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Topic 4 Chemical Changes By Bethan Poole (4.2 Reaction of acids (4.2.4 The…
Topic 4 Chemical Changes
By Bethan Poole
Reactivity of metals
4.1.2 The Reactivity Series
Determined by the metal reaction with oxygen in air, water and dilute acid
Metals react with acids to produce salts and hydrogen
Lithium, sodium and potassium
- very reactive, top of reactivity series, react vigorously with water to produce metal hydroxide solutions and hydrogen
Not safe to react with dilute acids
Calcium, magnesium, zinc and iron
- fairly reactive, react quicky with acids and slowly with water
Copper and gold
- very unreactive, don't react with water or acids
Includes
hydrogen
and
carbon
for comparison
4.1.3 Extraction of Metals and Reduction
The method of extraction depends on the metals reactivity
Unreactive metals (e.g. Gold) are found in there native state
Most are found as oxides or compounds
Metals less reactive than carbon can be extracted by heating with carbon
e.g.
iron oxide+ carbon-> iron + carbon dioxide
Iron loses oxygen so is reduced
Carbon gains oxygen so is oxidised
4.1.1 Metal Oxides
Metals react with oxygen to produce metal oxides.
4.1.4 Oxidation and Reduction in terms of electron (HT Only)
Oxidation
Substance
GAINS
oxygen
Metals react with oxygen to form
metal oxides
Reduction
Substance
LOSES
oxygen
Always occur together
4.3 Electrolysis
4.3.3 Using electrolysis to extract metals
Metals more reactive than carbon can be extracted from their ores using electrolysis
Electrolyis requires a lot of heat and energy so is
expensive
Aluminium
Obtained by the electrolysis of aluminium oxide
Aluminium oxide mixed with cryolite before to lower melting point to 1000°c
Aluminium forms at Cathode
Cathode:
2O2-> O2+4e-
Oxygen gas forms at Anode and reacts with carbon to produce CO2
This wears away the Anode- needs to be replaced regularly
Anode:
Al3+3e- -> Al
Overall:
2Al2O3->4Al+3O2
Aluminium oxide->auminium+oxygen
Step 1
Steel tank lined with graphite- Lining= -ve electrode. Graphite blocks hanging over tank +ve electrode
Step 2
Aluminium dissolved in molten cryolite
Al3+ and O2- ions free to move
Step 3
Current switched on molten aluminium forms at -ve electrode (bottom of tank)Gain electrons and become aluminium atoms- Reduction
Oxygen gas forms at +ve electrode- lose electrons and become atoms- Oxidation
4.3.4 Electrolysis of aqueous solutions
Water molecules break down to form OH- and H+ ions
Cathode
Hydrogen produced if metal
more reactive
than hydrogen
Metal produced if metal
less reactive
than hydrogen
Anode
Oxygen produced unless solution contains halide ions
Halide ions present , halogen produced
Electrolysis of sodium chloride solution
Hydrogen releases at negative electrode
Chlorine gas released at positive electrode
4.3.5 Representation of reactions at electrodes as half equations (HT Only)
At cathode
Positively charged ions gain electrons
Reduction reactions
At anode
Negatively charged ions lose electrons
Oxidation reactions
Example
Lead Bromine
Anode
2Br->Br2+2e-
Cathode
Pb2+2e- ->Pb
4.3.1 The process of electrolysis 4.3.2 Electrolysis of molten ionic compounds
The use of an electrical current to break down compounds containing ions into their elements
Electrolyte
When an ionic compound is melted or dissolved using inert (Chemically inactive) electrons in water, ions are free to move
They can conduct electricity
Passing a current through it causes ions to move towards electrodes
Electrodes
Cathode
Positive ions (Cations) move here - Metals
Negative
Anode
Negative ions (Anions) move here- Non-metals
Positive
Made from solids that conduct electricity
4.2 Reaction of acids
4.2.3 Soluble salts
Can be made by reacting acids with insoluble bases (e.g. metal oxides, metal hydroxides and metal carbonates)
Method
(1) Add the metal oxid or carbonate to the acid in excess
(2) Filter excess oxide or carbonate to leave salt solution
(3) Warm the salt solution so the water evaporates leaving chrystals of salt
4.2.4 The pH scale
Hydroxide ions (OH-) make solutions alkaline
Hydrogen ions (H+) make solutions acidic
Acids and Alkalis
pH less than 7 acidic- closer to zero=stronger
pH more than 7 alkaline- closer to 14= stronger
pH 7= neutral
Measure of the acidity or alkalinity of an aqueous solution
Measured using a pH probe or universal indicator
Indicators change colour depending or alkaline or acidic solutions
4.2.2 Neutralisation of acids and salt production
Acids neutralised by bases (alkalis)
Can also be neutralised by
metal oxides -
acid+ meatal oxide-> salt+water
metal carbonates-
acid + metal carbonate-> salt + water + carbon dioxide
Salt produced depends on acid used
Hydrochloric acid- Chloride salts
Nitric acid- nitrate salts
Sulfuric acids- sulphate salts
acid + metal hydroxide -> salt + water
When an acid reacts with an alkali the H+ and OH- ions react to make water H2O (pH 7)
Called neutralisation because (1) Acid neutralised by alkali (2) Remaining solution pH 7 (Neutral)
4.2.5 Titrations (Chemistry Only)
Method
(2) Place the conical flask on white tile so the colour can be seen clearly
(3) Place the acid in a burette and take a reading of the volume
(1) Measure a known volume of alkali using a pipette and place it in a conical flask and add a suitable indicator- phenolphthalein
(4)Carefully add the acid to the burette swirling the flask to thoroughly mix, continue until the indicator changes colour, this is the end point
(5)Take a reading of the volume of acid left in the burette and calculate the volume of acid added
Calculations
(1) Write down balanced equation- determines number of moles
(2) Calculate number of moles in the solution of known volume and concentration.
(3) Calculate concentration of other solution
Example
A titration is carried out and 0.04dm3 hydrochloric acid neutralises 0.08dm3 sodium hydroxide of concentration 1.00mol/dm3
(1) Calculate the concentraion of the hydrochloric acid: HCl + NaOH -> NaCl + H2O
(2) number of moles of NaOH= Volume x concentration = 0.08dm3 x 1.00mol/dm3= 0.08mol
(3) concentration of HCl= Number of moles of HCl/ Volume of HCl= 0.08mol/0.04dm3=
2.00mol/dm3
4.2.1 Reactions of atoms with metals
Acid react with some metals to produce salts and hydrogen
4.2.6 Strong and Weak Acids (HT Only)
Strong acids
Completely ionised in aqueous solutions
e.g. Hydrochloric Acid, Nitric Acis and Sulphuric Acid
Lower pH
Weak acids
Partially ionised in aqueous solutions
e.g. Ethanoic acid, citric acid and carbonic acid
Higher pH
A pH decrease of one indicates that the concentration of hydrogen ions as increased by a factor of 10
Dilute
Reducing the concentration of a solute in solution, usually simply by mixing with more solvent.
Concentrated
An acid with a concentration of 2mol/dm3 is more concentrated that an acid with a concentration of 0.5mol/dm3
