AQA A-Level Chemistry - Periodicity

Periodicity

pattern in change of properties in a row of elements and this pattern is repeated in the next row

Block

S block - the block that belongs to hydrogen and groups 1 and 2

p block - The block that belongs to the non-metals and metalloids

d block - The block that belongs to the transition metals

f block - The block that belongs to the lanthanides and actinides

elements in each block have their highest energy electron in that sub-shell

position in the periodic table is determined by its proton number

Physical properties of period 3 elements

Atomic radius trend - increases - increase in no. of protons - increased nuclear charge - increased force of attraction on e- - similar shielding - atomic radius decreases

First ionisation energy - increases but drop between Mg and Al and drop between P and S

Melting point of elements - Na - Al - increases - smaller ion -increased nuclear charge - stronger electrostatic force of attraction

Silicon - macromolecular - many strong colavlent bonds must be broken - peak

P-Ar - depends on size P4, S8, Cl2 and Ar

Larger the molecule - stronger Van der Waals between the molecules - more energy required to overcome the forces

N.B - Argon - small and single atoms with e- close to nucleus - cannot be polarised - electron cloud not easily distorted

Physical properties of Group 2 elements

Atomic radius - increases - extra shell - increased shielding - despite increase in nuclear charge

First ionisation energy - increased shell - increased nuclear distance - increased shielding - despite increase in nuclear charge

Melting point - decreases - increase in ion size - more shells - weaker electrostatic force of attraction - but blip in Mg due to a unique crystal structure

The reactions of Mg-Ba with water

produces X(OH)2

reactivity increases down the group

Atom gets larger - outershell e- furthe away from nucleus

e- more readily lost

OH- ions formed

Mg has lowest reactivity so might have to use steam

Solubilities

Solubility of hydroxides increase down the group

So more OH- ions formed down the group

So solutions are more alkaline as you go down the group

Mg(OH)2 is sparingly soluble

Solubility of sulphates decrease down the group. BaSO4 is insoluble.

Test the solubilities of group 2 hydroxides by mixing a soluble group 2 salt eg X(NO3)2 with NaOH.

Test the solubilities o f group 2 sulfates by mixing solutions of soluble group 2 salts eg X(NO3)2 with H2SO4

The use of magnesium in Titanium extraction

TiO2 (Titanium ore) converted to TiCl4 by heating with carbon in a steam of chlorine gas. It is difficult to extract from its ore because Titanium reacts with carbon to form Titanium carbide TiC.

The TiCl4 is then purified using distillation and then reduced by Magnesium at 1000C

It is quite expensive

Other used of group 2 compounds

Mg(OH)2 in medicine - antacid - reacts with excess acid in the body relieves digestion

Barium meals - insoluble so forms a suspension - soft tissue does not show up on X-ray but BaSO4 is insoluble and can be used to view organ structures. Other Ba compunds can't be used because Ba2+ ions are poisonous.

Ca(OH)2 in agriculture - neutralising acidic soils in agriculture

CaO or CaCO3 to remove SO2 from flue gases (SO2 is a pollutant) - CaO and CaCO3 can be mixed with water to produce an alkaline slurry which is sprayed on the flue gases. The SO2 reacts to produce the solid waste product Calcium Sulfite which must be disposed. It also produces CO2.

Acidified BaCl2 can be used to test for sulfate ions. This is because BaSO4 is insoluble. It must be acidified to remove the presence of any unwanted ions that could interfere with the results.

Why wash with deionised water

remove SOLUBLE impurities

filtrate - is the solution

Group 7(17) Halogens

Fluorine - pale yellow gas

Chlroine - green gas

Bromine - red-brown liquid

Iodine - Grey solid

electronegativity - decreases as you go down the group - outershell e- further away - increased shielding - reduced power of the nucleus to attract a pair of electrons

boiling point - low b.p because they are simple molecular structures but increases as you go down the group - size of molecules increases - stronger Van der Waals bet. the moelcules - more E required to overcome the forces of attraction

Oxidising ability

halogens are electron acceptors

decreases down the group

displacement reactions show oxidising ability

adding chlorine, bromine water or iodine solution to potassium halide solution

orange solution when bromine displaced

brown solution when iodine displaced

therefore can be used to identify which halide is in the solution

Reducing ability

halide ions are electron donors

increases down the group

solid sodium halide + sulfuric acid (proton donor, strong oxidising agent and removes water)

Bromide ions can reduce H2SO4 to H2S

Iodide ions can reduce H2SO4 to H2S to SO2

Observation

Hydrogen halide - misty fumes

Br2 - orange fumes

SO2 - choking gas (test potassium dichromate goes green)

I2 - black solid and purple gas

Sulphur - yellow solid

H2S - rotten egg smell

test for halide ions

acidified silver nitrate solution

AgF = soluble in water

AgCl = white ppt

AgBr = cream ppt

AgI = yellow ppt

ammonia solution can be used to distinguish the silver halide formed as the colour may be subjective

AgCl dissolves in dilute ammonia

AgBr dissolves in concentrated ammonia

AgI does not dissolve

The silver nitrate is acidified to remove ions that may interfere with the test results

Required practical 4 - carry out simple test tube reactions to identify ions

Group 2 metals

nichrome wire + HCl

dip in unknown solution

non-luminous flame to observe color change

- Ca2+ - brick red

- Sr2+ - red

- Ba2+ - pale green

NH4+ add dilute NaOH and gently heat mixture- NH3 gas is given off - it is alkaline - ammonia dissolves in water and turn damp red litmus paper blue

Group 7 halide ions (described above)

-OH- - alkaline - drop red litmus paper - turn blue

-CO32- - add dilute HCl - fizzes - CO2 gas produced - bubble the gas through a test tube of limewater - limewater turns cloudy

-SO42- - add dil HCl + BaCl2 - white ppt will form

Use of Chlorine and Chlorate(I) ions

the reaction of chlorine with water to form chloride ions and chlorate(I) ions

the reaction of chlorine with water in sunlight to form chloride ions and oxygen

the reaction of chlorine with cold dilute aqueous NaOH to produce sodium chlorate, sodium chloride and water.

Sodium Chlorate is bleach and it kills bacteria and exterts bleaching action by oxidation of organic compunds. Can be used as bleach or disinfectant.

The use of chlorine in water treatment

Chlorate(I) ions kills bacteria

so adding chlorine to water makes safer to drink and swim in

Benefits:

kills disease causing micro-organisms

prevents reinfection further down the supply as it persists

prevents growth of algae

eliminates bad tastes and smells

removes discolouration of organic compounds

only very low doses required

Risks:

gas is harmful and irritates the respiratory system if inhaled

liquid chlorine on skin or eye can cause severe burns

chlorine reacts with hydrocarbons in the water (from bacteria decay and decomposition of plants) to from chlorohydrocarbons which are carcinogenic but increase risk of cancer = small price for preventing chlolera epidemic

the benefits to health of water treatment by chlorine outweigh the toxic effects and chlorine is also used in small doses

society assesses the advantages and disadvantages when deciding if chemicals should be added to water supplies.

Amount of chlorination required also depends on the country as more is required in hotter countries where increased temp leads to increased growth of bacteria.