Lewis structures for elements in the same column (group) of the periodic table are similar to one another.R in the figures are either carbon or hydrogen.Boron: 3 bonds, 0 lone pairs (exception to the octet rule).Carbanion: C - has 3 bonds, 1 lone pair, formal charge -1.Carbocation: C + has 3 bonds, no lone pairs, formal charge +1.Hydrogen: 1 bond, 0 lone pair (exception to octet rule).N +: 4 bonds, 0 lone pair, formal charge of +1.O 1+: 3 bonds, 1 lone pair, formal charge of +1.O 1-: 1 bond, 3 lone pairs, formal charge of -1.It can either be 4 single bonds or two double bonds. Carbon: 4 bonds total (meaning 4 total bonds.All electrons in a bond are shared and can be used to satisfy the octet for both atoms on either side of the bond.Exceptions include the boron column (they form 3 bonds and have a six-tet), large elements (3rd row and below such as the 10-tet P in PO 4 3- and the 12-tet S in SO 4 2-), and radicals (compounds with an odd # total electrons that result in a single, unpaired electron). Formulas are drawn in such a way that an octet is achieved on each atom.A 'lone pair' is represented by two dots. Every line represents 1 bond (2 electrons). CO 2 has 2 double bonds and 0 unbonded electron pairs - it is linear.H 2O has 2 bonds, 2 unbonded electron pairs - it is bent.If you look up the table for VSEPR # = 4 and # unbonded electron pairs = 1, then you'll find that NH 3 is trigonal pyramidal. NH 3 has a vsepr number of 4 (3 bonds to H and 1 unbonded pair).When calculating the VSEPR number, always use the electron/bond configuration about the central atom.The VSEPR number is the total number of bonds + unbonded electron pairs.

Radicals also count as an electron pair.The shapes of molecules are determined by the molecular geometry.The VSEPR theory is used to predict the geometry of molecules.Valence shell electron pair repulsion and the prediction of shapes of molecules (e.g., NH3, H2O, CO2) Hybrid orbitals are most commonly used with carbon as the center atom.Sp: a hybrid between one s with one p orbital.Sp2: a hybrid between one s with 2 p orbitals.Sp3: a hybrid between one s with 3 p orbitals.Hybrid orbitals are produced by hybridizing (mixing) electron orbitals to produce geometries that facilitate bonding.Hybrid orbitals: sp3, sp2, sp and respective geometries They make up the second bond in a double bond, and both the second and the third bond in a triple bond. They also make up the first bond of double and triple bonds. The covalent bond results when there is a sharing of electrons between two atoms, resulting in the overlap of their electron orbitals.The big difference is that G is tiny compared to k, because gravitational force is weaker compared to the much stronger electrostatic force.G is analogous to k and m is analogous to q.Coulomb's law is analogous to the universal law of gravitation:.If r halves, the electrostatic force increase by a factor of 2 2 = 4.If q 1 doubles, the electrostatic force doubles.Opposite charges attract (negative F), same charges repel (positive F).Larger charge magnitudes + charges being closer together → greater electrostatic force.The lattice energy is proportional to the electrostatic attraction between the ions.

The larger magnitude of the lattice energy, the stronger the ionic bond and the harder it is to break.

Lattice energy is the energy required to break the ionic bond.

Lattice energy measures the ionic bond strength.

Ions that form strong ionic bonds have high charge density, that is, the charge to size ratio is high.

Strong ionic bonds are promoted by high charge magnitudes (q values) that are close together (small r value).

The greater the magnitude of electrostatic potential, the stronger the ionic bond.

Frequently, the negative sign is dropped and only the magnitude of the electrostatic energy is used.

Electrostatic energy is negative because q 1 and q 2 are opposite in charge (If q 1 and q 2 are not opposite in charge, then they would repel each other, and no ionic bond would form).

Electrostatic Energy = Electrostatic potential x charge = kq 1/r x q 2 = kq 1q 2/r.

The ionic bond forms when electrons transfer completely from one atom to another, resulting in oppositely charge species that attract each other via electrostatic interaction.

The ionic bond (electrostatic forces between ions)

Calculating formal charge