1. Molecules & Covalent Bonding
Chapter 5 & 15
Molecules & Covalent
Bonding

2. The Covalent Bond
When 2 hydrogen atoms combine to make H2, the atoms both share their 1 electron so each has a noble gas configuration like helium.

3. Electronegative Difference
The Chemical Bond
Why do atoms sometimes take and sometimes share to be come stable?
Electronegative Difference

5. Bond Type
Comparing the elctronegativities will help decide who wins, if anyone, in the tug of war for electrons!

6. Bond Types
Chlorine is strong enough to take the electron from sodium, but the hydrogen have equal electronegativities and therefore end up sharing the electrons.
Cl-
Na+ H

7. Determining the Bond Type
Electronegativity difference
Bond Type
Example
0.0 – 0.4
Nonpolar Covalent
H – H (0)
>0.4 – 2.0
Polar Covalent
H – Cl (0.9)
 2.0
Ionic
Na+Cl- (2.1)
Molecules are usually in this range

8. Bond Type
Let’s try a couple. What is the bond type for NaCl? Na = Cl = = 2.23 It is ionic (big surprise)!
.93
3.16

9. Bond Type What is the bond type in a water molecule, H2O?
Water has two H-O bonds!
H = O =
3.44 – 2.20 = 1.24
The H-O bonds are Polar Covalent
2.20
3.44

10. Covalent Vocabulary Covalent bond – sharing of a pair of electrons
Molecule – electrically neutral unit of a substance that retains the properties of the substance and is held together by two or more covalently bonded atoms.

11. Covalent Vocabulary Nonpolar covalent - equal sharing of electrons.
Polar covalent - unequal sharing of electrons.

12. Cl Cl Cl Covalent Vocabulary
Single covalent bond - sharing of 1 pair of e- between atoms.
Unshared pair - pair of electrons that is not being shared. Cl Cl Cl
Lewis Dot Structures

13. O N Covalent Vocabulary
Double covalent bond - sharing of 2 pairs of e- between atoms.
Triple covalent bond - sharing of 3 pairs of e- between atoms. O N

14. Diatomic Molecules Some elements commonly exist as molecules
Contain two atoms of the same element
H, O, N, F, Cl, Br, I

15. How do you draw molecules?
Space filling-model of HCl
H-Cl
Lewis dot of HCl
For the simple molecules we’ll be using, the key to drawing is to find the number of bonds.
For octet structures, an easy way to find the number of bonds is to use the N-A=S method !

16. Drawing Covalent Structures
First, find the number of bonds (shared e- pairs) using the following formula:
N – A = S
N = Total number of e- needed for all to have a full octet (Remember while most things need 8, H and He need 2)
A = Total number of e- available
(For polyatomic ions, it equals A - charge)
S = Shared electrons
# bonds =

17. Drawing Covalent Structures
Draw molecule without bonds.
Draw central atom(s). Usually the element there is only one of, otherwise use elements from the following families as the central atom
C  N → B  O (never H or halogen)
Draw outer atoms spaced around central atom(s). O O S
For SO3 O

18. Drawing Covalent Structures
Draw a single bond between all atoms.
If there are left over ones, then you will need to have multiple covalent bonds!
Multiple bonds will go between carbons (if there are 2) or between central atom and B, N, or O groups(usually O). S O
SO3 has 4 bonds!

19. Drawing Covalent Structures
Fill in dots (draw them in pairs) so all atoms have full octets.
If it’s a polyatomic ion, put the whole thing in brackets and write the charge in the upper right hand corner.
H-Cl +

20. Let’s do a few together:
8(1)+2(3)
5(1)+1(3)
14-8
6/2
NH3 14 8 6 3 O2
HCN
C2H4

21. Polyatomic Ions
Polyatomic ions are covalently bonded atoms that have gained or lost electrons to become stable.
Because they are covalent, we can use N-A=S to draw these structures.
The only difference is that the gain or loss of electrons has effected how we calculate A
Find A like before, then subtract the charge to find the actual value of A.
Draw NO2-
N =
A =
S =
8(3) = 24
5(1) + 6(2) = 17
- (-1) = 18
24 – 18 = 6

22. Resonance
When one Lewis structure does not correctly represent the molecule you have resonance.
This occurs when there is an equal choice for placing multiple bonds.
Draw all structures connected by double arrows. N O N O

23. Non-Octet Structures
When you have a structure that doesn’t follow the octet rule, you can’t use N-A=S!!!! Quickly decide on the number of bonds, and draw the structure!

24. Non-Octet Structures Non-Octet Structures Less than an octet
Beryllium
Stable with only 4 shared electrons (2 bonds)
Boron
Stable with only 6 shared electrons (3 bonds)
Expanded octets (more than an octet)
Central atom has more than 4 bonds
PCl5 and SF6
H—Be—H

25. Non-Octet Structures
BeCl2
PCl5
BH3

26. VSEPR Theory
Electron Dot Structures fail to reflect the 3 dimensional shapes of molecules
Ex: Draw Methane
What does it really look like?
The VSEPR Theory states that because electron pair repel, molecules adjust their shapes so that the valence electron pairs are as far apart as possible

28. Molecule Shape
Valence Shell Electron Pair Repulsion (VSEPR) Theory - electrostatic repulsion between valence electron pairs surrounding an atom cause these pairs to be oriented as far apart as possible.
Two atom molecules are always linear.

29. Molecule Shape A = number central atom B = number outside atoms
E = unshared
electron pairs
on the central
atom

30. Just like bonds can be polar or nonpolar, so can molecules.
Molecule Polarity
Just like bonds can be polar or nonpolar, so can molecules.
Molecular polarity starts with the bond type.
You can’t have a polar molecule unless you have polar bonds!
The problem is that polar bonds don’t always give you a polar molecule.
Perfect symmetry (symmetrical with all outside atoms the same element) can neutralize the polarity to give a nonpolar molecule.

31. Molecule Polarity
nonpolar
nonpolar
polar

32. Intermolecular Force
Intermolecular Force - forces of attraction between molecules.
Dipole-dipole: strong force of attraction between polar molecules.
Hydrogen bonding: very strong dipole-dipole caused between polar molecules containing hydrogen bonded to N, O, or F with an unshared pair of electrons.
Examples - HF, NH3, and H2O
Is it strong “H-NOF” (enough)
London dispersion forces: weak force of attraction between nonpolar molecules caused by momentary dipoles of constantly moving electrons.

33. Hydrogen Bonding
Hydrogen bonds: Attractive forces in which hydrogen that is covalently bonded to a very electronegative atom is also bonded to an unshared electron pair of an electronegative atom in the same molecule or nearby molecule

34. Hydrogen Bonding Strongest of the intermolecular forces.
Extremely important in determining the properties of water
High Surface Tension
Ice less dense than water
High Heat Capacity

35. Properties of Molecular Substances
A great variety of physical properties occurs among covalent compounds. This is because of the intermolecular attractions
Example: Most melting & boiling points are low compared to ionic compounds. Diamond does not melt, but vaporized at 3500o C and above.
Diamond is an example of a network solid in which all of the atoms are covalently bonded to each other

36. Characteristics of Ionic vs. Covalent Compounds
Ionic compounds
Covalent Compounds
Representative unit
Formula Unit
Molecule
Bond Formation
Electrons are transferred
Electrons are shared
Type of elements
Metal & Nonmetal
Nonmetals
Physical State
Solid
Solid, liquid, or gas
Melting Point
High (usually above 300o C)
Low (usually below 300oC)
Solubility in water
Usually high
Varies high to low
Electrical conductivity of aqueous solution
Good conductor
Poor to nonconducting