Bonding and structure

Electrostatic attraction between two oppositely charged ions

The shorter the distance between oppositely charged ions and the higher the charges, the stronger the electrostatic forces between them.

Positive ions from through the loss of electrons and negative ions form through the gain of electrons

The radii increases as there are more shells

The greater the number of protons the smaller the atomic radii as the electrons are pulled more closely to the nucleus

This is shown through x-ray diffraction as the electron density maps show the likelihood of finding electrons in a region

Strong electrostatic attraction between two nuclei and the shared pair of electrons between them

Generally, the shorter the bond, the stronger it is

This occurs by electron pair repulsion

Little - 180°

TurniPs - 120°

Never - 104.5°

Punch - 107°

Tiny - 109.5°

TriBes - 90° and 120°

Over - 90°

Each lone pair will reduce the bond angle by 2.5°

The ability of an atom to attract the bonding electrons in a covalent bond

0 - 0.4 = non-polar covalent

0.4 - 1.7 = polar covalent (permanent dipole)

1.7 - 4 = ionic bond

When they have an electronegativity between 0 and 0.4 and are symmetrical

forces

1. Occurs between polar molecules

2. It is stronger than London forces, so have higher b.p

3. They occur in asymmetrical molecules and have a bond with significant difference in electronegativity

1. Occurs with an H and a N/O/F which are the most electronegative and have a lone pair of electrons

2. There is a 180° between two molecules bonded by hydrogen bonds

3. Occurs in addition to London forces

1

1. High m.p and b.p- have both London (dipole-dipole) and hydrogen bonds

2. Ice has a lower density that water- when the hydrogen bonds are broken as the ice melts, the water molecules get closer together

HF → HI → HBR → HCl

This is due to the increasing electronegativity, meaning that HF can form hydrogen bonds while the others can’t

1. Ionic solids (giant ionic lattices)

2. Covalently bonded solids (giant covalent lattices)

3. Solid metals (giant metallic lattices)

They are simple molecular

1. It is macromolecular

2. 3 covalent bonds per atom

3. Delocalised electrons between layers

1. Macromolecular

2. Tetrahedral structure

3. 4 covalent bonds per atom

1. One layer of graphite

2. 3 covalent bonds per atom

Enthalpy change when 1 mole of substance is formed from its constituent elements in their standard states under standard conditions

Heat energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions

Heat energy required to remove 1 mole of electrons from 1 mole of gaseous +1 ions to form 1 mole of gaseous +2 ions

The enthalpy change when one mole of an ionic lattice is formed from its isolated gaseous ions

The enthalpy change when one mole of atoms is formed from its its element in its standard state under standard conditions

The enthalpy change when one mole of gaseous atoms acquires one mole of electrons to form one mole of gaseous negative ions (EXO)

The enthalpy change when one mole of gaseous negative ions acquires one mole of electrons to form one mole of gaseous 2- ions (ENDO)

1. Atomise metal

2. Ionise metal

3. Atomise non metal

4. Electron affinitise non metal

5. Lattice

fAt Ion At EaLing