Chemical Bonding
Ionic Bonding
Covalent Bonding
Metallic Bonding
Physical Properties
Polyatomic Ions
Under normal conditions, ionic compounds are usually solids with lattice structures.
The number of electrons lost or gained is determined by the electron configuration of the atom.
Positive ions (cations) form by metals losing valence electrons.
Negative ions (anions) form by non-metals gaining [electrons]
volatility low
conducts electricity when dissolved or molten, never conduct when it is a solid
Solubility: often dissolve in water and do not dissolve in solvents that are not water
Hydroxide 
Hydrogen Carbonate 
Ammonium 
formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.
Bonds
Single Bond: one shared pair of electrons
Double bond: two shared paired electrons
Triple Bond: three shared pair of electrons
Bond length decreases and bond strength increases as the number of shared electrons increases.
Bond polarity results from the difference in electronegativities of the bonded atoms.
Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.
The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.
However, some atoms like Be and B, might form stable compounds with incomplete octets of electrons.
Resonance Structure
occur when there is more than one possible position for a double bond in a molecule.
Shapes of species are determined by the repulsion of electron pairs according to VSEPR theory.
Linear
Trigonal Planar
Trigonal Pyramid
Tetrahedral
Bent
Carbon and silicon form giant covalent/network covalent structures.
Intermolecular Forces
Strength of intermolecular forces: hydrogen bond > dipole dipole> London Dispersion Force
electrostatic attraction between a lattice of positive ions and delocalized electrons.
The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.
Alloys usually contain more than one metal and have enhanced properties.
Properties
Electrical Conductivity: Bonding electrons are delocalised; Current flow occurs without displacement of atoms within the metal
Malleability: Can be hammered into thin sheets;
atoms capable of slipping with respect to one another;
Melting Point
Melting (and boiling points) increase across the three metals because of the increasing strength of the metallic bonds.
The ionic bond is due to electrostatic attraction between oppositely charged ions.
Nitrate 
Sulfate 
Phosphate 
Carbonate 
2 bonding pairs 0 lone pairs,180 degrees
3 bonding pairs, 0 lone pairs, 120 degrees
3 bond pairs, 1 lone pair, about 107 degrees
4 bond pairs, 0 lone pair, 109.5 degrees
2 bond pairs, 1 or 2 lone pairs, about 104 degrees
Physical Property
Poor electrical and thermal conductivity
High volatility
Covalent compounds that are non-polar are insoluble in polar solvents unless thay are able to form hydrogen bonds with water.
Shiny
Hydrogen Bond: attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Hydrogen usually bonds with oxygen, nitrogen, or fluorine to make a hydrogen bond.
Dipole-Dipole bond: force between an ion and a polar molecule. A hydrogen bond is an example of dipole-dipole force and is an attraction between a slightly positive hydrogen on one molecule and a slightly negative atom on another molecule.
London Dispersion Force: temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles
Properties of giant covalent compounds and their structures
Very high melting points as the strong covalent bonds must be broken
Diamond does not conduct electricity. Graphite, C60 Fullerene contains free electrons, so it does conduct electricity. Silicon is a very low conductor of electricity.
insoluble in both polar and non-polar solvents as they're not strong enough to break the super strong covalent bonds.
Diamond: Each C atom is covalently bonded to 4 others in a giant tetrahedral arrangement. This structure gives a very rigid, 3-dimensional structure (so diamond is very hard).
Graphite: Each C atom is covalently bonded to 3 others to make a layer structure. Each layer is a giant structure. Because the layers are held together by weak Van der Waals forces, is is relatively easy to make the layers slide across each other.
This is a spherical molecule composed of 60 carbon atoms. The structure is similar to a “rolled-up” graphite layer. Each C atom is covalently bonded to 3 other C atoms & the remaining unpaired electrons are delocalised.
This property of graphite makes it a good lubricant and also a good pencil “lead”
Graphene: Giant 2 dimensional sheet