1. Chemical Bonding Chapter 7 & 8

2. Chemical Bonding
Why do bonds form??? To complete their valence shell  “to satisfy their octet.”
Remember the Octet Rule: when elements bond they transfer/share electrons so both have a noble gas electron configuration.
Bonding only involves the valence electrons
(before we can discuss bond types we have to visualize electrons)

3. Lewis Dot Diagram Structures
Use dots to show all the valence electrons (outer electrons - same as group # except He)
For example, Arsenic – group 5
As = [Ar]4s2 3d10 4p s and p are the only outer e-

4. Lewis Dot Diagram Structures
The rules:
There are 4 sides around the symbol
No side can have more than 2 dots
When filling the sides of the element symbol, each side gets one dot before doubling up (except H and He)

5. YOU TRY! Write the Lewis Structure for
A) Xe
B) Br
C) Ba

6. Types of Chemical Bonds
Ionic bonds: attraction between positive and negative ions (e- are lost or gained)
metal + nonmetal (K2O)
Covalent bonds: sharing of electrons between two atoms
nonmetal + nonmetal (CO2)
Metallic bonds: metal atoms give up valence e- that are then free to move throughout the material
metal + metal (steel)
Goofy song 5 min

7. Bond Types
Ionic compounds conduct electricity when dissolved in water, because the dissociated ions can carry charge through the solution. Molecular compounds don't dissociate into ions and so don't conduct electricity in solution.

8. Lewis Structures…
(12 min)

9. Covalent Bonding
Bonds form when elements share electrons to end up with 8 valence electrons
Can have more than one bond
Single bond shares 1 pair of electrons
Double bond shares 2 pairs
Triple bond shares 3 pairs (very strong)

10. Lewis Structures & Covalent Bonds
Covalent Bonds – arrange elements so everyone has 8 valence electrons (except H and He want 2)
CH4
SO2
PCl3

11. Lewis Dot Diagram Rules
Hydrogen and Halogens are always terminal
LOWEST electronegativity is central molecule (often the oddball)
1) Find total # valence electrons by adding up group #’s of elements (add for negative charge, subtract for positive charge)
2) Place one pair of electrons between each pair of bonded atoms
3) Subtract from the total the number of bonds you just used
4) Place lone pairs about each terminal (except H) to satisfy octet rule. Left over pairs are assigned to central atom
5) If central atom is not yet surrounded by 4 electron pairs, convert to double bonds. BUT ONLY C, N, O, P, S!!
(HINTS: Elements like symmetry! Carbon is always central!)

12. YOU TRY! Draw Lewis Structures for: A) phosphorus trifluoride B) N2H2
C) CH2O

13. Strange bonding
Some bonds have a negative charge so an ADDITIONAL ELECTRON is added
(or a positive charge so an electron is eliminated)

14. YOU TRY Write the Lewis structures for the following compounds:
A) SO4 2-
B) ClO3 –
C) NOF

15. Exceptions to the Octet Rule
Beryllium (can share) and has a maximum of 4 valence electrons
Boron can bond with only 6 valence electrons
Fluorine will NOT double bond

16. Covalent Bonding But nothing is really equal…
In covalent bonding one element is more electronegative than another and the electrons move closer to one element.
(fluorine MOST electronegative)

17. Covalent Bonding Types of covalent bonds:
A. Nonpolar covalent: e- are shared evenly (H-H) (“NOT pushing’”)
B. Polar covalent: bonded atoms have unequal attraction for e- (H-Cl and H-O-H) (“pushing”)

18. Molecular Polarity
For a compound with only 1 bond, electronegativity is calculated
Calculate the electronegativity difference (p 403)
Below 0.50 = nonpolar
0.50 – 1.7 = polar (and one is slightly negative/positive so arrow goes from + to -)

19. Molecular Polarity For a compound Draw Lewis Dot Structures
Complete SYMMETRY (radial symmetry) = non polar
ASYMMETRICAL = POLAR

20. YOU TRY
Write the Lewis structures for the following molecules. Polar or nonpolar??
A) NF3
B) carbon monoxide
C) O3
D) NO3-

21. Ionic Bonding
Ions are formed when an atom, trying to move toward a stable electron configuration (octet rule), loses or gains electrons
Na : Na Cl : Cl 1-
Groups and their charges
1A- charge 1+ (Li) 5A- charge 3- (N)
2A- charge 2+ (Be) 6A- charge 2- (O)
3A- charge 3+ (B) 7A- charge 1- (F)
Nonmetals form anions (negative ion); metals form cations (positive ion)

22. Lewis Structures & Ionic Bonds
Ionic Bonds transfer electrons and a new charge is shown
the electronegativity difference (p 403) is greater than 1.7

23.
(Bozeman science 9 minutes)

24. YOU TRY Write the Lewis structures for the following ionic compounds:
A) MgBr2
B) NaI
C) K2S

25. Hydrogen “Bonds”
Hydrogen “bonds” are not real bonds and happen only between hydrogen and nitrogen/oxygen/fluorine
Let’s talk about water: H2O
INTRAmolecular force (covalent bond) keeps the 2 hydrogen atoms bonded to the oxygen atom (strong bond)
INTERmolecular force bonds one water molecule to a neighboring water molecule (weak bond)
(2min)