TEST 1  

Name______________________________________   Chem. 111-004A           9/18/2007

 

 

 

1. Briefly explain the following Laws & Theory

 

Law of conservation of mass

 

Mass is neither created nor destroyed

 

Law of definite proportions

 

A given compound always contains exactly the same proportion of elements by mass

 

Law of multiple proportions

 

When two elements form a series of compounds, the ratio of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers

 

Dalton’s Atomic Theory

 

1) All matter is composed of small indivisible particles called atoms. 2) All atoms of one element are identical; atoms of different elements differ in mass and other fundamental ways. 3) Chemical compounds form when different elements combine. 4) Chemical reactions involve reorganization of atoms; atoms do not change in a chemical reaction

 

 

 

 

2. The speed of the space shuttle in orbit is 7860 m/s. What is its speed in mi/hr if 1 in = 2.5400cm (exact), 1 ft = 12.00 inches (exact), and 1 mi = 5,280.00 ft (exact).

 

7860 m/s  x  60s/min  x  60 min/hr  x  100 cm/m  x  1 in/2.54 cm  x  1 ft/12 in

 x  1 mi/5280 ft  =  17,600 mi/hr  or 1.76 x 104 mi/hr

 

3. Convert the melting point of the greenhouse gas carbon dioxide (216 K) to

 

a)     oC ;  oC = K-273.15 = 216K-273.15 = -57

 

 

b)     oF; oF=9/5C  +32 = 9/5(-57oC)  +  32  = -71

 

 

 

 

 

 

 

 

  1. a. The density of pure copper is 8.96g/ml. If a block of pure copper  displaces 144.45 ml of water when placed in a container filled with water, what is the mass of the block of copper.

 

            144.45 ml H2O  x  8.96 g Cu/ ml H2O  = 1290 g Cu

 

 

 b. You received two gold looking rings from two different friends. One is pure gold, the other is an alloy of gold and silver.  What simple experiment/ determinations would allow you to determine which ring is pure gold. Assume the volume of each ring is different .

 

            Determine density of each ring by measuring its mass on a balance      and its volume by displacement of water. The ring that has the density of gold is pure; the other is an an alloy of silver and gold.

 

 

 

 

 

 

 

5. Fill in the following table of isotopes:

 

Symbol of isotope

# of Protons

# of neutrons

Mass number

# of electrons

Net charge

68Zn2+

 

30

38

68

28

2+

90

    Sr

38

52

90

38

0

32     3-

     P 

15

17

32

18

-3

 

 

 

 

6. Name each of the following

 

a)      Sr3N2 ____Strontium nitride___________________________

 

b)      V(NO2)3 _Vanadium III nitrite____________________

 

 

c)      S2Br2 ____Disulfur dibromide_________________________

 

 

d)      HNO2 _______Nitrous Acid_________________________

 

 

e)      HI(aq)___Hydroiodic acid_______________________________________

 

 

7. Write the formula for each of the following compounds

 

a)      Sodium sulfide__Na2S________________________________

 

 

b)      Strontium nitride ____Sr3N2__________________________________

 

 

c)      Chromium  III sulfate ____Cr2(SO4)3_____________________________

 

 

d)      Bromous Acid ____HBrO2_________________________________

 

 

e)      Chlorous acid _____HClO2___________________________________

 

 

8.    .    The freezing of water is a

               a)    chemical and physical damage          

               b)    physical change because the solid water is chemically the same as the liquid.

               c)    chemical change because heat is given off  when the process  occurs.

               d)    chemical change because solid water expands when freezing.

               e)    Physical change because the water merely turns to a solid.

.

            answer is b or e

9. Substances with constant composition that can be broken down into elements by chemical processes are called

               a)    Compounds

               b)    Mixtures

               c)    Solutions

               d)    Quarks

answer is a

 

10 .  A homogeneous  mixture. is also called a

               a)    pure mixture

               b)    heterogeneous mixture.

               c)    solution.

               d)    compound.

answer is c

 

 

11.   An example of a pure substance is  (an)

               a)    element.

               b)    compound.

               c)    pure table salt (sodium chloride).

               d)    carbon dioxide.

               e)    all of these

 

answer is e

 

12.       Generally observed behavior which can be formulated into a statement, sometimes mathematical in nature, is called a(n)

               a)    observation.

               b)    natural law.

               c)    theory.

               d)    measurement.

               e)    experiment.

 

answer is b

_________________________________________________________________________________________________ 

 

 

 TEST 2

 

Name_________KEY___________________________________ Chem111-004A 10/16/07

  1. (7 pts) a. How many molecules of SO 3 are in 5.00 g of SO 3?

 

5.00 g SO 3 x 1 mol SO 3/80.07 g SO 3 x 6.02 x 10 23 molecules SO 3/ 1 mol SO 3= 3.76 x 10 22 molecules SO 3

 

b. (7 pts) What mass is present in 1.85 mol of Na?

1.85 mol Na x 22.99 g Na/ 1 mol Na = 42.5 g Na

 

2. (7 pts) How many moles are present in 1.00 kg of Br 2?

1.00 kg Br 2 x 1000 g Br 2/1 kg Br 2 x 1mol Br 2/159.8 g Br2 = 6.26 mol Br 2

 

3. (7 pts) How does an empirical formula differ from a molecular formula? Give examples of each in your answer.

Empirical formula shows simplest ratio of elements in the formula. Example -CH

Molecular formula shows actual number of atoms of each element in the formula. Example-C 6H 6

 

4. (7 pts) If a protein contains 0.390% Cu, how many grams of Cu are present in 25.00 g of the protein?

25.00 g Protein x .390 g Cu/100 g Protein = 0.0975 g Cu

 

5. (7 pts) Balance the following reaction by inspection:

Balanced

Al(OH) 3(s) + 3HCl(aq) → AlCl 3 + 3H 2O

 

6. (10pts) For the following unbalanced equation:

 

P4(s) + F2(g) → PF3(g)

What is the theoretical yield if 72.50g of P 4 react with excess F 2

Balancing equation gives

P4(s) + 6F2(g) → 4PF3(g)

72.50 g P4 x 1 mol P4/123-88 g P4 x 4 mol PF3/1 mol P4 x 87.97g PF 3/1 mol PF3

= 205.9g PF 3

7. (10 pts) What mass of H 2(g) and I 2 (s) is required to prepare 22.5g of HI (g)

H 2 + I 2 + 2HI

22.5 g HI x 1mol HI/127.91 g HI = 0.176 mol HI

0.176 mol HI x 1 mol H 2/2 mol HI x 2.016 g H 2/ 1 mol H 2 = 0.177g H 2

0.176 mol HI x 1 mol I 2/2 mol HI x 253.8 g I 2/ 1 mol I 2 = 22.3 g I 2

 

8. (7 pts) What mass of NaOH is required to prepare 3.00L of 0.352M NaOH

VM = mol

(3.00L)(0.352 M) = 1.06 mol NaOH

1.06 mol NaOH x 40.00g NaOH/1 mol NaOH = 42.4 g NaOH

 

9. (7 pts) Write the net ionic equation when aqueous solutions of Iron III chloride and sodium hydroxide react.

Balanced equation for reaction

FeCl 3(aq) + 3NaOH(aq) → Fe(OH) 3(s) + 3NaCl(aq)

Net ionic equation

Fe 3+(aq) + 3OH -(aq) → Fe(OH) 3(s)

 

 

 

 

 

10. (7 pts) What is the concentration of all the ions that form when 0.100 mol of CrCl 3 is dissolved in water to a final volume of 125.00 ml.

CrCl3(s) → Cr 3+(aq) + 3Cl -(aq)

M= mol/L

0.100 mol CrCl 3/-o.12500L = 0.800 M CrCl3(aq)

From relationships in formula of CrCl 3 and/or above equation:

0.800 M CrCl3(aq) x 1 M Cr 3+/ 1 M of CrCl 3 = 0.800 M Cr 3+

0.800 M CrCl3(aq) x 3M Cl -/1M of CrCl 3 = 2.40 M Cl -

 

 

 

11. (7 pts) Complete the following acid-base reaction:

HClO 4(aq) + Cr(OH) 3(s) → Cr(ClO 4) 3 + 3H 2O

 

 

12. (4pts) a. Assign oxidation numbers to all atoms in the following reaction

 

3+ 5+ 2- 6+ 2- 4+ 2-

Cr 3+ + ClO 3 - --> Cr 2O 7 2- + ClO 2

 

b. (6 pts) Balance this oxidation-reduction reaction in acidic solution using the half reaction method.

1[7H 2O + 2Cr 3+ → Cr 2O 7 2- 14H + 6e -]

6[1e - + 2H + ClO 3 - → ClO 2 + H 2O]

 

Adding above gives

7H 2O + 2Cr 3+ + 6e - + 12H + + 6ClO 3 - → Cr 2O7 2- + 14H + + 6e - + 6ClO 2 + 6H 2O

Cancelling what is common to both sides gives:

H 2O + 2Cr 3+ + 6ClO 3 - → Cr 2O7 2- + 2H + + 6ClO 2

________________________________________________________________________________________________________

 

TEST 3

 

07chem111-004test3C5 &6.1-6.3

Multiple Choice

Identify the letter of the choice that best completes the statement or answers the question.

1. The molar masses of helium and oxygen are 4.0 g/mol and 16 g/mol, respectively. At the same temperature and pressure, 1 mole of helium will occupy

a.

the same volume as 1 mole of oxygen.

b.

four times the volume of 1 mole of oxygen.

c.

twice the volume of 1 mole of oxygen.

d.

half the volume of 1 mole of oxygen.

e.

one fourth the volume of 1 mole of oxygen.

Ans. a

2. Which conditions will cause the least deviation from the ideal gas law?

a.

100 atm and 500 K

b.

100 atm and 10 K

c.

0.001 atm and 500 K

d.

0.001 and 10 K

e.

0.001 and 273 K

Ans. c

3. All of the following statements are true EXCEPT

a.

the enthalpy change of an endothermic reaction is positive.

b.

at constant pressure the heat flow for a reaction equals the change in enthalpy.

c.

D H for a reaction is equal in magnitude but opposite in sign to D H for the reverse reaction.

d.

enthalpy is a state function.

e.

enthalpy change is dependent upon the number of steps in a reaction.

Ans. e

4. Water can be decomposed by electrolysis to hydrogen gas and oxygen gas. What mass of water must decompose to yield 48.0 L of oxygen gas at 1.00 atm and 25ºC?

2H 2O(l) gives 2H 2(g) + O 2(g)

Moles of O2; n=PV/RT; (1 atm)(48.0L)/(0.08206 L.atm/K.mol)(298) = 1.96 mol

1.96 mol O2 x 2 mol H2O/1 mol O2 x 18.02 g/1 mol H2O = 70.7 g H2O

 

 

 5. A balloon is filled with oxygen gas to a volume of 1.75 L at 36ºC. The balloon is then heated to 72ºC. What is the volume of the heated balloon?

V1/T1=V2/T2; V2=V1 (T2/T1)

V2 = 1.75 L (345K/309K) = 1.95 L

 

6. At 108ºC, the pressure in a 10.0 L flask is 874 mm Hg. How many moles of gas are in the flask?

n=PV/RT;

n=(10.0L)(874mm x 1 atm/760mm)/(0.08206L.atm/K.mol)(381K) = 0.0368 mol

7. A balloon is filled with 1.50 L of helium gas at sea level, 1.00 atm and 32ºC. The balloon is released and it rises to an altitude of 30,000 ft. If the pressure at this altitude is 228 mm Hg and the temperature is -45ºC, what is the volume of the balloon?

P1V1/T1=P2V2/T2; V2 =V1 (T2/T1)(P1/P2)

V2 = 1.50L(228 K/305K)(1.00 atm/228 mm Hg(1atm/760 mm Hg) = 3.74 L

 

8. The density of hydrogen gas in a flask is 0.147 g/L at 305 K. What is the pressure inside the flask?

P=dRT/MM

P=(0.147 g/L)(0.08206 L.atm/K.mol)(305K)/(2.028 g/mol) = 1.82 atm

9. A volume of 3.0 L of butane is burned in excess oxygen. Balance the chemical equation below and determine how many total liters of gases are produced? Assume that both the reactant and product temperature is 500 K and the pressure of the system remains constant at 1.0 atm.

C 4H 10(g) + O 2(g) gives CO 2(g) + H 2O(g)

 

C 4H 10(g) +13/2 O 2(g) gives 4 CO 2(g) + 5H 2O(g)

 

3.0 L C 4H 10(g) x 4.0L CO 2(g)/ 1L C 4H 10(g) = 12 L

 

3.0 L C 4H 10(g) x 5.0L H 2O (g)/ 1L C 4H 10(g) = 15 L

 

Total volume = 12L +15L = 27 L

10. A 10.0 L flask is used to collect 0.500 moles of N 2 and 0.180 moles of O 2 over water at 30ºC. What is the pressure in the flask? (vapor pressure H 2O(l) = 31.8 mm Hg)

P total = P N2 + P O2 + P H2O

n N2 + n O2 = 0.500 mol + 0.180 mol = 0.680 mol

P N2 + O2 = (0.680mol)(0.0806 L.atm/K.mol)(303K)/10.0L = 1.69 atm

P H2O = (31.8 mm Hg/1 atm/760 mm Hg) = 0.04 atm

P total = 1.69 + 0.04 = 1.73 atm

11. Aluminum has a specific heat of 0.902 J/g·ºC. How many joules of heat are required to change the temperature of 8.50 grams of aluminum from 25.0ºC to 93.4ºC?

q = Specific heat x mass x ∆T

q= 0.902 J/g oC x 8.50 g x 68.4 oC = 524 J

` 12. Determine the heat of reaction for the decomposition of one mole of benzene to acetylene

C 6H 6(l) gives 3C 2H 2(g)

given the following thermochemical equations:

2C 6H 6(l) + 15O 2(g) ® 12CO 2(g) + 6H 2O(g)

D H = -6271 kJ

2C 2H 2(g) + 5O 2(g) ® 4CO 2(g) + 2H 2O(g)

D H = -2511 kJ

 

½[2C 6H 6(l) + 15O 2(g) ® 12CO 2(g) + 6H 2O(g)]

D H = (1/2)-6271 kJ

-3/2[2C 2H 2(g) + 5O 2(g) ® 4CO 2(g) + 2H 2O(g)]

D H = (-3/2)-2511 kJ

∆Hrx = -3136 kJ + 3766kJ = 630 kJ

13. Consider three identical flasks filled with different gases.

Flask A: CO 2 at 760 torr and 0 oC

Flask B: N 2 at 760 torr and 0 oC

Flask C: H 2 at 760 torr and 0 oC

a. In which flask will the molecules have the greatest average kinetic energy? All Same

b. In which flask will the molecules have the least average velocity? Flask A

c. Which gas will effuse the fastest from a container with a small hole in it? Flask C

14. Explain the following :

a. Law of Conservation of Energy

Energy can be converted from one form to another, but can be

neither created or destroyed

b. ideal gas

A gas that follows the equation PV=nRT

c. enthalpy

H =E + PV or ∆H = q p

___________________________________________________________

TEST 4

Name________________________________________________________ 11/27/07

 

07fallchem111test4Chapters 6.4, 7.1-7.13.8.1-8.2

 

 

1.   What is the ground state electron configuration (1s 22s 2….) for the following atoms:? You may use the appropriate noble gas symbol to incorporate core electrons.

 

 

a.    50Sn [Kr} 5s24d105p 2

 

 

 

b .    24Cr [Ar] 4s13d 5

 

2.   Using molar enthalpies of formation, determine the heat of reaction for the combustion of 1.0000 mole of methanol.

 

 

2CH 3OH(l) +

3O 2(g) ®

2CO 2(g) +

4H 2O(l)

DHfº (kJ/mol)

-238.7

 

-393.5

-285.8

 

DHrxº = [2mol(-393.5kj/mol) + 4mol(-285.8kj/mol]-[2 mol(-238.7kj/mol) = 1452.8kj/2mol CH 3OH

 

                     = 764.2kj/mol

 

 

 

3.   Many hand held laser pointers emit 650 nm light. What is the frequency of this light?

 

 

n= c/l = 3.00 x 10 8 m/s / 650 x 10 -9 m = 4.62 x 10 14 s -1

 

 

4.   A helium-neon (or HeNe) laser emits light at 632.8 nm. What is the energy of a single photon from this laser?

 

 

DE = hc/l = (6.626 x 10 -34 J s) ( 3.00 x 10 8 m/s)/ 632.8 x 10 -9 m = 3.141 x 10 -19 J

 

 

 

  • Calculate the energy of a photon in the Balmer series that results from the transition n = 5 to n = 2.

(DE= -2.178 x 10-18 J (1/n final 2 - 1/n initial 2)

.

 

(DE= -2.178 x 10-18 J (1/2 2 - 1/5 2) = - 4.574 x 10 -19 J

 

6.   The Schrödinger wave equation

a.

proves electrons have positive and negative spins.

b.

calculates the precise position and momentum of an electron at any given time.

c.

is used to compute the wavelength of small particles.

d.

can be solved to find the probability of finding an electron in a region of space.

e.

proves that photons are particles.

 

Ans d

 

 

 

7.   Which of the following sets of quantum numbers refers to a 2s orbital?

a.

n = 1, = 2, m = 2, m s = + 1/2

b.

n = 1, = 2, m = 1, m s = + 1/2

c.

n = 2, = 2, m = 0, m s = + 1/2

d.

n = 2, = 1, m = -1, m s = + 1/2

e.

n = 2, = 0, m = 0, m s = + 1/2

 

Ans. E

 

 

8.   How many orbitals have the following quantum numbers: n = 6, = 2 , m = -2?

a.

0

b.

1

c.

3

d.

5

e.

6

 

Ans. b

 

  • What does the Heisenberg Uncertainty Principle tell us about what we can, or cannot know, about an electron in an orbital?

 

The limitation on how precision we can know the position and momentum of an electron in an orbital. As we know on with greater certainty, we know the other with less certainty. The equation which represents this principle is Dx x D (mv) = h/4p

 

 

 

 

 

 

 

 

 

10.   What is the correct orbital diagram for Ca 2+?

 

1s

2s

2p

3s

3p

4s

 

a.

(­¯)    (­¯)    (­¯)(­¯)(­¯)    

b.

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­ )

c.

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­¯)

d.

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­ )

e.

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)

 

Ans. c

 

11.   In general, ionization energies

a.

increase down a group and increase across a period.

b.

increase down a group and decrease across a period.

c.

decrease down a group and increase across a period.

d.

decrease down a group and decrease across a period.

e.

increase with atomic mass and increase with atomic radii.

 

Ans. c

 

 

12.   Place the following atoms in order of increasing size: Al, Cl, Mg, O, and P.

a.

Cl < O < P < Al < Mg

b.

Cl < P < Al < Mg < O

c.

O < Cl < P < Al < Mg

d.

O < Mg < Al < P < Cl

e.

none of the above

 

Ans. c

 

 

13.   The Pauli exclusion principle states that

a.

no two electrons from a given atom can have the same spin.

b.

no two electrons from a given atom can have the same four quantum numbers.

c.

two electrons can occupy an orbital if they have the same spin.

d.

two electrons can occupy an orbital if they have opposite spins.

e.

two electrons can occupy an orbital if they have opposite charges.

 

Ans. B

 

 

 

 

 

 

14.   What is the total capacity of electrons in n = 5, = 3?

a.

2

b.

6

c.

10

d.

14

e.

32

 

Ans. d

 

15.   Write of the chemical equation for the second ionization energy of a gaseous calcium atom?

 

Ca (g) + ®   Ca (g) 2+ + e -1

 

 

 

 

16.   Electronegativity increases

a.

moving down a group in the periodic table.

b.

moving from left to right across the periodic table.

c.

with increasing atomic mass.

d.

when electrons are paired.

e.

with increasing atomic radii.

 

Ans. b

 

17.   A chemical bond resulting from the electrostatic attraction between positive and negative ions is called a(n)

a.

covalent bond.

c.

charged bond.

b.

ionic bond.

d.

dipole bond.

 

Ans. b

 

 

18. Answer the following questions for an atom with the following orbital electron configuration:

 

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­¯)    

 

  • Does this represent a ground state or excited state? ___Ground____________________________

 

 

  • Assuming the atom has a 1+ charge, what element does this represent? ___Potassium__(K)__________

 

  • How many core electrons does the ground state of this ion have? ______18____________

 

  • How many valence electrons does the ground state electron configuration of the neutral atom have?

 

Ans. 1

 

  • Write a ground state orbital electron configuration for the 2+ ion?

 

(­¯)    (­¯)    (­¯)(­¯)(­¯)    (­¯)    (­¯)(­¯)(­ )    

 

 

 

 

  • Write a permissible set of quantum numbers for one of the electrons in a 2p orbital of the 1+ ion.

 

. [2,1,-1,+1/2]_others possibles

 

 ________________________________________________________________________________________________________

 TEST 5

Name______________________________________________ 12/11/07

 

 

chem 111004afallt5c8&9

 

Multiple Choice

Identify the choice that best completes the statement or answers the question.

 

 

Points (pts): 1-12 (6pts); 13 (8 pts); 14 (13 pts); 15 (9 pts); 16 (8 pts)

 

__e__ 1. Place the following ions in order of increasing size: Al 3+, F -, Mg 2+, and N 3-.

a.

F - < Mg 2+ < N 3- < Al 3+

b.

F - < N 3- < Al 3+ < Mg 2+

c.

F - < N 3- < Mg 2+ < Al 3+

d.

N 3- < F - < Mg 2+ < Al 3+

e.

Al 3+ < Mg 2+ < F - < N 3-

 

 

__b__ 2. Which of the following are correct resonance structures of SO 3. S is the central atom and all oxygens are bonded to th S.?

 

 

a.

(1) and (5)

b.

(2) and (4)

c.

(1), (2), and (4)

d.

(2), (3) and (4)

e.

(1), (2), (4), and (5)

 

 

__a__ 3. What is the hybridization of the carbon atoms in ethyne, C 2H 2 (HC ´ CH)?

a.

sp

b.

sp 2

c.

sp 3

d.

sp 3d

e.

sp 3d 2

 

 

 

 

 

__a__ 4. Which of the following is a correct Lewis structure for oxygen?

a.

b.

c.

d.

e.

 

 

__b__ 5. How many lone pairs of electrons are on the sulfur atom in sulfite ion, SO 3 2-?

a.

0

b.

1

c.

2

d.

3

e.

4

 

 

__c__ 6. What is the formal charge on each atom in CN -?

a.

C = 0, N = 0

b.

C = +1, N = -1

c.

C = -1, N = 0

d.

C = +2, N = -3

e.

C = +4, N = -5

 

 

__a__ 7. Use VSEPR theory to predict the molecular geometry of BH 3.

a.

triangular planar

b.

triangular pyramidal

c.

linear

d.

tetrahedral

e.

triangular bipyramidal

 

 

__a__ 8. Use VSEPR theory to predict the molecular geometry of HSeH.

a.

bent

b.

linear

c.

tetrahedral

d.

triangular planar

e.

triangular pyramidal

 

 

__c__ 9. What are the bond angles in CO 3 2-. (All oxygens bonded to C)?

a.

90º

b.

109.5º

c.

120º

d.

90º and 120º

e.

180º

 

 

__d__ 10. Which of the following molecules contain polar covalent bonds: CO, N 2, NH 3, and HCl?

a.

CO and HCl

b.

CO and NH 3

c.

CO, NH 3, and N 2

d.

CO, NH 3, and HCl

e.

All of the species contain polar covalent bonds.

 

 

__d__ 11. How many sigma and pi bonds are present in H 2CO (All hydrogen atoms and the oxygen atom are bonded to C) ?

a.

1 sigma bond and 3 pi bonds

b.

2 sigma bonds and 2 pi bonds

c.

2 sigma bonds and 1 pi bond

d.

3 sigma bonds and 1 pi bond

e.

none of the above

 

12. Using the trends in electronegativity in the periodic table, Which bond in each of the following is least polar.

 

I. a. Al-Br b. GaBr c. InBr d. Tl Br

 

Ans. d

 

II. a. C-H b. N-H c.O-H d. F-H

 

Ans. a

 

  • The SCN - ion has following resonance structures. Using formal charge calculations, evaluate each resonance form and indicate which is more preferred. Your answer must be justified by formal charge

calculations to receive credit.

 

 

 

3) S=0; C=0; N=1- (Preferred, low charges and more electronegative nitrogen atom has negative charge)

 

1) S= 1+; C=0; N=2-

2)Invalid carbon has 10 electrons

  • S=1-; C=0; N=0

 

 

 

 

 

 

 

 

14. a. Using the following molecular orbital diagram ,write the molecular electron configurations for F 2.

     
SEE TEXT: FIGURE 9.39, PAGE 410

 

 

 

 

b. Calculate the Bond Order (BO)

 

 

BO= ½(boding electrons –antibonding electrons)= (8-6)=1

 

 

 

c. Is the molecule paramagnetic or diamagnetic? Explain your answer

 

diamagnetic-no unpaired electrons

 

 

 

d. Is F 2 + paramagnetic or diamagnetic? Explain your answer

 

paramagnetic-one unpaired electron

 

 

 

 

15. Give the expected hybridization and bond angles for each lettered central atom in the following compound?

 

 

 

H2C a =

C bH -

O c -

C d(=O) -

C eH2 -

N fH2

Hybridization

sp 2

sp 2

sp 3

sp 2

sp 3

sp 3

 

 

 

 

 

 

 

Molecular Structure

Trigonal

Planar

Trigonal Planar

Bent

Trigonal Planar

Tetrahedral

Trigonal Pyrimidal

 

 

 

 

 

 

 

Approximate Bond Angle 0

120

120

<109

120

109

<109

 

 

 

16. Briefly explain the difference between ionic and covalent bonding.

 

Ionic Bonding: The electrostatic attraction between oppositely charged ions

 

Covalent Bonding: A type of bonding in which electrons are shared.

 

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