Topic 4 Inorganic Chemistry and the Periodic Table

• Group 2 elements are known as s-block elements because their valence electrons are in s orbitals.
• When Group 2 metals react, they lose two electrons to form 2+ ions, allowing them to achieve a full outer shell.
• The atomic radius of Group 2 metals increases down the group due to the additional electron shells.
• Reactivity increases down Group 2 since there is increased electron shielding and an increased atomic radius down the group, making it easier for the electrons to be lost.
• The first ionisation energy of Group 2 metals decreases down the group due to a greater atomic radius and increased shielding, making it easier for the electrons to be removed.
• Group 2 metals react with water in a redox reaction to produce a metal hydroxide and hydrogen. The metal hydroxide forms an alkaline solution.
• An example of a Group 2 metal reacting with water is magnesium. Magnesium reacts very slowly with liquid water, however steam can speed up this process (extra energy). Magnesium and water produces hydrogen and magnesium oxide (white powder).
• Group 2 metals react with chlorine gas to produce metal chlorides, which are white precipitates. As you move down the Group 2, the reactions become more vigorous (increased reactivity).
• Group 2 metals react with oxygen to form oxides. These reactions are vigorous. Strontium and barium can react with excess oxygen and heat energy to form metal peroxides.
• Group 2 metals react with dilute acids to produce bubbles of hydrogen gas and solutions of metal compounds.
• Group 2 hydroxides react with dilute acid to form a salt and water. This is known as a neutralisation reaction. Hydrochloric acid (HCl) forms chloride salts, sulfuric acid (H2SO4) forms sulfate salts, and nitric acid (HNO3) forms nitrate salts.
• The solubility of Group 2 hydroxides increases as you go down the group. Magnesium hydroxide, therefore, is the least soluble. Magnesium hydroxide has applications in medicine as an antacid since it is alkaline and can neutralise acids.
• Group 2 sulfates decrease in solubility down the group, meaning that magnesium sulfate is the most soluble. Barium sulfate (least soluble) has applications in medicine as a medical tracer for tissues and organs. It is toxic, but insoluble, so isn't absorbed by the blood.
• Barium chloride is used as a test for sulfate ions as it reacts to form barium sulfate which forms a white precipitate when sulfate ions are present.
• Group 2 carbonates and nitrates undergo thermal decomposition to produce solid metal oxides and gases. This is done by heating the substance in aerobic conditions.
• Group 2 carbonates produce carbon dioxide along with a metal oxide (white powder) when undergoing thermal decomposition.
• As you go down Group 2, more heat is required for thermal decomposition since the ions increase in size and the carbonates increase in thermal stability.
• The Group 2 nitrates produce nitrogen dioxide, oxygen and a metal oxide (white powder) when undergoing thermal decomposition.
• Group 1 carbonates and nitrates undergo thermal decomposition upon heating in aerobic conditions. In these reactions, lithium reacts the same as a Group 2 metals and produces the same products.
• All Group 1 metals besides lithium don't react like Group 2 because these reactions rarely achieve completion.
• Group 1 nitrates (aside from lithium) undergo thermal decomposition to produce a metal nitrate and oxygen.
• Group 1 metal carbonates won't thermally decompose in a lab because they require extremely high temperatures to achieve this.
• The elements of Group 1 and Group 2 form more stable carbonates and nitrates as you go down their groups, so they require more heat energy to thermally decompose.
• As you go down Group 1 and Group 2 , the ionic radius increases for the same overall charge. This means that smaller ions have a high charge density.
• The smaller ions at the top of Group 1 and Group 2 have a high charge density, so are able to polarise the negative carbonate and nitrate ions more. The more the negative ion is polarised, the less heat is required to separate the two ions. Therefore, smaller ions form less stable carbonates/nitrates and larger cations form more stable carbonates/nitrates.
• Group 1 and Group 2 elements can be identified via flame tests, since each metal has a unique flame colour.
• The procedure for flame tests involves firstly sterilising a nichrome wire and repeating this cleaning process until no colour is present in the Bunsen flame. Next, dip the wire into the unknown metal compound mixed with a few drops of HCl, and then place it into the flame. Observe the colour produced.
• In order to sterilise the nichrome wire, you should place it in a solution of concentrated HCl and then put it into a blue Bunsen flame.
• For Group 1 flame tests, lithium produces a red colour, sodium produces an orange/yellow colour and potassium produces a lilac colour.
• For Group 2 flame tests, magnesium produces no colour, calcium produces a brick red colour, strontium produces a crimson red colour, and barium produces an apple green colour.
• The formation of colours in flame tests can be explained by electron transitions.
• Electrons exist in orbitals, and when Bunsen heat is supplied, the energy is absorbed by the species. This causes some electrons to move to higher energy orbitals. When the electrons begin dropping down to their original energy orbital, they release energy.
• The colours produced by some elements in flame tests is a result of the electrons releasing energy. As they drop to their original level, energy is released and some is light energy. If a certain wavelength of light is emitted, it will fall within the visible light section of the electromagnetic spectrum, causing a colour to be produced.
• If no colour is produced during a flame test, then the energy released by the electrons transitioning doesn't fall within the visible light part of the electromagnetic spectrum.