1. Chemical Bonding

2. A chemical bond is a mutual attraction between nuclei and valence electrons of different atoms that binds atoms together.
As single particles, atoms have a high potential NRG. Nature favors a reduction in this energy-by bonding, they reduce their potential NRG.
Why do atoms form bonds?

3. The Octet Rule
Remember….for atoms to be most stable (i.e. have the lowest possible potential NRG) they must have their outermost shell full of electrons.
The Octet Rule states:
Atoms require 8 valence electrons to be stable
….There are exceptions to the octet rule. For example: H and He only need 2 electrons. Boron needs six and Beryllium only needs four.

4. Electron-Dot Notation
Used to show the number of valence electrons an atom

5. Periodic Table w/ Electron Dot Notations

6. Types of Bonds
Covalent
Ionic
Metallic

7. Covalent Bonding
Results from the sharing of electron pairs between two atoms.
Groups of atoms held together by covalent bonds are called molecules.

8. Properties of Covalent Bonds
Low melting and boiling points
Form between two nonmetals
Form by the sharing of electrons
Often liquids or gases at room temp.
Will either be nonconductors or poor conductors of electricity

9. Bonds: Single bonds- one pair of electrons shared Ex. Cl and Cl
Double bonds- sharing of two pairs of electrons
Ex. Ethene (C2H4)
Triple bonds- sharing of three pairs of electrons
Ex. N and N

10. Lewis Structures
Shows electrons being shared in covalent bonds as well as any unshared electrons (also called a lone pair) that are not involved in bonding.
Examples:
H and H
N and N
C and O

11. Covalent Bonds can be either:
Polar
Results from unequal sharing of electrons
Nonpolar
Results from equal sharing of electrons

12. Differences in electronegativity determine if a covalent bond will be polar or nonpolar-the bigger the difference-the more likely it will be polar

13. Predicting Bond Type Electronegativity Diff. Bond Type
Nonpolar Covalent
Polar Covalent
Ionic Bond

14. Predict the Bond Type: Bonding between sulfur and Hydrogen Cesium
Chlorine

15. Intermolecular Forces
Three types of intermolecular forces:
Dipole-Dipole (only present in polar molecules)
Hydrogen bonding (only present in polar molecules)
London dispersion forces (present in both polar and nonpolar)
These are forces that occur between
COVALENTY BONDED ELEMENTS

16. Dipole-Dipole
The strongest intermolecular forces occur between polar molecules because they act as tiny dipoles (created by equal but opposite charges separated by a short distance)
These oppositely charged areas are attracted to one another.

17. Hydrogen Bonding
Intermolecular force in which a H atom is bonded to a highly electronegative atom
This is a particularly strong type of dipole-dipole force

18. London Dispersion Forces
Any atom or molecule will experience weak intermolecular forces due to the continuous motion of electrons.
This movement creates uneven electron distribution which creates a temporary dipole.
London dispersion force- an attraction resulting from uneven electron distribution and the creation of temporary dipoles
These are very weak forces.

19. Remember….. In covalent bonds: Polar bonds are stronger than nonpolar
In reference to intermolecular forces:
H-bonding is strongest, then dipole-dipole,
then dispersion forces

20. Boiling Points and Intermolecular Forces
Boiling points can give information about bond strength.
Stronger bonds have higher boiling points and weaker bonds have lower boiling points.
Examples:
H2 is nonpolar covalent and boils at -253 C
NH3 is polar covalent and boils at -88 C
NaCl is ionic and boils at 1413 C

21. Molecular Geometry
VSEPR theory (valence-shell, electron pair repulsion)
Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

22. Molecules w/no unshared pairs that form linear molecules.
Formula Lewis Dot
BeF2 H2
HCl
Will have a sample formula of AB2.

23. Molecules w/ no unshared pairs that form trigonal planar molecules.
Formula Lewis Dot
BCl3
Will have a sample formula of AB3.

24. Molecules w/no unshared pairs that form tetrahedral molecules.
Formula Lewis Dot
CH4
Will have a sample formula of AB4.

25. Molecules w/unshared pairs of electrons.
VSEPR states that unshared pairs (lone pairs) of electrons will occupy a space around the atom just as bonding pairs do.
Example: water (H2O)
The unshared pair of electrons on the oxygen atom results in a “bent” molecule.
Has the formula AB2E2

26. More examples:
Formula Lewis Dot
NH3
Has the formula AB3E

27. PHET Simulations

28. sp3 Hybridization
Used to explain how the orbitals of an atom become rearranged when the atom forms covalent bonds
Hybridization is the mixture of two or more atomic orbitals of similar energies on the same atom to produce hybrid orbits of equal energy
Let’s look check out methane, CH4

29. Geometry of Hybrid Orbitals
Hybridization
# of Hybrid Orbitals
Geometry
s, p sp 2
Linear
s,p,p
sp2 3
Trigonal planar
s,p,p,p
sp3 4
tetrahedral

31. Ionic Bonding
Results from the electrical attraction between a cation and an anion
Bonds are formed when electrons are transferred.
By transferring electrons-their outermost shells become filled.

32. Two Types of Ions Monatomic -ions that form from one atom Ex. Na+1
Cl-1
Mg+2
S-2
Polyatomic
Ions that form from more than one atom
Ex. SO4-2 (sulfate)
NH4+1
(ammonium)

33. Oxidation Numbers Groups 1-18
Group Oxidation #
or -4
18 0
As electrons are gained or lost, the atom takes on either a negative or positive charge which is called an oxidation number.

34. Characteristics of Ionic Bonds
High melting pts.
High boiling pts.
Hard and brittle
can conduct electricity in solution
Form between a metal and a nonmetal

35. Metallic Bonding
Chemical bond that results from attraction between metal atoms and the surronding “sea” of electrons created by the overlapping of orbitals where electrons can move freely
Characteristics:
Good conductors of heat and electricity
Absorb a wide range of light frequencies
Malleable
Ductile
Form between two metals