Redox Reactions and Electrolysis PDF

Redox and electrolysis ​ Redox us short for ‘ reduction - oxidation’ ​ Electrolytic cell change electrical energy into chemical energy Uses -​ to extract reactive metals from their metal ores -​ Electroplating ( coating one metal with thin layer of other metal)...

Redox and electrolysis ​ Redox us short for ‘ reduction - oxidation’ ​ Electrolytic cell change electrical energy into chemical energy Uses -​ to extract reactive metals from their metal ores -​ Electroplating ( coating one metal with thin layer of other metal) -​ In refining ( purify ) copper -​ Electrolysis of brine Electrolysis - ending with -lysis : breaking down Redox reactions Reaction involving simultaneous oxidation and reduction -​ If a substance is reduced then another substance must oxidised Oxidation - gain of oxygen Reduction - loss of oxygen Eg. reduction of copper oxide with carbon 2CuO (s) + C(s) → 2Cu(s) + CO2 (g) Copper oxide - reduced Carbon - oxidised Roman numerals - uses to indicate oxidation number of an element eg. copper(ll)  ​ Not all redox reaction can be explain by the gain or loss of oxygen Eg. displacement of copper by magnesium Mg (s) + CuSO4(aq) → Cu(s) + MgSO4(aq) Magnesium - oxidised Copper - reduced Redox reaction are due to transfer of electrons in chemical reaction ​ Oxidation - loss of electrons ​ Reduction - gain of electrons Oxidation state - number assigned to an element showing number of electrons lost or gained in ions -​ Shows relative state of how much have been oxidised or reduce ( oxidation - positive) (reduction - negative) Eg. iron(ll) chloride - oxidation state of +2 iron(lll) chloride - oxidation state of +3 So iron(lll) chloride is more oxidised than iron (ll) chloride) ​ Oxidation cause an increase in oxidation state ​ Reduction causes a decrease in oxidation state Working with oxidation state ​ An element always have an oxidation state of 0 ​ For compounds ( CuO) the sum of oxidation state is 0 ​ For ions charge on the ion must equal to the sum of the oxidation state -​ Oxidation number of monatomic ion is the same as charge on the ion Eg. CuO Oxidation state : Oxygen : -2 Copper: +2 Sum of oxidation in compound = 0 Harder work example Show that manganese (VII) ion, MnO4- ( complex ions : contain charge ) has an oxidation state of +7 in the manganese ions Charge on oxygen : 4 4 x -2 ( oxygen’s oxidising state) = -8 Over all charge of ion is - (-1) Difference between -8 and -1 = 7 +7 is the oxidation state of manganese ( if there were two eg. Mn2 divide the difference by 2) During a reaction : ​ Oxidation cause an increase in oxidation state ​ Reduction causes a decrease in oxidation state Eg, MnO4- → Mn2+ +7 → +2 : change in oxidation state show that it is being reduced 2I- (iodine ions) → I2(iodine ) -1 → 0 : increase in oxidation number - is being oxidised Writing redox half equation 1.​ Using oxidation state, work out which substance is being oxidised which is being reduced 2.​ Construct a simple equation showing the change in each substance 3.​ Balance both equation ( atom) 4.​ Balance the reduction equation with electron of the left hand side Balance the oxidation equation with electron on the right hand side Eg. Mg(s) + CuSO4(aq) → Cu(s) + MgSO4 (aq) Mg is forming a compound so ion Mg2+ 0 → 2 : increase = oxidised Copper ( Cu2+) form an element 2→ 0 : is being reduced The 2+ charge can be obtain by knowing SO4^2-) Cu2+ + 2e- → Cu Mg → Mg2+ + 2e-  Reduction and oxidising agents ​ Oxidising agent : substance that oxidises another substance during a redox reaction, while itself is reduced ​ Reducing agent : is a substance that reduces another substance while it self oxidised Eg. Copper oxide + carbon → copper + carbon dioxide Copper oxide is being reduce : oxidising agent Carbon is being oxidised : reducing agent Potassium manganate (VII) and potassium iodine Potassium manganate (VII) : deep purple colour 7+ oxidation state Colorless when reduced to Mn+ Potassium iodine : colourless solution which can be oxidised to form brown iodine Electrolysis Principle of electrolysis ​ The breakdown of an ionic compound , molten or in aqueous solution by the passage of electricity -​ Carry out in a electrolytic cell 3 main component : ​ Battery ​ Two electrode -​ Connected to the negative terminal : cathode -​ Connected to the positive terminal : anode ​ An electrolyte -​ Aqueous or molten substance that conducts electricity -​ Is the substance being broken down ( usually ionic compound) -​ Must be molten or in aq solution - ion can move to opposite charge electrode Simple electrolytic cell  ​ Electrode often made out of -​ Carbon/graphite -​ Platinum Are inert ( unreactive ) and good conductor of electricity Main stages of electrolysis 1.​ Electrolyte is made molten or dissolve in water to give aqueous solution 2.​ Electrical current is pass through the cell 3.​ Cations move toward the negatively charge cathode -​ Hydrogen or metal ions 4.​ Anion moves towards positively charge anode -​ Anion Eg. electrolysis of NaCl Na+→ cathode → Na 2Cl- → anode → Cl2 PANIC - positive is anode, negative is cathode Reaction at electrode ( redox Electron flows from anode to cathode ( positive → negative) 1.​ Cathode Attract cation → move toward cathode -​ Causes reduction : gain electron 2.​ Anode attract anion → move toward anode -​ Cause oxidation : loses electrons Electrolysis of molten electrolytes electrolyte Produce at Produce at observations cathode (-) anode (+) Molten lead(ll) bromine Lead Bromine Silvery solid at cathode Brown gas at anode Pb2+(l) → 2e- → 2Br-(aq) → Pbs Br2+(g) + 2e- Concentrated aqueous hydrogen chlorine ​ Colourless gas at Sodium chloride ( brine) cathode which makes a splint light up and ‘ pop’ ​ Pale yellow green gas at anode - turn universal indicator red and then bleached to colorless ​ Also bleach red litmus paper - it is an acidic gas Diluted sulfuric acid hydrogen oxygen ​ Colorless gas at both electrodes ​ Cathode gas makes a lighted splint go ‘pop’ ​ Anode gas relight a glowing splint Molten only contains metal and non-metal ions -​ Product will always be metal (cathode) and nonmetal (anode) Direction of electron flow ( not current) And direction of ion flow Cathode loses electrons Anode gain electrons Electrolysis of aqueous electrolytes ​ More complex due to : -​ Ionisation : process of making ions by either gaining or losing electron -​ Dissociation : compounds splitting to form ions Of water ( H+) (OH-) Explaining with electrolysis of brine : Contains : Na+, Cl- ( from NaCl) and H+ and OH- (from dissociation of water) At cathode : -​ H+ ions better at gaining electron than Na+ ions -​ Hydrogen gas is form 2H+(aq) +2e- → H2(g) At anode : -​ Cl- ions are better at losing electrons so they will be oxidised in preference to OH- -​ Chlorine gas is form 2Cl-(aq) → Cl2(g) + 2e- ​ The Na+ and OH- stays in solution and form sodium hydroxide  ​ These are the basis of chlor-alkali industry Predicting the products ​ Product at cathode is based on reactivity series Less reactive metal ions ‘ leave’ the solution first so: ​ Cation below hydrogen -​ The metal will form ​ Cation above hydrogen -​ The H+ ions will form Eg. sodium chloride Cathode : H2 - sodium more reactive Anode : Cl2 ​ Product at anode is based on the electrochemical series -​ Electrochemical series : a list oh half equation showing the tendency of a reaction to want to be oxidised or reduced In general : If it is a halide ion : eg. Cl-, Br-, I- are present -​ Then the element will form Other negative anion eg. So4 2- or NO3- -​ Oxygen will form from dissociation of OH- ions -​ ( others are more reactive than OH- ions) Eg. copper sulfate Cathode : copper Anode : oxygen Electrolysis of copper sulfate ​ Product depends on the electrode used : copper and graphite Copper electrode is use for refining copper Graphite Cathode : copper Cu2+ +2e- → Cu Anode : oxygen 4OH- → 2H2O + O2 +4 e- Copper Cathode : copper Cu2+ +2e- → Cu Anode : copper Cu → Cu2+ + 2e- The effect of dilution If a diluted solution of aqueous sodium chloride solution electrolysed the product may change At cathode : hydrogen is still produce At anode : oxygen - due to having excess OH- Water molecules are presented in excess - decomposition of water leaving NaCl Application of electrolysis 1.​ Electroplating 2.​ Refining copper 3.​ Electrolysis of concentrated sodium chloride Electroplating ​ Coat one metal with another -​ Eg. coating nickel jewelry with gold/silver ( improve appearance) -​ Chromium plating on steel car bumpers ( protect from corrosion) Object to be electroplate : at cathode Coating layer metal : at anode Electrolyte : solution containing metal ion ( a salt ) of the metal providing electroplating layer Cathode : Ag+(aq) +e- → Ag(s) Anode : Ag(s) → Ag+(aq) + e- ​ Over time key coat with thin layer of silver ​ Anode get smaller ​ Silver anodes replace the silver ion in the electrolyte ( will be use up over time) Copper refining and purification ​ Impure copper - anode ​ Pure copper - cathode ​ Copper sulfate - electrolyte 1.​ At anode : copper loses electrons and go into solution as copper (ii) ions 2.​ The Cu2+ ions get attracted to the cathode ( is less reactive than H+) -​ The copper gain electron -​ A layer of pure copper build up at cathode 3.​ Anode slime ( impurities such as precious metals platinum, silver,gold) drop to the bottom 4.​ Anode loses mass, cathode gain mass Electrolysis of concentrated sodium chloride solution ​ Take place in a membrane cell Produce from mining underground rock salt deposit  ​ Product : hydrogen , chlorine, sodium hydroxide Uses: hydrogen Fuel cell, manufacturing of margarine and nylon. Manufacturing of hydrogen peroxide ( for ammonia) Chlorine Plastic, solvent, water purification, medical drugs Sodium hydroxide Soap and detergent, textile, paper, dye and medical drugs