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Let’s read! Pages 82 to 89. Objectives To know how to carry out electrolysis experiments. To work out what happens to ions at each electrode. To be able.

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Presentation on theme: "Let’s read! Pages 82 to 89. Objectives To know how to carry out electrolysis experiments. To work out what happens to ions at each electrode. To be able."— Presentation transcript:

1 Let’s read! Pages 82 to 89

2 Objectives To know how to carry out electrolysis experiments. To work out what happens to ions at each electrode. To be able to write successful half-equations. To predict the products of electrolysis. To do electrolysis calculations using the idea of a Faraday

3 Stick it in!

4 Electrolysis Lead bromide

5 Electrolysis Lead bromide Electrolysis is the break-down of a substance by electricity Electrolyte - a molten or aqueous solution through which an electrical current can flow. Copy please Electrode - an electrical conductor which carries charge to or from a liquid undergoing electrolysis.

6 Let’s try it!

7 Electrolysis experiments Electrolysis only happens in: - molten ionic liquids or - aqueous solutions containing ions. There must be a complete circuit. A lamp or ammeter shows that electricity is flowing around the circuit.

8 Moving charges Can you stick the sheet in your book?

9 Electrolysis of zinc chloride

10 1.Electrolytes contain positive and negative ions. 2.During electrolysis, positive and negative electrodes are put into the electrolyte. 3.The positive electrode is called the anode. 4.The negative electrode is called the cathode. 5.The negative ions (called anions) are attracted to the anode. 6.At the anode, the negative ions lose electrons to become atoms/molecules. 7.The positive ions (called cations) are attracted to the cathode. 8.At the cathode, the positive ions gain electrons to become atoms/molecules Draw the sentence

11 1.Electrolytes contain positive and negative ions. 2.During electrolysis, positive and negative electrodes are put into the electrolyte. 3.The positive electrode is called the anode. 4.The negative electrode is called the cathode. 5.The negative ions (called anions) are attracted to the anode. 6.At the anode, the negative ions lose electrons to become atoms/molecules. 7.The positive ions (called cations) are attracted to the cathode. 8.At the cathode, the positive ions gain electrons to become atoms/molecules Draw the sentence

12 At the electrodes Cathode (-) (negative electrode) Positive ions go here (cations). As metal ions are positive, they go to the cathode. Ions gain electrons. They are reduced and become neutral atoms. Anode (+) (positive electrode) Negative ions go here (anions). As non-metal ions are negative, they go to the anode. Ions lose electrons. They are oxidised and become neutral atoms (which react together to form molecules).

13 Common ions Li +, Na +, K +, Mg 2+, Ca 2+, Zn 2+ Cu +, Cu 2+, Fe 2+, Fe 3+, Al 3+ NH 4 + (ammonium ion) F -, Cl -, Br -, I - O 2-, S 2- OH - (hydroxide ion), CO 3 2-, NO 3 - (nitrate ion), SO 4 2- (sulphate ion)

14 Questions! Let’s try some easy questions

15 Half equations Show what happens at each electrode. Are balanced equations. Consider the electrolysis of copper chloride: Cu e -  Cu2Cl - - 2e -  Cl 2

16 Let’s read again! Page 92 to 95 (Chemistry for You)

17 Electrolysis of solutions – Cathode For solutions of highly reactive metals: Hydrogen gas, not the metal, is produced at the cathode.

18 Electrolysis of solutions – Anode The product at the anode depends on: The negative anions present in the solution.

19

20 Electrolysis of aluminium oxide CO 2

21 Electrolysis The Faraday

22 Objectives recall that one Faraday represents one mole of electrons calculate the amounts of the products of the electrolysis of molten salts and aqueous solutions

23 The Faraday A Faraday is one mole of electrons, and is equivalent to C (Coulombs) A current of 1A = 1C per second flowing For example, Cu e  Cu 1 mole of Cu 2+ ions reacts with 2 Faradays of electrons, to produce 1 mole of Cu

24 Quantity of electricity in coulombs = current in amps x time in seconds Q (C) = I (A) x t (s)

25 Example How much copper is deposited if a current of 0.2 Amps is passed for 2 hours through a copper(II) sulphate solution ?

26 current of 0.2 Amps is passed for 2 hours At the cathode: Cu 2+ (aq) + 2e-  Cu (s) Q = I x t = 0.2 x (2 x 60 x 60) =1440 Coulombs 1 mole electrons = Coulombs So, 1440 / = moles of electrons (Faradays)

27 Example moles of electrons passed through circuit = Cu 2+ (aq) + 2e-  Cu (s) From equation, it takes two moles of electrons to form one mole of copper moles copper = / 2 =

28 Example moles Cu = mass of Cu = moles x A r = x 64 = g of Cu deposited.

29 ‘How To” Guide 1.Write out relevant half equation 2.Work out coulombs of electrons flowing (Q = It) 3.Convert C into moles of electrons (Faradays) (Q/96500) 4.Work out moles of product using ratio from equation 5.Convert into mass (mass = moles x A r )

30 In an electrolysis of sodium chloride solution experiment a current of 2 A was passed for 2 minutes. –(a) Calculate the volume of chlorine gas produced. –(b) What volume of hydrogen would be formed? –(c) In practice the measured volume of chlorine can be less than the theoretical value. Why?

31 –Electrode equations: (-) cathode 2H + + 2e-  H 2 (+) anode 2Cl -  Cl 2 + 2e –(a) Calculate the volume of chlorine gas produced. Q = I x t, so Q = 2 x 2 x 60 = 240 C 240 C = 240 / = mol electrons this will produce / 2 = mol Cl 2 (two electrons/molecule) vol = mol x molar volume = x = 29.8 cm 3 of Cl 2

32 –(b) What volume of hydrogen would be formed? 29.8 cm 3 of H 2 because two electrons transferred per molecule, same as chlorine. –(c) In practice the measured volume of chlorine can be less than the theoretical value. Why? chlorine is moderately soluble in water and also reacts with the sodium hydroxide formed.

33 In the electrolysis of molten sodium chloride 60 cm 3 of chlorine was produced. –Calculate... –(a) how many moles of were chlorine produced? –(b) what mass of sodium would be formed? –(c) for how long would a current of 3 A in the electrolysis circuit have to flow to produce the 60cm 3 of chlorine?

34 (a) how many moles of chlorine produced? 60 / = mol Cl 2

35 (b) what mass of sodium would be formed? from the electrode equations 2 mol sodium will be made for every mole of chlorine so x 2 = mol sodium will be formed. A r (Na) = 23 mass = mol x atomic or formula mass = x 23 = 0.115g Na

36 (c) for how long would a current of 3 A in the electrolysis circuit have to flow to produce the 60cm 3 of chlorine? To produce mol of Cl 2 you need mol of electrons mol electrons = x coulombs = C Q = I x t, so = 2 x t, therefore t = / 3 = 161 s (to nearest second)

37 Brine? Brine is salty water (sodium chloride solution)

38 Electrolysis of brine

39 Hydrogen is produced at the cathode Chlorine is produced at the anode The solution remaining is sodium hydroxide

40 Electrolysis of brine Hydrogen is produced at the cathode Chlorine is produced at the anode The solution remaining is sodium hydroxide Cathy’s Ankles (CatHy’s AnCl)

41 Electrolysis of brine Hydrogen is produced at the cathode Chlorine is produced at the anode The solution remaining is sodium hydroxide Cathy’s Ankles (CatHy’s AnCl) Copy please!

42 Electrolysis of brine Cathode (-) 2H + (aq) + 2e - H 2(g) (SODIUM IS NOT FORMED (the sodium ion is more stable than the hydrogen ion in water H 2 OH + + OH - )) Anode (+) 2Cl - (aq) – 2e - Cl 2(g)

43 Electrolysis of brine Cathode (-) 2H + (aq) + 2e - H 2(g) (SODIUM IS NOT FORMED (the sodium ion is more stable than the hydrogen ion in water H 2 OH + + OH - )) Anode (+) 2Cl - (aq) – 2e - Cl 2(g) Copy please!

44 Chemicals from salt

45 Copy please!

46 Let’s try some questions!


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