1. Chapter 6.1-6.2 – Fundamentals of Chemical Bonding
CHM1111 Section 04
Instructor: Dr. Jules Carlson
Class Time: M/W/F 1:30-2:20
Monday, October 17th

2. p-block Elements - Correction
Properties of elements in the p-block vary more than in the s block.
Al, Ga, In, Sn, Tl, Pb, Bi have metallic propeties and form cations.
Al, Ga, In form 3+ cations Al(OH)3 and GaF3.
Can form anions if strongly electronegative (group 7 or 17-IUPAC halogens and group 6 or 16-IUPAC ).
Those less electronegative can form polyatomic oxoanions (eg. CO32- and NO3-), and halogens can too (eg. ClO4-).
Group 8 or 18-IUPAC atoms are noble gases and were believed to be inert. It has been found that Xe, Kr can react with halogens (eg. XeF6).

3. Bonding
A chemical bond is an attraction between atoms that allows for the formation of chemical substances that contain two or more atoms.
Atoms share electrons to different extents.
Ionic bond – electrostatic attraction between two oppositely charged ions, one atom formally has an electron(s) from the other atom.
Covalent bond – electrons are shared close to equally or equally between the two atoms.
Polar-covalent bond – electrons are shared but not equally.

4. Look at Covalent Bonds – H2
In covalent bonds, electrons are shared.
Two forces to consider:
Attraction between nuclei and electrons
Repulsion from nuclear-nuclear and electron-electron interactions
Consider the hydrogen atom (bond length = 74 pm):
Potential Well
At lengths > 300 pm, the total interaction energy is almost zero because of the large separation. At closer distances, the attraction between the electron of one atom and the nucleus of the other atom increases

5. Look at Covalent Bonds – F2
H2 is simple, only 1 electron to consider for each H, other atoms have many electrons.
F is 1s22s22p5, 2p electrons feel the strongest attraction, arrange in largest overlap.

6. Electron Sharing Can Be Unequal
In H2, each H nucleus has a charge of +1, both nuclei attract electrons equally, and both electrons have same orbital so offer equal repulsion.
Electron sharing is truly equal only between 2 identical atoms.
In HF, bonding electrons are attracted to the nucleus of the hydrogen and fluorine atoms – more strongly to F.
Degree of attraction depends upon size of nuclear charge and amount of screening.
The larger attraction of F results in partial +ve and –ve charges
Draw on electrons given by electronegativity.

7. Electronegativity
Electronegativity (χ) – the characteristic ability to attract bonding electrons.
Defined, and values calculated by Linus Pauling in 1932, calculated from dissociation energies relative to H (2.1)
Depend upon nuclear charge and electron screening – as do dissociation energies.
Atoms with higher electronegativity often have lower ionization energies and higher electron affinities, but the three processes are all different so are distinct properties.
Modern x-ray techniques can measure electron densities in bonds to confirm.

8. Electronegativity

9. Electronegativity Differences
The difference in electronegativity (Δχ) can tell you where a bond lies in the continuum of bond polarities.
F2, HF, and CsF are examples across this distribution.
F2 (Δχ = ) = Covalent bond (Δχ < 0.3)
HF (Δχ = ) = Polar-covalent bond (Δχ < 2.0)
CsF (Δχ = ) = Ionic bond (Δχ > 2.0)
Also, bond dissociation energy is larger for pairs of atoms with larger electronegativity differences as interaction energy is more negative (stable) for bonded molecule.

10. Drawing Problem
Draw depictions of Cl2, HCl, NaF, and CH2O showing electron density and orbital overlap.

11. I clicker Question Which of the following statements are true:
Electrons are shared more equally in a CH4 molecule than in a CF2 molecule.
The bond length and bond dissociation energy are larger for HI than for HF.
The C=O bonds in CO2 are polar-covalent bonds
Both (a) and (c) are true
All statements are true

12. Lewis Structures
Lewis structure: Drawing of a molecule showing how atoms are bonded together and reveals the distribution of bonding and non-bonding valence electrons.

13. Lewis Structure Conventions
Each atom is represented by its elemental symbol.
Only valence electrons are shown.
A line joining two elemental symbols represents one pair of electrons shared in a bond between two atoms (can have double, triple bonds too).
Dots placed next to an elemental symbol represent non-bonding electrons.

14. Bonding
Outer atoms only bond to one atom, and inner atoms bond to more than one atom.
Hydrogen atoms are always outer atoms
More electronegative atoms are outer atoms
Order of atom appearance in formula is often the bonding pattern (e.g HCN is H-C-N) – not true for oxoacids.

15. Rules for Building Lewis Structures
Count the valence electrons.
Assemble the bonding framework, placing 2 electrons per bond.
Place three non-bonding pairs of electrons on each outer atom except H.
Assign the remaining electrons to inner atoms to complete their octets.
Optimize electron configurations of the inner atoms.
Identify equivalent or near-equivalent Lewis structures.

16. First Example Dichloromethane CH2Cl2
Count the valence electrons. If anion, add one electron for each negative charge. If cation, remove one electron for each positive charge.
C – group 14: s2p2, H- group 1: s1
Cl – group 17: s2p5
So 7 for Cl x 2 Cl + 4 for C + 1 for H x 2 H = 20 e-
Assemble the bonding framework:
Remember H’s on outside and electronegative atoms.

17. First Example Continued
Place 3 non-bonding pairs on each outer atom except H.
Non-bonding pairs are also called lone pairs.
Assign the remaining valence electrons to inner atoms

18. Example Where Optimizing is Needed
Work through as we did with dichloromethane to step 4 for Formaldehyde H2CO.
Look at number of bonds to Carbon, how can we adjust this?