1. Define the following terms:

a. octet rule - main group elements tend to undergo reactions that leave them with eight valence electrons

b. antibonding molecular orbital - a type of molecular orbital that contains electrons that cannot occupy the region between nuclei and cannot contribute to bonding.

c. delocalization - the spreading of p electrons over the molecule

d. bond angle - the angle at which two adjacent bonds intersect

e. bonding electron pair - a pair of valence electrons in a covalent bond

f. bond length - the optimum distance between nuclei in a covalent bond

g. bond order - the number of electron pairs shared between two bonded atoms

h. covalent bond - a bond that occurs when two atoms share several (usually two) electrons

i. electronegativity - the ability of an atom in a molecule to attract the shared electrons in a bond

j. formal charge - the result of a method of electron bookkeeping that tells whether an atom in a molecule has gained or lost electrons compared to an isolated atom

k. hybrid atomic orbital - a set of wave functions derived by combination of atomic wave functions

l. Lewis structure - a representation of an atom that shows valence electrons as dots

m. lone pair electrons - a pair of valence electrons not used for bonding

n. molecular orbital theory - a quantum mechanical description of bonding in which electrons occupy molecular orbitals that belong to the entire molecule rather than to an individual atom

o. pi bond - a bond in which shared electrons occupy a region above and below a line connecting the two nuclei

p. polar covalent bond - a bond in which the bonding electrons are attracted somewhat more strongly by one atom than by the other

q. resonance hybrid - an average of several valid Lewis structures for a molecule

r. sigma bond - a bond in which the shared electrons are centered about the axis between the two nuclei

s. Coordinate covalent bond - a bond formed when one atom donates two electrons to another atom that has a vacant valence orbital

t. electron affinity – the energy change that occurs when an electron is added to an atom (or ion) in the gaseous state. 

u. valence bond theory – bonding occurs when an orbital of one atom overlaps with an orbital of another atom.

2. Write the complete and shorthand electronic configuration for the following atoms and ions.  If you suspect an anomalous configuration may be exhibited show that and explain your reasoning.

a. Si – 1s²2s²2p⁶3s²3p²

i) [Ne] 3s²3p²

b. Mg⁺² - 1s²2s²2p⁶

i) [Ne]

c. S^-2 - 1s²2s²2p⁶3s²3p⁶

i) [Ne] 3s²3p⁶

d. V - 1s²2s²2p⁶3s²3p⁶4s²3d³

i) [Ar] 4s²3d³

e. Ni⁺² - 1s²2s²2p⁶3s²3p⁶3d⁸

i) [Ar] 3d⁸

f. W - 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s¹5d⁵4f¹⁴

i) [Xe] 6s²d⁴f¹⁴

ii) Predict the following to give half filled s and d sublevels  [Xe] 6s¹5d⁵4f¹⁴

g. Am - 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s²5d¹⁰4f¹⁴6p⁶7s²6d¹5f⁶

i) [Rn] 7s²6d¹5f⁶

3. Arrange the following from left to right in order of increasing size:

a. As, Br, Ga, Kr

Kr<Br<As<Ga

b. Mg, Ba, Sr, Be

Be<Mg<Sr<Ba

c. F^-, Ne, F, O^2-

Ne<F<F^-2<O^-2

d. K, K⁺, Ca⁺², Sc⁺³

Sc⁺³<Ca⁺²<K+<K

4. Given the following ionization energy data:

a. Explain why there is an increase in ionization energy from Al to Ar.

Increased effective charge leads to increased ionization energy.

b. Explain the slight decrease in ionization energy from P to S.

S will have a half filled p sublevel if it loses 1 electron.  This is good therefore the ionization energy not that high.  P will lose half filled shell if it loses an electron therefore the ionization energy is high.

5. For each of the following pairs of ionic compounds circle the compound that has the higher lattice energy.  Explain your choice in each case.  (Remember that E_lattice = k Q⁺Q^-/d where k is a constant, Q’s are the ion charges and d is the distance between ions.)

a. LiF or LiBr

LiF has the higher lattice energy because F is smaller than Br.

b. Mg₃N₂ or NaCl

Mg₃N₂ has the higher lattice energy because Mg and N have higher charges than Na and Cl.

6. Draw the best resonance structure for each possible structure and assign formal charges to determine whether the connectivity in dinitrogen oxide is NNO or NON.  Identify the more stable structure and explain your reasoning.

7. Tabun is a chemical warfare agent

a.

i) How many pi  bonds are there in tabun?  3

ii) How many sigma bonds are there in tabun?  20

iii) Give the hybridization of each carbon, nitrogen and oxygen in the molecule.  The CºN C and N are both sp, the double bonded O is sp2, and all other non H atoms are sp3.

iv) Which is the shortest C-N bond in the molecule? CºN

v) What is the P-C-N bond angle?  180^o

vi) What is the C-N-C bond angle?  109^o

8. Complete the following table

BrF₅

Orbital geometry –

Octahedral

Molecular geometry

Square pyramidal

Hybridization of Br

sp³d²

SiO₃^-2

Orbital geometry –

Trigonal planar

Molecular geometry

Trigonal planar

Hybridization of Br

sp²

XeO₃

Orbital geometry –

Tetrahedral

Molecular geometry

Trigonal pyramidal

Hybridization of Br

sp³

PCl₃

Orbital geometry –

Tetrahedral 

Molecular geometry

Trigonal pyramidal

Hybridization of Br

sp³

9.  Draw the compounds XeF₃⁺ and CH₂O.  Explain the bonding in terms of valence bond theory.  That is show the atomic orbitals on the central atom, describe any electron promotion necessary, and show the orbitals involved in both sigma and pi bonding.

10. Nitrogen, N₂, can absorb an electron to give N₂^-.  Compare these species with regard to

a. magnetic character

N₂ as shown will have all paired electrons.  If one additional electron is added it will be in the P* level and will be unpaired. Thus N₂ is diamagnetic and N₂^-1 is paramagnetic.

b. bond order

N₂ has a bonding order of

10 bonding electrons – 4 antibonding electrons /2 = 3

N₂^-1 with an additional electron to be placed in an antibonding orbital has a bond order of

10 bonding electrons – 5 antibonding electrons /2 = 2.5

11. Answer the questions for the structure shown below.  Note that the molecule was drawn to fit the imagination of the instructor and may not faithfully represent the geometry of the molecule.  All bonds and lone pairs are shown.

a. What are the molecular and orbital geometries around Xenon?

Molecular ___square pyramidal_________      Orbital __octahedral___________

b. What is the hybridization of the nitrogen atom?        sp²

c. What is the formal change on the fluorine atom?                   +1       

d. What is the hybridization of carbon?            sp                     Charge on carbon?       0         

e. What are the orbital and molecular geometries around iodine?

Molecular ____T shaped_______        Orbital ______trigonal bipyramidal________

f. What is the hybridization of the bromine atom?                    sp³

g.

What are the orbital and molecular geometries around oxygen?

Molecular ___bent_____         Orbital _____tetrahedral_____

12. Calculate a lattice energy for CaH₂ in kJ/mol using the following information:

a. E_ea for H  =  -72.8 kJ/mol

b. Heat of sublimation for Ca  =  + 178.2 kJ/mol

c. E_i1 for Ca  =  +589.8 kJ/mol

d. Bond dissociation energy for H₂  =  +435.9 kJ/mol

e. E_i2 for Ca  =  +1145 kJ/mol

f. Net energy change for the formation of CaH₂ from its elements  =  -186.2 kJ/mol

Ca(s)  à  Ca(g)                                         + 178.2 kJ/mol

Ca(g)  à  Ca⁺¹  +  e^-1                                + 589.8 kJ/mol

Ca⁺¹(g)  __.  Ca⁺²(g)  +  e^-1                        + 1145 kJ/mol

H₂(g)  à  2 H(g)                                        + 435.9 kJ/mol

2 H(g)  + 2 e^-1 à  2 H^-1(g)             2(-72.8 kJ/mol)

2 H^-1(g)  +  Ca⁺²(g)  à  CaH₂(s)                lattice energy

Ca(s)  +  H₂(g)  à  CaH₂(s)                       -186.2 kJ/mol

+ 178.2 kJ +589.8 kJ + 1145 kJ + 435.9 kJ + 2(-72.8 kJ) + lattice energy = -186.2 kJ

lattice energy = -2389 kJ/mol