Chapter 2 Notes

Ionic Bonding

Ionic Bonding

Bond formed between atoms with a large difference in electronegativity, resulting in the complete transfer of electrons. Ionic compounds exist as crystals with high melting points due to strong attractive forces.

Covalent Bonding

Bond formed through the sharing of electrons. Can be covalent or polar covalent. Covalent bonds are directional and determine orbital shapes.

Lewis Structures

A method to represent the valence electrons of atoms in a molecule. Each atom shares valence electrons to attain the electron configuration of a noble gas atom in the same period.

Formal Charges

Calculation of the formal charge of an atom in a molecule by subtracting non-bonding electrons and half of the bonding electrons from the number of valence electrons. The best structure has minimized formal charges and is consistent with the electronegativities of the atoms.

Octet Rule

Atoms cannot overfill their valence shells and hydrogen must be surrounded by 2 electrons. Atoms tend to bond in a way that achieves a stable electron configuration with 8 valence electrons.

Resonance Structures

Different representations of a molecule that differ in the position of electrons. The actual structure is an average of the resonance structures, known as a resonance hybrid.

Bond Order

The number of bonds between two atoms. Single bond has a bond order of 1, double bond has a bond order of 2, and triple bond has a bond order of 3.

Line Diagrams

A simplified way to represent molecules, where solid lines represent covalent bonds, carbon atoms are at the intersection of two lines, and hydrogen atoms surround the carbon atom.

VSEPR Theory

Provides information about the 3D shape of molecules. The electronic arrangement is determined by the lone pairs and bonded atoms, while the molecular arrangement ignores lone pairs.

Molecule Polarity

Molecules with a dipole moment greater than 0 are polar, while molecules with a dipole moment of 0 are non-polar. Polar molecules have asymmetrical shapes, while non-polar molecules have symmetrical shapes.

Valence Bond Theory

Covalent bonds are formed by the overlap of atomic orbitals. Electrons are localized in bonds, and bonds can be formed through head-on or side-to-side overlap of orbitals.

Hybridization

Fixed combinations of two or more valence atomic orbitals. Hybrid orbitals were invented to make Valence Bond theory consistent with observed molecular shape. Different hybrid orbitals have different numbers of regions of electron density.

Boron Hydrides

Compounds containing boron-hydrogen bonds called boranes. Boron has an electron-deficient bond and can only form B2H6 (diborane). 3-centre, 2-electron bonds are formed where one electron pair is shared between three atoms.

Molecular Orbital Theory

Molecular orbitals are formed by combining the wavefunctions of atomic orbitals. Constructive combination results in bonding orbitals, while destructive combination results in antibonding orbitals. The number of molecular orbitals formed is equal to the total number of atomic orbitals combined.

MO Bond Order & Stability of Molecules

Species with higher bond orders are more stable. A bond order of 0 means there are no net electrons holding the molecule together.

Paramagnetism & Diamagnetism

Paramagnetic molecules have unpaired electrons and exhibit magnetism, while diamagnetic molecules have paired electrons and do not exhibit magnetism.

Diatomic (Homonuclear) Second Row Elements

Valence electrons in the n=2 orbital form 8 molecular orbitals. π MOs are formed from constructive combination, while π* MOs are formed from destructive combination. Relative energies of the MOs vary depending on the molecule.

Semiconductors & Band Theory

Explains the conducting, semiconducting, and insulating properties of materials. Conductors have a conducting band, insulators have a large band gap, and semiconductors have a small band gap.

Intermolecular Forces

Forces between molecules that hold them together. Types include Van der Waals forces (dispersion), dipole-dipole interactions, and hydrogen bonding. Boiling and melting points are influenced by the strength of intermolecular forces.