Presentation is loading. Please wait.

Presentation is loading. Please wait.

Higher Unit 3 Electrolysis. After today’s lesson you should be able to:  Use ‘Q = I x t’ to calculate - the value of faraday - the mass of product at.

Similar presentations


Presentation on theme: "Higher Unit 3 Electrolysis. After today’s lesson you should be able to:  Use ‘Q = I x t’ to calculate - the value of faraday - the mass of product at."— Presentation transcript:

1 Higher Unit 3 Electrolysis

2 After today’s lesson you should be able to:  Use ‘Q = I x t’ to calculate - the value of faraday - the mass of product at an electrode - the time or current used in the experiment - the volume of gas given off at an electrode - the charge on an ion

3 Electrolysis  Electrolysis is a chemical reaction which occurs when an ionic compound is broken down into its elements using electricity.  The ionic compound must be molten or in solution.  A d.c. supply is used to separate the products to different electrodes.  Positive ions are attracted to the negative electrode where they undergo reduction.  Negative ions are attracted to the positive electrode where they undergo oxidation.

4  The current, symbol ‘I’, is the speed at which the electrical current flows round the circuit.  Current is measured in Amperes (A).  The current is kept constant by using a variable resistor in the circuit.  The electrical charge, symbol ‘Q’ is measured in Coulombs (C).  The time the current has been flowing round the circuit, symbol ‘t’, is measured in seconds (s).  The total charge is calculated using the equation Q = It

5 Electrons and Faradays  The charge is carried through the wires of the circuit (the external circuit) by electrons.  1 mole of electrons = 1 Faraday = 96500C.  The ion-electron equation for the element being calculated gives the number of moles of electrons and hence, the number of Faradays required to deposit 1 mole of the element.  e.g. Cu e - → Cu. 1 mole of copper is deposited using 2 moles of electrons = 2 x 96500C = C.

6 Example 1 – Calculating the Faraday experimentally  A solution of copper(II) nitrate was electrolysed for 20 mins using a current of 0.2A. A mass of 0.08g of copper was deposited on the negative electrode. Calculate the value of the Faraday.

7 Step 1 – calculate the total charge using the equation ‘Q = It’ Q = ? I = 0.2A t = 20 mins = (20 x 60)s = 1200s Q = 0.2 x 1200 = 240C

8 Step 2 – calculate the number of moles of the element deposited using the equation n = m ÷ gfm. n = ? m = 0.08g gfm Cu = 63.5g n = 0.08 ÷ 63.5 = mol

9 Step 3 – calculate the value of Faraday mol Cu is deposited by 240C ∴ 1 mol Cu → 240 = C

10 From the equation Cu e - → Cu 1 mole of Cu is deposited using 2 moles of electrons but 1 mole of electrons = 1 Faraday ∴ Faraday = = 95238C 2

11 Exercise P118 of ‘Test your Higher Chemistry Calculations’ Q15.1 – 15.5

12 Example 2 – calculating the mass deposited What mass of nickel is deposited in the electrolysis of nickel(II) sulphate solution if a current of 0.4A is passed for 120mins.

13 Step 1 – calculate the total charge using the equation ‘Q = It’ Q = ? I = 0.4A t = 120 mins = (120 x 60)s = 7200s Q = 0.4 x 7200 = 2880C

14 Step 2 – calculate the number of moles of electrons that this charge represents C →1 mole ∴ 1C → ∴ 2880C→ 1 x = 0.03 mol

15 Step 3 – calculate the mass of the element deposited From the equation Ni e - → Ni 1 mole of Ni is deposited using 2 moles of electrons i.e. 2 moles e - → 1mole Ni 2 mol e - → 58.7g Ni 1mol e - → mol e - → 58.7 x 0.03 = g 2

16 Exercise P of ‘Test your Higher Chemistry Calculations’ Q15.21 – 15.35

17 Example 3 How long must a current of 0.25A flow in the electrolysis of molten aluminium oxide to cause the deposition of 1.08g of aluminium at the negative electrode?

18 Step 1 – calculate the number of moles of element deposited n = ? m = 1.08g gfm = 27.0g n = 1.08 ÷ 27.0 = 0.04 mol

19 Step 2 – calculate the number of moles of electrons involved. From the equation Al e - → Al 1 mole of Al is deposited using 3 moles of electrons i.e. 1mole Al → 3 moles e - 1mole Al → 3 moles e mol Al → 0.04 x 3 = 0.12mol

20 Step 3 – calculate the charge this represents 1 moles e - → 96500C 0.12 mol e - → x 0.12 = 11580C = 11580C Step 4 – calculate t. t = Q I = = 46320s OR 772mins = = 46320s OR 772mins

21 Exercise P of ‘Test your Higher Chemistry Calculations’ Q15.46, , – 15.60

22 Example 4 What volume of hydrogen would be given off at the negative electrode in the electrolysis of an aqueous solution if a current of 0.25A flowed for 2hrs. (molar volume of hydrogen = 24 l mol -1 )

23 Step 1 – calculate the total charge using the equation ‘Q = It’ Q = ? I = 0.25A t = 2hrs = (2 x 60 x 60)s = 7200s Q = 0.25 x 7200 = 1800C

24 Step 2 – calculate the number of moles of the electrons used C →1 mole ∴ 1C → ∴ 1800C→ 1 x = mol

25 Step 3 – calculate the number of moles of hydrogen given off From the equation 2H + + 2e - → H 2 1 mole of H 2 is deposited using 2 moles of electrons i.e. 2 moles e - → 1mole H 2 2 moles e - → 1mole H 2 1 mole e - → ½ mole H 2 1 mole e - → ½ mole H moles e - → X ½ = mole H 2 = mole H 2

26 Step 4 – calculate the volume of hydrogen given off 1 moles H 2 → 24L 1 moles H 2 → 24L mole H 2 → x mole H 2 → x L OR 228cm L OR 228cm 3

27 Exercise P of ‘Test your Higher Chemistry Calculations’ Q15.36 – 15.45

28 Example 5 A molten iron compound is electrolysed using a current of 4.73A for 30mins during which 1.64g of iron is deposited. Calculate the number of positive charges on each iron ion.

29 Step 1 – calculate the total charge using the equation ‘Q = It’ Q = ? I = 4.73A t = 30 = (30 x 60)s = 1800s Q = 4.73 x 1800 = 8514C

30 Step 2 – calculate the number of moles of the electrons used C →1 mole ∴ 1C → ∴ 8514C→ 1 x = mol

31 Step 3 – calculate the number of moles of iron deposited n = ? n = ? m = 1.64g gfm Fe = 55.8g n = 1.64 ÷ 55.8 = 0.03mol

32 Step 3 – calculate the charge on the ion 0.03 mol Fe is deposited by 0.088mol e - ∴ 1 mol Fe → = 2.93 ≃ 3 mol = 2.93 ≃ 3 mol The ion-electron equation must be: Fe e - → Fe

33 Exercise P133 of ‘Test your Higher Chemistry Calculations’ Q15.61 – 15.65


Download ppt "Higher Unit 3 Electrolysis. After today’s lesson you should be able to:  Use ‘Q = I x t’ to calculate - the value of faraday - the mass of product at."

Similar presentations


Ads by Google