Topic 4: Inorganic Chemistry and the Periodic Table

Reactivity increases as you go down the group because atomic radius increases due to more shells. So less energy is required down the group to remove the outermost 2 e

You see a bright white flame and the formation of a white solid, with more vigorous reactions down the group. If the burning metal is placed in a glass jar of 02 then the same reaction occurs but more vigorously. For the reaction to start the element must be heated, but even w/o heat hay slow reaction entre the air and the metal, and an oxide coating prevents further reaction. Ba is so reactive that it must be stored in oil. general equation: 2M+ O2——-> 2MO

G2 oxides= basic oxides, meaning they react w/ water to form alkalies. You observe solids reacting to form colourless solutions G2 oxides react w/ acids to form salt+water (neutralisation). You observe a white solid reacting to form a colourless solution, and the reaction is exothermic.

SAME AS G2 OXIDES: G2 hydroxides react w/ acids to form salt+water (neutralisation). You observe a white solid reacting to form a colourless solution, and the reaction is exothermic.

solubility of G2 hydroxides increases down the group. Solubility of G2 sulphates increases up the group

K= lilac Na= orange Li= orange/red Cu= blue/green Ba= pale green Ca= brick red Sr= crimson Mg= no colour

-Weigh an empty boiling tube and place 1cm depth of metal carbonate into it. -reweigh the boiling tube, and clamp it at a 45 degree angle. -fit a bung w/ a delivery tube at the top of the boiling tube. The delivery tube should contain a test tube of limewater. -Strongly heat the boiling tube for 5 mins and record when the limewater is cloudy. -After 5 mins remove the the delivery tube from the limewater, leave it to cool and reweigh it. -Repeat with other metal carbonates. -Heat causes air in the test tube to expand, so the first bubbles you observe in the tube are expanded air, NOT CO2. CaCO3 —–> CaO + CO2

Dip a nichrome wire in HCL to avoid any impurities. Dip wire into the pure solid metal. Hold it to the flame and observe the colour change

Group 7 elements must gain an electron. As atomic radius increases down the group it becomes harder to attract an electron as the positive attraction of the nucleus is weakened by additional shielding. Therefore reactivity decreases down the group FIND INFO ABOUT ATOMISATION ENERGY AND E AFFINITY

Most halogens react with metals to produce metal halides. All halogens react with Na: 2Na + X2 ———> 2NaX. You observe a bright yellow flame and white solid forming of the sodium halide. All halogens except iodine react with Fe to produce solid Fe (iii) halides: 2Fe + 3X2 ———> 2FeX3. Iodine isn’t a powerful enough oxidising agent to oxidise Fe from 0 to +3 but it can oxidise it to +2. Fe + I2 ———> FeI2 Overall halogens are stronger oxidising agents at the top of the group

Add silver nitrate. Add nitric acid to prevent other precipitates forming. Add the halide sample, and then ammonia. Chloride ions = white precipitate that dissolves in dilute NH3 Bromide ions= Cream precipitate that dissolves only in concentrated NH3 Iodide ions= Yellow precipitate that does not dissolve in ammonia

Hydrogen halide + ammonium hydroxide –> salt + water HX + NH4OH ———> NH4X + H2O Hydrogen halide + water —> X⁻ + H3O⁺ Hydrogen halide + ammonia –> NH4X HX + NH3 —> NH4X

M+ Cl2——-> MCL2 Reactions get more vigorous down the group

Magnesium + water is a very slow reaction, but down the group there is an increase in fizzing. Metal hydroxides form Mg+ steam rapidly forms MgO (white solid) and H gas. The H gas is burned when it leaves the test tube for safety to prevent a highly flammable gas entering a lab

Add concentrated H2SO4 to a solid sample. Also soak a splint in ammonia solution to test for the hydrogen halide gas that will be produced: Floride and chloride ions: white misty fumes of HF and HCl Bromide: misty acidic fumes contaminated with brown bromine vapour Iodide: some steamy fumes of HI but lots of red and purple vapour

Each of the halogens could take electrons from other species to make their ions, so they can be oxidising agents. Whenever one of these halogens oxidise something IN SOLUTION, the halogen ends up as halide ions with water molecules attached

It is such a powerful oxidising agent. F oxidises water to oxygen so it’s impossible to do simple solution reactions with it.

Oxidising ability falls down the group. So a halogen higher in the group can oxidise all those lower down

Cl2 + 2I- ———> 2Cl- +I2 The iodine appears as a red solution if you use little chlorine, or a dark grey precipitate if chlorine is used in excess

Br2 + 2I- ——-> 2Br +I2 A red solution of iodine is formed until the bromine is in excess, the u get a dark grey precipitate

The halogen starts as diatomic molecules which may be a solid liquid or gas. These must be split apart to make individual atoms. Those atoms gain an electron, and isolated ions become wrapped in water molecules to form hydrated ions

The energy needed to produce 1 mole of isolated gaseous atoms starting from an element in its standard state, (eg gas for Cl2, liquid for Br2 etc)

For a gas like Cl2, it’s half the bond enthalpy because breaking the Cl-Cl bond produces 2 atoms, not 1. For a liquid or solid like Br2 or I2, it also includes the energy needed to convert them into gases

The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions X(g) + e- –> X- (g)