1. Ch. 8 Covalent Bonding
College Chemistry

2. Review Remember an ionic bond?
Holding together of cations and anions by electrostatic forces, or the “giving and taking” of electrons
So if an ionic bond is the giving and taking of electrons, what would a covalent bond be?
Molecule held together by sharing electrons
Another way to achieve stability and have both electrons achieve a full octet

3. Rules
1. Decide which element goes in the “middle” (usually least electronegative)
What are the trends again?
Increase from bottom left to upper right
2. Draw valence electrons around each atom
Count how many valence electrons each element has (remember to add/subtract for ions!)
3. Draw covalent bonds between unpaired electrons.
4. Remember: C, N, O, S can double bond (a select few others can rarely)
C, N, and O NEVER disobey the octet rule
5. Disobey the octet rule only if you have to!
Place extra electrons on the central atom – it will be the one that disobeys the octet rule

4. Electron Dot Diagrams Remember these?
What would the electron dot diagram for C be?
For Cl?
Why were electron dot diagrams useful for ionic compounds?
Show which electrons are accepted/ donated and shows bonding

5. Electron Dot Diagrams
So why do you think electron dot diagrams are useful for covalent (molecular) compounds?
To show which electrons are shared!
Let’s try Hydrogen!
Electron dot diagram:
Molecular compound:
This shared pair of
electrons is called a
pair or shared electrons

6. Electron Dot Diagrams Electron dot diagram of Fluorine:
Molecular Compound:
These valence electrons not shared (and not in bonding are referred to as…
Unshared electrons (or pair)

7. Let’s try another
How would we do water?

8. Let’s try another
How many shared pairs: 2
Unshared:

9. Let’s try another
How would we do methane?

10. Let’s try another
Shared pairs: 4
Unshared pairs:
None!

11. Formal Charge
Formal charge – electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to an atom in the Lewis structure
In a lone atom, the # of electrons associated with the atom is simply the # of valence electrons (ex: Na  1)
In a molecule, we have bonding electrons and non-bonding electrons, to worry about…this can give us a “formal” charge, a charge you wouldn’t normally expect that atom to have
This also helps us define resonance

12. Formal Charge How to figure out Formal charge:
1. (valence electrons) – (electrons assigned to atom) = formal charge
For molecules  formal charges must equal zero
For cations, the formal charge must equal the overall positive charge
For anions, the formal charge must equal the overall negative charge

13. Formal Charge Example Valence: 6 6 6 e- assigned: 6 5 7
C H H H
Valance:
e- assigned:
Formal charge:

14. What about one not bonding with H?
Let’s try BF3
Formal Charge:
0 on all
Shared pairs: 3
Non shared
On B – none
On F - 9

15. Here’s a tough one Let’s try making O2
Remember: all electrons MUST be paired
What did you notice?
A double bond needed to be formed
Will a double bond be more or less stable than a single bond?

16. Here’s a tough one
Double bond – a bond that shares two pairs of electrons

17. Here’s another
Let’s try N2
A triple bond

18. Here’s one more:
CO2
Atoms don’t have to be the same type of element to participate in double or triple bond
Hydrogen will never participate. Why?
Only 1 electron!

19. Another: Let’s try CO Both electrons were shared from the O
Coordinate covalent bond - When one atoms contributes both bonding electrons

20. With Ions
Let’s try NH4+
Remember to take away one electron for the +1 charge!
Formal Charge:
N: +1
H: 0

21. With Ions
Let’s try OH-1
Remember to add one electrons for the -1 charge!
Formal Charge:
O: H: 0

22. Another… PCl5 Forced to disobey the octet!
Octet rule cannot be satisfied in molecules who have an odd number of valence electrons

23. One More BF3 (again) What do you notice?
It cannot obey the octet rule because the “middle” atom does not have enough electrons!

24. Ozone O3 Where do you put the double bond?
It doesn’t matter!
Resonance structure – occurs when it is possible to draw 2 or more valid electron dot structures
We are moving ELECTRONS, NOT ATOMS!

25. Another example of resonance
Formal Charges:
O: N: O: 0

26. Review - electronegativity
Remember electronegativity?
Increases across, decreases down
Electronegativity is important in differentiating between covalent and ionic bonds
Covalent bonds – sharing of electrons
Ionic bonds – give/take of electrons
Let’s consider NaCl
Which atom is more electronegative?
Cl (3.0) vs. Na (0.9)
By a lot or a little?
A lot
Do you think this would be covalent or ionic? Why?
Ionic – big difference between electronegativities

27. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally
Polar covalent bonds – covalent bond between atoms, electrons shared unequally
Ionic bonds – complete giving/taking of electrons
Will the more or less electronegative atom attract the electrons better?
more electronegative atoms – slight/partial negative charge
less electronegative atom – slight/partial positive charge

28. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally
Polar covalent bonds – covalent bond between atoms, electrons shared unequally
Ionic bonds – complete giving/taking of electrons

29. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally (less than 0.4)
Polar covalent bonds – covalent bond between atoms, electrons shared unequally ( )
Ionic bonds – complete giving/taking of electrons (greater than or equal to 2.1)

30. Bond Polarity
Special “arrows” are used to show the most electronegative side of the bond OR
“delta” signs denotes the partial positive or negative charge

31. Bond Polarity Let’s try one: CO2 EN of C: EN of O: Draw structure: 2.5
3.5
Draw structure:

32. Molecular Polarity
If a polar bond is an unequal sharing of electrons, what would a polar molecule be?
One end of the molecule is more negative, one end is positive
Nonpolar molecule?
The molecule has no polarity between bonds
OR molecule has polarity between bonds, but the molecule is symmetrical!

33. Molecular Polarity Let’s look at CO2 EN of C: EN of O:
2.5
EN of O:
3.5
Does it have polar bonds?
yes
Is it symmetrical?
Polar or nonpolar molecule?
Non polar

34. Molecular Polarity Let’s look at N2 EN of N: Does it have polar bonds?
3.0
Does it have polar bonds? no
Polar or nonpolar molecule?
Non polar

35. Molecular Polarity Let’s look at H2O EN of H: EN of O:
2.1
EN of O:
3.5
Does it have polar bonds?
yes
Is it symmetrical ?
No, look at the Lewis structure otherwise you will think WRONG
Polar or nonpolar molecule?
polar

36. Molecular Polarity (click me) Polar bonds? Yes Symmetrical?
Yes – polar molecule
No – non polar molecule No
Non polar
(click me)

37. Molecular Polarity
Its’s hard to tell if a molecule is symmetrical or not.
Here’s some hints
Tetrahedral, linear, trigonal planer, trigonal bipyramidal, octahedral and square planer will always be nonpolar if the elements off the central atom are the same
Do dipoles cancel out?

38. Molecular Polarity Let’s look at BF3 EN of F: EN of B:
4.0
EN of B:
2.0
Does it have polar bonds?
yes
Is it symmetrical ?
Polar or nonpolar molecule?
Non-polar

39. Ionic vs. Covalent

40. Ionic vs. Covalent
Why are ionic compounds good conductors and covalent poor?
Hint: electricity flows through ions
Ionic compounds have ions, covalent do not
Why are ionic compounds usually more soluble than covalent compounds in water?
Hint: think of electronegativities in water
Water is very polar, ionic compounds have ions, ions dissociate in water easily