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"; Using The Principles Of Stoichiometry, Determine The Second Product From The Reaction - Chem Homework Help
Using the principles of stoichiometry, determine the second product from the reaction

1. Using the principles of stoichiometry, determine the second product from the reaction: 2 C7H6O3 + C4H6O3 -> 2 C9H8O4 + ___________.

a. H2O

b. 2 H2O

c. 1/2 H2O

d. nothing

e. cannot be determined

2. For the previous reaction, if we begin with 1.10 moles of C7H6O3 and 0.850 moles of C4H6O3, what is the theoretical number of moles of the organic product, C9H8O4?

a. 2.11 moles

b. 2.55 moles

c. 4.00 moles

d. 1.10 moles

e. 0.85 moles

3. If the reaction in 19 yields 0.73 moles of C9H8O4, what is the percent yield (rounded)?

a. 66%

b. 42%

c. 34%

d. 85%

e. 11%

4. How does one calculate a mole?

a. Grams product divided by grams starting material

b. Grams of a material divided by molecular weight of same material

c. 6.023 molecules divided by the temperature

d. Grams of product divided by 6.023 molecules

e. Weight of the sample times the volume

5. What is the value of calculating a mole?

a. It allows someone to equate different materials in terms of numbers of units of each

b. It is a way of determining the amount of product possible from starting reagents

c. It is a method for determining the reaction conditions

d. It tells you how hot the reaction is.

e. It allows someone to equate different materials in terms of numbers of units of each and it is a way of determining the amount of product possible from starting reagents

6. Balancing a chemical equation tells you:

a. how much of each material can be present in a reaction.

b. the limiting reagent.

c. what kind of atoms are present and in what amounts.

d. all of the choices apply.

7. What is the empirical formula of the sugar glucose: C6H12O6?

a. C2H2O2

b. CH2O

c. C6H12O6

d. CH4O

e. H2O

8. Calculate the molecular weight of glucose: C6H12O6.

a. 180.16

b. 180.16g/mol

c. 18.0g/mol

d. 182.36g/mol

e. 184.52g/mol

9. If one wishes to prepare any sugar from the elements, which elements need not be present?

a. Hydrogen

b. Oxygen

c. Carbon

d. Water

e. Bromine

10. A mass spectrometer tells you:

a. the elemental composition of a sample.

b. the empirical formula of a sample.

c. how much energy is in the sample.

d. the elemental composition of a sample and the empirical formula of a sample.

e. the elemental composition of a sample and how much energy is in the sample.

11. The term, empirical formula, is:

a. the number of atoms in a compound.

b. the relative number of atoms in a compound.

c. the number of molecules in a mole.

d. the relative number of compounds in a sample.

e. none of the choices apply.

Module 3 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. Which of the following compounds is a strong electrolyte?

a. H2O

b. CH3OH

c. CH3CH2OH

d. NaF

2. Which of the following compounds is a weak electrolyte?

a. HNO3

b. NaNO3

c. HNO2

d. NaNO2

3. Based on the solubility rules, which one of the following compounds should be insoluble in water?

a. NaCl

b. MgBr2

c. FeCl2

d. AgBr

4. Based on the solubility rules, which one of the following compounds should be insoluble in water?

a. CaCO3

b. (NH4)2CO3

c. Na2CO3

d. K2CO3

5. Based on the solubility rules, which of the following will occur when a solution containing about 0.1g of Pb(NO3)2(aq) is mixed with a solution containing 0.1g of KI(aq) /100 mL?

a. KNO3 will precipitate; Pb2+ and I– are spectator ions.

b. No precipitate will form.

c. Pb(NO3)2 will precipitate; K+ and I– are spectator ions.

d. PbI2 will precipitate; K+ and NO3– are spectator ions.

6. Which of the following is the correct net ionic equation for the reaction that occurs when solutions of Pb(NO3)2 and NH4Cl are mixed?

a. Pb(NO3)2(aq) + 2NH4Cl(aq)  NH4NO3(aq) + PbCl2(s)

b. Pb2+(aq) + 2Cl–(aq)  PbCl2(s)

c. Pb2+(aq) + 2NO3– (aq) + 2NH(aq) + 2Cl–(aq)  2NH(aq) + 2NO3– (aq) + PbCl2(s)

d. NH4+(aq)+ NO3– (aq)  2NH4NO3(s)

7. The common constituent in all acid solutions is:

a. H2.

b. H+.

c. OH–.

d. H2SO4.

8. Identify the major ions present in an aqueous LiOH solution.

a. Li2+, O– , H–

b. Li+, OH–

c. LiO–, H+

d. Li+, O2–, H+

9. What is the correct formula of the salt formed in the neutralization reaction of hydrochloric acid with calcium hydroxide?

a. CaO

b. CaCl2

c. CaH2

d. CaCl

10. The oxidation number of N in NaNO3 is:

a. +6.

b. +5.

c. +3.

d. –3.

11. Select the compound in which sulfur has its highest possible oxidation number.

a. H2S

b. SO2

c. SCl2

d. Na2SO4

12. In the following redox reaction, which element is oxidized and which is reduced? 4NH3 + 3Ca(ClO)2  2N2 + 6H2O + 3CaCl2

a. H is oxidized and N is reduced.

b. N is oxidized and Cl is reduced.

c. N is oxidized and O is reduced.

d. Cl is oxidized and O is reduced.

13. Identify the reducing agent in the following chemical reaction: Cd + NiO2 + 2H2O  Cd(OH)2 + Ni(OH)2.

a. Cd

b. NiO2

c. H2O

d. Cd(OH)2

14. What element is oxidized in the following chemical reaction? NiO2 + Cd + 2H2O  Ni(OH)2 + Cd(OH)2.

a. Ni

b. Cd

c. O

d. H

15. Which of the following represents a metal displacement reaction?

a. 2NaN3(s)  2Na(s) + 3N2(g)

b. Fe2O3(s) + 2Al(s)  2Fe(s) + Al2O3(s)

c. 3NO2(g) + H2O(l)  2HNO3(aq) + NO(g)

d. 2P(s) + 3Cl2(g)  2PCl3(g)

16. Which of the following represents an acid-base neutralization reaction?

a. 2Al(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2(g)

b. SO2(g) + H2O(l)  H2SO3(g)

c. LiOH(aq) + HNO3(aq)  LiNO3(aq) + H2O(l)

d. 2KBr(aq) + Cl2(g)  2KCl(aq) + Br2(l)

17. What mass of K2CO3 is needed to prepare 200 mL of a solution having a potassium ion concentration of 0.150 M?

a. 4.15g

b. 10.4g

c. 13.8g

d. 2.07g

18. A 20.00 mL sample of 0.1015 M nitric acid is introduced into a flask, and water is added until the volume of the solution reaches 250.00 mL. What is the concentration of nitric acid in the final solution?

a. 1.27 M

b. 8.12 x 10–3 M

c. 0.406 M

d. 3.25 x 10–2 M

19. 34.62 mL of 0.1510 M NaOH was needed to neutralize 50.0 mL of an H2SO4 solution. What is the concentration of the original sulfuric acid solution?

a. 0.0229 M

b. 0.218 M

c. 0.0523 M

d. 0.209 M

20. If 73.5 mL of 0.200 M KI(aq) was required to precipitate all of the lead (II) ion from an aqueous solution of lead (II) nitrate, how many moles of Pb2+ were originally in the solution?

a. 7.35 x 10–3 moles of Pb2+

b. 73.5 x 10–3 moles of Pb2+

c. 73.5 x 10–2 moles of Pb2+

d. 7.35 x 10–2 moles of Pb2+

21. The pressure of a gas sample was measured to be 654 mmHg. What is the pressure in kPa? (1 atm = 1.01325 x 105 Pa)

a. 87.2 kPa

b. 118 kPa

c. 6.63 x 104 kPa

d. 8.72 x 104 kPa

22. Which of these properties is/are characteristic(s) of gases?

a. High compressibility

b. Relatively large distances between molecules

c. Formation of homogeneous mixtures regardless of the nature of gases

d. All of the choices apply

23. A sample of a gas occupies 1.40 x 103 mL at 25°C and 760 mmHg. What volume will it occupy at the same temperature and 380 mmHg?

a. 2,800 mL

b. 2,100 mL

c. 1,400 mL

d. 1,050 mL

24. The gas pressure in an aerosol can is 1.8 atm at 25°C. If the gas is an ideal gas, what pressure would develop in the can if it were heated to 475°C?

a. 0.095 atm

b. 0.717 atm

c. 3.26 atm

d. 4.52 atm

25. The temperature of an ideal gas in a 5.00 L container originally at 1 atm pressure and 25°C is lowered to 220 K. Calculate the new pressure of the gas.

a. 1.0 atm

b. 1.35 atm

c. 8.8 atm

d. 0.738 atm

26. At what temperature will a fixed mass of gas with a volume of 125 L at 15°C and 750 mmHg occupy a volume of 101 L at a pressure of 645 mm Hg?

a. –73°C

b. 10.4°C

c. 2°C

d. 34°C

27. Calculate the number of moles of gas contained in a 10.0 L tank at 22°C and 105 atm. (R = 0.08206 Latm/Kmol)

a. 1.71 x 10–3 mol

b. 0.0231 mol

c. 1.03 mol

d. 43.4 mol

28. Calculate the volume occupied by 35.2g of methane gas (CH4) at 25°C and 1.0 atm. (R = 0.08206 Latm/Kmol)

a. 0.0186 L

b. 4.5 L

c. 11.2 L

d. 53.7 L

29. Calculate the density, in g/L, of CO2 gas at 27°C and 0.50 atm pressure.

a. 0.89 g/L

b. 1.12 g/L

c. 9.93 g/L

d. 46.0 g/L

30. Determine the molar mass of chloroform gas if a sample weighing 0.389g is collected in a flask with a volume of 102 cm3 at 97°C. The pressure of the chloroform is 728 mmHg.

a. 187 g/mol

b. 121 g/mol

c. 112 g/mol

d. 31.6 g/mol

31. A mixture of three gases has a total pressure of 1,380 mmHg at 298 K. The mixture is analyzed and is found to contain 1.27 mol CO2, 3.04 mol CO and 1.50 mol Ar. What is the partial pressure of Ar?

a. 0.258 atm

b. 301 mmHg

c. 356 mmHg

d. 5,345 mmHg

32. What volume of CO2 gas at 645 torr and 800 K could be produced by the reaction of 45g of CaCO3 according to the equation? CaCO3(s)  CaO(s) + CO2(g)

a. 0.449 L

b. 22.4 L

c. 25.0 L

d. 34.8 L

33. Which statement is false?

a. The average kinetic energies of molecules from samples of different “ideal” gases is the same at the same temperature.

b. The molecules of an ideal gas are relatively far apart.

c. All molecules of an ideal gas have the same kinetic energy at constant temperature.

d. Molecules of a gas undergo many collisions with each other and the container walls.

34. Which of these gas molecules have the highest average kinetic energy at 25°C?

a. H2

b. O2

c. N2

d. All the gases have the same average kinetic energy

35. Deviations from the ideal gas law are greater at:

a. low temperatures and low pressures.

b. low temperatures and high pressures.

c. high temperatures and high pressures.

d. high temperatures and low pressures.

4

+

4

+

Module 4 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. Radiant energy is:

a. the energy stored within the structural units of chemical substances.

b. the energy associated with the random motion of atoms and molecules.

c. solar energy; i.e., energy that comes from the sun.

d. energy available by virtue of an object’s position.

2. Thermal energy is:

a. the energy stored within the structural units of chemical substances.

b. the energy associated with the random motion of atoms and molecules.

c. solar energy; i.e., energy that comes from the sun.

d. energy available by virtue of an object’s position.

3. Chemical energy is:

a. the energy stored within the structural units of chemical substances.

b. the energy associated with the random motion of atoms and molecules.

c. solar energy; i.e., energy that comes from the sun.

d. energy available by virtue of an object’s position.

4. Potential energy is:

a. the energy stored within the structural units of chemical substances.

b. the energy associated with the random motion of atoms and molecules.

c. solar energy; i.e., energy that comes from the sun.

d. energy available by virtue of an object’s position.

5. Heat is a measure of:

a. temperature.

b. the change in temperature.

c. thermal energy.

d. thermal energy transferred between two bodies at different temperature.

6. An endothermic reaction causes the surroundings to:

a. warm up.

b. become acidic.

c. condense.

d. decrease in temperature.

7. An exothermic reaction causes the surroundings to:

a. warm up.

b. become acidic.

c. condense.

d. decrease in temperature.

8. When 0.7521g of benzoic acid was burned in a calorimeter containing 1,000g of water, a temperature rise of 3.60°C was observed. What is the heat capacity of the bomb calorimeter, excluding the water? The heat of combustion of benzoic acid is –26.42 kJ/g.

a. 15.87 kJ/°C

b. 4.18 kJ/°C

c. 5.52 kJ/°C

d. 1.34 kJ/°C

9. Naphthalene combustion can be used to calibrate the heat capacity of a bomb calorimeter. The heat of combustion of naphthalene is –40.1 kJ/g. When 0.8210g of naphthalene was burned in a calorimeter containing 1,000g of water, a temperature rise of 4.21°C was observed. What is the heat capacity of the bomb calorimeter excluding the water?

a. 32.9 kJ/°C

b. 7.8 kJ/°C

c. 3.64 kJ/°C

d. 1.76 kJ/°C

10. When 0.56g of Na(s) reacts with excess F2(g) to form NaF(s), 13.8 kJ of heat is evolved at standard-state conditions. What is the standard enthalpy of formation (H°f) of NaF(s)?

a. 24.8 kJ/mol

b. 570 kJ/mol

c. –24.8 kJ/mol

d. –7.8 kJ/mol

11. The combustion of butane produces heat according to the equation: 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(l), H°rxn = –5,314 kJ/mol. What is the heat of combustion per gram of butane?

a. –32.5 kJ/g

b. –45.7 kJ/g

c. –91.5 kJ/g

d. –2,656 kJ/g

12. Ethanol (C2H5OH) burns according to the equation: C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(l), H°rxn = –1367 kJ/mol. How much heat is released when 35g of ethanol is burned?

a. 1,797 kJ

b. 1,367 kJ

c. 9.61  10–4 kJ

d. 1,040 kJ

13. According to the first law of thermodynamics,:

a. energy is neither lost nor gained in any energy transformations.

b. perpetual motion is possible.

c. energy is conserved in quality, but not in quantity.

d. energy is being created as time passes. We have more energy in the universe now than when time began.

14. A gas is compressed in a cylinder from a volume of 20.0 L to 2.0 L by a constant pressure of 10.0 atm. Calculate the amount of work done on the system.

a. 1.01  104 J

b. –180 J

c. 1.81 104 J

d. –1.81  104 J

15. Which of the following processes always results in an increase in the energy of a system?

a. The system loses heat and does work on the surroundings.

b. The system gains heat and does work on the surroundings.

c. The system loses heat and has work done on it by the surroundings.

d. The system gains heat and has work done on it by the surroundings.

16. A calorimeter measures:

a. the change in temperature of a process/reaction.

b. the change in volume of a process/reaction.

c. the change in pressure of a process/reaction.

d. the change in randomness of a process/reaction.

e. none of the choices apply.

17. Most chemical processes are:

a. spontaneous.

b. random.

c. endothermic.

d. exothermic.

e. man-made.

18. In the natural world, most chemical processes occur in contact with the Earth’s atmosphere at:

a. a constant pressure.

b. a constant temperature.

c. a constant volume.

d. a constant amount.

e. none of the choices apply.

19. Which of the following processes do not involve energy changes?

a. Plants converting carbon dioxide to oxygen, water, and sugars

b. Animals digesting food

c. Automobile burning fuel

d. Animals breathing in oxygen

e. None of the choices apply

20. The idea gas law relates:

a. volume and pressure in a predictable way.

b. temperature and volume in a predictable way.

c. pressure and temperature in a predictable way.

d. quantities and pressure in a predictable way.

e. all of the choices apply.

Module 5 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. What is the wavelength of radiation that has a frequency of 2.10  1014 s –1?

a. 6.30  1022 m

b. 7.00  102 nm

c. 7.00  105 m

d. 1.43  10–6 m

2. What is the energy in joules of one photon of microwave radiation with a wavelength 0.122 m? (c = 2.9979  108 m/s; h = 6.626  10–34 Js)

a. 2.70  10–43 J

b. 5.43  10–33 J

c. 1.63  10–24 J

d. 4.07  10–10 J

3. What is the binding energy (in J/mol or kJ/mol) of an electron in a metal whose threshold frequency for photoelectrons is 2.50  1014 /s?

a. 99.7 kJ/mol

b. 1.66  10–19 J/mol

c. 2.75  10–43 J/mol

d. 7.22  1017 kJ/mol

4. Complete this sentence: Atoms emit visible and ultraviolet light as:

a. electrons jump from lower energy levels to higher levels.

b. the atoms condense from a gas to a liquid.

c. electrons jump from higher energy levels to lower levels.

d. they are heated and the solid melts to form a liquid.

5. Calculate the frequency of the light emitted by a hydrogen atom during a transition of its electron from the n = 6 to the n = 3 principal energy level. Recall that for hydrogen, En = –2.18  10–18 J(1/n2).

a. 1.64  1015 /s

b. 9.13  1013 /s

c. 3.65  1014 /s

d. 2.74  1014/s

6. If a hydrogen atom and a helium atom are traveling at the same speed,

a. the wavelength of the hydrogen atom will be about 4 times longer than the wavelength of the helium atom.

b. the wavelength of the hydrogen atom will be about 2 times longer than the wavelength of the helium.

c. the wavelength of the hydrogen atom will be roughly equal to the wavelength of the helium atom.

d. the wavelength of the helium atom will be about 2 times longer than the wavelength of the hydrogen atom.

7. What is the maximum number of electrons in an atom that can have the following set of quantum numbers? n = 4 l = 3 ml = –2 ms = +1/2

a. 0

b. 1

c. 2

d. 6

8. Electrons in an orbital with l = 3 are in a(n):

a. d orbital.

b. f orbital.

c. g orbital.

d. p orbital.

9. The maximum number of electrons that can occupy an energy level described by the principal quantum number, n, is:

a. n.

b. n + 1.

c. 2n.

d. 2n2.

10. How many orbitals are allowed in a subshell if the angular momentum quantum number for electrons in that subshell is 3?

a. 1

b. 3

c. 5

d. 7

11. Which element has the following ground-state electron configuration?: 1s22s22p63s2

a. Na

b. Mg

c. Al

d. Si

12. Which element has the following ground-state electron configuration?: [Kr]5s24d105p3

a. Sn

b. Sb

c. Pb

d. Bi

13. How many electrons in a ground-state tellurium atom are in orbitals labeled by l = 1?

a. 4

b. 10

c. 12

d. 22

14. Which of the following ground-state atoms is diamagnetic?

a. Ca

b. As

c. Cu

d. Fe

15. When the electron in a hydrogen atom falls from the n = 3 excited energy level to the ground state energy level, a photon with wavelength, , is emitted. An electron having this same wavelength would have a velocity of:

a. 7.10  103 m/s.

b. 2.93  106 m/s.

c. 2.93  103 m/s.

d. 7.10 m/s.

16. In order to fill an orbital, electrons must be:

a. paired.

b. parallel.

c. ejected.

d. promoted.

e. round.

17. Unpaired electrons are found in:

a. Hund’s rule.

b. paramagnetic compounds.

c. magnetic compounds.

d. Hund’s rule and paramagnetic compounds.

e. Hund’s rule and magnetic compounds.

18. The Aufbau Principle states that:

a. in the ground state, electrons are all excited.

b. each successive electron naturally gets placed in the next most stable orbital.

c. electrons which are placed, are stable.

d. an electron’s location and speed cannot be known.

e. electrons are bound to atoms.

19. The structure of the periodic table was created by:

a. Mendeleev.

b. Pauli.

c. Aufbau.

d. Einstein.

e. Bohr.

20. If we somehow knock a proton off of a Fluorine atom, we are left with:

a. the nucleus of a carbon atom.

b. the nucleus of an oxygen atom.

c. the nucleus of a fluorine atom.

d. the nucleus of a neon atom.

e. the nucleus of a boron atom.

21. The element with 32 protons can only be:

a. Germanium.

b. Silicon.

c. Bismuth.

d. Gadolinium.

e. Arsenic.

22. A valence electron is:

a. any electron in the shell of an atom.

b. all those of the highest principle quantum number, plus any in partially-filled orbitals.

c. all electrons in the atom.

d. the core electrons.

e. the majority of electrons.

23. A core electron:

a. is extremely difficult to remove from the atom.

b. is close to the center of the atom.

c. is tightly held.

d. can be removed with very high energy beams.

e. all of the choices apply.

24. When an electron is gained or lost, a(n) ________ is formed.

a. cation

b. ion

c. anion

d. all of the choices apply

25. Electrons repel each other,:

a. unless they are paired.

b. unless they are in different orbitals.

c. unless Hund’s rule fails.

d. unless they reach equilibrium.

e. unless they are lost.

26. An atom with all electrons paired, not attracted by magnets, is called:

a. paramagnetic.

b. diamagnetic.

c. magnetic.

d. geomagnetic.

e. unmagnetic.

27. An atom with all electrons unpaired, and attracted by magnets, is called:

a. paramagnetic.

b. diamagnetic.

c. magnetic.

d. geomagnetic.

e. unmagnetic.

28. Most of the volume in the atom is occupied by:

a. the valence electrons.

b. the core electrons.

c. the nucleus.

d. the quanta.

e. the neutrons.

29. Electron affinity is defined as:

a. the tendency of an atom to repel an electron.

b. the energy change when an electron is added to an atom.

c. the ionization energy.

d. the work required to add an electron to a system.

e. the first electron to leave an atom.

30. The size of an atomic radius is determined by:

a. how tightly that atom holds its electrons.

b. how the atoms fit together.

c. how the atoms repel each other.

d. how much the electrons try to leave.

e. how the atom boosts its energy through repulsion.

Module 6 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. As opposed to early periodic tables based on the law of octaves, modern periodic tables arrange the elements in order of increasing:

a. nuclear binding energy.

b. number of neutrons.

c. atomic mass.

d. atomic number.

2. The elements in Group 7A are known by what name?

a. Transition metals

b. Halogens

c. Alkali metals

d. Alkaline earth metals

3. The elements in Group 2A are known by what name?

a. Transition metals

b. Halogens

c. Alkali metals

d. Alkaline earth metals

4. The alkali metal elements are found in __________ of the periodic table.

a. Group 1A

b. Group 2A

c. Group 3A

d. Period 7

5. Which one of the following elements is a transition element?

a. Sr

b. Pb

c. As

d. Fe

6. Consider the element with the electron configuration, [Kr]5s24d7. This element is a(n):

a. representative element.

b. transition metal.

c. nonmetal.

d. actinide element.

7. How many valence electrons does an oxygen atom have?

a. 2

b. 4

c. 6

d. 7

8. What is the charge on the monatomic ion that calcium forms in its compounds?

a. +2

b. +1

c. –1

d. –2

9. Which two electron configurations represent elements that would have similar chemical properties?: (1) 1s22s22p4 (2) 1s22s22p5 (3) [Ar]4s23d5 (4) [Ar]4s23d104p5

a. (1) and (2)

b. (1) and (3)

c. (2) and (3)

d. (2) and (4)

10. The electron configuration of a copper(I) ion is:

a. [Ar]4s23d8.

b. [Ar]4s13d9.

c. [Ar]3d10.

d. [Ar]4s23d64p2.

11. Arrange the following ions in order of increasing ionic radius: K+, P3– , S2– , Cl– .

a. K+ < Cl– < S2– < P3–

b. K+ < P3– < S2– < Cl–

c. P3– < S2– < Cl– < K+

d. Cl– < S2– < P3– < K+

12. Arrange the following ions in order of decreasing ionic radius: Al3+, Mg2+, Na+, O2–.

a. Al3+ > Mg2+ > O2– > Na+

b. Al3+ > Mg2+ > Na+ > O2–

c. Na+ > Mg2+ > Al3+ > O2–

d. O2– > Na+ > Mg2+ > Al3+

13. For which of the following reactions is the enthalpy change equal to the second ionization energy of nitrogen?

a. N2+(g)  N3+(g) + e–

b. N2+(g) + e–  N+(g)

c. N(g)  N2+(g) + 2e–

d. N+(g)  N2+(g) + e–

14. Which of the elements listed below has the highest first ionization energy?

a. He

b. Ne

c. Ar

d. Kr

15. Which of the elements listed below has the smallest first ionization energy?

a. C

b. Ge

c. P

d. O

16. Which of the following elements has the greatest electron affinity (largest positive value)?

a. Mg

b. Al

c. Si

d. S

17. Which of the following elements has the greatest electron affinity (largest positive value)?

a. K

b. Br

c. As

d. Ar

18. Which pair of elements from different groups resembles each other the most in their chemical properties?

a. Be and B

b. Al and Si

c. Li and Be

d. Al and Be

19. In a surprisingly large number of their properties, beryllium resembles aluminum and boron resembles silicon. Such a relationship is called:

a. amphoterism.

b. an allotropic relationship.

c. a diagonal relationship.

d. the periodic law.

20. Which of the following elements has the greatest metallic character?

a. Br

b. F

c. Ge

d. Sc

221. Which is not a transition metal?

a. Ge

b. Ni

c. Po

d. Te

22. Which element is least likely to form a bond with another atom?

a. Kr

b. Xe

c. Fr

d. Ac

23. In the old theory of triads, which element could not be paired with Au?

a. Ag

b. Pt

c. Cu

d. Tc

24. How many different versions of the periodic table are known?

a. 700

b. 2

c. 10

d. 1

25. What is one discovery which was predicted by the periodic table?

a. Which elements give off light

b. High temperature superconductors

c. Better batteries

d. The incandescent light bulb

26. Most of the elements can be found in which form?

a. Solid

b. Liquid

c. Gas

d. Plasma

27. The energy necessary to remove an electron from an element or compound is called:

a. ionization energy.

b. effective nuclear charge.

c. electronegativity.

d. the electron affinity.

28. The change in energy when an electron is added to a neutral atom is called:

a. electronegativity.

b. electron affinity.

c. effective nuclear charge.

d. The Bohr radius.

29. The upper right part of the periodic table (not including the noble gasses) has the elements which are the most:

a. electropositive.

b. electronegative.

c. neutral.

d. unreactive.

30. The periodic table is a tool that is most useful for:

a. determining the relative properties of atoms around one which is known.

b. organizing elements according to their absolute properties.

c. determining which elements will bond with other elements.

d. scientists.

Module 7 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. Which one of the following is most likely to be an ionic compound?

a. CaCl2

b. CO2

c. CS2

d. SO2

2. Which one of the following is most likely to be an ionic compound?

a. ClF3

b. FeCl3

c. NH3

d. PF3

3. Which one of the following is most likely to be a covalent compound?

a. Rb2O

b. BaO

c. SrO

d. SeO2

4. Which one of the following is most likely to be a covalent compound?

a. KF

b. CaCl2

c. SF4

d. Al2O3

5. The Lewis dot symbol for the lead atom is:

a.

b.

c.

d.

6. The Lewis dot symbol for the S 2– ion is:

a.

b. 2–

c. S2–

d.  2–

7. Which one of the following ionic solids would have the largest lattice energy?

a. NaCl

b. NaF

c. CaBr2

d. CaCl2

8. Which of the following ionic solids would have the largest lattice energy?

a. SrO

b. NaF

c. CaBr2

d. CsI

9. Which of the following solids would have the highest melting point?

a. NaF

b. NaCl

c. NaBr

d. NaI

10. Which of the following solids would have the lowest melting point?

a. KI

b. KBr

c. KCl

d. KF

11. Which of the elements listed below has the greatest electronegativity?

a. Na

b. As

c. Ga

d. Cs

12. Which of the elements listed below is the least electronegative?

a. Sr

b. V

c. Ni

d. P

13. Which of the bonds below would have the greatest polarity (i.e., highest percent ionic character)?

a. Si  P

b. Si  S

c. Si  Se

d. Si  Cl

14. The Lewis structure for CS2 is:

a.

b.

c.

d.

15. Assuming the octet rule is obeyed, how many covalent bonds will a nitrogen atom form to give a formal charge of zero?

a. 0

b. 1

c. 2

d. 3

16. Which of the following is a useful guideline for the application of formal charges in neutral molecules?

a. A Lewis structure in which there are no formal charges is preferred.

b. Lewis structures with large formal charges (e.g., +2,+3 and/or -2,-3) are preferred.

c. The preferred Lewis structure is one in which positive formal charges are on the most electronegative atoms.

17. For which of these species does the best Lewis structure have two or more equivalent resonance structures?

a. HCO2–

b. SCN–

c. CNO–

d. N3–

18. Which one of the following compounds does not follow the octet rule?

a. NF3

b. CF4

c. PF5

d. AsH3

19. Which one of the following compounds does not follow the octet rule?

a. NF3

b. CO2

c. CF4

d. NO

20. Use the bond enthalpy data given to estimate the heat released when 6.50g of nitrogen gas reacts with excess hydrogen gas to form ammonia at 25°C. BE(NN) = 941.4 kJ/mol, BE(H–H) = 436.4 kJ/mol, BE(N–H) = 393 kJ/mol

a. 228 kJ

b. 340 kJ

c. 107 kJ

d. 24.9 kJ

21. Give the number of lone pairs around the central atom and the molecular geometry of CBr4.

a. 0 lone pairs, square planar

b. 0 lone pairs, tetrahedral

c. 1 lone pair, square pyramidal

d. 1 lone pair, trigonal bipyramidal

22. Give the number of lone pairs around the central atom and the molecular geometry of SCl2.

a. 0 lone pairs, linear

b. 1 lone pair, bent

c. 2 lone pairs, bent

d. 3 lone pairs, bent

23. According to the VSEPR theory, the geometry of the SO3 molecule is:

a. pyramidal.

b. tetrahedral.

c. trigonal planar.

d. distorted tetrahedron (seesaw).

24. According to the VSEPR theory, the molecular geometry of ammonia is:

a. linear.

b. trigonal planar.

c. bent.

d. trigonal pyramidal.

25. The bond angle in SCl2 is expected to be:

a. a little less than 109.5°.

b. 109.5°.

c. a little more than 109.5°.

d. 120°.

26. The bond angle in Cl2O is expected to be approximately:

a. 90°.

b. 109.5°.

c. 120°.

d. 145°.

27. Complete the following sentence: The PCl5 molecule has:

a. nonpolar bonds, and is a nonpolar molecule.

b. nonpolar bonds, but is a polar molecule.

c. polar bonds, and is a polar molecule.

d. polar bonds, but is a nonpolar molecule.

28. Which one of the following molecules has a non-zero dipole moment?

a. BeCl2

b. Br2

c. BF3

d. IBr

29. Which one of the following molecules has a zero dipole moment?

a. CO

b. CH2Cl2

c. SO3

d. SO2

30. Indicate the type of hybrid orbitals used by the central atom in CCl4.

a. sp

b. sp2

c. sp3

d. sp3d

31. Indicate the type of hybrid orbitals used by the central atom in SF6.

a. sp

b. sp2

c. sp3

d. sp3d2

32. What is the hybridization on the central atom in NO3– ?

a. sp

b. sp2

c. sp3

d. sp3d

33. Which of the following is not true of molecular orbitals?

a. The number of molecular orbitals formed is always equal to the number of atomic orbitals combined.

b. A molecular orbital can accommodate up to two electrons.

c. When electrons are added to orbitals of the same energy, the most stable arrangement is predicted by Hund’s rule.

d. For any substance, the number of electrons in molecular orbitals is equal to the sum of all the valence electrons on the bonding atoms.

34. Which of the following correctly lists species in order of increasing bond length?

a. O2 < O2+ < O2–

b. O2– < O2 < O2+

c. O2+ < O2 < O2–

d. O2– < O2+ < O2

35. The N – N – H bond angles in hydrazine N2H4 are 112°. What is the hybridization of the nitrogen orbitals predicted by valence bond theory?

a. sp3

b. sp2

c. sp4

d. sp6

Pb

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Pb

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:

S

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S

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::

S

·

·

CSS

······

······

=-

SCS

······

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::

SCS

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SCS

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:

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Module 8 Check Your Understanding 1

Directions: Please highlight or bold the correct answer.

1. Alkanes have the general formula,:

a. CnH2n–4.

b. CnH2n–2.

c. CnH2n.

d. CnH2n+2.

2. Alkenes have the general formula,:

a. CnH2n–4.

b. CnH2n–2.

c. CnH2n.

d. CnH2n+2.

3. Alkynes have the general formula,:

a. CnH2n–4.

b. CnH2n–2.

c. CnH2n.

d. CnH2n+2.

4. How many structural isomers are there of C4H10?

a. 4

b. 6

c. 2

d. 8

5. Which of these species is an aromatic compound?

a. C2H2

b. C6H12

c. C6H4Br2

d. C5H10

6. Which of these is the systematic name for the compound represented below?

alt

a. 2-ethylbutane

b. 3-methylpentene

c. 3-methyl-1-pentene

d. 3-methyl-1-hexene

7. The group of atoms that is responsible for the characteristic properties of a family of organic compounds is called a(n):

a. reaction center.

b. functional group.

c. binding site.

d. enzyme.

8. Which one of the following functional groups is found in alcohols?

untitled.bmp

a. A

b. B

c. C

d. D

9. Which one of the following functional groups is found in carboxylic acids?

CHM101-8.bmp

a. A

b. B

c. C

d. D

10. The reaction of an alcohol and a carboxylic acid yields a(n):

a. hydrocarbon.

b. ester.

c. ether.

d. aldehyde.

11. The reaction of Cl2 with CH4 to produce methyl chloride is an example of a(n):

a. free radical reaction.

b. addition reaction.

c. reduction reaction.

d. ester hydrolysis.

12. Which of these statements describes a condensation reaction?

a. Addition of H2O to a double bond

b. Linking an acid and an alcohol to make an ester and water

c. Addition of H2 to an alkene

d. Oxidation of ethanol to acetaldehyde

13. Bromination of benzene (C6H6), an aromatic compound,:

a. occurs by substitution rather than addition.

b. occurs by addition rather than substitution.

c. occurs more rapidly than bromination of a nonaromatic compound.

d. results in formation of 1,2,3,4,5,6-hexabromocyclohexane.

14. Which functional group, when present in a compound that is allowed to stand in air, poses a danger of slowly yielding explosive peroxides?

a. Ether

b. Alcohol

c. Carboxylic acid

d. Ketone

15. The molecule shown below is chiral; i.e., not superimposable on its mirror image.

alt

a. True

b. False

16. The term, “hybridization,” refers to:

a. a change in the orbital configuration of electrons about an atom which has undergone bonding.

b. a way to describe elemental reactions.

c. changes in shapes of atoms.

d. a way to describe elements.

e. none of the choices apply.

17. When carbon atoms undergo bonding interactions,:

a. the hybridization of the atoms always changes.

b. the electronic state of the atoms change.

c. the reaction is always reversible.

d. none of the choices apply.

18. Enantiomers are:

a. sets of molecules with the same connectivity, atom to atom.

b. molecules with different orientation of substituents, about one or more atoms.

c. molecules which behave differently in biological systems.

d. non-super imposable mirror image isomers.

e. all of the choices apply.

19. Carbon chemistry is vastly diverse as compared to non-carbon chemistry because:

a. a large diversity of compounds are stable under normal conditions.

b. inorganic compounds cannot make large structures.

c. organic compounds do not make salts.

d. the chemistry of life is based on carbon.

e. a large diversity of compounds are stable under normal conditions and the chemistry of life is based on carbon.

20. It is not safe to carry out chemical experiments without:

a. proper ventilation.

b. chemical-resistant gloves.

c. someone else present in case of an emergency.

d. firm knowledge of what you’re doing, and the consequences.

e. all of the choices apply.

21. A(n) __________ is composed of only hydrogen and carbon.

a. halocarbon

b. salt

c. hydrocarbon

d. organic compound

e. reagent

22. The products of the complete combustion of a hydrocarbon are:

a. water and carbon dioxide.

b. water, carbon dioxide, and carbon monoxide.

c. carbon dioxide and carbon monoxide.

d. ozone.

e. water and ozone.

23. Boiling points of alkanes with un-branched changes:

a. increase with decreasing chain length.

b. increase with increasing chain length.

c. increase with temperature.

d. decrease with size.

24. Alkanes and cycloalkanes are soluble in water.

a. True

b. False

c. Not possible to know

25. The name of a parent compound with alkanes comes from:

a. The shortest sub-unit.

b. The longest chain of carbon atoms.

c. The octane value.

d. The number of other atoms in the chain.

26. Reaction of bromine with a double bond:

a. decolorizes because of evaporation.

b. decolorizes because of reaction.

c. decolorizes because the bromine is no longer present.

d. causes the solution to fume.

e. decolorizes because of reaction and decolorizes because the bromine is no longer present.

Correct = e

Hint = Page 388

LO = 9A8

27. Formation of acetic acid from ethanol is:

a. a process which occurs only in a lab.

b. a process which occurs naturally in the human body.

c. a process which should be carried out only under proper supervision.

d. a process which is an oxidation.

e. a process which occurs naturally in the human body and a process which is an oxidation.

28. The Lewis dot structure of a carbon atom in a molecule must:

a. have 10 electrons around the carbon.

b. have 8 electrons around the carbon.

c. have 6 electrons around the carbon.

d. have 4 electrons around the carbon.

29. Which atom in the following list does not adhere to the Lewis electron dot rule?

a. H

b. Xe

c. N

d. O

e. F

30. Bonding between two carbon atoms is always:

a. covalent.

b. ionic.

c. non-polar.

d. polar.

e. distributed equally.

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