Molecular Structure, Bonding, and Intermolecular Forces

Iodine (I) has seven valence electrons.
For a molecule with three iodine atoms and a -2 charge ( $I_3^{-2}$ ), there are a total of 21 valence electrons from the iodine atoms plus 2 electrons from the charge, totaling 23 electrons.
- This total number of electrons needs to be organized in the Lewis structure for the molecule.
Iodine can have an expanded octet because it is in or after the third period and thus has access to d orbitals.

The central iodine atom in $I_3^{-2}$ has five electron groups (or regions) around it, which may consist of bonded or non-bonded electron pairs.
Based on the number of electron groups, the electron arrangement is trigonal bipyramidal.
However, the molecular shape is determined by the arrangement of the atoms, not just the electron groups.

Carbon-containing molecules are examined using Valence Shell Electron Pair Repulsion (VSEPR) theory to determine molecular geometry.
For a carbon atom with three electron groups and no lone pairs, the molecular shape is trigonal planar with bond angles of 120 degrees.

Dipole moment is related to the electronegativity difference between two atoms in a bond.
In hydrogen fluoride (HF), fluorine is more electronegative than hydrogen, resulting in a polar covalent bond and a partially negative charge on fluorine (δ-) and a partially positive charge on hydrogen (δ+).
The dipole moment is represented by a crossed arrow pointing towards the more electronegative atom.

In the absence of an electric field, polar molecules are randomly arranged.
When an electric field is applied, polar molecules align with their partially positive ends towards the negative plate and partially negative ends towards the positive plate.

A polar diatomic molecule consists of two different elements with differing electronegativities.
A nonpolar diatomic molecule consists of two identical elements, such as O2, with an electronegativity difference of zero.

For molecules with three or more atoms, the dipole moment is determined by the polarity of the bonds and the molecular geometry.
In carbon dioxide (CO2), there are two polar bonds between carbon and oxygen. However, because the molecule is linear and symmetrical, the dipole moments cancel each other out, making the molecule nonpolar overall.
If carbon dioxide were bent, the dipole moments would not cancel out, resulting in a polar molecule with a net dipole moment.
The overall dipole moment is the vector sum of individual bond dipole moments.

Diatomic Molecules: Atoms have significant differences in electronegativity in polar molecules.
Bond Dipole Moments: Represented using δ+ and δ-.
Polar Molecules: Dipoles do not cancel out; asymmetrical molecular symmetry; uneven charge distribution.
Nonpolar Molecules: Dipoles cancel out; even distribution of charges; symmetrical.
Solubility: Polar molecules dissolve in polar solvents (e.g., ethanol in water).
Intermolecular Forces: Dipole-dipole forces and hydrogen bonding in polar molecules; London dispersion forces in nonpolar molecules.

Ammonia (NH3) has nitrogen more electronegative than hydrogen, so the bond dipoles point towards nitrogen, thus the resulting total dipole moment points upwards.
Nitrogen trifluoride (NF3) has fluorine more electronegative than nitrogen, so the bond dipoles point towards fluorine. The resulting total dipole moment is smaller than that of ammonia because the dipoles partially cancel each other.

Naming Binary Compounds: Name the first element, then name the second element with an "-ide" ending (e.g., bromine chloride).
Bromine chloride (BrCl) is polar because chlorine is more electronegative than bromine.
Bromine trifluoride (BrF3) is trigonal planar.
Carbon tetrachloride (CCl4) structure is tetrahedral.
- Even with polar bonds, the symmetry of the molecule causes the dipole moments to cancel, resulting in a nonpolar molecule.

Cis-dichloroethylene has chlorines on the same side of the molecule, resulting in a polar molecule.
Trans-dichloroethylene has chlorines on opposite sides of the molecule, causing the bond dipoles to cancel and resulting in a nonpolar molecule.

Valence bond theory explains how atoms form covalent bonds through the overlap of atomic orbitals.
Potential energy decreases as atoms approach each other and form a bond. The most stable state occurs when the potential energy is at its minimum.

Hybridization involves mixing two or more atomic orbitals to form new hybrid orbitals.
Example: Beryllium (Be) typically has an electron configuration of $1s^2 2s^2$ . To bond with two hydrogen atoms, one electron from the 2s orbital is excited to the 2p orbital.
This results in two hybrid sp orbitals.