1. Bond Polarity and Electronegativity
1.1. Bond Polarity
1.1.1. Uneven distribution of electrical charge across a chemical bond between two atoms
1.1.1.1. Polar
1.1.1.1.1. Electrons are shared unequally between atoms; they have a north and south pole, similar to a magnet
1.1.1.1.2. There is one region where the negative electrons spend more time (negative) and a region where they do not (positive)
1.1.1.1.3. Partial charges on atoms
1.1.1.2. Nonpolar
1.1.1.2.1. Electrons are shared equally between atoms; there are no magnetic poles in the bond
1.1.1.2.2. No charges on atoms
1.2. Electronegativity
1.2.1. An atom's tendency to attract electrons shared in a bond, or strip another atom of its valence electrons
1.2.1.1. High electronegativity have stronger holds on their own electrons, and attract electrons belonging to other atoms
2. Bond Strength (Bond Enthalpy) and Multiple Bonds
2.1. Bond Strength
2.1.1. Measured by the energy required to break it; the energy necessary to separate the bonded atoms.
2.1.1.1. The stronger the bond, the greater the energy required to break it.
2.2. Bond Enthalpy
2.2.1. The energy required to break a specific covalent bond
2.2.1.1. Methane: C-H bond enthalpy is 438 kJ/mol
2.2.1.2. Triflouromethane: C-H bond enthalpy is 446 kJ/mol
2.3. Multiple Bonds
2.3.1. Two or more electron pairs are shared between two atoms.
2.3.1.1. Double Bonds
2.3.1.1.1. Four bonding electrons participate in the bond rather than two electrons in a single bond
2.3.1.2. Triple Bonds
2.3.1.2.1. Six bonding electrons
3. The more strongly an atom attracts the electrons in its bonds, the larger it's electronegativity. The larger the difference in electronegativity between two atoms in a bond, the more polar the bond.
3.1. Oxygen has a higher electronegativity than hydrogen, so the shared electrons in the O-H bond are pulled more towards the oxygen atom, making the O-H bond polar.
4. Triple Bonds are much stronger than double or single bonds.
5. Chemical Bonds
5.1. Metallic Bonds
5.1.1. Group fo atoms share a cloud of valence electrons
5.1.1.1. Magnesium (Mg)
5.1.1.2. Sodium (Na)
5.2. Ionic Bonds
5.2.1. Oppositely charged ions are attracted to one another
5.2.1.1. Potassium Chloride (KCL)
5.2.1.2. Calcium Floride (CaF2)
5.3. Covalent Bonds
5.3.1. Valence electrons are shared between atoms, forming electron pairs
5.3.1.1. Methane (CH4)
5.3.1.2. Carbon Monoxide (CO)
6. Lewis Symbols and the Octet Rule
6.1. Lewis Symbols
6.1.1. Display the valence elctrons in outermost energy levels
6.2. The Octet Rule
6.2.1. The preference to have a total of eight electrons in their outermost energy levels.
6.2.1.1. Nobel Gases
6.2.1.1.1. Helium (He)
6.2.1.1.2. Neon (Ne)
6.2.2. There are exceptions to the Octet Rule.
6.2.2.1. Too few valence electrons
6.2.2.1.1. Borane (BH3)
6.2.2.2. Too many valence electrons
6.2.2.2.1. Sulfur Hexaflouride (SF6)
7. Lewis Structures and Resonance Structures
7.1. Lewis Structures
7.1.1. Representation of the valence shell electrons
7.1.1.1. Electrons are shown as dots.
7.1.1.2. Bonding electrons are shown as lines between the two atoms.
7.2. Resonance Structures
7.2.1. Group of two plus Lewis Structures. These are typically used when a single Lewis structure cannot fully explain the bonding because of partial charges in it
7.2.1.1. Ozone (O3)
8. Molecular Shapes and VSEPR Model
8.1. 3-dimensional arrangement of atoms in space; used to predict how the molecule will form
8.2. VSEPR (Valence Shell Electron Pair Repulsion): Electrons prefer to be as far as possible from one another
8.2.1. Linear
8.2.1.1. 2 outside atoms bond to a cental atom
8.2.1.1.1. Beryllium Hydride (BeH2)
8.2.2. Trigonal Pyramidal
8.2.2.1. 3 outside atoms bonded to a central atom, electrons forming a triangle.
8.2.2.1.1. Boron Trichloride (BCl3)
8.2.2.1.2. Ammonia (NH3)
8.2.3. Tetrahedral
8.2.3.1. 4 bonds on a central atom, has bond angles of exactly 109.5 degrees
8.2.3.1.1. Methane (CH4)
8.2.4. Trigonal Bipyramidal
8.2.4.1. 5 atoms around a central atom, three in a pane and two on oppostie sides of the molecule
8.2.4.1.1. Phosphorous Pentachloride (PCl5)
8.2.5. Octahedral
8.2.5.1. 6 atoms around the central atom, with bond angles of exactly 90 degrees
8.2.5.1.1. Sulfur Hexaflouride (SF6)
9. Orbitals
9.1. Hybrid Orbitals
9.1.1. Combination of two or more atomic orbitals; results in a different shape and energy
9.1.1.1. sp, sp2, sp3, sp3d, an sp3d2
9.1.1.1.1. sp: Beryllium Flouride (BeF2)
9.2. Molecular Orbitals
9.2.1. Combination of two or more different atoms; spatial distribution of electrons in a molecule associated with a particular orital energy
9.2.1.1. Bonding
9.2.1.2. Antibonding