SECTION I: Multiple Choice

Please mark the best choice for each question.

Note: For all questions involving solutions and/or chemical equations, assume that the system is in pure water at room tem-

perature unless otherwise noted.

1. A combination of sand, salt, and water is an

example of a ___________________.

a. Homogeneous mixture d. Pure substance b. Heterogeneous mixture e. Solid c. Compound

2. Which state or states of matter are significantly

compressible?

a. gases d. liquids and gases b. liquids e. solids and liquids c. solids

3. If matter is uniform throughout and cannot be

separated into other substances by physical

means, it is _____.

a. a compound d. a heterogeneous mixture b. either an element or a

compound e. an element

c. a homogenous mixture

4. Of the following, only _______________ is a

chemical reaction.

a. Melting of lead d. Crushing of Stone b. Dissolving sugar in water e. Placing a penny in water c. Tarnishing of Silver

5. Of the following, _____ is smallest in mass.

a. 25 kg d. 2.5 x 109 fg b. 2.5 x 10-2 mg e. 2.5 x 1010 ng c. 2.5 x 1015 pg

6. Precision refers to _______________.

a. How close a measured value is to other measured numbers b. How close a measured number is to the true value c. How close a measured number is to the calculated value d. How close a measured number is to zero e. How close a measured number is to the determined mean

7. A scientific __________ is a concise statement or

an equation that summarizes a broad variety of

observations. It has been extensively tested.

a. law d. trend b. hypothesis e. pattern c. theory

8. Which one of the following is NOT one of the

postulates of Dalton’s atomic theory?

a. Each element is composed of tiny, indivisible particles

called atoms b. All atoms of a given element are identical to each other

and different from those of other elements c. During a chemical reaction, atoms are changed into atoms

of different elements d. Compounds are formed when atoms of different elements

combine e. Atoms of an element are not changed into different types

of atoms by chemical reactions

9. The charge on an electron was determined in the

______.

a. cathode ray tube by JJ Thompson b. Rutherford gold foil experiment c. Millikan oil drop experiment d. Dalton atomic theory e. radioactive decay of Carbon by James Chadwick

10. Which combination of protons, neutrons, and

electrons is correct for the 6329Cu isotope of

Copper?

a. 29 protons, 34 neutrons, and 29 electrons b. 29 protons, 29 neutrons, and 63 electrons c. 63 protons, 29 neutrons, and 63 electrons d. 34 protons, 29 neutrons, and 34 electrons e. 34 protons, 34 neutrons, and 29 electrons

11. In the Rutherford nuclear-atom model,

a. The heavy subatomic particles, protons and neutrons,

reside in the nucleus b. The three principle subatomic particles (protons, neutrons,

and electrons,) all have essentially the same mass. c. The light subatomic particles, protons and neutrons, reside

in the nucleus d. Mass is spread essentially uniformly throughout the atom e. The three principle subatomic particles(protons, neutrons,

and electrons) all have essentially the same mass and mass

is spread essentially uniformly throughout the atom.

12. In the periodic table, the elements are arranged in

_____________.

a. Alphabetical order b. Order of increasing atomic number c. Order of decreasing metallic properties d. Order of increasing neutron content e. Chemically families based on bit physical and chemical

properties

13. Which pair of elements would you expect to

exhibit the greatest similarities in the physical and

chemical properties?

a. Oxygen and Sulfur d. Hydrogen and Helium b. Carbon and Nitrogen e. Silicon and Phosphorous c. Potassium and Calcium

14. The elements in groups 1A, 6A, and 7A are called

_______________, respectively.

a. Alkaline earth metals, halogens, and chalcogens b. Alkali metals, chalcogens, and halogens c. Alkali metals, halogens, and noble gasses

d. Alkaline earth metals, transition metals, and halogens

15. Which one of the following does not occur as a

diatomic molecule in elemental form?

a. oxygen d. hydrogen b. nitrogen e. bromine c. sulfur

16. When a metal and a nonmetal react, the ________

tends to lose electrons and the _______ tends to

gain electrons.

a. Metal, metal d. Nonmetal, metal b. Nonmetal, nonmetal e. None of the above, these

elements share electrons c. Metal, nonmetal

17. The correct formula for molybdenum (IV)

hypochlorite is ____________.

a. Mo(ClO3)4 d. Mo(ClO4)4 b. Mo(ClO)4 e. MoCl4 c. Mo(ClO2)4

18. When the following equation is balanced, the

coefficients are ____________.

NH3 + O2 → NO2 + H2O

a. 1, 1, 1, 1 d. 1, 3, 1, 2 b. 4, 7, 4, 6 e. 4, 3, 4, 3 c. 2, 3, 2, 3

19. When the following equation is balanced, the

coefficient of Al is __________.

2 3 2Al (s) + H O (l) Al(OH) (s) + H (g)

a. 1 d. 4 b. 2 e. 5 c. 3

20. Predict the product in the combination

reaction below.

Al (s) + N2 (g) → ____________

21. Pentacarbonyliron (Fe(CO)5) reacts with

phosphorous trifluoride and hydrogen to

release carbon monoxide. The reaction of 5.0

mol of Fe(CO)5, 8.0 mol of PF3, and 6.0 mol of

H2 will release ________ mol of CO.

Fe(CO)5+2PF3+H2→Fe(CO)2(PF3)2H2+3CO

a. 15 d. 6.0 b. 5.0 e. 12 c. 24

22. When H2SO4 is neutralized by KOH in

aqueous solution, the net ionic equation is:

a. SO42- + 2K+ → K2SO4 (aq)

b. SO42- + 2K+ → K2SO4 (s)

c. H+ + OH- → H2O (l) d. H2SO4 + 2OH- → 2H2O (l) + SO4

2-

e. 2H+ + 2KOH → 2H2O (l) + 2K+

23. When aqueous solutions of ___________ are

mixed, a precipitate forms.

a. NiBr2 and AgNO3 d. KOH and Ba(NO3)2

b. NaI and KBr e. Li2CO3 and CsI

c. K2SO4 and CrCl3

24. Which one of the following is a correct

expression for molarity?

a. mol solute / L solvent d. mol solute / kg solvent

b. mol solute / mL Solvent e. μmol solute / L solution

c. mmol solute / mL

solution

25. How many moles of Co2+ are present in

0.200L of a 0.400 M solution of CoI2

a. 2.00 d. 0.0800

b. 0.500 e. 0.0400

c. 0.160

26. There are _____ paired and _______ unpaired

electrons in the Lewis Symbol for a

phosphorous atom. a. 4, 2 d. 4, 3

b. 2, 4 e. 3, 2

c. 2, 3

27. The halogens, alkali metals, and alkaline

earth metals have ____ valence electrons,

respectively.

a. 2, 4, and 6 d. 7, 1, and 2

b. 1, 5, and 7 e. 2, 7, and 4

c. 8, 2, and 3

28. Elements from opposite sides of the periodic

table tend to form ___________.

a. Covalent compounds

b. Ionic compounds

c. Compounds that are gases at room temperature

d. Diatomic compounds

e. Metallic compounds

Questions 29-32

(A) Heisenberg uncertainty principle

(B) Pauli exclusion principle

(C) Hund's rule (principle of maximum multiplicity)

(D) Shielding effect

(E) Wave nature of matter

29. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic C

30. Explains the experimental phenomenon of electron diffraction E

31. Indicates that an atomic orbital can hold no more than two electrons B

32. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron A

Questions 33-35 refer to the phase diagram below of a pure substance.

(A) Sublimation

(B) Condensation

(C) Solvation

(D) Fusion

(E) Freezing

33. If the temperature increases from 10° C to 60° C at a constant pressure of 0.4 atmosphere, which of the processes occurs? A

34. If the temperature decreases from 110° C to 40° C at a constant pressure of 1.1 atmospheres, which of the processes occurs?

B

35. If the pressure increases from 0.5 to 1.5 atmospheres at a constant temperature of 50° C, which of the processes occurs? B

Questions 36-38 refer to the following diatomic species.

(A) Li2

(B) B2

(C) N2

(D) O2

(E) F2

36. Has the largest bond-dissociation energy E

37. Has a bond order of 2 D

38. Contains 1 sigma (σ) and 2 pi (π) bonds C

39. In a molecule in which the central atom exhibits sp3d2 hybrid orbitals, the electron pairs are directed toward the corners of

(A) a tetrahedron

(B) a square-based pyramid

(C) a trigonal bipyramid

(D) a square

(E) an octahedron

40. Based on the information below, what is the standard enthalpy change for the following reaction?

Na2O(s) + H2O(l) 2 NaOH(s)

H2(g) + 1/2 O2(g) H2O(l) ∆H° = x

2 Na(s) + 1/2 O2(g) Na2O(s) ∆H° = y

Na(s) + 1/2 O2(g) + 1/2 H2(g) NaOH(s) ∆H° = z

(A) x + y + z

(B) x + y - z

(C) x + y - 2z

(D) 2z - x - y

(E) z - x - y

41. Which of the following sets of quantum numbers (n, l, ml, ms) best describes the valence electron of highest energy in a

ground-state gallium atom (atomic number 31) ?

(A) 4, 0, 0, 1/2

(B) 4, 0, 1, 1/2

(C) 4, 1, 1, 1/2

(D) 4, 1, 2, 1/2

(E) 4, 2, 0, 1/2

42. CH3CH2OH boils at 78 °C and CH3OCH3 boils at - 24 °C, although both compounds have the same composition. This differ-

ence in boiling points may be attributed to a difference in

(A) molecular mass

(B) density

(C) specific heat

(D) hydrogen bonding

(E) heat of combustion

43. Based on concepts of polarity and hydrogen bonding, which of the following sequences correctly lists the compounds below

in the order of their increasing solubility in water?

X = CH3-CH2-CH2-CH2-CH3

Y = CH3-CH2-CH2-CH2-OH

Z = HO-CH2-CH2-CH2-OH

(A) Z < Y < X

(B) Y < Z < X

(C) Y < X < Z

(D) X < Z < Y

(E) X < Y < Z

44. Concentrations of colored substances are commonly measured by means of a spectrophotometer. Which of the following

would ensure that correct values are obtained for the measured absorbance?

I. There must be enough sample in the tube to cover the entire light path.

II. The instrument must be periodically reset using a standard.

III. The solution must be saturated.

(A) I only

(B) II only

(C) I and II only

(D) II and III only

(E) I, II, and III

45. The data below were gathered in order to determine the density of an unknown solid. The density of the sample should be re-

ported as

Mass of an empty container = 3.0 grams

Mass of the container plus the solid sample = 25.0 grams

Volume of the solid sample = 11.0 cubic centimeters

(A) 0.5 g/cm3

(B) 0.50 g/cm3

(C) 2.0 g/cm3

(D) 2.00 g/cm3

(E) 2.27 g/cm3

46. Which of the following pairs of compounds are isomers?

47. It is suggested that SO2 (molar mass 64 grams), which contributes to acid rain, could be removed from a stream of waste gas-

es by bubbling the gases through 0.25-molar KOH, thereby producing K2SO3. What is the maximum mass of SO2 that could

be removed by 1,000. liters of the KOH solution?

(A) 4.0 kg

(B) 8.0 kg

(C) 16 kg

(D) 20. kg

(E) 40. kg

48. I2(g) + 3 Cl2(g) 2 ICl3(g)

According to the data in the table below, what is the value of ∆H° for the reaction represented above?

Bond Average Bond Energy

(kilojoules / mole)

I---I 150

Cl---Cl 240

I---Cl 210

(A) - 870 kJ

(B) - 390 kJ

(C) + 180 kJ

(D) + 450 kJ

(E) + 1,260 kJ

49. The electron-dot structure (Lewis structure) for which of the following molecules would have two unshared pairs of electrons

on the central atom?

(A) H2S

(B) NH3

(C) CH4

(D) HCN

(E) CO2

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

(A) C6H6 (benzene)

(B) NO

(C) SO2

(D) NH3

(E) H2S

51. _____ Fe(OH)2 + _____ O2 + _____ H2O _____ Fe(OH)3

If 1 mole of O2 oxidizes Fe(OH)2, how many moles of Fe(OH)3 can be formed?

(A) 2

(B) 3

(C) 4

(D) 5

(E) 6

Questions 52-53nrefer to the following types of energy

A) Activation energy

B) Free energy

C) Ionization energy

D) Kinetic energy

E) Lattice energy

52. The energy required to convert a ground-state atom in the gas phase to a gaseous positive ion. C

53. The energy released when gas phase ions bond to form a crystalline solid. E

Questions 54-57 refer to the following descriptions of bonding in different types of solids.

A) Lattice of positive and negative ions held together by electrostatic forces

B) Closely packed lattice with delocalized electrons throughout

C) Strong single covalent bonds with weak intermolecular forces

D) Strong multiple covalent bonds (including π-bonds) with weak intermolecular forces

E) Macromolecules held together with strong polar bonds

54. Cesium chloride, CsCl (s) A

55. Gold, Au (s) B

56. Carbon dioxide, CO2 (s) D

57. Methane, CH4 (s) C

58. When the following equation is balanced, the coefficient of 3 8 3C H O is __________.

3 8 3 2 2 2C H O (g) + O (g) CO (g) + H O (g)

A) 1

B) 2

C) 3

D) 7

E) 5

59. A compound contains 40.0% C, 6.71% H, and 53.29% O by mass. The molecular weight of the compound is 60.05 amu. The

molecular formula of this compound is __________.

A) 2 4 2C H O

B) 2CH O

C) 2 3 4C H O

D) 2 2 4C H O

E) 2CHO

60. What mass of copper is produced when 0.050 00 mol of CuCl2 (aq) is reacted completely with excess Zn(s)?

A) 3.178 g

B) 7.868x10-4 g

C) 6.723 g

D) 3.270 g

E) 6.355 g

61. When a solution of sodium chloride is vaporized in a flame, the color of the flame is _____

A) blue

B) yellow

C) green

D) violet

E) white

62. The phase diagram for a pure substance is shown at right. Which point on the diagram corresponds to the equilibrium be-

tween the solid and liquid phases at the normal melting point?

A) A

B) B

C) C

D) D

E) E

63. Types of hybridization exhibited by the C atoms in propene, CH3 CHCH2, include which of the following?

I. sp

II. sp2

III. sp3

A) I only

B) III only

C) I and II only

D) II and III only

E) I, II, and III

64. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the minimum number of

moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl- as AgCl(s)? (Assume that AgCl is

completely insoluble)

A) 0.10 mol

B) 0.20 mol

C) 0.30 mol

D) 0.40 mol

E) 0.60 mol

65. The ionization energies for element X are listed in the table below. Based on this data, element X is most likely to be

Ionization Energies for Element X (kJ mol-1)

First Second Third Fourth Fifth

580 1 815 2 740 11 600 14 800

A) Na

B) Mg

C) Al

D) Si

E) P

66. Of the following molecules, which has the largest dipole moment?

A) CO

B) CO2

C) O2

D) HF

E) F2

67. Which of the following techniques is most appropriate for the recovery of solid KNO3 from an aqueous solution of KNO3?

A) Paper chromatography

B) Filtration

C) Titration

D) Electrolysis

E) Evaporation to dryness

68. In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic radius?

A) It remains constant

B) It increases only

C) It increases, then decreases

D) It decreases only

E) It decreases, then increases

69. According to the balanced chemical equation, how many moles of HI would be necessary to produce 2.5 mol of I2, starting

with 4.0 mol of KMnO4 and 3.0 mol H2SO4?

10 HI + 2 KMO4 + 3H2SO4 → 5 I2 + 2MnSO4 + K2SO4 + 8 H2O

A) 20.

B) 10.

C) 8.0

D) 5.0

E) 2.5

70. A yellow precipitate forms when 0.5 M NaI (aq) is added to a 0.5M solution of which of the following ions?

A) Pb2+ (aq)

B) Zn2+ (aq)

C) CrO42-(aq)

D) SO42- (aq)

E) OH- (aq)

71. For the reaction of ethylene, ∆Hrxn = -1 323 kJ mol-1. What is the value of ∆H if the combustion produced liquid water H2O

(l), rather than water vapor, H2O (g)?

2 C2H2 (g) + 5 O2 (g) → 4 CO2 (g) + 2 H2O (g)

∆Hvap = -44 kJ mol-1 H2O (g) → H2O (l)

A) -1 235 kJ

B) -1 279 kJ

C) -1 323 kJ

D) -1 367 kJ

E) -1 411 kJ

72. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl (aq) in order to prepare a 0.500 M HCl (aq) so-

lution is approximately

A) 50.0 mL

B) 60.0 mL

C) 100. mL

D) 110. mL

E) 120. mL

Section II: Free Response

YOU MAY USE YOUR CALCULATOR

CLEARLY SHOW THE METHOD USED AND THE STEPS INVOLVED IN ARRIVING AT YOUR ANSWERS. It is to

your advantage to do this, since you may obtain partial credit if you do and you will receive little or no credit if you do not.

Attention should be paid to significant figures.

1. Polystyrene is often used as plastic foam insulation in disposable cups and plates. Polystyrene, like other plastics and like

diamonds and silica, is a network solid. Its carbon and hydrogen atoms are connected by huge networks of covalent bonds. It

isn't made up of identical small molecules, so it doesn't really have a molar mass. But like any compound, polystyrene's

atoms are combined in definite proportions, so it does have an empirical formula. To learn more about the composition of

polystyrene, combustion analysis was performed on a 2.57 g sample. In the combustion analysis, 8.67 g of carbon dioxide

and 1.77 g of water were produced.

a. How many moles of carbon were present in the original sample of polystyrene?

b. How many moles of hydrogen were present in the original sample of polystyrene?

c. What is the empirical formula of polystyrene?

d. Write a balanced chemical reaction for the combustion of Polystyrene.

e. What is the molecular formula of polystyrene if it had a molecular mass of 195 g/mol?

2. One commercial method used to peel potatoes is to soak them in a solution of NaOH for a short time, remove them from the

NaOH, and spray off the peel. The concentration of NaOH is normally in the range of 3 to 6 M. The NaOH is analyzed

periodically. In once such analysis, 45.7mL of 0.500M H2SO4 is required to neutralize a 20.0mL sample of NaOH solution.

a. Write a balanced chemical equation for the mentioned neutralization process.

b. How many moles of H2SO4 are contained in the 45.7mL sample?

d. What volume of 18M solution of H2SO4 is required to prepare the 45.7mLsample that is required for neutralization?

e. What was the concentration of the NaOH solution?

3. A sample of 70.5 mg of Potassium Phosphate is added to 15.0mL of a 0.050 M Silver Nitrate solution, resulting in the

formation of a precipitate.

a. Write a balanced chemical equation for this reaction.

b. How many grams of each compound were put into this reaction?

c. Which compound is the limiting reagent?

d. Calculate the theoretical yield of the precipitate that forms.

e. What was the % yield if you only obtained 9.85 x 10-2 g in the lab?

4. Answer all three portions in this part. Give the formulas to show the reactants and the products for the three following chemi-

cal reactions. Each reaction occurs in aqueous solution unless otherwise indicated. Represent the substances in solution as

ions if the substance is extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. In

all cases a reaction occurs. You need to balance the chemical equation. Then answer a question about the reaction.

Example: A strip of magnesium is added to a solution of silver nitrate

a. Solutions of tin(II) chloride and iron(III) chloride are mixed.

b. Solutions of cobalt(II) nitrate and sodium hydroxide are mixed.

c. Ethene gas is burned air.

d. Solid calcium sulfite is heated in a vacuum.

e. Solid sodium oxide is added to distilled water.

f. A strip of zinc is added to a solution of hydrobromic acid.

Mg +2 Ag+ --> Mg2+ +2 Ag

5. You are given a solution containing barium nitrate, mercury(II) nitrate, and magnesium nitrate. Using only sodium chloride,

sodium fluoride, and sodium hydroxide you need to separate the barium, mercury(II), and magnesium ions. How would you

go about separating these ions? Discuss your experimental procedure and defend your answer.

Using the solubility rules, chloride compounds are generally soluble, but mercury is an exception, so mixing all three

solutions with sodium chloride would mean that only one solution would result in precipitate, which is the mercury solution.

To distinguish between barium and magnesium ions, react with sodium hydroxide. Magnesium is insoluble (precipitate will

form) in hydroxides whereas barium is soluble, so whichever solution forms the precipitate is the magnesium and the soluble

solution is barium.

6. Describe how you would separate a mixture of water, cyclohexanol, and lead shavings. (Water boils at 100° C and

cyclohexanol boils at 161°C.) Classify each mixture and the components at each step of the separation protocol as a

heterogeneous mixture, a homogeneous mixture, an element, or a compound.

The initial mixture of water, cyclohexanol, and lead shavings is a heterogeneous mixture. First, separate the lead shavings by

filtration. The subsequent mixture of water and cyclohexanol is homogeneous (both compounds are polar). To separate the

mixture, distill based on boiling points.

7. A Lewis dot structure can be drawn for boron trifluoride, BF3, in which the octet rule is satisfied for all the atoms.

a. How many total valence electrons are there in this compound? 24

b. Draw the Lewis dot structure for the resonance structure is followed.

c. How many total sigma bonds are there? How many total pi bonds?

3 sigma, 1 pi

d. Write the formal charge of each atom in the Lewis representation next to the symbol for the atom.

Formal Charge = # of valence electrons – # of non-bonding electrons – ½ # of bonding electrons

e. Why is this structure of BF3 unreasonable from the standpoint of electronegativities? The Pauling electronegativity

of boron is 2.0, and that of fluorine is 4.0

The most reasonable Lewis structure is the one with all formal charges closest to zero. If there are charges, then the

most negative formal charge would be assigned to the most electronegative atom. Therefore, it would be more

likely for fluorine (4.0) to be assigned the negative formal charge because it is more electronegative than boron

(2.0).

f. What would a stable resonance structure look like?

8. Use the principles of atomic structure and/or chemical bonding to explain each of the following. In each part, you answer

must include references to both substances.

a. The atomic radius of Li is larger than that of Be.

Both Li and Be have their outer electrons in the same shell (and/or they

have the same number of inner core electrons shielding the valence

electrons from the nucleus). However, Be has four protons and Li has

only three protons. Therefore, the effective nuclear charge experienced

(attraction experienced) by the valence (outer) electrons is greater in Be

than in Li, so Be has a smaller atomic radius.

b. The second ionization energy of K is greater than the second ionization energy of Ca.

The second electron removed from a potassium atom comes

from the third level (inner core). The second electron

removed from a calcium atom comes from the fourth level

(valence level). The electrons in the third level are closer to

the nucleus so the attraction is much greater than for electrons

in the fourth level.

c. The carbon-carbon bond energy in C2H4 is greater than it is in C2H6.

C2H4 has a double bond between the two carbon atoms, whereas C2H6 has a carbon-carbon single

bond. More energy is required to break a double bond in C2H4 than to break a single bond in

C2H6; therefore, the carbon-to-carbon bond energy in C2H4 is greater. d. The boiling point of Cl2 is lower than the boiling point of Br2.

Both Cl2 and Br2 are nonpolar, and the only intermolecular attractive forces are London dispersion

forces. Since Br2 has more electrons than Cl2, the valence electrons in Br2 are more polarizable.

The more polarizable the valence electrons, the greater the dispersion forces and the higher the

boiling point.

9. Answer the following questions about the element selenium, Se (atomic number 34).

a. Samples of natural selenium contain six stable isotopes. In terms of atomic structure, explain what these isotopes

have in common, and how they differ.

b. Write the complete electron configuration (e.g., 1s2 2s2 … etc.) for a selenium atom in the ground state. Indicate

the number of unpaired electrons in the ground-state atom, and explain your reasoning.

c. In terms of atomic structure, explain why the first ionization energy of selenium is

i. less than that of bromine (atomic number 35), and

ii. greater than that of tellurium (atomic number 52).

d. Selenium reacts with fluorine to form SeF4. Draw the complete Lewis electron-dot structure for SeF4 and sketch

the molecular structure. Indicate whether the molecule is polar or nonpolar, and justify your answer.

10. The longest wavelength of light with enough energy to break the Cl-Cl bond in Cl2(g) is 495 nm.

a. Calculate the frequency, in s-1, of the light.

b. Calculate the energy, in J, of a photon of the light.

c. Calculate the minimum energy, in kJ mol-1, of the Cl-Cl bond.

11. A certain line in the sp ectrum of atomic hydrogen is associated with the electronic transition in the H atom from the sixth

energy level (n = 6) to the second energy level (n = 2).

a. Indicate whether the H atom emits energy or whether it absorbs energy during the transition. Justify your answer.

b. Calculate the wavelength, in nm, of the radiation associated with the spectral line.

c. Account for the observation that the amount of energy associated with the same electronic transition (n = 6 to n = 2)

in the He+ ion is greater than that associated with the corresponding transition in the H atom.

3d

3d

12. Refer to the atomic orbitals shown below for the following questions:

A)

B)

C)

D)

E)

b. Represents an atom that is chemically unreactive E

c. Represents an atom in an excited state C

d. Represents a violation of Hund’s rule B

e. Represents an atom that has four valence electrons A

f. Represents an atom of a transition metal D

g. Use the quantum numbers to identify the location of the following electron from Choice A

n = 2, l = 0, ml = 0, ms = -1/2

h. Use the quantum numbers to identify the location of the following electron from Choice D

n = 3, l =2, ml =2,1,0,-1,or -2, ms = ½

i. Use the quantum numbers to identify the location of the following electron from Choice E

n =2, l =1, ml = -1, 0, or 1, ms = -1/2

13. When a 2.000-gram sample of pure phenol, C6H5OH(s), is completely burned according to the equation above,

64.98 kilojoules of heat is released. Use the information in the table below to answer the questions that follow.

C6H5OH(s) + 7 O2(g) → 6 CO2(g) + 3H2O(l)

Substance Standard Heat of Formation,

ΔH°f, at 25°C (kJ/mol)

C(graphite) 0.00

CO2(g) -395.5

H2(g) 0.00

H2O(l) -285.85

O2(g) 0.00

C6H5OH(s) ?

a. Calculate the molar heat of combustion of phenol in kilojoules per mole at 25°C.

b. Calculate the standard heat of formation, ΔH°f, of phenol in kilojoules per mole at 25°C.

14. In an experiment, a sample of an unknown, pure gaseous hydrocarbon was analyzed. Results showed that the sam-

ple contained 6.000 g of carbon and 1.344 g of hydrogen.

a. Determine the empirical formula of the hydrocarbon.

In another experiment, liquid heptane, C7H16(l), is complete combusted to produce CO2(g) and H2O(l), as represented by

the following equation. The heat of combustion, ΔHcombo, for one mole of C7H16(l) is –4.85 x 103 kJ.

C7H16(l) + 11 O2(g) 7 CO2(g) + 8 H2O(l)

b. Using the information in the table below, calculate the value of ΔHfo for C7H16(l) in kJ mol-1.

Compound AHfo (kJ mol-1)

CO2(g) -393.5

H2O(l) -285.8

c. A 0.0108 mol sample of C7H16(l) is combusted in a bomb calorimeter.

i. Calculate the amount of heat released to the calorimeter.

ii. Given that the total heat capacity of the calorimeter is 9.273 kJ/oC-1, calculate the

temperature change of the calorimeter.

15. The reaction between silver ion and solid zinc is represented in the following equation:

2 Ag+ (aq) + Zn (s) → Zn2+ + 2Ag (s)

a. A 1.50 g sample of Zn is combined with 250 mL of 0.110 M AgNO3 at 25o C. Identify the limiting reactant.

Show calculations to support your answer.

b. On the basis of the limiting reactant that you identified in part (a), determine the concentration of Zn2+ after the

reaction is complete. (Assume volume is conserved).

16. The heat of solution of potassium permanganate can be determined in a coffee-cup calorimeter. 10.0 g of solid

KMnO4 is dissolved in 500.0 mL of water. Water has a density of 1 g/mL. The temperature of the solution changes

from 24.8 oC to 23.5 oC. The specific heat of the solution is equal to the specific heat of water, 4.18 J/gK.

a. Is energy being absorbed by the system or released to the surroundings in this reaction?

Absorbed (temp decreases) = endothermic reaction

b. Is the process endothermic or exothermic?

c. Calculate the heat transferred for this process.

q = mCT

q = 500g x 4.184 x (23.5-24.8)

q = -2179.6J

d. Determine the enthalpy of this process, ∆Hsoln.

+2179.6J

e. What is the molar heat of solution for KMnO4?

10g KMnO4 = .063mol (use molar mass)

-2179.6J/.063mol = -34597J = -34.6kJ

17. Why is it that the aluminum atom is larger than the phosphorus atom, but the aluminum ion is smaller than the

phosphorus ion? Defend your answer using periodic trends.

As you go down a group the number of energy levels increase, so atomic radius increases. As you go across a

period effective nuclear charge increases as the number of electrons increases with the increasing pull of the nucleus

since the number of protons are also increasing.

With negative ions, the number of electrons are increasing yet the pull of the positive protons in the nucleus is

decreasing as there are more electrons than protons so the effective nuclear charge is decreased. With positive ions,

the effective nuclear charge increases because there are now less electrons than protons, so the positive pull of the

nucleus is greater than the neutral atom.

18. Neils Bohr deduced his “planetary model” of the atom from the data of the line spectrum of a hydrogen atom. It was

very successful in justifying the empirical relationships of data but did not work for atoms larger than hydrogen.

Schrödinger’s wave function equation modified the Bohr model to account for the splitting of energy levels. We

now have quantum numbers to describe the placement of electrons in an atom. Please explain the advantages in the

modern quantum mechanical model over the Bohr model. Use the following concepts in your answer: Heisenberg

Uncertainty Principle, quantum numbers, orbital, sublevel, Hund’s Rule, Pauli Exclusion Principle.

The Bohr model represents electrons as orbiting the nucleus in set energy levels whereas the Quantum model repre-

sents electrons as having both wave-like and particle-like characteristics with an orbital being a 3-D region with a

high probability of locating electrons. The Bohr model did not consider electrons as wave-like. The Bohr model

can also not explain spectra for larger atoms in explaining energy level transitions.

19. Consider the hydrocarbon pentane, C5H12 (molar mass 72.15 g/mol)

a. Write the balanced equation for the combustion of pentane.

b. The structural formula of one isomer of pentane is shown below. Draw the structural formulas for two other

isomers of pentane. Be sure to include all atoms of hydrogen and carbon in your structures.