The periodic table is split into groups, which go vertically in columns along the table and periods, which go horizontally along the rows.
There is a line on the right hand side of the periodic table, which separates metals from non-metals.
As a general rule, metals are good conductors and non-metals are poor conductors. However, this is not conclusive. Metals and non-metals can be absolutely distinguished by the pH of their oxides. Metals make basic oxides, whereas non-metal oxides are acidic.
Elements in the same group in the periodic table have similar chemical properties because the have the same number of outer shell electrons, which is what determines chemical properties.
The noble gases (group 0) are a family of inert gases – they are unreactive due to their electronic configuration. They all have a full outer shell of electrons.
Group one metals react vigorously with cold water to form a metal hydroxide and hydrogen. Potassium burns the most violently, with a purple flame, then sodium with an orange/yellow flame. Lithium burns with a red flame, however it is less reactive than potassium or sodium and therefore does not always light. Sodium and potassium move quickly around the surface of the water and melt, and all the group one metals fizz and the metal disappears as it reacts.
The group one metals get more reactive going down the group as the outer shell gets further away from the nucleus and therefore the attraction between the nucleus and the electron gets weaker.
In lithium, the outer shell electrons are closest to the nucleus and so more energy is needed to take the electron, whereas in potassium, it is much further away.
At room temperature, the halogens are all quite obviously different. Fluorine is a yellow gas. Chlorine is a green gas. Bromine is a brown liquid. Iodine is a dark grey solid.
The halogens get denser and darker in colour as they go down the group. They also get less reactive as they go down the group because the outer shell of electrons gets further from the nucleus and the halogens want to gain an electron. There is less pull from the nucleus because the outer shell is further away.
Hydrogen chloride gas is a neutral gas, whereas hydrochloric acid, which is the gas, dissolved in water, is acidic.
Hydrogen chloride is acidic in water because H2O dissociates into H+ and OH- ions. When HCl dissolves in water, this results in it also splitting into H+ and Cl- ions, which makes it acidic (as it contains H+ ions) whereas when it dissolves in an organic, non-polar solvent, such as methylbenzene, it remains as the HCl molecule, which is not acidic.
Group 7 elements get less reactive as you go down the group. Fluorine is the most reactive group 7 metal.
E.g. 2KI (aq) + Cl2 (aq) → 2KCl (aq) + I2 (s) The chlorine displaces iodine, as it is less reactive. A grey solid (iodine) is formed as a precipitate.
These displacement reactions are REDOX reactions. The iodine is oxidized, as it loses electrons, and the chlorine is reduced.
There are a number of gases that make up our air:
You can find out the percentage of O2 in the air by doing a simple experiment. 100cm3 of air is put into a closed system that is comprised of two gas syringes and a middle section containing a metal that reacts with oxygen, in this case, copper, in excess. The system is then heated until all the oxygen has been used up, i.e. the gas volume is not changing and the copper is no longer reacting. The system is then allowed to cool down; otherwise it will not be an accurate reading, as gases take up more space when they are hot. The volume of gas left at the end is subtracted from your original amount, which should leave a value of around 20-21%.
2H2O2 →(using MnO2) → 2H2O + O2 Hydrogen peroxide breaks down into hydrogen and oxygen, which is sped up by the use of manganese dioxide as a catalyst.
Magnesium, sulphur and carbon all burn in air, a reaction that gives out heat and light.
CaCO3 + 2HCl → CO2 + H2O + CaCl2
Carbon dioxide is prepared in a laboratory using calcium carbonate and hydrochloric acid. This reaction occurs in a conical flask and the CO2 is collected downwards.
CuCO3 → CuO + CO2
Under heat, copper carbonate breaks down into copper oxide and carbon dioxide. This is a thermal decomposition.
Carbon dioxide is denser than air and will dissolve in water under pressure. These qualities mean that CO2 has certain industrial uses.
Carbon dioxide is used in fire extinguishers due to the property of it being denser than air. This means that it makes a barrier between the fuel and the oxygen and prevents oxygen getting to the fire. Without oxygen, the fire cannot continue to burn and therefore it goes out.
Carbon dioxide is also used in fizzy drinks because it dissolves in water under pressure, which is what makes the drinks fizzy, as the CO2 tries to escape.
Carbon dioxide is also a greenhouse gas, which means that it contributes the greenhouse effect, which in turn contributes to global warming.
Metal + acid → salt + hydrogen
Hydrogen reacts with oxygen to form water:
The water is formed as a vapour because the reaction is exothermic and therefore the water can only exist as a gas.
A chemical test for the presence of water is that anhydrous copper sulphate changes from white to blue. Cobalt chloride paper also changes from blue to pink in the presence of water.
A physical test for pure water is its boiling point, as water boils at 100 degrees Celsius.
The metal reactivity series
Only the more reactive metals react with cold water – with the most reactive of those being the most vigorous. Magnesium reacts with warm water but zinc and any further down the reactivity series do not react. Their reactivity can be determined by reactions with dilute acids. Silver and gold do not react.
Precise differences in reactivity can be found using displacement experiments. More reactive metals (and their salts) will displace those of less reactive metals.
Oxidation and reduction can also be described as the loss and addition of oxygen. Oxidation is the addition of oxygen, whereas reduction is the removal.
Redox reactions – a reaction in which both oxidation and reduction occur.
Oxidising agent – Gives oxygen (or takes electrons)
Reducing agent – Takes oxygen (or gives electrons)
Iron oxide forms in the presence of oxygen and water. It is a brown solid that is commonly known as rust.
Rust can be prevented by many methods:
Sacrificial protection means putting zinc or another more reactive metal onto the iron, which will react with the air more readily than iron, which prevents rusting.
Tests for Cations
Flame Tests Lithium compounds – red flame Sodium compounds – yellow/orange flame Potassium compounds – lilac flame Calcium compounds – brick red flame
Method: Clean a wire by dipping it into hydrochloric acid and then putting it into a flame and repeating the process until no colour comes from the wire. Then, put a sample of the compound being tested onto the wire. Place this over a roaring flame of a Bunsen burner, and see what colour appears in the flame.
Testing for NH4+ NH4+ + OH- → NH3 + H2O NH4Cl + NaOH → NaCl + NH3 + H2O Usually in this reaction, you can smell the ammonia, but a definitive test is to test for it with damp red litmus, which will turn blue
Testing for Cu2+ Fe2+ and Fe3+ Add aqueous sodium hydroxide (NaOH) to a solution of these metal ions and a coloured precipitate of metal hydroxide will form. Cu2+ + 2OH- → Cu(OH)2 Blue Precipitate Fe2+ + 2OH- → Fe(OH)2 Green Precipitate Fe3+ + 2OH- → Fe(OH)3 Brown Precipitate If you leave the Fe(II) solution for two long, it will reoxidise in the air and become Fe(III).
Tests for Anions
Testing for Cl- Br- and I- Add enough dilute nitric acid to make the solution acidic, which will remove any impurities, and then add some silver nitrate solution. Ag+ + Cl- → AgCl White Precipitate Ag+ + Br- → AgBr Cream Precipitate Ag+ + I- → AgI Yellow Precipitate
Test for SO42- (sulphate ion) Add some hydrochloric acid, followed by a few drops of barium chloride. A white precipitate of barium sulphate appears, which is insoluble in water. Ba2+ + SO42- → BaSO4
Testing for CO32- (carbonate ion) Add nitric acid, which will produce CO2. Test for the CO2 using limewater, which should go cloudy. ZnCO3 + 2HNO3 → Zn(NO3)2 + CO2 +H2O
Hydrogen Collect: Over water or upwards into a test tube
Test: Produces a squeaky pop when held to a lighted splint
Oxygen Collect: Over water
Test: Relights a glowing splint
Carbon dioxide Collect: Over water or downwards into a test tube
Test: Turns limewater milky – forms a white precipitate of CaCO3: Ca(OH)2 + CO2 → CaCO3 + H2O
Ammonia Collect: Upwards into a test tube
Test: Turns damp red litmus blue
Chlorine Collect: Downwards into a test tube or over salt solution
Test: Bleaches damp litmus