Chemistry Unit 1

There are 4 main steps:

• Ionization: An electron gun fires at the sample knocking out an outer electron, leaving a postive ion. 
• Acceleration: The ions pass through an electric field formed using negatively charged plates.
• Deflection: They are then deflected by the strong magnetic field, lighter ions being deflected more than heavier ones.
• Detection: A negatively charged plate detects tiny electric currents.

First ionization energy is energy required to remove 1 electron from each atom in 1 mole of gaseous atoms forming 1 mole of ions with a single positive charge.
Ionization energy is influenced by:

• the number of protons.
• the distance of the outermost electron from the nucleus (shielding).

As a result of this, first ionization energy decreases as you go down the group.

The first ionization energy across period 3 increase because the force of attraction between the nucleus and electrons increases and the electrons are the same distance from the nucleus. However, there are two points where this does not apply.

• Between Magnesium and Aluminium there's a a decrease because the outer electron is now in the p sub-level.
• Between Phosphorus and Sulfur as the two outer electrons are paired and repel each other.

Empirical formula: the simplest whole number ratio of the atoms of each element in a compound.

Molecular formula: the actual number of atoms of each element in one molecule of the compound. It is a whole number multiple of the empirical formula. 

mass = Mr x n

n = c x v
     1000

pV = nRT

p = pressure in Pascals (Pa)
V = volume in m^3
n = moles
R = gas constant = 8.31 JK^-1mol^-1
T = temperature in Kelvin (K)

1 cm^3 == 10^-6 m^-3
1 kPa = 10^3 Pa
0 degrees C = 273 K 

Ionic bonding occurs when a metal and a non-metal react to form a compound.

• Metal atoms lose electrons to form cations.
• Non-metal atoms gain electrons forming anions.
• Ions are formed by the transfer of electrons.

 Ionic compounds are formed of lattice structures:

• Each ion is surrounded by ions with the opposite charge.
• Ionic compounds are said to have giant structures.

You cannot write molecular formula for an ionic compound, you must instead write an empirical formula.

Covalent bonds involve the sharing of electrons between two atoms.

Non-metal atoms can form double or triple covalent bonds by sharing more than one pair of electrons. These are stronger than single bonds.

A lone pair of electrons is a pair of electrons not involved in bonding.

Co-ordinate bonds 
Lone pairs of electrons form co-ordinate bonds with atoms that have vacant orbitals. 

• a co-ordinate bond is a shared pair of electrons where both electrons come from one of the atoms.
• they can be shown in displayed formula by an arrow.

Electronegativity is the ability of an atom to withdraw electron density from a covalent bond.

• If two atoms in a bond have different electronegativities then the most electronegative has a greater share of the electrons.
• Non-metals at the top of groups 5,6 and 7 are the most electronegative.

Electronegativity increases across the periodic table, as the number of protons increases.
Electronegativity decreases down the periodic table, as the amount of shielding increases.

Polar bonds are those in which the electrons are not shared equally.

• The more electronegative element has a partial negative charge.
• The less electronegative element has a partial positive charge .

Permanent dipole-dipole forces

When two polar molecules come near each other in space there will be an attraction between them.
Dipole-dipole forces exist between two molecules that have permanent dipoles.

Hydrogen bonding

Strong type of permanent dipole-dipole which exists between molecules that contain very electronegative elements.
For a hydrogen bond to occur:

• The molecule must contain an O,N or F atom bonded to a H.
• The molecule to which it is attracted must contain an O,N or F atom.
• Common examples include: water, ammonia and hydrogen fluoride.

Ice's unusual properties are due to its hydrogen bonds. 

Van der Waals' forces

VdWs is a force of attraction between a temporary dipole on one molecule and an induced dipole on another.
They also form between polar and non-polar molecules but aren't the dominant force.

Temporary dipoles occur because electron clouds aren't static and therefore the distribution of electrons at any moment may not be even.

Induced dipoles are caused by the presence of a temporary dipole nearby. These can then induce dipoles in other atoms or molecules.

Van der Waals' strength depends on a) the size of the atom/molecule and b) the area of contact.
a) temporary dipoles form more readily in large molecules and they get stronger down groups.
b) decreasing the area of contact (by branching etc.) reduces the strength of the VdWs.

Metallic bond is the electrostatic attraction between metal ions and the delocalized electrons in a metallic lattice.
In a metallic bond:

• Each metal atom forms a positive ion (cation).
• The positive ions are arranged in lattice structure.
• The atoms lost become delocalized.
• Delocalized electrons aren't attracted to any particular ion so are free to move through the metal.

Metallic bonds don't all have the same strength. The strength is dependent on:

• the charge of the ions.
• the number of electrons.

The cations formed arrange themselves in a hexagonal close-packed structure.

The atomic radius of the elements decreases across the period, this is because the number of protons increases and the extra electron is in the same shell so the pull of the nucleus increases slightly.

Sodium - metal - s block - metallic bonding,
Magnesium - metal - s block - metallic bonding,
Aluminium - metal - p block - metallic bonding,
Silicon - non-metal - p block - giant covalent bonding,
Phosphorus - non-metal - p block - simple covalent bonding,
Sulfur - non-metal - p block - simple covalent bonding,
Chlorine - non-metal - p block - simple covalent bonding,
Argon - non-metal - p block - monatomic bonding,

Argon is monatomic because it has a full outer shell and so doesn't generally form chemical bonds.

Alkanes have a single covalent bond. 
They have the general formula CnH2n+1.
If the chain is branched it is said to have a side chain, this is named depending on the number of carbons it contains.

Alkenes have a double covalent bond between two carbon atoms. 
Alkenes have the general formula CnH2n. 

Halkoalkanes are compounds containing carbon, hydrogen and a halogen atom.

Isomerisms occur when two or more molecules have the same molecular formula but different arrangements of atoms.

Structural isomerisms occur when they have different structural formulae.

Chain Isomerism: these occur when there is a different arrangement of the carbon chain. Typically they involve straight and branched hydrocarbons.
Example: butane and methylpropane.

Positional Isomerism: these occur when a functional group can be in more than one position on the carbon chain.
Example 1: but-1-ene and but-2-ene. 
Example 2: 1-bromopropane and 2-bromopropane.

C5H10 has five possible isomers:
pent-1-ene, pent-2-ene, 2-methylbut-1-ene, 2-methylbut-2-ene, 3-methylbut-1-ene.

Cracking is a decomposition reaction which involves the breaking of C-C bonds. This forms shorter chain alkanes and alkenes.

Short chain alkanes are more useful as fuels as they are more volatile (evaporate more easily) and burn more easily. The alkenes are used as raw materials for making polymers.

There are two ways of cracking:

• Thermal cracking.
• Catalytic cracking.

We use the cracking process because the amount of petrol produced, by fractional distillation, is insufficient and too much naptha is made instead.

Petrol is a mixture of hydrocarbons, it must vaporize readily enough to ignite engines but not so easily that excess vapour enters the engine. It also musn't ignite too early.

Octane rating tells us how easily the petrol ignites, a lower octane rating indicates the fuel will more readily auto-ignite (which causes a knocking sound and can lead to engine damage).

For complete combustion to occur:

• An unlimited supply of oxygen is required.
• Alkane + oxygen ---> carbon dioxide + water + release of energy.

In a limited supply of oxygen, incomplete combustion occurs:

• Water is still formed.
• Carbon only undergoes partial oxidation to form carbon monoxide (toxic).
• Unburned carbon particles are released (soot).

Global warming:
Greenhouse gases are good at absorbing infrared radiation and re-emitting in a different direction. This warms the atmosphere - the greenhouse effect.

The greenhouse effect is needed to keep the planet warm enough to sustain life. However, processes (like burning fossil fuels) have released large amounts of greenhouse gases which are heating up the Earth. This is global warming.

Acid rain is caused by sulfur dioxide, sulfur is present in fossil fuels. Acid rain can cause chemical weathering and acidification.