A regionaround the nucleus that can hold up to two electrons with opposite spins.
s- orbitals: spherical shaped
p- orbitals: dumbbell shaped
No. of electron pairs in orbitals
s : 1
p : 3
d : 5
f : 7
Ionic bonding
Electrostaticattraction between positive and negative ions
Giant ionic lattices: Oppositelychargedionsstronglyattracted in alldirections
High melting and boiling points
Requires more energy to overcome strong forces of electrostatic attraction.
Electrical conductivity
Solid: Nofree mobileelectrons;ions held in place by strong electrostatic forces of attraction
Molten: Ions are free to move; free mobile charge carriers.
Aqueous: Ions disassociate in water; freemobilecharge carriers.
Covalent bonds
Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom.
Single covalent bonding: 2 atoms contribute to and share a pair of electrons
Multiple covalent bonding: Atoms contribute and sharemultiplepairs of electronsbetween them
Dative covalent (coordinate) bonding: A bond in which one atom singly sharesallelectrons with another atom
Average bond Enthalphy: The meanenergyrequired to break1mole of bonds in gaseousmolecules.
A measurement of covalent bond strength
Bond angles
Trigonal planar: 120
Tetrahedral: 109.5
Trigonal bipyramidal:
Axial: 90
Equatorial: 120
Octahedral: 90
Electronegativity
The ability of an atom to attract the bonding electrons in a covalent bond.
Polar bond and permanent dipole
Within molecules containing covalently-bonded atoms with different electronegativities
Polar molecules require polar bonds with dipoles that do not cancel due to their direction.
Permanent and induced dipole-dipole forces can both be referred to as Van Der Waal‘s forces.
Induced dipole-dipole forces = London Dispersion Forces
Hydrogen bonding
Intermolecular bonding between molecules containing N. O or F and the H atom of -NH, -OH, or HF.
Electronegativity
Ability of an atom to attractbonding electrons in a covalent bond.
Increases to F in the periodic table
Electron Pair Repulsion Theory
Pairs of electrons around a nucleus repeleach other so the shape that a moleculeadopts has the electronspositioned as far apart as possible.
Strength of Repulsion
LP - LP -> LP - BP -> BP - BP
Orbitals are filled in order of increasing energy, with orbitals of the same energyoccupied singlybeforepairing.
Energy level
The shell that an electron is in.
Shell
The orbit that an orbital is in around the nucleus of an atom.
Permanent Dipole
A permanent uneven distribution of charge
Polar bond
A covalent bond that has a permanent dipole due to the different electronegativities of the atoms that make up the bond.
Non-linear (bent)
Central atom has 2 bonding pairs and 2 lone pairs
Octahedral
Central atom has 6 bonding pairs.
London (Dispersion) Forces
Induced dipole-dipole interactions caused when the random movement of electrons creates a temporary dipole in one molecule inducing a dipole in a neighbouring molecule.
Macroscopic Properties
Properties of a bulk material rather than the individual; atoms/molecules that make up the material
Polar Molecule
A molecule that contains polar bonds with dipoles that do not cancel out due to their direction (asymmetrical)
Pyramidal
Central atom has 3 bonding pairs and 1 lone pair
Simple Molecular Lattice
Solid structure made of covalently bonded molecules attracted by intermolecular force.
Tetrahedral
Central atom has 4 bonding pairs
Trigonal bipyramidal
Central atom has 5 bonding pairs.
Trigonal Planar
Central atom has 3 bonding pairs
Strength of metallic bonds
directly proportional to the number of delocalised electrons
Inversely proportional to the radius of the atoms.
Boiling points are based on strength of van der Waals’ interactions.
Affected by the number of electrons
Greater opportunity for more extensive temporary dipoles.