GCSE Chemistry Unit 3

Sun causes evaporation of sea water

Water vapour rises & condenses, creating clouds
Droplets get too big -fall as rain
Returns to the sea, dissolving different minerals from rocks on / under the ground

Water is a solvent - it dissolves ionic compounds;

Nitrates,
Chlorides, (except for silver and lead)
Sulfates, (except for barium and lead)
Sodium, Potassium and Ammonium salts 

Water molecules are polar - positive hydrogen side, negative oxygen side, this attracts the different ions. 

Other compounds that are soluble in water - carbon dioxide, sulphur dioxide, salts, fertilisers and gases. 

Hard water is caused by calcium and magnesium ions picked up from the rocks and soil.

Adv. of hard water:
Good for teeth and bones (calcium)
Scale in pipes acts as a protective coating, stopping metal ions getting into drinking water, and protects pipes from rust

Disadv. of hard water:
Doesn't form a lather, but scum, when you add soap, so you need more soap
Forms scale (Calcium carbonate) on pipes, kettles etc. This causes problems that cost to fix

Why does hard water get scum?

Calcium sulfate (sol./water) + Sodium stearate (sol./soap) = Calcium stearate (insol./scum) + Sodium sulfate (sol.)
Sol - soluble Insol - insoluble
IT'S A PRECIPITATION REACTION

1. Heat it - temporary hard water

Calcium hydrogen carbonate is known as temporary hard water. It contains calcium ions that make the water hard, but when heated it thermally decomposes into calcium carbonate, which doesn't contain calcium ions.

2. Adding sodium carbonate (washing soda)

Na2CO3(aq) + CaSO4(aq) = CaCO3(s) + Na2SO4(aq)

Sodium carbonate + Hard water = Limestone + Soft water

As the limestone is insoluble, there are no more calcium ions in the water, and so it is soft.

3. Ion exhange

Hard water is put into the column, which contains Na+ ions. The Ca2+ and Mg2+ stays in the column (resin prefers to bond with 2+ than +) and the Na+ leaves in the water, making it soft. 

Water from reservoirs gets taken to water treatment works to be cleaned.

1. Screening - the water is passed through a net so that any big objects like twigs will be removed
2. Ozone treatment - microorganisms in the water are killed
3. Removal of salts and microorganisms - chemicals cause solids and microorganisms to stick together and fall to the bottom
4. Filtration - the water is passed through gravel beds to remove all solids
5. pH correction - ensures the water is neutral
6. Chlorination - kills off any remaining microorganisms that could cause harm

Solute - solid that you dissolve

Solvent - liquid you dissolve something in
Saturated solution - solvent can't hold any more solute

Increase temperature, increase solubility of solids, decrease solubility of gases
Increase pressure, increase solubility of gases (coke can)

Measured in g/100cm3 of water

To find the solubility of a substance:
Chuck the solute into the solvent until it is saturated (solute no longer dissolves)
Weigh the solution
Evaporate the water, leaving the crystals
Weigh the crystals

Mass of solution - Mass of crystals = Mass of water

x or / until it's in g/100cm3

Flame tests
Some metal compoounds burn with coloured flames
Li+ - Lithium - crimson red
Na+ - Sodium - yellow/orange
K+ - Potassium - lilac
Ca2+ - Calcium - brick red
Ba2+ - Barium - lime green

Precipitation tests
Many metal hydroxides are insoluble. Mixing a metal solution with sodium hydroxide (NaOH) creates a precipitation reaction as the metal bonds with the hydroxide, forming a coloured precipitate.

Ca2+ - Calcium - white (flame test)
Al3+ - Aluminium - white (redissolves in excess NaOH)
Cu2+ - Copper - blue
Fe2+ - Iron(II) -  dark green
Fe3+ - Iron(III) - reddish brown
Mg2+ - Magnesium - white (flame test) 

Testing for Carbonates

Carbonates fizz when added to acids, and form CO2. This can be tested by passing the gas produced through limewater - if the limewater turns milky, CO2 has been produced.

Some carbonates change colour when they are heated. Copper carbonate turns black and stays black, zinc carbonate turns from white to yellow, and then back as it cools.

Testing for Sulfates

Sulfates form white precipitates when added to barium chloride.

Testing for Nitrates

When added to aluminium powder and sodium hydroxide and then heated, nitrates form ammonia. Test for this using damp red litmus paper, which should turn blue.

Add dilute nitric acid and silver nitrate solution

Each one gives a silver precipitate;

Silver chloride - white
Silver bromide - cream
Silver iodide - yellow

If you can't tell the difference, add more nitric acid - the silver precipitate will eventually dissolve and the solution will become clear; first the chloride, then the bromide, and eventually the iodide (though iodide never goes completely clear)

Nearly all organic compounds contain carbon.

When burnt in air, hydrocarbons:
- produce a yellowy-orange and/or blue flame
- produce water and carbon dioxide
- char (go black) on the surface

To test for C=C double bonds, put the organic compound in bromine water, which should become colourless. 

Finding the empirical formula of organic compounds:

CxHy --> H2O + CO2

(unknown hydrocarbon burns to give water and carbon dioxide)

Masses:

CxHy - 0.4g
H2O - 0.9g
CO2 - 1.1g

0.9/18 = 0.05
1.1/44 =  0.025

C = 25 = 1
H = 50 = 2 - This is then doubled because there are 2 hydrogens in water
so H=4

CH4

In the early 1800's the periodic table was arranged in order of atomic mass.

1864 - Newlands noticed that every eighth element had similar properties, and so listed the known elements in rows of seven. Unfortunately, transition metals messed it up on the third row. 

This theory became known as Newlands' Law of Octaves.

It was criticised because:
- there were no gaps for undiscovered elements
- metals and non-metals were mixed up
- the groups contained elements that didn't have similar properties

1869 - Mendeleev arranged around 50 elements into his table of elements. He also ordered them by atomic mass, but left gaps to keep elements with similar properties in the same vertical groups. This left a massive gap on the third row - the transition metals. The gaps allowed people to predict the properties of undiscovered elements. 

At first, scientists didn't think the periodic table was important, but when electrons, protons and neutrons were discovered in the late 19th century, the periodic table was rearranged in order of atomic numbers, and all of the elements were put into groups. You can use the periodic table to find out the electrons in an atom, and from this predict it's properties. 

Electrons are set out in shells, or energy levels.

The maximum number of electrons that can occupy an energy level is given by the formula 2x(NxN), where N is the number of the energy level. Eg, energy level 1 has a max of 2 electrons (2x1x1), energy level 2 has a max of 8 electrons (2x2x2) and energy level 3 has a max of 18 electrons (2x3x3).

The group number tells how many electrons are on the last shell, and the period number tells how many shells there are. 

Properties.
All Group I metals: have one outer electron which they try to get rid of; form 1+ ions; only form ionic compounds; react with water to produce hydrogen gas. Li, Na, and K all float on water, and are soft, silvery metals.

Trends.
As you go down the group: the atoms get bigger because there is 1 extra full shell of electrons; the density increases because the mass does; the melting and boiling point lowers; the elements beome more reactive.

They become more reactive for 2 reasons;
1. Electrons are held in place because of the attraction between the nucleus and the electrons. The more electron shells, the further away the outer electron is from the nucleus, and so the attraction is less. This makes it more reactive because the electron is more easily lost.
2. As the atoms get bigger, there are more electrons between the nucleus and the outer electrons, reducing the electrostatic attraction. This is called shielding. As the nucleus is sheilded, it is easier for the outer electron to be lost. 

Properties.
All Group VII halogens: form in pairs (eg Br2); form 1- ions; bond ionically or covalently; and react with metals to form salts. More reactive halogens will displace less reactive ones. They all have distinctive colours - Flourine is a colourless gas with a yellow tinge; Chlorine is a green gas; Bromine is a red-brown liquid that turns to a orange-brown gas; and Iodine is a dark grey solid that turns to a purple gas. All the gas forms are poisonous.

Trends.
As you go down the group: the elements become less reactive; the melting and boiling points increase.

Transition metals are found inbetween Group II and Group III. They have the traditional metal properties; good conductors of heat and electricity, dense, strong, and shiny. They are less reactive than Group I metals, and they are denser, stronger, and harder than Group I metals. They also have much higher melting points (except mercury).

Transition metals often form more than one ion, eg Fe2+ and Fe3+, Cu+ and Cu2+, Cr 2+, Cr3+ and Cr6+.  These differnet ions usually from different coloured compounds - Fe2+ makes green compounds, Fe3+ makes red/brown compounds.

The compounds of transition metals are very colourful eg Potassium manganate(VII), Potassium Chromate(VI) and Copper(II) sulfate.

Transition metals and their compounds make good catalysts, eg Iron is the catalyst in the Haber process, nickel is the catalyst for turning oils into fats.

Their properties are caused by the way their electron shells fill. As you get further away from the nucleus, the closer the energy levels are. Eventually they start to overlap. This first happens between level 3 and 4. Eg Potassium has 19 electrons, but electron 19 goes on to the 4th level not the third, giving the arrangement 2,8,8,1 as opposed to 2,8,9

Arrhenius - 1880's

All acids release H+ ions when dissolved in water.
All alkalis form OH- ions when dissolved in water.

But not accepted as only works for those that dissolve in water, and scientists didn't believe it was possible - charged subatomic particles not discovered yet.

Lowry and Bronsted

They came up with definitions that works for soluble and insoluble bases (the problem with Arrhenius' theory)

Acids release H+ ions - they are proton donors
Bases accept H+ ions - they are proton acceptors

In acidic solutions the H+ ions are released, and become hydrated. These make the acid acidic.
In basic solutions the base takes H+ ions from the water, causing the water to release H+ ions, leaving OH- ions behind. These make the base basic. 

Acids can be strong, weak, concentrated or dilute;

Strong - ionise completely in water
Weak - ionise partially in water
Concentrated - more acid than water
Dilute - more water than acid

Titration
Shows how much acid is needed to neutralise an amount of alkali (or vice versa)

1.Put some alkali in a flask, with indicator
2.Add the acid a bit at a time using a burette
3.The indicator changes colour when the alkali is neutralised e.g phenolphthalein is pink in alkalis but colourless in acids
4.Record how much acid was used to neutralise the alkali 

To find out the concentration of a solution, use the equation
N=CxV where N=moles, C=concentration and V=volume

The concentration of a solution is it's molarity. If you have a 1 molar solution, then 1dm3 is equal to 1 mole

Titration can be used to find out the concentration of a solution.

E.g A student titrated 25cm3 of unknown concentration NaOH whith 13.9cm3 of 0.1 molar HCl
What's the NaOH concentration?

N=CxV
N=0.1x(13.9/1000)
  =0.00139 moles of HCl

The balanced equation means that NaOH and HCl are in the ratio 1:1 so

0.00139=Cx0.025 so C=0.00139/0.025=0.0556 molar

All reactions are either exothermic or endothermic

Exothermic reaction - gives out energy to the surroundings, usually as heat. Can be measured as a  rise in temperature.

Endothermic reaction - takes in energy from the surroundings, usually as heat. Can be measured as a drop in temperature.

Energy transfer in reactions can be measured. Measure the temperature of each reactant separately. Then mix them together and record the temperature every 30 seconds, until the temperature stops increasing.

In a chemical reaction, old bonds are broken and new bonds are formed. Energy is needed for bonds to break, and energy is released when bonds form. 

In an endothermic reaction, the energy released by new bonds is less than that of the energy needed to break the old bonds.

In an exothermic reaction, the energy released by new bonds is greater than that of the energy needed to break the old bonds.

Measuring the amount of energy produced when a fuel is burnt;

Burn the fuel. Use the flame to heat some water. This is called calorimetry. It uses a metal container, normally made of copper.

Method:

1. Put 50g of water in the copper can and record its temp.
2. Weigh the spirit burner and lid
3. Put the spirit burner underneath the can, and light the wick. Heat the water, stirring constantly, until the temperature reaches about 50°C
4. Put out the flame using the burner lid, and measure the final temperature of the water
5. Weigh the spirit burner and lid again

Disadvantages of burning fuels;
 - global warming, climate change - expensive to counter these effects
 - using fuel up means that it's going to get more expensive, including anything that needs fuel to be transported 

Mass of spirit burner + lid before heating = 68.75g
Mass of spirit burner + lid after heating = 67.85g
Mass of spirit burnt = 0.90g

Temp. of water before heating = 21.5°C
Temp. of water after heating = 52.5°C
Temp. rise of 50g of water due to heating = 31.0°C

So 0.90g of fuel produces enough energy to heat up 50g of water by 31°C

It takes 4.2 joules of energy to heat up 1g of water by 1°C (this will be included in the exam)

Energy produced in the experiment = 4.2 x 50 x 31 = 6510 joules
0.9g of fuel produces 6510 joules of energy, meaning 1g of fuel produces 6510/0.9 = 7233J

(example from CGP GCSE AQA Chemistry The Revision Guide)

Every chemical bond has a bond energy. This varies slightly depending what compound the bond occurs in. Known bond energies can be used to calculate the overall energy change for a reaction.

Eg
H2 + Cl2 = 2HCl

The bond energies you need are: H–H = +436 kJ/mol, Cl–Cl = +242kJ/mol, H–Cl = +431kJ/mol

1. Breaking one mole of H–H and one mole of Cl–Cl bonds requires 436 + 242 = +678kJ
2. Forming two moles of H–Cl bonds releases 2 x 431 = 862 kJ
3. Overall more energy is released than is used to form the products = 862-678 = 184kJ/mol released
4. Since this is energy released, if we wanted to show ∆H we'dneed to put a negative sign in front of it to indicate that it's an exothermic reaction, like this: ∆H = -184 kJ/mol 

Joules is the standard unit of energy, it replaced an older unit, the calorie.

1 calorie = amount of energy needed to raise the temperature of 1g of water by 1°C
1 calorie = 4.2 joules

Information on food labels is always written kilocalories, however they write it as 'calorie' not 'Calorie'.

1 Calorie (capital C) = amount of energy needed to raise the temperature of 1kg of water by 1°C
1 Calorie = 4200 joules 

We get our energy from food. Foods high in fat and oil produce large amounts of energy and carbohydrates produce some energy. Protein also contains energy, however we don't use it. 

Energy is released from food in respiration. Excess energy (food) is stored as fat. The fat stores are then used up if you do not have enough energy (food).