Chemistry

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

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

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.

ELECTROLYSIS:

An electric current is the flow of ions (e.g. electrons)

Covalent compounds cannot conduct electricity as they have no ions meaning there can be no potential difference. Ionic compounds conduct electricity when molten or dissolved as it means the ions in the compound can now move around freely which means that it is possible for current to flow as the ions can now carry the current.

Experiment to distinguish between electrolytes and non-electrolytes: 

If when the circuit is turned on (the switch is closed) the lamp turns on, then the substance is an electrolyte.

Electrolysis of molten salts (e.g.PbBr₂)

1. Bubbling (effervescence) at the anode (positive) of Br gas
2. Lead fired at the cathode (negative)

Electrolysis of an aqueous solution with carbon electrodes:

Sodium Chloride:

The experiment is set up as before except that it is flipped and two test tubes are filled with water placed above the electrodes. This is to collect the gas.

At Cathode:

For positive ions, the lower down in the reactivity series the more easily it accepts electrons. Therefore H+ in the water discharges rather than Na+. Therefore, hydrogen gas is released.

2H⁺ + 2e^- —> H₂(g)

The sodium reacts with the remaining OH- in the water to form an alkaline sodium hydroxide solution  - Na⁺(aq) + OH^-(aq) —> NaOH(aq)

At anode:

The chlorine and hydroxide ions are at a similar position in the reactivity series. But there are more chlorine ions in the solution so it is mainly these that are discharged.

2Cl^- —> Cl₂(g) + 2e^-

Copper Sulphate:

At cathode:

Copper is lower in the reactivity series compared to hydrogen so copper is formed.

Cu²⁺ + 2e^- —> Cu

At anode:

The hydroxide in the water discharges, releasing oxygen.

4OH^- —> 2H₂O(l) + O₂(g) + 4e^-

Sulphuric Acid:

At cathode:

Hydrogen ions discharges.

2⁺ + 2e^- —> H₂(g)

At anode:

Once again the hydroxide ions discharge, releasing oxygen.

4OH^- —> 2H₂O(l) + O₂(g) + 4e^-

Calculations involving electrolysis:

1 Faraday = 1 Mole of electrons = 96500 Coulombs

Coulomb = current * time

A current of 0.2 amp is passed through copper sulphate for 2 hours, how much copper is formed?

Cu²⁺ + 2e^- —> Cu

2 hours = 2 * 60 * 60 = 7200

7200 * 0.2 = 1440

Therefore, there are 1440 coulombs

1440/96500 = 0.015

Therefore, there is 0.015 mole

0.015 / 2 = Cu²⁺

0.0075 = Cu²⁺

Ar of Cu = 63.5

63.5 * 0.0075 = 0.48

Therefore, answer is 0.48g

CHEMISTRY OF THE ELEMENTS (PART 1)

The Periodic Table

Metals - 

1. High melting and boiling points
2. Good conductors of electricity and heat
3. Have oxides that tend to be basic

Non-metals -

1. Generally have lower melting and boiling points
2. Don’t usually conduct electricity
3. Form acids with oxides. Sometimes they can also be neutral

Elements in the same group have similar chemical properties because their atoms have the same number of electrons in their outermost energy shell/level. Therefore, the atoms give or take the same number of electrons during a reaction if they are in the same group. This means that atoms in the same group have similar chemical properties.

Noble gases (group 0) are inert gases that generally do not react very well. This is because their outermost shell is full, they, therefore, do not need to take or give electrons as they are are already at their most stable configuration. 

Group 1 Elements:

All group 1 elements react with water to form hydrogen and a metal hydroxide solution. The reaction is also exothermic meaning it gives off heat.

Lithium - when placed in water the metal reacts with the water. Because the hydrogen released is not released symmetrically, the metal moves around and over the water. A white trail is formed of LiOH.

Sodium - the same reaction as lithium except that it melts as the heat produced is enough to melt sodium.

Potassium - enough heat is produced to burn the hydrogen released with a lilac flame.

Rubidium, Caesium - these metals are so reactive that the reaction created is explosive.

As you can see the further down you go in terms of the period of the group 1 elements (i.e. how big the element is), the higher the reactivity. This is because as the atoms of the elements become bigger, the electron in the outer shell moves further away, this means that it is easier for it to escape from the attractions of the protons in the nucleus. This, therefore, means reaction occurs faster.

Group 7 - the Halogens:

Properties:

1. Boiling and melting point increases as the period goes up
2. Salt - producing
3. Diatonic
4. Poor conductors
5. Colour becomes darker as the period goes up

Hydrogen chloride gas is an anhydrous compound in the gas state while hydrochloric acid is an aqueous solution of hydrogen chloride

Hydrochloric acid is also acidic because of the H⁺ ions in the water.

Hydrogen chloride is an ionic compound meaning it only dissolves in polar solvents such as water and not in organic substances such as methylbenzene. Dissociation of HCl because of the water creates an acidic solution as there will be H⁺ ions in the water.

As the period increases for the halogens, the reactivity decreases. Because of this, the larger the group 7 atom, the less volatile it is. This is why iodine is solid while fluorine is a gas. Also, as the period decreases, the colour lightens. 

Adding a solution of chlorine to a potassium chomped with a less reactive element (i.e. lower period) will cause the solution to become coloured as the least reactive element with be displaced.

e.g. 2KBr + Cl₂ —> 2KCl + Br⁺

These reactions are called redox reactions.

Oxygen and Oxides:

Nitrogen - 78.1%

Oxygen - 21%

Argon - 0.9%

CO₂ - 0.04%

Proving the volume of O₂ in the air:

With Copper:

1. Oxygen reacts with copper
2. Attach a syringe with 100cm³ of air to a tube packed with copper
3. Use a bunsen burner to heat the copper, allowing it to react with oxygen faster
4. Move the burner as you do to make sure only fish copper is burnt
5. Keep on doing this until the copper stops turning black
6. See how much of the air has been used and take it as a percentage

With the rusting of iron:

1. Place the iron wool in an inverted test tube filled with air and then place it in water
2. Mark the original level of water in the test tube with a rubber band
3. Wait a week for the iron to react with the oxygen and water (rust)
4. Mark the higher water level with another rubber band
5. Fill the tube up to the rubber band and pour the water into a measuring cylinder. Do this for the second rubber band level too
6. First volume / second volume = amount of oxygen as a %

Oxygen from hydrogen peroxide:

This is a catalytic decomposition reaction with manganese (IV) oxide as the catalyst. The oxygen is collected through the displacement of water.

Oxygen burning in…

1. Magnesium - burns with a bright white flame, with white powdery ash (MgO), which when dissolved forms a basic solution.
2. Carbon - burns with a yellow flame produces colourless CO₂ gas, which the dissolved forms an acidic solution.
3. Sulphur - burns with a bright blue flame, produces colourless sulphur dioxide, which when dissolved forms an acidic solution.

CHEMISTRY OF THE ELEMENTS (PART 2)

Properties of CO₂:

How to collect CO₂:

1. Calcium carbonate (marble chips) placed in a flat-bottom flask
2. Hydrochloric acid poured through a thistle funnel
3. CO₂ collected through upward displacement of air as CO₂ is more dense than air

Thermal decomposition of metal carbonates:

Copper (II) carbonate (green powder) to Copper oxide (black powder):

CuCo₃(s) —> CuO + CO₂

Carbon dioxide is a colourless and odourless gas that is denser the air and slightly soluble int water. It is used in carbonated drinks. It is more soluble in higher pressures so when you open the can, as the pressure falls the gas bubbles out. It is also used in fire extinguishers as it is more dense than oxygen and so prevents oxygen from reaching, thereby stopping the fire.

Acid rain is formed when acidic air pollutants such as sulfur dioxide and nitrogen dioxide dissolve in rain water. This increases the rate of erosion to marbles and limestone, destroys marine life and causes stunted growth as they important nutrients.

Hydrogen and Water:

1. Magnesium + Hydrochloric acid —> Magnesium chloride + Hydrogen

    Magnesium + Sulphuric acid —> Magnesium sulphate + Hydrogen

     Mg + 2HCl —> MgCl₂ + H₂

     Mg + H₂SO₄ —> MgSO₄ + H₂

      Rapid effervescence (gas releasing), exothermic, H₂ released, acidic solution left

2.  Aluminium + Hydrochloric acid —> Aluminium chloride + Hydrogen 

     Aluminium + Sulphuric acid —> Aluminium Sulphate + Hydrogen

     2Al + 6HCl —> 2AlCl₃ + 3H₂

     2Al + 3H₂SO₄ —> Al₂(SO₄)₃ + 3H₂

     Effervescence, H₂ released

3. Same occurs with zinc only slower

4. Copper does not react with acid as it is less reactive than hydrogen meaning it cannot displace     it.

Combustion of Hydrogen:

2H₂ + O₂ —> 2H₂O

Anhydrous Copper (II) Sulphate (white powder) reacts with water to create Hydrous Copper (II) Sulphate (blue powder). This is one of the tests for water.

The only water to test if water is pure, however, is to see if it boils at 100°C and freezes at 0°C.

Reactivity series:

Potassium - K

Sodium - Na

Lithium - Li

Calcium - Ca

Magnesium - Mg

Aluminium - Al

Carbon - C

Zinc - Zn

Iron - Fe

Tin - Sn

Lead - Pb

Hydrogen - H

Copper - Cu

Mercury - Hg

Silver - Ag

Gold - Au

Platinum - Pt

The order of the reactivity series is deduced on how violently they react with water and dilute acid. This can be seen from the speed bubbles are produced. If it does not react with cold water then it’s below magnesium. If it reacts in acid when warmed then it is above hydrogen.

A more reactive substance will displace a lesser one in a displacement reaction. this can be used to deduce the order using metal oxides with metals and metal salts with metals.

Redox reactions:

Oxidation - the gaining of oxygen, or the loss of an electron

Reduction - the losing of oxygen, or the gain of an electron

Redox - a reaction involving both oxidation and reduction 

e.g. Mg + CuO —> MgO + Cu

Oxidising agent - something that causes oxidation. CuO, in this case, is the oxidising agent.

Reducing agent - something that causes reduction. Mg in this case.

Rusting:

Iron rusts in the presence of oxygen and water. It may be sped up by electrolytes such as salts.

To prevent rusting, you can:

1. Coat the iron with paint, oil, grease or plastic to prevent oxygen or water from reaching the iron.
2. Galvanising - covering the iron with a layer of zinc. This means that even if a hole is created, the iron would not be touched the zinc would have to corrode first since zinc is more reactive.
3. Sacrificial Protection - this is the same as galvanising except with other metals such as aluminium.

Testing for ions and gases:

Flame test:

Place a small amount of the metal on a wire and burn using a bunsen burner. The colour of the flame should indicate what metal it is.

Lithium - red

Sodium - orange

Potassium - lilac

Calcium - brick - red

Using sodium hydroxide:

Add some sodium hydroxide to the solution. The colour of the solution formed should indicate what the substance.

Copper - blue gelatinous precipitate

Iron (III) - orange-brown gelatinous precipitate

Iron (II) - green gelatinous precipitate

Ammonium - solution does not have colour, a smell of ammonia, gas turns damp litmus paper red

The above were all tests for cations (positive ions). The following are for anions (negative ions).

Testing for chlorides bromides and iodides:

1. Add nitric acid to remove other substances apart from the substance you are testing for.
2. Add silver nitrate
3. Chlorides - white precipitate

      Bromides - cream precipitate

      Iodides - yellow precipitate

Testing for sulphates:

1. Add HCl to remove other substance such as carbonates
2. Add barium chloride
3. Sulphates - white precipitate

Testing for carbonate:

1. Add nitric acid
2. If CO₂ is formed then carbonate is present

The following are tests for gases…

1. Hydrogen - squeaky pop heard when a it splint is placed in the gas
2. Oxygen - relights a glowing splint 
3. CO₂ - turns limewater from colourless to cloudy
4. Ammonia - turns damp litmus paper blue
5. Chlorine - turns damp litmus paper red

PHYSICAL CHEMISTRY

Acids, Alkalis and Salts:

Universal indicator colours:

Red - acidic

Orange - weak acid

Green - neutral 

Blue - weak alkaline

Purple - alkaline

Acid - contains an excess of H⁺

Alkalis - contains an excess of OH^- ions

Metal + acid —> salt + hydrogen

Metal oxide + acid —> salt + water

Metal carbonate + acid —> salt + carbon dioxide + water

General rules for the solubility of salts in water:

• All common sodium, potassium and ammonium salts are soluble. 
• All nitrates are soluble
• Common chlorides except silver chloride and lead chloride are soluble
• Common sulphates except barium, lead and calcium sulphates are soluble
• Common carbonates except sodium, potassium and ammonium carbonates are insoluble

Preparing soluble salts:

1. Ready an acid with an excess of a solid metal (until there is some left at the bottom). Filter and head unless there are magnesium or carbonates. Place in an evaporating basin to form crystals.
2. Titration - only use this method if the metal is soluble in acid. Place 25cm³ of metal hydroxide in a conical flask using a pipette and add some phenolphthalein. Add the appropriate acid using a burette until the solution turns colourless. Repeat the experiment without phenolphthalein, heat gently to make the solution more concentrated and leave to cool to form crystals.

Preparing an insoluble salt:

Precipitate reactions - mix two solutions, one with the correct positive ion and the other with the correct negative ion. Filter out the spectator ions and water. Wash the residue with pure distilled water to remove any remaining spectator ions. After obtaining the purer residue of the salt, leave the salt to dry in a warm place.

Energetics:

Exothermic - a reaction where heat energy is given out

Endothermic - a reaction where heat energy is taken in

Measure the temperature at the beginning of the experiment and at the end. The amount of heat energy that has been released or taken in is the calorimetry of the reaction.

∆H = the energy released or taken per mole of the substance

∆H = temperature change * mass * heat capacity (J/g) / number of moles

Activation energy - energy required to start the reaction 

Breaking bonds require energy, this takes in heat and so is endothermic.

Making of bonds releases energy, this releases heat and so is exothermic

Calculating enthalpy change:

C-H = 413kJ/mol O=O = 495kJ/mol C=O = 803kJ/mol H-O - 463kJ/mol

CH₄ + 2O₂ —> CO₂ + 2H₂O

4*C-H + 2*O=O —> 2*C=O + 4*H-O

4*413 + 2*497 - 2*803 - 4*463 = -812kJ

This means that the reaction is exothermic. If the result was positive then it would be endothermic.

This also means that 812kJ energy is released for every mole of methane

Rates of reaction:

Factors affecting rate of reaction in terms of particle collision theory

1. Surface area - higher surface area means particles collide more frequently as there is more contact between the reactants.
2. Concentration / pressure - a higher concentration or pressure will mean there is a higher chance of particles colliding as there are more particles
3. Temperature - particles receive more kinetic energy meaning they move faster and so collide more frequently. 
4. Catalyst - provides an alternative pathway which requires a lower activation energy.

Equilibria:

Some reactions happen both ways with reactants becoming products and products becoming reactants. The symbol ⇌ is used to show this.

e.g. CuSO₄.5H₂O ⇌ CuSO₄ + 5H₂O

       NH₄Cl ⇌ HCl + NH₃

Dynamic Equilibrium is when in a reversible reaction, the forward and reverse reactions are at the same rate 

Le Chatelier’s Principle states that a reaction will move in such a way as to minimise the effect of a change to the system in order to return to its dynamic equilibrium

Because of this, for a reversible endothermic reaction, increasing the temperature will cause more products to be produced as a forward reaction occurs. However, for an exothermic reversible reaction when the temperature is increased fewer products will be made. This is because if fewer products are produced in an exothermic reaction then less heat is produced, this balances out the increased temperature. 

e.g. H₂ + I₂ ⇌ 2HI     ∆H = -10kJ (so exothermic)

Increasing the temperature here favours the endothermic reaction so more HI decomposes while when the temperature is decreased more HI will be formed.

The same principle applies to changing the pressure.

e.g. N₂ + 3H₂ ⇌ 2NH₃

      4 moles of gas - 2 moles of gas

So under high pressure, more ammonia will form and less nitrogen and hydrogen. The opposite happens when the pressure is decreased.

CHEMISTRY IN SOCIETY

Extraction and uses of metals:

If the metal is below carbon in the reactivity series than carbon could be used to react with the metal oxide to create a pure version of the metal. Above zinc, for metals such as aluminium, electrolysis has to be used. For some metals, such as titanium, its oxide is reacted with a more reactive metal. 

Using carbon - cheapest

Using electrolysis - expensive

Using other metals - most expensive

Extraction of aluminium from aluminium oxide:

Electrolysis is used as aluminium does not react with carbon. For this to work the ions must be free to move. Molten aluminium oxide only occurs at high temperature so it is dissolved in molten cryolite, another aluminium compound that melts at a lower temperature of 100°C and is a better conductor of electricity. 

The electrodes used for the electrolysis is made of graphite (carbon) and needs to be replaced regularly to prevent CO₂ from being formed, this applies to the anode. 

The amount of electricity involved is huge as it needs to create enough heat to keep the cryolite molten. Therefore, it is very costly.

At cathode (negative) electrode :

Al³⁺ + 3e^- —> Al

At anode (postive) electrode:

2O^2- —> O₂ + 4e^-

Extraction of iron from iron ore:

1. Hematite (a type of iron ore), coke (impure carbon) and limestone are placed in a blast furnace
2. C + O₂ —> CO₂ occurs
3. CO₂ + C —> 2CO occurs as the CO₂ is reduced by the coke at such a high temperature
4. Fe₂O₃ + 3CO —> 2Fe + 3CO₂ - the CO (reducing agent) reacts with hematite to form iron
5. Fe₂O₃ + 3C —> 2Fe + 3CO - this also happens at the hot parts of the furnace
6. Iron flows to the bottom of the furnace and can be tapped off
7. The limestone is added to remove impurities. Limestone undergoes thermal decomposition - CaCO₃ —> CaO + CO₂ - the calcium oxide (a basic oxide) reacts with acidic oxides such as silicon oxide and forms a molten slag which floats on top of the iron and can be tapped off separately.

Uses for…

Aluminium:

1. Conducts heat and electricity - used in planes (so that lightning conducts through the body of the plane), also used in cooking pans.
2. Shiny appearance - used in pans
3. Light (low density) - planes, electric cables
4. Resists corrosion (because of the layer of Al₂O₃) - aeroplanes, pans, electric cables
5. Strong when made into an alloy

Iron:

• Cast iron (4% carbon) - doesn't shrink much when solidified and so it is used for making castings. It is also used for manhole covers, guttering and drainpipes as it is very hard and cheap.
• Mild iron (<0.75% carbon) - nails, car bodies, girders. Very hard and relatively expensive
• Wrought iron (0% carbon) - decorates gates and raising as it is fairly soft
• High ( carbon steel 0.25 - 1.5% carbon) - cutting tools, masonry nails. The difference between measuring nails and mild steel nails is that the later would bend if miss-struck while the former would break as more carbon, means more strength but also means the material is more brittle.
• Stainless steel - an alloy of iron with chromium and nickel. Forms an oxide layer like aluminium so does not corrode. Used in kitchen utensils.

Crude Oil:

Crude oil is a mixture of hydrocarbons.

Fractional distillation is used to separate crude oil. A fractionating column is cooler at the top than at the bottom. Say a hydrocarbon boils at 120°C, the temperature as it enters the bottom of the column is high enough to make it rise as a gas. It rises until the temperature is low enough for it to condense, it can then be tapped off.

Crude oil contains:

1. Refinery gases - contains ethanes and is used as LPG (liquefied petroleum gas) for domestic heating and cooking.
2. Gasoline (petrol) - used in cars as fuel
3. Kerosene - used as fuel for jet aircrafts, for domestic heating oil and as paraffin for small heaters
4. Diesel gas - buses, lorries, some cars, oil railway engines
5. Fuel oil - ship boilers, industrial heating
6. Bitumen - tarmac

Viscosity and boiling point increases from 1 - 6 as the crude oil fraction becomes large in length.

Incomplete combustion may occur when burning hydrocarbons.

e.g. CH₄ + 3O₂ —> CO + 4H₂O

Carbon monoxide is bad for humans as it combines with haemoglobin in the bloodstream, preventing it from carrying oxygen, which can cause death.

In car engines, temperatures are high enough to allow nitrogen and oxygen in the air to react, forming nitrogen oxide.

The problem with fractional distillation of crude oil is that too much of the long chain hydrocarbons such as fuel oil is produced and too little of small chain hydrocarbons such as those of petrol are contained in crude oil.

To solve this, catalytic cracking may be used where long chain hydrocarbons such as fuel oil using silicon dioxide or aluminium aside at 600-700°C undergoes thermal decomposition to produce a random group/mixture of shorter chained hydrocarbons which can then be separated by fractional distillation.

e.g. C₁₃H₁₈ —> C₂H₄ + C₃H₆ + C₈H₁₈

Synthetic polymers:

Monomers are the base units of a polymer. A polymer is a long chain molecule made up of many monomers. 

Nylon is a condensation polymer where small molecules from both monomers are lost to allow them to join together. This molecule lost is usually water or hydrogen chloride.

Monomers of nylon are basically hydrocarbons with either COOH (carboxylic acid) or Cl(chloride) at both ends.

The Industrial Manufacture of Chemicals:

Ammonia is made by reacting nitrogen from the air and hydrogen from natural gas or from the cracking of hydrocarbons.

The Haber process:

In the Haber process, nitrogen and 3 hydrogen atoms react at 450°C, a pressure of about 200 atmospheres and an iron catalyst. 15% of it turns into ammonia. The unused gas is then recycled and reused after the ammonia has turned to liquid via cooling and condensation.

Ammonia is used to create nitric acid and is put into fertiliser as is contains nitrate ions.

Sulphuric acid is made by reacting sulphur (found in rocks and some natural gases) and oxygen (from the air)

Contact process:

1. Burn sulphur in air - S + 2O₂ —> SO₂
2. React oxygen with sulphur dioxide at 450°C, a pressure of about 2 atmospheres and a Vanadium (V) oxide catalyst. SO₂ + O —> SO₃
3. Reacts with sulphuric acid - H₂SO₄ + SO₃ —> H₂S₂O₇
4. Then water is added - H₂S₂O₇ + H₂O —> 2H₂SO₄

Sulphuric acid is used in detergents, fertiliser and paint.

Electrolysis of sodium chloride solution creates the products of chloride at the anode, hydrogen at the cathode and sodium hydroxide in the solution. Chloride is prevented from coming into contact with sodium chloride which creates bleach by using a diaphragm cell.

At anode:

2Cl^- —> Cl₂ + 2e^-

At cathode:

2H⁺ + 2e^- —> H₂

Uses of sodium hydroxide:

1. Making bleach by reacting it with chlorine
2. Making paper by helping break the wood down into pulp
3. Making soap by reacting with vegetable fats and oils to make compounds such as sodium stearate, which is present in soap

Uses of chlorine:

1. Sterilised water
2. Making bleach (as above)
3. Making hydrochloric acid

So that's all there is to IGCSE Chemistry. Once again now all you have to do is learn these notes, memorise them and then you're done.

Good luck in your exams!

PS. I've attached a complete copy in case you're too lazy to read it from the blog (it has a mighty 21 pages).

Click here