Chemical Bonding

I. Types of Chemical Bonds
A. Bond energy is the energy required to break the bond. A stable bond is achieved by reaching the lowest possible energy state. Ionic bonding is the electrostatic attraction of oppositely charged ions. The energy of interaction is defined as E=2.31e-19 J nm*(Q1Q2/r). This is a modification of Coulumb's law.
B. Covalent bonding is where the atoms share the electrons. Polar covalent bonding is where there is an unequal sharing of electrons. There is a partial charge on each atom of this polar covalent bond.

II. Electronegativity
A. The ability of an atom to attract shared electrons to itself. Electronegativity increases from left to right and bottom to top on the periodic table.

III. Ions: Electron Configuration and Sizes
A. Bonding involves atoms attaining the number of valence electrons of the nearest noble gas. To predict the formula of the ionic compound, simply balance the compound so that it is neutral.
B. The ration of protons to electrons greatly influences the size of the ions. The greater number of protons the smaller the ion will be.

IV. Formation of Binary Ionic Compounds
A. The strength of an ionic bond is indicated by the lattice energy or the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. Lattice energy = k*(Q1Q2/r). The greater the charges of the ions the greater the lattice energy will be.

V. The Covalent Chemical Bond
A. Bonds result from the tendency of a system to seek its lowest possible energy.

VI. The Localized Electron Bonding Model
A. A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Pairs that are localized on the atom are called lone pairs while pairs that are in between atoms are called bonding pairs.

VII. Lewis Structures
A. Any atom really wants to achieve a noble gas configuration. Every atom wants to have eight electrons around it as described by the octet rule, except for the Hydrogen atom that only wants two.
B. Exceptions: Boron and Beryllium is electron-deficient while Sulfur, Xenon and other third-row elements can be electron rich.
C. Resonance involves several acceptable Lewis structures for a molecule. The variation in these structures is usually the double bond in different places. The actual structure is actually a mixture of all the resonance structures.
D. Formal charge is used to determine if one structure is more stable than the other is. Formal charge is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule (=number of valence electrons on free atom-(number of lone pair electrons + number of bonds). A molecule with an overall neutral charge will be favored.

VIII. Molecular Structure: The VSEPR Model
A. The basis of this model is the tendency of electrons to distance themselves from each other due to repulsion. It is important to remember that lone pairs of electrons determine the structure but naming the arrangement is based on the atoms. For example NH3 has a tetrahedral shape with all of its electron pairs but the shape of the atoms is trigonal pyramidal.

IX. Hybridization and the LE Model
A. Central atoms hybridize (combine, mix) the orbitals of their electrons to bond with other atoms.
B. sp hybridization involves the combination of a single s and single p orbital. This produces a linear bond.
C. sp2 hybridization involves the combination of a single s and two p orbitals. This produces a trigonal pyramidal arrangement.
D. sp3 hybridization involves the combination of a single s and three p orbitals. This produces a tetrahedral arrangement.
E. dsp3 hybridization involves the combination of a singe s, three p, and one d orbital. This produces a trigonal bipyramidal arrangement.
F. d2sp3 hybridization involves the combination of a single s, three p, and two d orbitals. This produces a octehedral arrangement.
G. A double bond always consists of a s bond (hybridized orbitals) and a p bond (un-hybridized p orbitals).