Post on 26-Dec-2015
Types of Chemical Bonds
A. Ionic bond: Results from the electrostatic
attraction between positive and negative ions
B. Covalent bond: Resulting from the sharing of
electrons
A+ : B-
1. A + B
A : B
2. Chemical bonds between unlike atoms are never really ionic
or covalent, fall somewhere in between
3. How do we determine whether a bond is ionic or covalent?
a. Compare electronegativities
(ability of an atom to attract electrons)
b. Electronegativity difference
5. Nonpolar covalent:
a. share electrons equally
b. No charge
6. Polar covalent:
a. share electrons unequally
b. Have a partial positive (δ+) and a partial negative charge (δ-)
7. Examples:
Using the following chart classify bonds between the following,
which bond will be more negative?
Element EN value
Sulfur 2.5
Hydrogen 2.1
Cesium 0.7
Chlorine 2.5
a) H-S b) Ce-S c) S-Cl
C.Metallic bonding
1. happens with solids and liquids
2. metal atoms give electrons
3. electrons slide through the material (not transferred to a negative)
Covalent bonding and molecular compounds
1. Share electrons
2. Still using valance electrons
3. Share 1 pair of electrons: a single bond
4. Two shared pairs: a double bond
5. Three shared pairs: a triple bond
6. Examples:
a) H2 H +H = H:H
H-H
7. Octet rule: Chemical compounds tend to form so that each
atom, by gaining, losing, or sharing electrons, has an octet
of electrons in its highest occupied energy level
a. He is happy with 2 electrons
b. exceptions to the octet rule
Structures
1. Molecular structure: shows the types and numbers of atoms
PH3
2. Lewis structure:
a. Atomic symbols represent nuclei and inner-shell electrons
b. Dot-pairs or dashes between two atomic symbols
c. Will have shared pairs and unshared pairs
ex. Cl-Cl
3. Structural formula
a. Indicates the kind, number, arrangement, and bonds of
the atoms in a molecule
4 Step Method 1. Add total valence e-
2. Draw skeleton structure
3. Subtract bonded e- from total valence e-
4. Distribute remaining e- (in pairs) to satisfy octet rule
• Resonance
1. Refers to bonding in molecules that cannot be correctly represented
by a single Lewis structure
2. SO3
2. Lattice: The energy released when one mole of an ionic crystalline compound is formed
from gaseous ions.
Metallic Bonding
1. metallic bond: resulting from the attraction between
positive ions and surrounding mobile electrons
The VSEPR Theory
(Valence-Shell Electron Pair Repulsion Theory)
The pairs of electrons in the valence shell will repel each other until they are as far away from other electron pairs as they can be.
Molecular Atoms bonded Lone pairs Type Formula Formula
Shape to central atom of electrons of molecule example Structure
around central
atom
Linear NA NA AB HCl H – Cl
Linear 2 0 AB2 CO2 O = C = O
Bent 2 1 or 2 AB2E H2O O
H H
Trigonal 3 0 AB3 AlCl3 Cl
Planar Al
Cl Cl
Trigonal 3 1 AB3E NH3 N
Pyramidal H H H
Molecular Atoms bonded Lone pairs Type Formula Formula
Shape to central atom of electrons of molecule example Structure
around central
atom
Tetrahedral 4 0 AB4 CH4 H
C
H H
H
E: stands for unshared electron pair
***Make sure you put dots around the elements where they needed them.
B. CS2 Is it AB2 or AB2E
To move or not to move: 1. Polarity a. Nonpolar molecules
1. electrons shared equally
2. molecule doesn’t move
b. Polar molecules 1. electrons shared unequally
2. molecule would move in the direction of
strongest pull
Hybridization
1. The blending of two or more orbitals to make a new
orbital with a new shape
2. Helps explain the bonding and geometry of some Group
15 and Group 16 elements
3. Hybridization rules
How to find the (-) areas:
a. each bond type counts as 1 area of – charge
b. Each unshared pair of electrons count
c. NEVER COUNT HYDROGEN
Bond length
A. The average distance between two bonded atoms, minimum potential energy
1. Longest - single bonds
2. Middle - double bonds
3. Shortest - triple bonds
B. Bond energy
1. Energy required to break a chemical bond and form
neutral atoms
2. can be related to bond length
a. Long bond length needs low energy to break it apart
b. Short bond length will need high energy
Other properties of polarity
1. Boiling point a. Temp at which a liquid turns to a
gas
b. The greater attraction of charges, the higher the boiling point
c. Ionic - highest
NaCl + -
- +
d. Polar covalent - middle
&+ &-
&- &+
e. Nonpolar covalent - lowest
2. Intermolecular forces
a. Dipole-dipole
1. Polar molecules
2. The attraction between the + pole of one molecule
and the - pole of a different molecule
3. Has a middle boiling point
b. London dispersion force 1. Nonpolar 2. The attraction between the
temporary pole of one molecule with an adjacent molecule
3. Low boiling point
c. hydrogen bonding
1. attraction between a hydrogen atom and an unshared pair of electrons on a strongly electronegative atom (fluorine, oxygen, nitrogen
2. EN difference between hydrogen and fluorine makes them highly polar
3. highest boiling points
Melting points/freezing points
A. ionic bonds
1. Strongest attraction
2. Highest melting/freezing point
B. Polar bonds
1. Somewhat strong bonds
2. Middle values
C. Nonpolar bonds
1. Weakest attractions
2. Lowest melting/freezing points
D. Ex. Rank from lowest to highest
1) boiling point
2) freezing /melting points
H2O NaCl F2