Bonding

Ionic bond

Electrostatic force of attraction between oppositely charged ions

Covalent bond

A shared pair of electrons

Dative covalent bond

The shared pair of electrons in the covalent bond come from only one of the bonding atoms

Direction of arrow when drawing dative covalent bond

From atom that provides lone pair to atom that is deficient

Metallic bond

Electrostatic force of attraction between positive metal ions and delocalised electrons

3 factors affecting strength of metallic bond

Nuclear charge, number of delocalised electrons per atom, size of ion

Shape formed by ionic bonding

Giant ionic lattice

Shapes formed by covalent bonding

Simple molecular and giant covalent

Shape formed by metallic bond

Giant metallic lattice

Ionic boiling and melting points

High (because of giant lattice of ions with strong electrostatic forces between oppositely charged ions)

Ionic solubility in water

good

Ionic conductivity when solid

Poor (ions can’t move/in fixed lattice)

Ionic conductivity when molten

Good (ions can move)

Simple molecular boiling and melting points

Low (weak intermolecular forces)

Simple molecular solubility in water

Generally poor

Simple molecular conductivity when solid

Poor (no ions)

Simple molecular conductivity when molten

Poor (no ions)

Macromolecular boiling and melting points

High (strong covalent bonds)

Macromolecular solubility in water

Insoluble

Macromolecular conductivity when solid

Diamond and sand poor, graphite good

Macromolecular conductivity when molten

Poor

Metallic boiling and melting points

High (strong electrostatic forces)

Metallic solubility in water

Insoluble

Metallic conductivity when solid

Good (delocalised electrons can move through structure)

Metallic conductivity when molten

Good

Metallic - malleable description

Planes of ions can slide over each other easily as ions are identical

Linear (b + l pairs, bond angle, example)

2, 0, 180, CO₂

Trigonal planar (b + l pairs, bond angle, example)

3, 0, 120, AlCl₃

Tetrahedral (b + l pairs, bond angle, example)

4, 0, 109.5, NH₄⁺

Trigonal pyramidal (b + l pairs, bond angle, example)

3, 1, 107, PF₃

Bent (b + l pairs, bond angle, example)

2, 2, 104.5, H₂O

Trigonal bipyramidal (b + l pairs, bond angles, example)

5, 0, 120 and 90, PCl₅

Octahedral (b + l pairs, bond angle, example)

6, 0, 90, SF₆

What shape of molecule?

Linear

What shape of molecule?

Trigonal planar

What shape of molecule?

trigonal pyramidal

What shape of molecule?

Tetrahedral

What shape of molecule?

Bent

What shape of molecule?

Trigonal bipyramidal

What shape of molecule?

Octahedral

Electron pair repulsion hierarchy

lone pair-lone pair > lone pair-bond pair > bond pair-bond pair

Electronegativity

The tendency of an atom in a covalent bond to attract electrons in the covalent bond to itself

Scale that measure electronegativity

Pauling Scale (0 to 4)

4 most electronegative atoms

F, O, N, Cl

Why does electronegativity increase across a period

The number of protons increases and the atomic radius decreases because electrons in the same shell are pulled in more

Why does electronegativity decrease down a group

The distance between the nucleus and the outer electrons increases and the shielding of the inner shell electrons increases

how large does difference in electronegativity have to be to form an ionic compound

More than 1.7

When does a polar covalent bond form

When atoms in the bond have different electronegativity (from 0.3 to 1.7)

How is a charge separation in a polar covalent bond produced

Unequal distribution of electrons (forming dipole ends)

Why symmetrical molecules can never be polar

If polar bonds are present, the individual dipoles ‘cancel out’ so there is no net dipole moment

how induced dipoles are formed

In any molecule electrons move about randomly which can form areas of high/low electron density so parts of molecule become more/less negative. This can cause dipoles to form in neighbouring molecules

Main factor affecting Van der waals forces

The number of electrons (the more electrons, the higher the chance that temporary dipoles will form, making VdW stronger between molecules)

What types of molecules do permanent dipole-dipole forces occur between

Polar molecules

Where do hydrogen bonds form

In compounds that have a hydrogen atom attached to on of the 3 most electronegative atoms (N, O and F - which must have an available lone pair of electrons)

Hierarchy of intermolecular forces

Hydrogen > permanent dipole-dipole > Van der Waals

What type of bonding structure is shown?

Giant ionic lattice (alternate positive and negative ions)

What type of bonding structure is shown?

Giant metallic lattice (close packing metal ions)

What type of bonding structure is shown?

Simple molecular structure (regular arrangement held together by weak vdw forces)

What type of bonding structure is shown?

Simple molecular - ice (central water molecule with 2 ordinary covalent bonds and 2 hydrogen bonds in a tetrahedral arrangement)

What type of bonding structure is shown?

Macromolecular - diamond

What type of bonding structure is shown?

Macromolecular - graphite (planar arrangement of carbon atoms in layers, 3 covalent bonds per atom in each layer, 4th outer electron is delocalised and is between layers)