C3 (Chemistry Extension)

Water Cycle

Continuous process by which water is transferred between the surface of the Earth and the atmosphere.

Sea water contains more ions than fresh water. To change sea water into fresh water use simple distillation (evaporation and condensation)

Purifying water

1. River water is screened to remove large debris.
2. The screened water and water from aquifers have ozone added, this removes algae, pesticides and kills 99% of bacteria.
3.Filter through layers of sand and gravel to remove smaller insoluble particles.
4. Filter through beds of activated carbon, this removes and remaining pesticides.
5. Add chlorine, killing any remaining bacteria.
6. Water is stored in water towers or covered reservoirs, with a small amount of chlorine to keep it fresh.
7. Water is sent to homes or factories for use. 

Solubility

Solute - solid that dissolves in a solvent to form a solution

Solvent - liquid in which a solute dissolves to form a solution

Solubility - number of grams of solute that dissolve in 100g of water (or other solvent) at a given temperature.

Saturated solution -  solution in which no more solute will dissolve at a given temperature.

Solubility

Gases
*Temperature - the higher the temperature, the less soluble gases become.
*Pressure - the higher the pressure, the more soluble gases become.

Solids
*Temperature - the higher the temperature, the more soluble solids become.
*Pressure - pressure has no effect on the solubility of solids. 

Solubility

Solubility curves can be used to
* Find the mass of a solute that will dissolve in a different amount of water.
* Find the minimum amount of water needed to dissolve a certain amount of solute at a given temperature.
* Predict when solutes will crystallise out of solution. 

Hard water

Hard water contains dissolved calcium or magnesium ions. It forms a lot of scum with soap and is difficult to lather.

Soft water doesn't contain dissolved calcium or magnesium ions. It easily forms lather with soap and not much scum.

Formation of hard water

Temporary hard water
Calcium carbonate + Water + Carbon dioxide = Calcium hydrogencarbonate (or the magnesium form)  

Permanently hard water
Calcium sulphate + Water = Dissolved Calcium sulphate (or the magnesium form)

+ Soap
Calcium sulphate + sodium stearate = sodium sulphate + calcium stearate (or the magnesium form)

Removing hardness from water

Temporary only
Heat it
Limescale forms, but the water is no longer hard

Both types
+ washing soda (Sodium carbonate)
Limescale forms (water is no longer hard) 

Both types
Ion exchange column
Calcium and magnesium are exchanged for sodium ions.
However the water should not be drunk in large quantities as high sodium content causes high blood pressure. The column also needs to be regenerated when all ions are used up (strong sodium solution)

Flame tests

Method of identifying metal ions from groups 1 and 2 in compounds (white)

Lithium = Crimson (red)

Sodium = Orange/ yellow

Potassium = Lilac

Calcium = Brick red

Barium = Lime green

Copper = Blue/ green
Some metals have no flame colour, and many colours are similar 

Adding Sodium hydroxide

This forms a coloured precipitate

Aluminium = White (redissolves if more is added)

Calcium = White (forms more precipitate if more is added)

Copper = Blue

Iron (II) = Green

Iron (III) = Orange

Magnesium = White (forms more precipitate if more is added)

Ammonium ions

Ammonium salt + sodium hydroxide = Ammonia + chemical salt + water

Ammonia is a alkali gas so it turns red litmus paper blue

Carbonates

Carbonate + Acid = Salt + Water + Carbon dioxide
Carbon dioxide makes limewater cloudy 

Heating carbonates

Copper carbonate = Black when heated (so it's really copper oxide)

Zinc carbonate = Yellow when hot, white when cold (but it's really zinc oxide)

Calcium carbonate = White when heated (really calcium oxide)

*(The lower the metal is in the reactivity series the easier it is to decompose)

Halides (Halogen ions)

+ Silver nitrate

Forms a coloured precipitate

Chloride = White (soluble in excess dilute ammonia solution)

Bromide = Cream (soluble in excess concentrated ammonia solution)

Iodide = Yellow (insoluble in excess concentrated ammonia solution)

Sulphates

+ Barium chloride to a solution

If a white powdery precipitate is formed the solution is a compound containing sulphate ions.

Nitrates

+ Sodium hydroxide to a solution

Heat it and + Aluminium

Ammonia gas should be given off if the solution contains nitrate ions

This turns red litmus paper blue.

Theory of acids and alkalis
Strong acid
Strong acids (such as hydrochloric) fully ionise in water producing many H+ ions.
Weak acids
Weak acids (such as ethanoic)  only partially ionise in water producing few H+ ions.
Strong alkalis
Strong alkalis (such as sodium hydroxide) fully ionise in water producing many OH- ions
Weak alkalis
Weak alkalis (such as ammonium hydroxide) partially ionise in water producing few OH- ions. 

Theory of acids and alkalis

Acids are proton donors, meaning they 'give' H+ ions (or a proton) to the solution.

Alkalis are proton acceptors, meaning they 'take' H+ ions (or a proton) from the solution.

Empirical formulae by combustion analysis

This is the method by which you can work out the formulae of a unknown compound (containing only hydrogen, carbon and oxygen), by burning it.

To find the mass of Carbon, take the mass of carbon dioxide produced by burning and multiply it by 12/44 (mass of carbon in carbon dioxide)
To find the mass of Hydrogen, take the mass of water produced by burning and multiply it by 2/18 (mass of hydrogen in water)
To find the mass of Oxygen, take the total mass of compound before burning and subtract the mass of hydrogen and the mass of carbon.

With your masses you can now work out the empirical formulae of the compound. 

Alkanes and Alkenes

Alkanes (such as ethane)
*Burn easily
*Do not react with bromine water
*Do not react with acidified potassium permanganate

Alkenes (such as ethene)
*Burn easily
*Cause orange bromine water to lose its colour
*Cause purple potassium permanganate to lose its colour. 

Other methods of identifying substances
Elements only
*Atomic line spectroscopy - each element has its own line emission spectrum
Compounds only
*Infrared (I.R.) spectroscopy - this identifies specific bonds in compounds
*Ultraviolet spectroscopy (U.V.) - detects nitrate and phosphate ions in water (commonly used in fertilisers)
*Nuclear magnetic resonance spectroscopy (NMR) - detects organic molecules
*Gas/liquid chromatography - separate gases
Both
*Mass spectroscopy - measures Ar and Mr
 

The alkali metals (group 1)
The alkali metals react with air and water, creating a alkali product.
As you travel down group 1
*They get softer (you can cut them with a knife)
*They get denser (Although the first two float on water)
*The melting point decreases (many are melted by their own reactions)
*The boiling point decreases
*The reactivity of the group increases, this is because the alkali metals reacting means they lose 1 electron from the outer shell. As you descend the group there are more shells between the outer shell and the nucleus, so the attractive force between the two is weak. This means it is easier for the atom to lose the electron and so react. (Basically the alkali metal at the bottom of group 1 is the most reactive)

Transition metals

Transition metals are the group between groups 2 and 3. They have 3 special properities
*They produce coloured compounds (this is the reason precious stones are coloured)
*They can be used as catalysts to change the reaction rate
*They form ions of different charges such as Fe (2+) and Fe (3+)

They have these properties due to their electronic structure. Across each row of transition metals a inner shell of electrons is being filled (not the usual outer shell) 

Halogens (group 7) (saltmakers)

As you go down group 7
*Melting point and boiling point increase (at the top the halogens are gas at room temperature, at the bottom they are solid)
*Density increases
*Reactivity decreases, this is because to react halogens must gain 1 electron, but as you go down the table there are more shells between the outer shell and the nucleus, so the attractive force is less, so it is harder to gain electrons and so react. (Basically the most reactive halogen is at the top). 

Periodic table
The periodic table as we know it today is mostly the work of John Newlands and Dmitri Mendeleev.

John Newlands (1864)
Newlands noticed that if the elements were arranged in atomic mass number their properties repeated at intervals. However this only worked for the first few elements.

Dmitri Mendeleev (1869)
Mendeleev copied Newland's table, except he left gaps for undiscovered elements and predicted their properties (accurately). He also put elements in atomic number order (without realising it). However again Mendeleev's table only worked for the first few elements, as the noble gases were yet to be discovered. 

Solution calculations

1. Calculate the number of moles of the substance whose quantity and concentration is known. (Moles = concentration*volume)
2. Work out the number of moles of the substance you need (reactant or product) 
3. Calculate the concentration, volume or mass of the substance as required. (Mass = Mr*moles) 

Choosing indicators for titrations

A indicator must change colour at the end point (neutralisation)
Strong acid and Strong alkali
Any indicator (but NEVER universal indicator)
Strong acid and Weak alkali
Methyl Orange
Weak acid and Strong alkali
Phenolphthalein
Weak acid and Weak alkali
No suitable indicator

Other indicators
Litmus
Bromothymol Blue
 

Energy changes in reactions

Exothermic
Chemical energy into heat energy
Reactants have more energy than products

Endothermic
Heat energy into chemical energy
Products have more energy than reactants 

Bond Energies

Bond energy is the amount of energy required to break or make a bond. Bond energy can be used to calculate the energy change in reactions.

1. Calculate energy required to break bonds in reactants.
2. Calculate energy required to make bonds in products.
3. Energy change = energy required to break bonds- energy required to make bonds
4. Negative value = exothermic
    Positive value = endothermic 

Energy in fuels/food

Heat energy = mass of water * specific heat capacity * change in temperature

Reactions in solution

1. Write balanced symbol equation
2. Calculate heat change
3. Calculate number of moles reacting
4. Calculate heat release/absorb for 1 mol