Redox Chemistry & Groups 1, 2, 7

• An oxidising agent is a species which oxidises another species by removing one or more electrons.
• A reducing agent is a species which reduces another species by adding one or more electrons.
• The oxidation number of oxygen is -2, except for in peroxides where it is -1.
• In groups 1 & 2, IE decreases down the group, due to an larger atomic radius and increased shielding effect (leading to decreased experienced nuclear attraction), which all result in less energy being required to remove electrons. Although nuclear charge also increases down the group, the effect of this is outweighed by the other factors.
• There is a general increase in reactivity down groups 1 and 2, as ionisation energy decreases.
• Alkali metals react with O2 in the air to form metal oxides, which is why group 1 metals tarnish when exposed to air.
• A metal oxide is a dull coating which covers the surface of the metal.
• lithium + water: relatively slow reaction; effervescence seen and heard
• sodium + water: lots of heat released, sodium melts; hydrogen released catches fire and causes the melted ball of sodium to dash on the surface
• potassium + water: lots of heat released, the hydrogen released burns with a lilac flame, the potassium melts into a shiny ball which dashes around the surface
• group 1 + 2 metals react with chlorine to form metal chlorides. group 1 chlorides are white solids at room temperature.
• When Sr and Ba react with chlorine, they can also form a peroxide.
• Magnesium reacts very slowly with cold water. The solution formed - Mg(OH)2, is weakly alkaline as magnesium is only slightly soluble in water.
• When magnesium is heated with steam, it reacts vigorously with the steam to make MgO and hydrogen gas, which is burned as it leaves the tube.
• metal oxide + HCl --> metal chloride + water
• metal oxide + H2SO4 --> metal sulfate + water
• A base is a proton receiver, which increases concentration of OH- ions.
• An alkali is a soluble base.
• An acid is a proton donor, which increases concentration of H+ ions.
• Group 1 metal oxides react with water to give a colourless alkaline solution (metal hydroxide).
• Group 1 metal hydroxides act like an alkali when added to dilute acid. Therefore, this is a neutralisation reaction, and forms a salt and water.
• All group 2 oxides are basic, apart from BeO which is amphoteric.
• CaO reacts vigorously with water, releasing a lot of energy. In this reaction, some of the water boils off as the solid lump of CaO seems to expand and crack open.
• When a group 2 oxide + H2SO4 forms an insoluble sulfate, it forms at the surface of the oxide, so the solid oxide beneath it cannot react with the acid. This is prevented by using the oxide in powder form and stirring so that neutralisation takes place.
• Group 2 metal hydroxide + dilute acid --> colourless solution of metal salt.
• Sulfates decrease in solubility going down the group. This trend is seen by the fact that BaSO4 is a white precipitate.
• Group 2 hydroxides increase in solubility down the group, so become more alkaline.
• We test for CO2 by bubbling it through limewater. The limewater will go cloudy/milky if present, and a white precipitate of calcium carbonate forms.
• Milk of magnesia is magnesium hydroxide, sold as an indigestion remedy. It neutralises HCl in the stomach which relieves symptoms.
• Sulfate ions can be tested for by adding solution containing barium ions. Sulfate ions will react with barium ions to form a white precipitate of BaSO4. There must be H+ ions present to stop other ions from forming a precipitate, so HNO3 or HCl are usually added.
• There are no trends for the solubility of group 1 hydroxides/sulfates, as they are all soluble in water, so no precipitates will form. We must use the flame tests.
• Thermal decomposition is the breakdown of a compound into 2 or more different substances using heat.
• Barium ions are poisonous, but hospital patients are sometimes fed barium meals as the dense white precipitate makes soft tissue more visible on X-rays. Not poisonous, as it is highly insoluble, so ions aren't free to move.
• We test for carbonates by reacting with acid. If carbonates are present, carbon dioxide will be produced, and the limewater test will be positive.
• Group 1 and 2 carbonates decompose when heated to form a metal oxide and carbon dioxide.
• Group 1 and 2 nitrates decompose to form the metal oxide, oxygen, and nitrogen dioxide gas.
• Thermal stability of Group 1 and 2 carbonates and nitrates increase down the group. This is because smaller positive ions at the top of the groups will polarise the anions more than the larger ions at the bottom of the group. The small positive ion attracts the delocalised electrons in the nitrate/carbonate towards itself. The higher the charge, and the smaller the ion, the higher the polarising power. Therefore, the more polarised they are, the greater the ease of thermal decomposition, as the bonds in the carbonate and nitrate ions become weaker.
• We test for NO2 by dissolving it in water, and if present, it would give an acidic solution which would turn blue litmus paper red.
• Metal ions produce a colour if heated strongly in a flame, as the heat excites the outer electron and causes it to move to a higher energy level. The electron is unstable at this energy level so falls back down. As it drops back down, energy is emitted in the form of visible light energy with the wavelength of the observed light.
• Flame test colours: Li+ = scarlet red, Na+ = yellow, K+ = lilac, Rb+ = red, Cs+ = blue, Ca2+ = brick red, Sr2+ = red, Cu2+ = green, Ba2+ = green.