Bonding Activities (topic 4/14):
Lesson 1 – Ionic Bonding and structure & Covalent Bonding
1. Name one ionic compound that contains only non-metallic elements.
2. Name the following ions:
a. NH4+
b. OHc. NO3d. HCO3e. CO3-
f.
SO42-
g. PO43-
3. Write Lewis dot symbols for the following ions: Li+, Cl-, S2-, Sr2+, N3-
4. Use Lewis dot symbols to show the formation of aluminum oxide.
5. Why are solid ionic compounds rather poor conductors of electricity? Why does conductivity increase when an
ionic compound is melted or dissolved in water?
6. What is lattice energy and what role does it play in the stability of ionic compounds?
7. Specify which compound in the following pairs of ionic compounds has the higher lattice energy: (a) KCl or
MgO; (b) LiF of LiBr; (c) Mg3N2 or NaCl. Explain your choice.
8. What is Coulomb’s law and what are its implications in determining lattice energies?
9. Explain, according to Coulomb’s law, why MgCl2 has a much stronger attraction than NaCl.
10. Define ionic character as it relates to binary compounds and how is it used to distinguish types of bonding?
11. Explain why ions with charges greater than 3 are seldom found in ionic compounds.
4.2 Covalent Bonding
1. How does electronegativity relate to electron affinity and ionization energy?
2. Considering electronegativity values, which of the following bonds is covalent, which is polar covalent, and
which is ionic?
(a) the bond in CsCl
(b) the bonds in H2S
(c) the NN bond in H2NNH2
3. List the following bonds in order of increasing ionic character;
a. lithium-to-fluorine bond in LiF
b. potassium-to-oxygen bond in K2O
c. nitrogen-to-nitrogen bond in N2
d. sulfur-to-oxygen bond in SO2, the chlorine-to-fluorine bond in ClF3
4. What is the relationship between bond length and bond strength
5. What is bond enthalpy and how does it relate to compound stability?
6. Bond Enthalpy and Lattice Energy Exercise
Objective:
 Understand the difference of bond strength in the different types of bonds of a covalent compound
 Compare bond energies in covalent compounds to the energy involved in an ionic structure
Directions:
I. Research the average bond enthalpies for bonds in polyatomic molecules and calculate the energy needed to
break the bonds indicated in the structures below
1. The C-H bond in methane (draw Lewis structure of the molecule)
2. The C-C bond in ethane (draw Lewis structure of the molecule)
3. The C-C bond in ethene (draw Lewis structure of the molecule)
4. The C-C bond in ethyne (draw Lewis structure of the molecule)
5. The C-C bond in benzene (draw Lewis structure of the molecule)
6. How do bond energies compare between single, double, and triple bonds? How do these compare to a
resonant bond?
II. For the structures above (1-5), calculate the amount of energy needed to break every bond in the structure. Redraw the structure on the space provided to the right of each number and report your answers on the line.
1. ___________________________
2. ____________________________
3. ____________________________
4. ____________________________
5. ______________________________
III. Draw the structure of TNT and calculate the amount of energy that a 2 kg bag of this explosive will release
after breaking every single bond.
IV. Apply your understanding of the energy involved in breaking/forming bonds to calculate the enthalpy of the
following reaction and predict whether the reaction will release energy (exothermic) or absorb energy
(endothermic):
6CO2 + H2O  C6H12O6 + O2
V. Find the lattice energy for NaCl and compare it to the bond energy in the C-Cl bond of
CCl4.
7. What is bond polarity? How does this relate to the understandings of electronegativity and electron affinity?
4.3 Covalent structures
1. Write Lewis structure for NF3, CS2, HNO3, HCOOH, CO3-2, NO2-, OF2, N2F2, Si2H6, OH-, CH2ClCOO-, CH3NH3+.
For each structure, identify the electron domain geometry and molecular geometry.
2. For each of the following molecular geometries, provide the bond angles associated with the geometry.
a. Linear
b. V-shaped with one lone pair
c. V-shaped with two lone pairs
d. Trigonal Planar
e. Trigonal Pyramidal
f.
Tetrahedral
3. Explain why the magnitude of repulsion between electrons decreases in the following order: lone pair-lone pair >
lone pair-bonding pair > bonding pair-bonding pair.
4. Write Lewis structures for the following ions: (a) O22-, (b)C22-, (c) NO+, (d) NH+
5. The following Lewis structures for (a) HCN, (b) C2H2, (c) SnO2, (d) BF3, (e) HOF, (f) HCOF, and (g) NF3 are
incorrect. Explain what is wrong with each one and give a correct structure for the molecule. (Relative positions
of atoms are shown correctly.)
6. What is a coordinate covalent bond? When will a compound exhibit coordinate covalent bonds?
7. What is meant by the dipole moment of a molecule? Why are some molecules with polar bonds nonpolar?
Return to question 4 of this section and add the resulting dipole moment?
8. How are delocalized electrons in covalent compounds associated with resonant bond structures?
9. What is a resonant hybrid?
10. Define bond length, resonance, and resonance structure. What are the rules for writing resonance structures?
11. Draw the resonance structures for the molecule nitrous oxide, N2O. Do the same for SCN-, the chlorate ion, O3,
and OCN-.
12. Draw the Lewis electron-dot structures for CO32-, CO2, and CO, including resonance structures where appropriate.
Which of the three species has the shortest C-O bond length? Explain the reason for your answer.
13. Which elements are the most notable for forming giant molecular structures? What do these elements have in
common?
14. How do the structures of these giant molecular compounds determine the properties of these molecules?
15. Why does graphite conduct electricity and not heat, while diamonds do not conduct electricity but will efficiently
conduct heat, when both are simply giant molecular structures?
4.4 Intermolecular Forces
Compounds
Cl2 and CBr4
I2 and NO3-
CH3OCH3 and
CH4
HCOOH and
C6H14
Br2 and ICl
SiH4 and SnH4
Molecular Representation:
Intramolecular dipole AND
Net dipole (use different
colors to distinguish
between the two)
Research:
Actual boiling
point for each
compound
Conclusion:
Compare prediction to actual
data and explain results in
terms of intermolecular forces
PF3 and NaCl
SO2 and CO2
HF and HI
NH3 and PH3
NH3 and NCl3
HCl and Cl2
Br2 and Cl2 and F2
1. Which elements allow for the highest strength of dipole-dipole intermolecular forces?
2. What form of intermolecular forces can exist between non-polar molecules? Explain why only this form of
intermolecular force is possible.
3. Explain what is meant by the term “induced dipole”. What factors can affect the strength of this temporary dipole
attraction?
4. Is hydrogen bonding a bonding force or an intermolecular force of attraction between molecules? Explain.
4.5 Metallic Bonding
1. Why do alloys have different properties from the elements that are mixed to form the alloys?
14.1 Further aspects of covalent bonding and structure
Exceptions to the Octet Rule
1. What is meant by the term expanded octet
2. What VSEPR shapes are associated with molecules where the central atom has 5 electron domains? What
feature determines the shapes of these molecules? What bond angles are associated with each of these
shapes?
3. What VSEPR shapes are associated with molecules where the central atom has 6 electron domains? What
feature determines the shapes of these molecules? What bond angles are associated with each of these
shapes?
4. How many electron domains exist for the central atom in molecules with the following geometries:
Square planar
Octahedral
Square pyramidal
Trigonal pyramidal
Linear
5. What is meant be the term reduced octet?
6. Which elements are noted for their ability to form covalent compounds instead of ionic compounds and
do so by forming incomplete octets?
7. Several covalent compounds can form where there is an odd number of electron surrounding the central
atom. Describe the general stability of these molecules. What term is used to describe the majority of
these compounds?
8. The AlI3 molecule has an incomplete octet around Al. Draw three resonance structures of the molecule in
which the octet rule is satisfied for both the Aluminum and Iodine atoms. Show formal charges.
9. Of the noble gases, only krypton, xenon, and radon are known to form a few compounds with O and/or F.
Write Lewis structures for the following molecules: (a) XeF2, (b) XeF4, (c) XeF6, (d) XeOF4,
(e) XeO2F2. In each case Xe is the central atom.
10. Write Lewis structures for SeF4 and SeF6. Is the octet rule satisfied for Se?
11. Write the Lewis formula for gaseous beryllium chloride, a covalent compound.
12. Write the Lewis formula for the following: SF4, ClO4-, SO32-, PCl3. Include representation of the net
dipole for each.
13. None of the following is known to exist as covalent compounds. What is wrong with each one?
Formal Charge & Resonance
14. Write the full formula for formal charge as it appears in the text (do not use the shorthand version).
15. Why are only half of the bonded pair electrons subtracted from the number of valence electrons?
16. Why are all of the electrons in lone pairs subtracted from the number of valence electrons?
17. How is this equation used to then determine which resonant structure is the most stable structure?
18. What is meant by the term “preferred structure”? Why would this structure be preferred?
19. If the formal charge for two resonant structures is the same, how is the preferred structure identified?
Why?
20. Draw three resonance structures for the molecule nitrous oxide, N2O. Indicate formal charges. Do the
same for SCN-, ClO3-, and OCN-. Identify the preferred structure for each.
21. Draw the Lewis electron-dot structures for CO32-, CO2, and CO, including resonance structures where
appropriate. Which of the three species has the shortest C-O bond length? Explain the reason for your
answer.
22. Read carefully the discussion of the dissociation of oxygen and ozone. Provide a summary of this
material including the relationship between bond strength, Planck’s Law, and why the oxone layer is
essential for life as we know it on Earth.
23. Write the catalytic reactions for the depletion of ozone using NOx and Chlorofluorocarbons (CFC’s).
Demonstrate using Hess’s Law how these reactions result in the reduction of ozone. Also explain, in the
context of the reactions, why NOx and CFC’s are catalysts in the reactions.
24. Explain why these reactions are important by reference to the sources of NOx and CFC’s and how
depletion of ozone will impact living systems.
Pi & Sigma Bonding
25. Describe the difference between sigma bonds and pi bonds. Include in the description the type and how
many orbitals are overlapping along with the orientation of the overlap.
26. For the following molecular structures, provide the total number of pi and sigma bonds.
a) H3C – CH3
b) H3C – CH = CH2
c) CH3 – C = C – CH2OH
d) CH3CH = O
e) CH3COOH
f) CO
g) CO2
h) CN-
14.2 Hybridization
1. Out of the following choices BF3, NH3, NaCl, CF4, and PH3,
a) Which ones have a central atom that forms sp2 hybrid orbitals?
b) Which ones have a central atom that forms sp3 hybrid orbitals?
c) Which ones have a central atom that forms sp hybrid orbitals?
2. Which form of hybridization can form molecules with shapes that are either trigonal pyramidal or tetrahedral?
(a)sp, (b) sp2, (c) sp3
3. In which of the following species does the central atom NOT form sp2 hybrid orbitals:
(a) SO2 (b) BF3, (c) NO3-, (d) SO3, (e) PCl3
4. What are the hybrid orbitals of the carbon atoms in the following molecules?
a) H3C – CH3
b) H3C – CH = CH2
c) CH3 – C = C – CH2OH
d) CH3CH = O
e) CH3COOH
f) CO
g) CO2
h) CN-