1. CHEMICAL BONDING

2. ***Occurs when atoms of elements combine together to form compounds.*****

3. Formation of compounds
involve adjustments in the position of one or more valence electrons.
PE is lower in bonded atoms.
Attractive force that develops is called "chemical bond“
process is called "bonding" and occurs during chem. reactions

4. Two (or three) methods Ionic bonding - attraction of ions
Covalent bonding - shared pairs of electrons
Metallic bonding - alloys/metals - not compounds

5. Covalent Bond
An attractive force that develops between atoms that are sharing pairs of electrons
∆EN < 1.7
Ex: Hydrogen – H2
dot diagram H• + H• → H:H
Structural formulas use a dash H - H

6. Bond energy Energy required to break a bond.
Bonds form to lower PE, so breaking bonds will increase PE.
Depends on the size of atoms and number of electrons shared.

7. Breaking Bonds is always endothermic
Energy is required.
ALWAYS.

8. Bond Length

10. Why not ionic? The difference in electronegativity is less than 1.7
Electrons are not pulled away from either atom.
They are shared.

11. Fluorine – F2

12. Ammonia -NH3

13. Special Types of Covalent Bonds
Multiple bonds – Occasionally atoms share more than one pair of electrons
Double bond – two shared pairs
Ex. O O=O or

14. Triple bonds - two atoms share three pairs of electrons. Ex. N2
N N or

15. Coordinate Covalent Bonds
Occurs when one atom donates both of the shared electrons in a pair
Is most commonly seen in polyatomic ions. Ex. NO3-1

16. How to write Lewis Structures
Set out the atoms – think symmetry.
Count all the valence e-
Insert single bonds first, then fill rest.
Usually, the e- are paired.
each nonmetal atom requires an octet.
H only requires 2 e-.
Multiple bonds may be needed.
Readily formed by C, N, O, S, and P.

17. Exceptions to the octet rule

18. Resonance – more than one Lewis diagram is possible.+ - - + •• •• •• •• •• •• O O O •• •• O O O
Frequently seen in polyatomic ions •• •• •• ••

19. Expanded valence – a few elements in the third and fourth period have the ability to have more than 4 pairs of electrons around the central atom.
Ex. PCl5 or SF6 S F •• •• •• Cl P Cl •• •• P •• •• Cl Cl •• •• •• •• ••

20. Polarity
An unequal sharing of electrons due to difference in electronegativity.
Polar bond – Any bond with ∆EN
Polar molecule – Has a positive end and a negative end. Called dipoles.
Occurs if there are polar bonds and the molecule is asymmetrical.
Causes intermolecular attraction to increase.

21. An ionic bond -(electrovalent)
Definition - An attraction that forms between oppositely charged ions
 + ions ions → neutral compound
∆EN > 1.7

22. An ionic bond -(electrovalent)
c. dot diagrams of ions 1) formation
A + IE → positive ion (A+) + e-
B + e- → negative ion (B-) + EA
2) bonding A B- → AB

23. Ionic Solids Made from + and - ions Compound is neutral.
Have strong attractions in all directions
Metal-nonmetal (or polyatomic ions)
High MP and BP
Tend to be hard, solid, and brittle
Some dissolve readily
Non conductors as solids, but will conduct when molten or dissolved

24. Ionic solids are brittle
due to Crystal lattice + -

25. Ionic solids are brittle
Strong Repulsion breaks crystal apart. + - + - + - + -

26. Characteristics of Covalent Compounds
Nonmetal-nonmetal combinations
Can be gases, liquids, or solids
Low to med. MP and BP
Insulators/Nonconductors (except for acids)
Molecular (a few are crystalline)
Generally not soluble (some polar exceptions)

27. Lattice Energy (ionic)
More correct than bond energy for crystals.
Lattice Energy = k(Q1Q2 / r)
k is a constant that depends on the structure
of the crystal.
Q’s are charges.
r is internuclear distance.
Lattice energy is greater with more highly charged ions.

28. Molecular Substances
Covalently bonded substances – show more variety in phases and properties. Tend to be insulators(nonconductors).
Nonmetal-nonmetal combinations
Nonpolar – tends to be gases at room temp. – have only dispersion (Vanderwaals) forces. Have low MP and BP and high VP
Polar –(dipoles) tend to be liquids or solids at room temp. Have ↑MP and BP and↓VP
H- bonding – very strong type of polar

29. Metallic Bonding
In metals the valence electrons are not bonded to any specific atom. (delocalized)
Able to move freely over the positive centers.
Referred to as a “sea of electrons”
Moving charge allows current, malleability, ductility,etc.

30. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

31. Malleable +

32. Malleable
Electrons allow atoms to slide by. + + + + + + + + + + + +

33. Shapes of Molecules 3 or more atoms
Results from VSEPR and/or hybridization of orbitals.
Can affect properties of the molecule (polarity)

34. VALENCE SHELL ELECTRON PAIR REPULSION
VSEPR
VALENCE SHELL ELECTRON PAIR REPULSION
A theory which describes the shapes of molecules based on the idea that pairs of electrons will repel each other as much as possible.

35. H H C H H VSEPR Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
Farthest apart is at 109.5º. H H C H H

36. H C H H H 4 atoms bonded Basic shape is tetrahedral.
A pyramid with a triangular base.
Same shape for everything with 4 pairs. H
109.5º C H H H

37. N H N H H H H H 3 bonded - 1 lone pair
Still basic tetrahedral but you can’t see the electron pair.
Shape is called pyramidal. N H N H H H
<109.5º H H

38. O H O H H H 2 bonded - 2 lone pair
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is called bent. O H O H
<109.5º H H

39. H H C C O H O H 3 atoms no lone pair
The farthest you can get the electron pair apart is 120º.
Shape is flat and called trigonal planar. H
120º H C C O H O H

40. 2 atoms no lone pair
With three atoms the farthest they can get apart is 180º.
Shape called linear.
180º O C O

41. VSEPR Table Example #of BP #of LP shape bond angle BeH2 2 0 linear 180
BH trigonal
(planar)
CH tetrahedral
NH pyramidal
(unshared pairs take up more space)
H bent
See table 6-5 in text for more info

42. Polar molecules are asymmetrical and have polar bonds.
Nonpolar molecules are symmetrical and may or may not have polar bonds.

43. Intermolecular Attractions (Vanderwaals forces)
Attractive forces between molecules
Happens to covalent molecules.
Strongest - hydrogen bonds
Medium dipoles/ polar
Weakest - dispersion or London forces
( All molecules have London forces)

44. Hydrogen bonding
In some highly polar compounds, a H atom is attracted to, and forms a weak bond with, an adjacent molecule.
Only occurs in compounds where there is:
H-F (strongest)
H-O
H-N (weakest)

45. Hydrogen Bonding H O d+ d- H O d+ d-

46. Hydrogen bonding H O H O H O H O H O H O H O

47. H-bonding Gives water its very unusual properties. High MP and BP
Holds the DNA molecule together.
Provides stability and shapes for proteins, enzymes, etc.
Strongest type of intermolecular attraction.

48. London dispersion (Van der waals) Forces
Weakest type of intermolecular attraction.
Develops between nonpolar molecules due to momentary shifts in the electron positions.
Strength of attraction is directly proportional to the number of electrons
(wax is a nonpolar molecule – but really big)

49. Network Solids
Large arrays of covalently bonded crystals. Do not conduct, very high MP and BP, hard solids
Examples: diamond, graphite, SiO2 (very few – easiest to just memorize)
More on this later

50. Bond length and energy
Bond length depends on the size of the atoms/ions and the number of bonds between them
C-C is longer than C=C is longer than CC
Bond energy is related to bond length.
Longer bonds tend to have lower BE

51. Breaking Bonds is always endothermic
Energy is required.
ALWAYS.

52. Endothermic reactions

53. Forming bonds is exothermic
Energy is released.
ALWAYS

54. Exothermic reactions