Chemistry - EDEXCEL IGCSE - Ionic, covalent and metallic compounds

IONIC COMPOUNDS:

Ions are formed through the adding (reduction) or removal (oxidation) of electrons.

Ionic Bond - a strong electrostatic attraction between oppositely charged ions

Properties:

• Structure - a 3D lattice structure held together by ionic bonds.
• Melting and boiling points - Very high (Note - the higher the ionic charge the higher the mp/bp)
• Conductivity - only conducts when molten or dissolved
• Solubility - dissolves in water but not organic substances

Ionic compounds have high melting and boiling points because of the strong lattice structure which means the compounds are held together more strongly. This means more energy is needed to break apart the bonds and so more heat. Furthermore, the electrostatic attraction between ions is very strong.

Ionic compounds do not conduct as a solid as the ions are not free to move and so cannot carry the charge. When molten or dissolved, the ions are now free to move and so the compounds conducts.

COVALENT COMPOUNDS:

Covalent bond: the attraction between positive nuclei for a shared pair of electrons.

                    

Covalently bonded molecules are held together by weak intermolecular forces so are mostly volatile liquids or gases as they have very low melting and boiling points.

Covalent substances also cannot conduct electricity as they have no charged particles and so cannot carry the charge. 

Substances with a giant covalent structure such as diamond high a high melting and boiling point as they are held together by strong covalent bonds which hold them together very strongly.

Diamond = carbon joined to 4 others

Graphite = carbon joined to 3 others

Diamond is hard as it has very many covalent bonds holding it together. This makes it very useful for cutting.

Graphite is built in layers of carbon which can slide over each other making it slippery and, therefore, a good lubricant. It has a relatively high melting and boiling point as it is held together by strong covalent bonds. Furthermore, it is also a good conductor as the fact that only 3 carbons are joined to each other means that there is always 1 free electron to carry the charge.

METALLIC CRYSTALS

Metals are a giant structure of positive ions surrounded by a ‘sea’ of delocalised electrons.

They are malleable, this is because they can ‘slide’ over each other. This is because they have a regular shape/structure and also because the ‘sea’ of electrons means the positive ions are always held together even when they are being moulded.

Metals also conduct as the delocalised electrons are free to move and can, therefore, carry the charge as they are ions.

Chemistry - EDEXCEL IGCSE - Separation techniques and calculations involving moles

SEPARATION TECHNIQUES

Solids and Liquid:

1. If solid has not dissolved then use filtration. The liquid (filtrate) passes through leaving the solid (residue). This works because of different sizes of particles.
2. If the solid in a solution is required use evaporation.
3. If both are needed then simple distillation is required. The solution is heated and the liquid evaporates and passes into the condenser where it cools and condenses. The solid remains in the flask as residue, the liquid (distillate) can be collected.

Liquid and Liquid:

1. For immiscible liquids (liquids that don’t mix) use a separating funnel. This works because of differences in densities.
2. For miscible liquids then use fractional distillation. This is where a temperature gradient is created in a fractionating column meaning the liquid with the lower boiling point condenses at the top and vice versa. This occurs as the liquid with the lower boiling point boils first.

Solid and Solid:

1. Chromatography:

• This takes advantage of different solubility. 
• Draw a pencil line on paper.
• Place a concentrated dot of the mixture on the paper.
• Place the paper in a beaker of solvent which dissolves the mixture.
• Allow the solvent to move up the paper and dry the paper.
• The solids will have been separated.

Retention factor = distance travelled by the solid / distance travelled by the solvent.

RELATIVE FORMULA MASS AND MOLAR VOLUMES OF GASES

Relative formula mass (Mr) - relative atomic mass of a formula. e.g. the Mr of H2O is 18

A mole is one of chemistry’s counting unit. One mole is equal to 6.022 * 10

For molar volume of gas, one mole always equates to 24dm3 or 24,000 cm3 at room temperature and pressure.

Example question: 6g of a substance is 4.8dm3, what is its molar mass?

1 mole = 24dm3

0.2 mole = 4.8dm3

-

0.2 moles = 6g

1 mole = 30g

Molar mass = 30g/mol

Chemical formulae and equations:

How to test if a formulae is correct:

1. Weigh substance (A)
2. Remove one element through a reaction
3. Weigh again (B)
4. Start weight (A-B) = weight of element removed (C) 
5. Repeat until all the elements have been separated
6. Divide weight C by the atomic mass of the element removed
7. The ratio tells you the formula

Calculating empirical formulae:

1. Find the mass of all the different elements
2. Divide the masses by the molar mass of the element
3. Divide through by the number of moles which is the lowest
4. Multiply to get the smallest whole numbers 

Calculating molecular formulae:

1. Work out the empirical formula
2. Do molecular mass / formula mass
3. Times the result by the empirical formula

Calculations with reacting masses:

What mass of oxygen is needed to burn 3kg of propane, C₃H₈?

C₃H₈ + 5O₂ —>  3CO₂ + 4H₂O

The relative formula mass of propane = 3*12+8*1 = 44

So the molar mass = 44g/mol

3000 / 44 = 68.2 moles

5O₂ = 68.2 * 5 = 341 moles of O₂

O₂ = 32g/mol

341*32 = 10.9kg

Therefore answer = 10.9kg

Calculating percentage yield:

Percentage yield = total amount obtained * 100 / maximum theoretical amount 

2.8g of Fe reacts with S to form 4.1g of FeS, what is its percentage yield?

Atomic mass of Fe = 56

2.8 / 56 = 0.05

0.05 moles of FeS = 0.05*88 = 4.4g

4.1*100/4.4 = 93.2(3s.f.)

Therefore answer = 93.2% 

Molar Concentration (molarity) = Number of moles / Volume

Chemistry - EDEXCEL IGCSE - States of matter, atoms and atomic structure

STATES OF MATTER

ATOMS

Evidence:

1. Dilution of coloured (e.g. copper sulphate solution)

• Use serial dilution (step-wise)
• Each stage shows the same colour but a lighter shade
• Provides evidence of small ‘particles’ which spread out

     2. Diffusion of potassium manganate and water

• Water becomes a purple colour
• Shows evidence of particles which move.

Definitions:

Atom – the basic unit of chemistry and matter it is made of electrons, protons, neutrons.

Molecule – where atoms join together covalently.

Element – a substance made of only one type of atom.

Compound – a substance made up of 2 or more types of atoms chemically bonded and in a fixed ratio.

Mixture – a group of substances not chemically bonded together and not in a fixed ratio.

Atomic number - number of protons

Mass number - number of protons and neutrons in the nucleus

Isotopes - atoms of an element with differing mass numbers but the same atomic number.

Relative Atomic Mass - the weighted average relative to a Carbon - 12 atom.

Formula fo relative atomic mass (Ar):

Ar = ((x * %x) + (y * %y)) / 100

Where x, y = mass of the isotope 

And %x, %y = abundance of the isotope 

ATOMIC STRUCTURE