilibrium:

a) [SO2] = [SO3] d) [SO2] < [O2]

b) 2[SO2] = 2[SO3] e) [SO2] > [O2]

c) [SO2] = [O2]

Questions 3 & 4 refer to the following:

At a given temperature, 0.300 mole NO, 0.200 mol Cl2 and 0.500 mol ClNO were placed in a 25.0 Liter container. The following equilibrium is established: 2ClNO(g) [pic] 2NO(g) + Cl2(g)

3. At equilibrium, 0.600 mol of ClNO was present. The number of moles of Cl2 present at equilibrium is

a) 0.050 d) 0.200

b) 0.100 e) 0.250

c) 0.150

4. The equilibrium constant, Kc, is:

a) 4.45 x 10-4 d) 0.167

b) 6.67 x 10-4 e) 1500

c) 0.111

5. At 985(C, the equilibrium constant for the reaction,

H2(g) + CO2(g) [pic] H2O(g) + CO(g)

is 1.63. What is the equilibrium constant for the reverse reaction?

a) 1.63 d) 0.613

b) 0.815 e) 1.00

c) 2.66

6. What is the relationship between Kp and Kc for the reaction, 2ICl(g) [pic] I2(g) + Cl2(g)?

a) Kp = Kc(RT)-1 d) Kp = Kc

b) Kp = Kc(RT) e) Kp = Kc(2RT)

c) Kp = Kc(RT)2

7. For the reaction 2NO2(g) [pic] N2O4(g), Kp at 25(C is 7.3, when all partial pressures are expressed in atmospheres. What is Kc for this reaction? [R=0.0821 L(atm(mol-1(K-1]

a) 4270 d) 179

b) 0.0119 e) 2.06

c) 0.291

8. 0.200 mol NO is placed in a one liter flask at 2273 K. After equilibrium is attained, 0.0863 mol N2 and 0.0863 mol O2 are present. What is Kc for this reaction?

2NO(g) [pic] N2(g) + O2(g)

a) 9.92 d) 39.7

b) 3.15 e) 0.576

c) 0.0372

9. N2O4(g) [pic] 2 NO2(g)

At 25(C, 0.11 mole of N2O4 reacts to form 0.10 mol of N2O4 and 0.02 mole of NO2. At 90(C, 0.11 mole of N2O4 forms 0.050 mole of N2O4 and 0.12 mole of NO2. From these data we can conclude

a) N2O4 molecules react by a second order rate law.

b) N2O4 molecules react by a first order rate law.

c) the reaction is exothermic.

d) N2O4 molecules react faster at 25(C than at 90(C.

e) the equilibrium constant for the reaction above increases with an increase in temperature.

10. For the equilibrium system

H2O(g) + CO(g) [pic] H2(g) + CO2(g)

(H = -42 kJ/mol

Kc equals 0.62 at 1260 K. If 0.10 mole each of H2O, CO, H2 and CO2 (each at 1260 K) were placed in a 1.0-Liter flask at 1260 K, when the system came to equilibrium…

| |The temperature would |The mass of CO would |

|a) |decrease |increase |

|b) |decrease |decrease |

|c) |remain constant |increase |

|d) |increase |decrease |

|e) |increase |increase |

11. For the reaction system,

N2(g) + 3H2(g) [pic] 2NH3(g) + heat

the conditions that would favor maximum conversion of the reactants to products would be

a) high temperature and high pressure

b) high temperature, pressure unimportant

c) high temperature and low pressure

d) low temperature and high pressure

e) low temperature and low pressure

12. Solid HgO, liquid Hg, and gaseous O2 are placed in a glass bulb and are allowed to reach equilibrium at a given temperature.

2HgO(s) [pic] 2Hg(l) + O2(g) (H = +43.4 kcal

The mass of HgO in the bulb could be increased by

a) adding more Hg.

b) removing some O2.

c) reducing the volume of the bulb.

d) increasing the temperature.

e) removing some Hg.

[pic]

17 ( Acid-Base Equilibria

PRACTICE TEST

1. What is the [H+] when [OH-] = 8.1 x 10-5?

a) 8.1 x 10-5 M d) 3.6 x 10-6 M

b) 1.0 x 10-7 M e) 8.1 x 10-5 M

c) 1.2 x 10-10 M

2. What is the [H+] when [OH-] = 3.3 x 10-9?

a) 3.0 x 10-6 M d) 6.6 x 10-5 M

b) 1.0 x 10-7 M e) 3.3 x 10-9 M

c) 3.3 x 10-5 M

3. What is the [H+] in a 0.0025 M HCl solution?

a) 1.0 x 10-7 M d) 3.6 x 10-5 M

b) 4.0 x 10-12 M e) need more info

c) 2.5 x 10-3 M

4. What is the [OH-] in a 0.0050 M HCl solution?

a) 5.0 x 10-3 M d) 6.6 x 10-5 M

b) 1.0 M e) 2.0 x 10-12 M

c) 1.0 x 10-7 M

5. A solution in which [H+] = 10-8 has a pH of ___ and is _______.

a) 8, acidic d) -8, neutral

b) 6, basic e) 8, basic

c) -6, basic

6. What is the pH of a 0.00030 M HNO3 solution?

a) 8.11 d) 4.48

b) 3.00 e) none of these

c) 3.52

7. What is the pH of a 0.0060 M KOH solution?

a) 5.12 d) 8.88

b) 2.22 e) 7.00

c) 11.78

8. A sample of lemon juice is found to have a pH of 2.55. What is the H+ concentration of the juice?

a) 0.0035 M d) 0.0080 M

b) 0.0028 M e) 355 M

c) 11.6 M

9. A sample of milk is found to have a pH of 6.60. What is the OH- concentration of the milk?

a) 2.5 x 10-21 M d) 4.0 x 10-8 M

b) 1.0 x 10-7 M e) 2.5 x 10-7 M

c) 5.0 x 10-7 M

10. What is the concentration of OCl- in a 0.60 M solution of HOCl? Ka = 3.1 x 10-8.

a) 1.8 x 10-4 M d) 1.4 x 10-4 M

b) 7.1 x 10-11 M e) 1.1 x 10-4 M

c) 0.40 M

11. What is the pH of a 0.020 M solution of hydrosulfuric acid, a diprotic acid?

Ka1 = 1.1 x 10-7 Ka2 = 1.0 x 10-14

a) 7.00 d) 4.33

b) 9.67 e) 3.05

c) 7.84

12. What is the concentration of CO32- in a 0.010 M solution of carbonic acid? The relevant equilbria are,

H2CO3 [pic] H+ + HCO3- Ka1 = 4.3 x 10-7

HCO3- [pic] H+ + CO32- Ka2 = 5.6 x 10-11

a) 6.6 x 10-5 M d) 7.5 x 10-7 M

b) 5.6 x 10-11 M e) 7.9 x 10-7 M

c) 6.7 x 10-11 M

13. What is the S2- concentration in a saturated solution (0.10 M) of H2S, in which the pH has been adjusted to 6.00 by the addition of HCl? For H2S, Ka1 = 1.1 x 10-7 and Ka2 = 1.0 x 10-14.

a) 1.1 x 10-16 M d) 3.2 x 10-8 M

b) 1.1 x 10-10 M e) 3.2 x 10-6 M

c) 1.0 x 10-2 M

14. Which of the following salts will result in a basic solution when it is dissolved in water?

a) KCl d) MgBr2

b) NH4I e) none of these

c) NaCN

15. What is the pH of a 0.50 M solution of NaNO2? For HNO2, Ka = 4.5 x 10-4.

a) 12.18 d) 8.52

b) 5.48 e) 7.00

c) 1.82

16. What is the pH of a 1.0 M solution of NaOCl? For HOCl, Ka = 3.1 x 10-8.

a) 10.75 d) 10.25

b) 3.25 e) 7.00

c) 3.75

21 ( Electron Transfer Reactions

PRACTICE TEST

1. Which of the following is the correct cell notation for the reaction

Hg22+ + Cd(s) ( Cd2+ + 2Hg(l)

a) Cd2+ | Cd | | Hg22+ | Hg

b) Cd2+ | Hg22+ | | Cd | Hg

c) Cd | Cd2+ | | Hg22+ | Hg

d) Cd2+ | Hg | | Hg22+ | Cd

e) Hg | Cd | | Hg22+ | Cd2+

2. Consider an electrochemical cell where the following reaction takes place:

3Sn2+(aq) + 2Al(s) ( 3Sn(s) + 2Al3+(aq)

Which of the following is the correct cell notation for this cell?

a) Al | Al3+ | | Sn2+ | Sn

b) Al3+ | Al | | Sn | Sn2+

c) Sn | Sn2+ | | Al3+ | Al

d) Sn | Al3+ | | Al | Sn2+

e) Al | Sn2+ | | Sn | Al3+

Standard Reduction Potentials at 25(C E( (volts)

F2(g) + 2e- ( 2F-(aq) +2.87

Au3+ + 3e- ( Au(s) +1.50

Cl2(g) + 2e- ( 2Cl-(aq) +1.36

O2(g) + 4H3O+(aq) + 4e- ( 6H2O(l) +1.23

Br2(l) + 2e- ( 2Br-(aq) +1.08

Ag+(aq) + e- ( Ag(s) +0.80

Hg22+(aq) + 2e- ( 2Hg(l) +0.79

I2(s) + 2e- ( 2I-(aq) +0.535

Cu2+(aq) + 2e- ( Cu(s) +0.337

Sn4+(aq) + 2e- ( Sn2+(aq) +0.15

Sn2+(aq) + 2e- ( Sn(s) -0.14

Cd2+(aq) + 2e- ( Cd(s) -0.40

Zn2+(aq) + 2e- ( Zn(s) -0.763

2H2O(l) + 2e- ( H2(g) + 2OH-(aq) -0.828

Al3+(aq) + 3e- ( Al(s) -1.66

K+(aq) + e- ( K(s) -2.93

Li+(aq) + e- ( Li(s) -3.045

3. Given the two half reactions and their potentials, which net reaction is spontaneous?

Ni2+(aq) + 2e- ( Ni(s) E° = -0.25 V

Mg2+(aq) + 2e- ( Mg(s) E° = -2.37 V

a) Ni(s) + Mg2+(aq) ( Mg(s) + Ni2+(aq)

b) Ni2+(aq) + Mg(s) ( Mg2+(aq) + Ni(s)

c) Ni(s) + Mg(s) ( Mg2+(aq) + Ni2+(aq)

d) Mg2+(aq) + Ni2+(aq) ( Mg(s) + Ni(s)

e) Mg2+(aq) + Mg(s) ( Ni(s) + Ni2+(aq)

4. Calculate E° for the following reaction:

Sn4+(aq) + 2K(s) ( Sn2+(aq) + 2K+(aq)

a) +6.00 V d) +2.78 V

b) -3.08 V e) -2.78 V

c) +3.08 V

5. Calculate E° for the following reaction:

2Al3+(aq) + 3Cd(s) ( 2Al(s) + 3Cd2+(aq)

a) -2.06 V d) -4.52 V

b) +4.52 V e) -1.26 V

c) +2.06 V

6. Using data from the reduction potential table and the reaction

2Ag(s) + Pt2+(aq) ( Pt(s) + 2Ag+(aq) E° = 0.38 V

calculate the standard reduction potential of the half-reaction

Pt2+(aq) + 2e- ( Pt(s)

a) -1.18 V d) 1.18 V

b) -0.40 V e) 2.00 V

c) 0.40 V

7. Using data from the reduction potential table, predict which of the following is the best oxidizing agent.

a) F2 d) Ag+

b) Ag e) Al3+

c) Sn4+

8. An electrochemical cell of notation Pd | Pd2+ | | Cu2+ | Cu has an E° = -0.65 V. If we know that the standard reduction potential of Cu2+/Cu is E° = 0.34 V, what is the standard reduction potential for Pd2+/Pd?

a) -0.99 V d) 0.62 V

b) -0.31 V e) +0.99 V

c) +0.31 V

9. The standard cell potential for

3Sn4+(aq) + 2Al(s) ( 3Sn2+(aq) + 2Al3+(aq)

is E° = 1.81 V. What is Ecell when

[Sn4+] = 1.0,

[Sn2+] = 1.0 x 10-2, and

[Al3+] = 1.5 x 10-3 at 298 K.

a) 1.70 V d) 1.86 V

b) 1.76 V e) 1.93 V

c) 1.81 V

10. Predict the product at the anode when electric current is passed through a solution of KI.

a) I2(l) d) K(s)

b) K+(aq) e) O2(g)

c) H2(g)

11. If electric current is passed through aqueous LiBr, the product at the cathode would be __________ and the product at the anode would be __________.

a) H2O(l), Li+(aq) d) Br2(l), H2(g)

b) Br2(l), Li(s) e) H2(g), Br2(l)

c) Li(s), Br2(l)

12. How long would it take to deposit 1.36 g of copper from an aqueous solution of copper(II) sulfate by passing a current of two amperes through the solution?

a) 2070 sec d) 736 sec

b) 1.11 x 10-5 sec e) 1030 sec

c) 2570 sec

13. If a current of 6.0 amps is passed through a solution of Ag+ for 1.5 hours, how many grams of silver are produced?

a) 0.60 g d) 3.0 g

b) 36 g e) 1.0 g

c) 0.34 g

14. How is aluminum currently produced in industry?

a) by reduction of Al3+ in Al2O3 with Na(s)

b) electrochemical reduction of pure Al2O3 to give Al and O2

c) electrolysis of AlF3 to give Al and F2

d) electrolysis of a mixture of Al2O3 and Na3AlF6 to give Al and O2

e) by reduction of Al3+ in Al2O3 with CO(g)

15. How was aluminum originally made?

a) the Hall-Heroult process

b) Al2O3 mixed with cryolite is electrolyzed

c) electrolysis of molten Al2O3

d) mining and purifying directly

e) reducing AlCl3 with sodium

16. Under acidic conditions the bromate ion is reduced to the bromide ion. Write the balanced half-reaction for this process.

a) BrO3- + 6H+ + 6e ( Br- + 3H2O

b) 2BrO3- + 6H+ ( Br2- +6H2O + 3e

c) Bro3- + 6H2O + 10e ( Br2- + 12H+ + 3 O2

d) 2BrO3- + 6H2O ( 2Br- + 12H+ + 6 O2 + 8e

e) 2BrO3- + 6H+ ( Br2- + 3H2O + 3e

17. Balance the following redox equation which occurs in acidic solution.

N2H4(g) + BrO3-(aq) ( Br-(aq) + N2(g)

a) 3N2H4 + BrO3- ( 3N2 + Br- + 3H2O + 6H+

b) N2H4 + BrO3- + 2H+ ( 2Br- + N2 +3H2O

c) 3N2H4 + 2BrO3- + 12H+ (

3N2 + 2Br- + 6H2O + 12H+

d) N2H4 + 2BrO3- + 8H+ ( 2Br- + N2 + 6H2O

e) 3N2H4 + 2BrO3- ( 3N2 + 2Br- + 6H2O

18. Which of the following reactions is NOT a redox reaction?

a) 2HgO(s) ( 2 Hg(l) + O2(g)

b) H2(g) + Br2(g) ( 2HBr(g)

c) 2HCl(aq) + Zn(s) ( H2(g) + ZnCl2(aq)

d) H2CO3(aq) ( H2O(l) + CO2(g)

e) 2KClO3 ( 2KCl(s) + 3 O2(g)

1.C 2.A 3.B 4.C 5.E 6.D 7.A 8.E 9.E 10.A

11.E 12.A 13.B 14.D 15.E 16.A 17.E 18.D

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download