Topic 4 Chemical Changes By Bethan Poole

Reactivity of metals

4.1.2 The Reactivity Series

4.1.3 Extraction of Metals and Reduction

4.1.1 Metal Oxides

4.1.4 Oxidation and Reduction in terms of electron (HT Only)

4.3 Electrolysis

4.2 Reaction of acids

4.2.3 Soluble salts

4.2.4 The pH scale

4.2.2 Neutralisation of acids and salt production

4.2.5 Titrations (Chemistry Only)

4.2.1 Reactions of atoms with metals

4.2.6 Strong and Weak Acids (HT Only)

4.3.3 Using electrolysis to extract metals

4.3.4 Electrolysis of aqueous solutions

4.3.5 Representation of reactions at electrodes as half equations (HT Only)

4.3.1 The process of electrolysis 4.3.2 Electrolysis of molten ionic compounds

Metals react with oxygen to produce metal oxides.

At cathode

The use of an electrical current to break down compounds containing ions into their elements

Electrolyte

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

Oxygen gas forms at Anode and reacts with carbon to produce CO2

This wears away the Anode- needs to be replaced regularly

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

Anode: Al3+3e- -> Al

Cathode: 2O2-> O2+4e-

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

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

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

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

Anode

Positive ions (Cations) move here - Metals

Negative ions (Anions) move here- Non-metals

Made from solids that conduct electricity

Oxidation

Reduction

Substance GAINS oxygen

Always occur together

Substance LOSES oxygen

Metals react with oxygen to form metal oxides

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

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

Calcium, magnesium, zinc and iron- fairly reactive, react quicky with acids and slowly with water

Not safe to react with dilute acids

Copper and gold- very unreactive, don't react with water or acids

Includes hydrogen and carbon for comparison

Strong acids

Completely ionised in aqueous solutions

e.g. Hydrochloric Acid, Nitric Acis and Sulphuric Acid

Weak acids

Partially ionised in aqueous solutions

e.g. Ethanoic acid, citric acid and carbonic acid

Lower pH

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

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

Measure of the acidity or alkalinity of an aqueous solution

Measured using a pH probe or universal indicator

pH 7= neutral

Indicators change colour depending or alkaline or acidic solutions

Acids neutralised by bases (alkalis)

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)

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 react with some metals to produce salts and hydrogen

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

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

Positive

Negative