Bonding and structure

Topic 2 (no 20,21,22,26,27), Topic 13A definitions

Define ionic bonding

Electrostatic attraction between two oppositely charged ions

Describe the effects that ionic radius and ionic charge have on the strength of ionic bonding

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

How are ions formed

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

Describe the trend of ionic radii down the groups

The radii increases as there are more shells

Describe the trend of ionic radii with isoelectronic ions

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

Describe the evidence for the existence of ions

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

Define covalent bond

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

Describe the relationship between bond length and strength

Generally, the shorter the bond, the stronger it is

How is the shape of a simple molecule determined?

This occurs by electron pair repulsion

State the main shapes of molecules and their angles

Little - 180°

TurniPs - 120°

Never - 104.5°

Punch - 107°

Tiny - 109.5°

TriBes - 90° and 120°

Over - 90°

How do lone pairs affect bond angles

Each lone pair will reduce the bond angle by 2.5°

Define elctronegativity

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

State the ranges of the electronegativity scale and the bonding it corresponds to

0 - 0.4 = non-polar covalent

0\.4 - 1.7 = polar covalent (permanent dipole)

1\.7 - 4 = ionic bond

When will molecules be non polar

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

Describe London forces

\
forces

Describe permanent dipole 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

Describe hydrogen bonding

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

Describe H2O, NH3 and HF in terms of their hydrogen bonds

1

Explain the two anomalous properties of water

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

Explain the trends in boiling temperatures of the hydrogen halides

HF → HI → HBR → HCl

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

Where are giant lattices present

1\. Ionic solids (giant ionic lattices)

2\. Covalently bonded solids (giant covalent lattices)

3\. Solid metals (giant metallic lattices)

What is the structure of iodine and water

They are simple molecular

Describe the structure of graphite

1\. It is macromolecular

2\. 3 covalent bonds per atom

3\. Delocalised electrons between layers

Describe the structure of diamonds

1\. Macromolecular

2\. Tetrahedral structure

3\. 4 covalent bonds per atom

Describe the structure of graphene

1\. One layer of graphite

2\. 3 covalent bonds per atom

Define enthalpy change of formation

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

Define 1st ionisation energies

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

Define 2nd ionisation energies

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

Define lattice formation energy

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

Define enthalpy of atomisation

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

Define 1st electron affinity

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

Define 2nd electron affinity

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

State the steps of the Born-Haber process

1\. Atomise metal

2\. Ionise metal

3\. Atomise non metal

4\. Electron affinitise non metal

5\. Lattice

f==At== @@Ion@@ At %%E%%a^^L^^ing