2.38 describe tests for the anions...

Halide ions
To test for halide ions, add dilute nitric acid (HNO3) followed by silver nitrate solution (AgNO3). A precipitate will be produces, the colour of this precipitate will determine what ions are present.
Cl -  > white precipitate
Br- > cream precipitate
I-   > yellow precipitate

NOTE: The silver nitrate solution determines which halide ions are present. The dilute nitric acid is added to 'get rid' of an carbonate or sulphite ions (as these would react with the silver nitrate).


Sulphate ions
To test for sulphate ions (SO42- ), add dilute HCl and then barium chloride solution (BaCl2). If sulphate ions are present, a note precipitate would form (this precipitate is barium sulphate)

NOTE: The HCl is used to 'get rid' of any carbonate ions, and these may impede results if present as they would also produce a precipitate).


Carbonate ions
To test for carbonate ions (CO32- ), add dilute HCl to the sample you are testing. If a gas is produced, collect it and bubble it through limewater. If carbonate ions are present, the limewater will go cloudy as carbon dioxide will be released.

2.32 understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidation is the gain of oxygen, reduction is the loss of oxygen.

NOTE: don't get this confused with the loss and gain of electrons... oxidation is loss of electrons and reduction is gain (O.I.L.R.I.G). Think of it as OXidation is gain ox OXygen (therefore, reduction must be loss)

2.31 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

Any metal higher in the reactivity series will displace one lower down from its oxide . For example...

to find out whether magnesium or copper is more reactive (therefore, higher in the series) just add magnesium to copper(II) oxide. The magnesium will displace the copper, meaning the magnesium is more reactive than the copper. Alternatively, you could add copper to magnesium oxide, No reaction would take place. This is because the copper is less reactive that magnesium (therefore, lower down in the series) as it will not displace the magnesium.

This must mean magnesium is above copper in the reactivity series.

It's the same thing with metals and a solution of their salt (the more reactive metal will displace a less reactive metal). For example...

To find out whether zinc or copper is more reactive, add zinc to a solution of copper (II) sulphate. The zinc will displace the copper (as it is more reactive), meaning it is positioned higher than copper in the reactivity series.

NOTE: in this particular reaction, the blue colour of the copper (II) sulphate solution fades as colourless zinc sulphate solution is formed. 

2.39 describe tests for the gases

Hydrogen
Squeaky pop test - put a lit splint into a test tube of gas. If it makes a 'squeaky pop' sound, hydrogen is in the tube.

Oxygen
Oxygen will relight  glowing splint.

Carbon Dioxide
If bubbled through limewater the limewater will go cloudy if carbon dioxide is present.

Ammonia
Damp red litmus paper will turn blue if ammonia is present.

Chlorine
Turns damp blue litmus paper red, then white (as it bleaches).

2.37 describe the tests for the cations...

Flame tests

To do a flame test, dip a platinum wire loop in dilute HCL then hold it in a flame, take it out once the flame burns without a colour (this means the platinum is clean). Dip the loop in the sample your testing and put it back in the flame, the colour of the flame will tell you what metal ion is in the substance...

Li+     > burns with a crimson red flame
Na+   > burns with a yellow-orange flame
K+     > burns with a lilac flame
Ca2+ > burns with a brick red flame

Sodium hydroxide solution and identify the ammonia evolved (NH4+)
You can check for ammonia gas using damp red litmus paper (will turn from red to blue is ammonia is present).

To test for ammonium ions in an unknown substance, add a few drops of sodium hydroxide solution to a solution of the unknown substance (in a test tube). Hold a piece of litmus paper near the top of the test tube, if ammonia is being given off, ammonium ions are present in the unknown substance.

Sodium hydroxide

Metal hydroxides are insoluble and precipitate out of solution when formed.

For this test add a few drops of sodium hydroxide solution to a solution the 'mystery compound' (what your testing). This will form an insoluble hydroxide. Some of these hydroxides have characteristic colours, this can be used to tell which metal ions are present in the unknown solution...

Cu2+ > blue
Fe2+  > sluggish green (yes, they are genuinely the words used)
Fe3+  > reddish brown

2.36 understand the sacrificial protection of iron in terms of the reactivity series

The sacrificial method involves placing a more reactive metal (such as zinc) with the iron. Water and oxygen then react with the sacrificial metal rather than the iron (as its more reactive than the iron).

Zinc is often used as it is more reactive than iron, so zinc will be oxidised instead of iron. A coating of zinc could be sprayed onto the iron object (galvanising). Another method is to bolt big blocks of zinc onto the iron (e.g. a ships hull of under groups iron pipes).

2.35 describe how the rusting of irony be prevented by grease, oil, paint, plastic and galvanising

These all create a 'barrier' around the iron, stopping water and/or oxygen from reaching it.

Paint/plastic can also be decorative and can be used on big or small structures (its versatile) and can be decorative (:

Grease/oil can only be used on moving parts e.g. bike chains

2.34 describe the conditions under which iron rusts

Rusting ONLY occurs when iron is in contact with both water and oxygen (from the air). The chemical reaction that is taking place is oxidation of iron to form iron(III) oxide (oxidation reaction), water then bonds to the iron(III) oxide and forms hydrated iron(III) oxide - this is rust.

NOTE: I'm not sure if we need the word equation so here it is just incase...

iron + oxygen ---> iron(III) oxide

iron(III) oxide + water ---> hydrated iron(III) oxide

2.33 understand the terms redox, oxidising agent, reducing agent

The oxidising agent is the element that gets reduced, the reducing agent is the element that gets oxidised. Reactions where oxidation and reduction take place are known as 'redox reactions' (REDuctionOXidation)

2.30 describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

Metals high up in the series (potassium, sodium, lithium and calcium) react very vigorous with water.

Metals in the middle (magnesium, zinc and iron) react with steam but don't react with cold water.

Copper won't react with either steam or water.

The more reactive the element is with water and dilute acid, the further up the series the element is positioned.

2.29 understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

The 'reactivity series' lists metals in order of their reactivity with other substances. The most reactive is at the top, while the least reactive is at the bottom. This is the order...

Potassium

Sodium

Lithium

Calcium

Magnesium

Aluminium

Zinc

Iron

Copper

Silver

Gold

KEY:

Super reactive

Fairly reactive

A little bit reactive

Not really at all reactive (basically inert)

NOTE: we need to learn this :( The best way is with an mnemonic. I use Peter.Says.Lions.Cats.Mice.And.Zebras.In.Cages.Suffer.Greatly but if you have another please comment!

2.28 describe a physical test to show whether water is pure

If the substance you are testing boils at 100ºC and/or freezes at 0ºC, it is pure water

2.26 describe the combustion of hydrogen

If hydrogen is burnt in air, water is produced (initially as water vapour, as it is hot, but later condenses to water if cooled)

The equation for this reaction is  2H₂+ O₂ → 2H₂O

2.25 describe the reactions of dilute hydrochloric acid and dilute sulphuric acids with magnesium, aluminium, zinc and iron

Acid + metal ---> salt + hydrogen


Magnesium

- Reacts vigorously with cold dilute acids
- Produces lots of bubbles


Aluminium

- Little reaction with cold dilute acids (as it has a protective aluminium oxide layer)
- Reacts vigorously with warm dilute acids and produces lots of bubbles


Zinc

- React slowly with dilute acids but more strongly if you heat them up


Iron

Same reactions as zinc.

NOTE: The more reactive the metal, the faster the reaction. For example, very creative metals (such as sodium) will react explosively, as the reaction is very quick. The speed of reaction is given off by the amount of bubbles that are given off.

2.24 understand the carbon dioxide is a greenhouse gas and may contribute to climate change

Gases in the atmosphere (such as carbon dioxide, methane, water vapour) act as natural insulators, trapping heat from the sun inside Earths atmosphere. This is because they absorb most of the heat that would normally be mediated out into space. They reradiate this trapped heat in all directions, including back towards Earth.

The level of carbon dioxide can increase due to many factors, some as a result of humans. For example, deforestation (fewer CO2 is removed by photosynthesis) and burning fossil fuels (CO2 that was 'locked' in these fuels is being released).

Although the Earth's temperature naturally varies (interglacial and glacial periods), there is reason to believe that extra CO2 has caused the average temperature to increase.

2.23 explain the uses of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

The stuff that makes your fizzy drinks fizzy is the carbon dioxide. As CO2 is slightly soluble in water and dissolves in drinks when user high pressure, a slightly acidic solutions forms due to the formation of carbonic acid. It eventually goes flat because, once you open the bottle, the CO2 is escaping (when it is flat, its all gone)

CO2 is more dense than air, this makes it perfect for fire extinguishers. The CO2 will sink onto flames and 'suffocate' them (stop the oxygen getting to the flames). As fire needs oxygen to burn, the fire will go out as no oxygen can get to it.

NOTE: CO2 fire extinguishers are only used when water extinguishers aren't safe. For example, when putting out an electrical fire.

2.22 describe the properties of carbon dioxide, limited to its solubility and density

Carbon dioxide is more dense than air (which is why it can be collected using the downward delivery method). It is also soluble in water at high pressure.

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate

The thermal decomposition (heating)  of metal carbonates will produce CO2 as, in thermal decomposition, the substance being heated will break down into simpler substances.

Method

- Put some copper(II) carbonate (its a green powder, if your wondering) into a test tube and insert a bung with a delivery tube at the top
- Clamp the test tube at a 90º angle and insert the delivery tube into another test tube (that is positioned vertically)
- Heat the copper(II) carbonate with a bunsen burner.

NOTE: Because CO2 is denser than air, the downward delivery method can be used.

Equations

Copper(II) carbonate ---> copper oxide + carbon dioxide

CuCO3(g) ---> CuO(s) + CO2(g)

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

Dilute HCL will react with calcium carbonate (aka marble chips) to produce calcium chloride, water and carbon dioxide...

Method

- Put marble chips at the bottom of a flask
- Fill the flask with hydrochloric acid, it doesn't have to be full, just covering the marble chips
- immediately attach a bung with a delivery tube into an upturned test tube in water
- the carbon dioxide will collect in this test tube


Equations

hydrochloric acid + calcium carbonate ---> calcium chloride + water + carbon dioxide

2HCL(aq) + CaCO3(s) ---> CaCl2(aq) + H2O(l) + CO2(g)

2.19 describe the reactions of magnesium, carbon and sulphur with oxygen in air, and the acid-base character of the oxides

When anything is burnt, it reacts with oxygen in the air to form oxides (which can have wither acidic or basic character). 

Magnesium
When magnesium burns in air, it produces a bright white flame and a white powder is formed (this is magnesium oxide). Magnesium oxide is slightly alkali when dissolved in water.

The equation for the reaction (burning Mg in air) is 2Mg_(s) + O_2(g) ---> 2MgO_(s)

NOTE: This question often comes up in exams so make sure you know it!

Carbon
Carbon will only burn in air if it is very strongly heated. It burns with a yellowy-orangey flame and produces carbon dioxide (as a gas). Carbon dioxide is slightly acidic when dissolved in water.

The equation for the reaction (burning C in air) is C_(s) + O_2(g) ---> CO_2(g)

Sulfur
Sulphur burns (in air) with a pale blue flame and produces sulfur dioxide which is acidic when dissolved in water.

The equation for this reaction (burning S in air) is C_(s) + O_2(g) ---> SO_2(g)