1. Bonding: General Concepts
Chapter 8
Bonding: General Concepts

2. Chemical Bonds Bond Energy - energy required to break a bond
Types: ionic and covalent

3. Ionic Bonding
Electrostatic attractions between tightly packed, oppositely charged ions
Between metals (cations) and nonmetals (anions)…can be polyatomic
High melting/boiling points
Ex: NaCl

4. Coulomb’s Law E = (2.31 X 10-19 J * nm)(Q1Q2/r)
Calculates the energy between two ions
E -> Energy in Joules
R -> distance between ion centers in nm
Q1 and Q2 are ions’ charges
When E is negative, the pair of ions has a lower energy than the ions separated (ions are attracted)
When E is positive, the repulsive energy is greater (occurs with two like-charged ions)
In nature, a system wants the lowest possible energy

5. Covalent Bonding Molecules where electrons are SHARED by nuclei
Equal sharing of electrons occurs between the two atoms
Nonequal sharing = polar covalent

6. Identical Atoms Diatomic atoms… Systems of energy favor LOWER energy
If a system can lower energy by forming bonds, it will happen

7. Bond Lengths
Distance at which the system has minimum energy

8. Polar Covalent Bond
Covalent bond where one atom pulls more on an electron than the other
Ex: H-F (fluorine slightly negative)
+ and - are used to represent slight charges
Water is another example
Due to electronegativity

9. Electronegativity
(electron affinity)…ability to attract shared electrons
Higher electronegative atom will have the slight negative charge
Higher difference in electronegativities, more likely to be an ionic bond

10. Bond Polarity
Dipolar/dipole moment occurs when a molecule has a slight positive and slight negative charge.
Arrow points towards the slight negative charge with a plus sign on the other end of the arrow
Electrostatic potential diagram also used (red = electron-rich, blue = electron-poor)

11. Polar Bond without Dipole Moment
Polar bonds where slight charges cancel each other out
No lone electrons

12. Ion Electron Configuration and Sizes
Atoms of stable compounds (ionic and covalent), have noble gas electron configurations
Covalent: electrons are shared so that the valence electron configurations of both nonmetals attain noble gas e- configurations
Ionic: Nonmetal achieves e- configuration of the next noble gas atom and the metal’s valance electrons are emptied so both ions achieve noble gas e- configurations

13. States of Ionic Compounds
Usually “ionic compound” means a solid state where ions are close to one another/interacting, minimizing the - - and + + repulsions and maximizing the attractions

14. Trend in Ions’ Sizes
Positive ion = loss of e- (cation is smaller than original atom)
Negative ion = gain of e- (anion is larger than neutral atom)
Ion size increases going down a group
Horizontally is more complicated (metals vs. nonmetals)…..

15. Isoelectronic Ions
Ions that have the same electron configuration (b/c have the same number of electrons)
O2-, F-, Na+, Mg2+, Al3+
Same number of electrons, but different numbers of protons
From O2- to Al3+ attraction to nucleus (increase in protons) increases so smaller ions

16. Bonds Single Bonds: one pair of electrons shared
Double bond: two pairs of electrons shared
Triple bond: three pairs of electrons shared
More bonds = shorter bond length
More bonds = stronger bonds
Bond energies - amount of energy required to break bond

17. Breaking Vs. Forming Bonds
To break bonds, energy must be added (required, endothermic, energy is positive)
To form bonds, energy must be removed (given off, exothermic, energy is negative)
Change in enthalpy: DH = sum of the energies required to break old bonds (positive signs) plus sum of energies released in formation of new bonds (negative signs)

18. Bond Energies
Energy stored in bonds can be used to determine energy accompanying various chemical reactions.
Bond dissociation energies = energy required to break bonds (positive kJ/mol, endothermic)… reverse for forming bonds. Table 8.4 in book
DH = all bonds broken – all bonds formed
Must use a balanced equation!

19. CH4(g) + 4F2(g)  CF4(g) + 4HF(g)
EXAMPLE
Use data 8.4 in your book to determine DH for the following reaction.
CH4(g) + 4F2(g)  CF4(g) + 4HF(g)
Answer: kJ

20. Localized Electron Bonding Model…Highlights
Defined: A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bonded atoms
Localized to one of the atoms (LONE PAIR)
Localized to the space between atoms (BONDING PAIR)

21. LE Model Parts
1. Description of the valence electron arrangement in the molecule using Lewis structures
2. Prediction of the geometry of the molecule using the valence shell electron-pair repulsion (VSEPR) model
3. Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs

22. LE Model Part 1: Lewis Structures
All atoms want to have noble gas electron configurations
Only valence electrons are included
Determine total valence electrons
Determine layout
Determine bonds
Place remaining valence electrons

23. Lewis Dot Structure Examples
Ammonia: NH3
Water: H2O
Acetylene: C2H2
Carbon Tetrachloride: CCl4
Dihydrogen Selenide: H2Se

24. Exceptions to Octet Rule
If valence electron total is odd, the octet rule doesn’t work
Some atoms do not require all 8 valence electrons (or have more than 8)
These molecules can exist in stable form
Boron: forms compounds where boron has less than 8 electrons (ex: BF3)
More than 8 only happens with elements in period 3 and beyond ex: SF6
See pg. 371 purple box for rules if needed

25. Exception Examples Draw the Lewis dot structure for the following:
ClF3
XeO3
RnCl2
BeCl2
ICl4-

26. Resonance Molecules with more than one possible electron dot structure
Do not switch back and forth
Molecules exist as a mixture (hybrid) of the resonance forms
Use double headed arrow to signify
Example:

27. Resonance Example
Draw the Lewis dot (Localized Electron Model) structure(s) of CO32-

28. # valence e- assigned = (# lone pair e-) + (½ # shared e-)
Formal Charges
Can’t use oxidation numbers because electrons are not shared evenly between atoms (electronegativities)
Atoms can be assigned formal charges using the following:
Formal Charge = (# valence e- on atom) – (# valence e- assigned to the atom in the molecule)
# valence e- assigned = (# lone pair e-) + (½ # shared e-)
Atoms want to have formal charges close to zero
Negative formal charges should be on the most electronegative atom
***ESTIMATES of charges (not exact charges)

29. Example Assign formal charges to each atom in CO2
Which is more likely?
Answer: two double bonds b/c all formal charges are zero
Draw all resonance structures and show the most stable one for SCN-
Answer: two double bonds b/c N more electronegative than S

30. LE Model Part 2: VSEPR Model
Molecular structure shows 3-D arrangement of molecules
Based on minimizing electron-pair repulsions (bonding and nonbonding electrons will be placed as far apart as possible)

31. No Unshared Pairs
Linear (180°)
BeCl2
Trigonal Planar (120°)
BCl3

32. No Unshared Pairs
Tetrahedral (109.5°)
CClF3
Square Planar (90°)

33. With Unshared Pairs
Bent (104.5°)
H2O
Pyramidal (107°)
NH3

34. More to Consider Tigonal bipyramidal See-saw T-shaped Octahedral 
Deal with atoms that do not obey the octet rule

35. Examples
H2S
CO2
PCl3
OCS
H2CO
CH4
N2H4