ELECTROLYSIS The Divorcing of a Compound PART ONE

ELECTROLYSIS The Divorcing of a Compound PART ONE ADEng. PROGRAMME Chemistry for Engineers Department of Pure and Applied Sciences Portmore Community College Prepared by M. J. McNeil, MPhil. 1

SPECIFIC OBJECTIVES I 2

SPECIFIC OBJECTIVES II 3

WHAT DO ALL THE ITEMS IN THE PHOTOGRAPH HAVE IN COMMON? 4

WHAT DO ALL THE ITEMS IN THE PHOTOGRAPH HAVE IN COMMON? The items could be solid silver, bronze or chrome. They are only plated - coated with a thin layer of the valuable metal. They have silver s, bronze s, etc dull shine and color, but are less costly to make and buy. What process binds these metals to a surface? How is the process related to the batteries in a radio? 5

WHAT IS ELECTROLYSIS? Electrolysis is a chemical process in which bonded elements and compounds are decomposed by passing an electric current through them. Any directed flow of charged particles - electrons or ions is an electric current. Two common examples of electrolysis: - changing of an automobile battery - silver plating a metal The electrolysis of water: 2H 2 O + energy 2H 2 + 2O 2 What in this room is a product of electrolysis? 6

ELECTRICAL CONDUCTION CONDUCTORS A substance that permits the passing of an electric current through it. E.g. all metals, graphite (carbon) and solutions of ionic and polar substances, e.g. iron, dilute NaCl solution and dilute H 2 SO 4. What are semi-conductors? NON-CONDUCTORS (INSULATORS) Do not allow the passage of electricity. E.g. non-metals and plastics. Classify each of the following as conductors and nonconductors: kerosene, Oil, iron, sulphur paper. 7

ELECTRICAL CONDUCTION METALLIC CONDUCTION Metallic conduction pertains the flow of a sea of free (mobile) electrons through a metal. All metals (and graphite) are classified as electrical conductors. They can conduct in the solid and liquid states. The metal remains chemically unchanged. ELECTROLYTIC CONDUCTION Ionic substances do not conduct electric current in the solid phase. They conduct only when molten or in aqueous solution eg. molten PbBr 2 and dilute H 2 SO 4, respectively. Mobile ions move randomly and move in one direction if attached to a power source (electric current). Chemical changes occur when they conduct an electric current. 8

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ELECTROLYTES Strong electrolytes are substances which are completely ionised when dissolved or molten. Electrolytes are molten substances or solutions that allows electricity to pass through causing it to decompose, e.g. molten salts, aqueous solutions of acids, alkali and salts. There are two types of electrolytes based on the extent of ionisation and conductivity when in the liquid state. Eg. Strong acids, salts such as NaCl. They have strong conductivity. NaCl (s) Na + (aq) + Cl - (aq) Weak electrolytes are substances which are partially ionised when dissolved or molten. That is, they contain few ions. E.g. weak acids, weak alkali and pure water. CH 3 COOH (s) CH 3 COO - (aq) + H + (aq) (vinegar/acetic acid) What are non-electrolytes and e.g.s? wax, glucose, molten covalent substances. 10

TYPES OF CELLS 11

TYPES OF ELECTROCHEMICAL CELLS Electrolytic Cell - Electrical energy is used to drive a nonspontaneous redox reaction. Voltaic/Galvanic Cell - Energy released from spontaneous redox reaction can be transformed into electrical energy (part two) 12

THE ELECTROLYTIC CELL Cells are containers of liquid with electrodes: Cell + Source or use of electricity Electrode Molten or aqueous chemicals In electrolytic cells, electricity is used to force chemicals 13 to undergo a redox reaction

SCHEMATIC OF ELECROLYIC CELL (I) 14

THE ELECTROLYTIC CELL Electrolysis involves the use of electrodes, which are pieces of metals connected to a battery that carry current into and out of the electrolyte. There are two types of electrodes: Anode: positive electrode connected to the (+) terminal of battery Cathode: negative electrode connected to the ( ) terminal of battery. The electrolyte contains: cations and anions. When the circuit is closed, the electrons in the battery and wires begin to move, causing the anode to acquire a positive charge and the cathode to acquire a negative charge. During electrolysis, the following occurs in the solution: Anions (-) are attracted and move towards the anode (+) terminal where they lose electrons to become neutral atoms or molecules. Cations (+) are attracted and move towards the cathode (-) terminatl where they gain electrons and become neutral atoms or molecules. Electrons flow anode to cathode and electric current flows in the opposite direction. 15

ELECTROLYSIS The end result of electrolysis is the breakdown of a compound into its constituent elements which are discharged or liberated at the electrodes. Discharge is the process by which ions gain or lose electrons and become atoms or molecules. Ions lose their charge during discharge. 16

ELECTROLYSIS OF MOLTEN ELECTROLYES 17

ELECTROLYSIS OF AQUEOUS SOLUTIONS An aqueous solution usually contains at least two different cations and two different anions due to the presence of H + and OH - ions from the water. However, only one type of ion of each charge is usually discharged at each electrode. Three factors determine which ions are preferentially dicharged (1) The position of the ions in the electrochemical series. (See mnemonic in course notes). The lower an ion in the electrochemical series, the more likely it is to be discharged from solution. (2) Relative concentration of the ions The greater the concentration of an ion, the more likely it is to be discharged, especially solutions containing halide ions (Cl -, Br -, I - ) e.g electrolysis of conc. HCl and NaCl solution. (3) Type of electrode - only non-inert (active) electrodes participates in electrolysis, e.g. electrolysis of CuSO 4 using copper anode; electrolysis of sodium chloride solution using mercury 18 (Hg) electrode.

ELECTROLYSIS OF CONCENTRATED SODIUM CHLORIDE 19

ELECTROLYSIS OF WATER The electrolysis of water produces hydrogen gas at the cathode (on the right) and oxygen gas at the anode (on the left). 20

ELECTROLYSIS OF SOME AQUEOUS SOLUTIONS 21

ELECTROLYSIS OF SOME AQUEOUS SOLUTIONS 22

PRODUCTS OF COMMON e.g. of ELECTRLOYSIS 23

QUANTITATIVE ELECTROLYSIS - MICHAEL FARADAY FIRST LAW Faraday s law of electrolysis pertains the quantity of an element measured in moles, formed during electrolysis. Faraday s first law of electrolysis: The mass of a substance produced at, or dissolved from, an electrode during electrolysis is directly proportional to the quantity of electricity passing through the electrolyte. 24

QUANTITAIVE ELECTROLYSIS- MICHAEL FARADAY The quantity of electricity depends on the: - length of time a steady current is flowing through the electrolyte - measured in seconds. - size of that steady current measured in amperes (A or amps). - The charge on the ion. Quantity of electricity (Q) = current (A) time (s) x Q (in coloumbs) = I x t. That is, one coloumb is one ampere flowing for one sec. 25

QUANTITAIVE ELECTROLYSIS- MICHAEL FARADAY 96 500 C = 1 Faraday. When 96 500 C of electricity have flowed through a circuit, one mole of electrons or 6.2 x 10 23 electrons have flowed. Since one mole of electrons is required to discharge one mole of an ion with a single charge. A - (1 mole) - e- A (1 mol) or C+ ( 1 mol) + e- C (1 mol) This indicates that one faraday is the quantity of electricity required to discharge one mole of an ion with a single charge. 96 500 Cmol -1 is called the FARADAY CONSTANT.****** 26

FARADAY S CALCULATIONS (1 st LAW) Calculate the quantity of electricity passed by: (a) a current of 0.5 A flowing for 10 mins. Ans. = 300 Coloumbs (C). (b) a current of 10 A flowing for 2 hours. Ans. = 72 000 C. Use Q = I x t 27

TESTING UNDERSTANDING I 28

TESTING UNDERSTANDING II 29

ELECTROLYSIS CALCULATIONS 1. Calculate the mass of magnesium liberated during the electrolysis of fused (molten) MgCl 2 if a current of 1.93 A is passed for 16 mins, 40 sec. 2. Calculate the mass of silver and the volume of oxygen (s.t.p) liberated during electrolysis by 9650 C (Ag = 108); (1 mol of gas at s.t.p. = 22.4 dm 3.) 30

SPECIFIC OBJECTIVES III 31

APPLICATIONS OF ELECTROLYSIS Electrolysis has many crucial applications in industry. Some of these are: Metal/Non-metal extraction Electrorefining of metals Electroplating Anodising etc. 32

METAL/NON-METAL EXTRACTION Many elements are made by electrolysis. Pb Al Zn Na K Li H 2 Cl 2 F 2 I 2 O 2 Pb 2+ + 2e - Pb (s) 2Cl - Cl 2(g) + 2e - This is sometimes called electrowinning- the element is won from its ion. The metallic compound along with other impurities is referred to as the ore of the metal. For e.g. bauxite is the ore for aluminum, containing aluminum oxide with silicon and iron(iii) oxide among the impurities. 33

METAL EXTRACTION & USES K Na Li Ca Mg Al C Zn Fe Pb H Cu Ag Au REACTIVITY SERIES Metals above carbon require the extraction by electrolysis of their molten compounds. Metals below carbon in the reactivity series can be extracted from their ores (oxides, carbonates and sulphides) using carbon in a method called chemical processes. Cu, Ag and Au found free in nature. Copper is extracted by displacement or by electrolysis of aqueous solution, crushing and washing. You are required to know: Sodium from sodium chloride (Down s cell) Aluminum from aluminum oxide (bauxite) 34

MNEMONIC FOR ACTIVITY SERIES 35

ELECTROLYSIS OF ALUMINUM OXIDE (ALUMINA) Most abundant metallic element in the earth s crust (8%). Found in minerals, bauxite, crylolite, mica and clay. Al is a very reactive metal and is thus very difficult to extract from its ore. Metals bind tightly to their ores, but can be extracted from them by electrolysis. Bauxite Mining in Australia 36

EXTRACTION OF ALUMINA FROM BAUXITE Ore is bauxite = hydrated aluminum oxide (Al 2 O 3.2H 2 O) (i) The mined bauxite is first treated with caustic soda (NaOH) solution to remove impurities of silica (SiO 2 ) and haematite (Fe 2 O 3 ). (ii) Aluminum oxide being amphoteric dissolves in the alkali leaving behind solid impurities which are filtered off. (iii) The resulting solution of sodium aluminate is allowed to crystallize then the crystals are washed. (iv)the anhydrous oxide, alumina (Al 2 O 3 ) is dissolved in molten cryolite (3NaF. AlF 3 ) which lowers the melting point of the electrolyte from 2050 to 900 o C. (v) Oxygen is evolved at the carbon anodes which slowly burns away to form CO 2. (vi) Molten aluminum is collected on the floor of the cell which is lined with graphite cathode. Cathode: 4Al 3+ + 12e- Anode : 6O 2 - -12e - 3O 2 4Al 37

EXTRACTION OF ALUMINA FROM BAUXITE 38

ELECROREFINING Most metals are too reactive to be found free (uncombined) in nature. The metals that are found free in nature are the unreactive ones - copper, silver and gold. All the other metals have to be extracted from their compounds. Electrolytic or electrorefining is a process of obtaining pure metal by the process of electrolysis to render them more useful. In this method, an electric current is passed through the electrolytic solution that contains dissolved metals. This causes the dissolved metals to deposit onto the negatively charged cathode. 39

ELECTROREFINING OF COPPER Anode half-reaction: Cu(s) -2e - Cu 2+ (aq) Copper atoms leave the anode and enter the skeleton copper ions. Cathode half-reaction: Cu 2+ (aq) + 2e- Cu(s). Electrolyte: mixture of CuSO 4 and H 2 SO 4 Overall, Cu leaves the anode and is deposited on the cathode. 40

ELECTROREFINING OF COPPER 41

NICKEL (Ni) PLATING In nickel plating, the mixtures correspond to the plated intended articles is to be used. Plating solution: Ni(II) sulphate or Ni(II) chloride solution to which boric acid and a chemical wetting agent is applied. Anode: Ni Cathode: the metal to be plated or coated. Nickel-plated objects are then often plated with chrome or silver. 42

SILVER (Ag) PLATING Nickel objects are often coated with silver which make the object more expensive than they really appear. Anode: Ag Cathode: the metal to be plated or coated. 43

+ve -ve stainless steel or Au AuCN 44

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Au plated 46

Copper Ring Gold Plated 47

ANODIZING Anodizing is a method of coating a metal with an oxide layer to make them resistant to corrosion (rusting). E.g. Anodizing aluminum sheets by the electrolysis of dilute H 2 SO 4. Aluminum readily forms a coating of Al 2 O 3, which renders the metal unreactive and thus resistant to corrosion. Anodized aluminum sheets can be dyed or painted because the oxide layer is very absorbent. 48

ANODIZED CIRCULON POTS 49

USES OF ELECTROLYSIS ELECTROPLATING (i) The object to be plated is made the cathode. (ii) The object is plated with a metal. ANODISING (i) The object to be anodised is made the anode. (ii) The object anodised with an oxide layer. Advantage of recycling metals (a) Reduces the difficulty of disposing solid waste. (b) Saves on valuable resources of metal ores. N.B. It is not feasible to anodise iron as the coating forms rust and is not adhesive. 50

OTHER APPLICATIONS OF ph meters - ELECTROLYSIS 51

USE OF ELECTROLYSIS 52

THE PROCESS of hair removal 53

ELECTROCHEMISTRY PART TWO I. Galvanic or Voltaic Cells: a) Anode/Cathode/Salt Bridge b) Cell Notations c) Determining Cell Potential/Cell Voltage/Electromotive force (emf) II. Relating Cell Potential to Spontaneity III. Effect of Concentration on Cell Potential 54

GALVANIC OR VOLTAIC CELL Salt Bridge allows current to flow Connected this way the reaction starts. Stops immediately because charge builds up. Salt bridge/porous disk: allows for ion migration such that the solutions will remain neutral etc. 55

GALVANIC CELL In turns out that we still will not get electron flow in the example cell. This is because charge buildup results in truncation of the electron flow. We need to complete the circuit by allowing positive ions to flow as well. We do this using a salt bridge which will allow charge neutrality in each cell to be maintained. 56

SPONTANEOUS REDOX REACTION Zn(s) + Cu 2+ (aq) -> Cu(s) + Zn 2+ (aq) Zn Cu 2+ time Zn 2+ Cu 57

CELL POTENTIAL Oxidizing agent pulls the electron. Reducing agent pushes the electron. The push or pull ( driving force ) is called the cell potential E cell Also called the electromotive force (emf) Unit is the volt (V) = 1 joule of work/coulomb of charge Measured with a voltmeter. 58

STANDARD REDUCTION POTENTIAL TABLE (SEE DATA BOOKLET) Most reference tables for electrochemistry are written as reductions in a half-reaction format, with the most negative reduction on the top and the most positive on the bottom. 59

E 0 is for the reaction as written. The half-cell reactions are reversible. The sign of E 0 changes when the reaction is reversed. Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0 The more positive E 0 the greater the tendency for the substance to be reduced. 60

REDUCTION POTENTIAL More negative Eº more easily electron is added More easily reduced Better oxidizing agent More positive Eº more easily electron is lost More easily oxidized Better reducing agent 61

Gets Smaller -> <- Gets Larger 62

Voltmeter GALVANIC CELL Salt Bridge 63

ELECTROCHEMICAL CELL The difference in electrical potential between the anode and cathode is called: cell voltage electromotive force (emf) cell potential E 0 Cell E 0 oxidation E 0 reduction UNITS: Volts Volt (V) = Joule (J) Coulomb, C 64 19.2

RULES FOR USING STANDARD REDUCTION POTENTIALS 1) Read the half reactions as written. 2) The more POSITIVE the reduction potential, the greater the tendency is for the substance to be reduced and therefore the better the oxidizing agent. (keep the half-reaction with the more POSITIVE cell potential as written in the table - that is the reduction reaction) 3) The half-cell reactions ARE reversible. IF you need to reverse, you MUST change the sign of the ξ 0 cell. 4) If you change the stoichiometric coefficients, ξ 0 cell remains the same 65

PREDICTING GALVANIC CELLS Given two 1/2 cell reactions, how can one construct a galvanic cell? Need to compare the reduction potentials of the two half cells. Turn the reaction for the weaker reduction (smaller E 1/2 ) and turn it into an oxidation. This reaction will be the anode, the other the cathode. 66

PREDICTING GALVANIC CELLS Example. Describe a galvanic cell based on the following: Ag + + e - Ag E 1/2 = 0.80 V Fe +3 + e - Fe +2 E 1/2 = 0.77 V Weaker reducing agent turn it around Ag + + Fe +2 Ag + Fe +3 E cell = 0.03 V E cell > 0.cell is galvanic 67

EXAMPLE For the following reaction, identify the two half cells, and use these half cells to construct a galvanic cell. 3Fe 2+ (aq) Fe(s) + 2Fe 3+ (aq) +2 0 +3 oxidation reduction Fe 2+ (aq) + 2e- Fe(s) E = -0.44 V Fe 3+ (aq) + e- Fe 2+ (aq) E = +0.77 V 68

EXAMPLE weaker reduction - turn it around OR?? Fe 2+ (aq) + 2e- Fe(s) E = -0.44 V 2 x Fe 3+ (aq) + e- Fe 2+ (aq) E = +0.77 V 2Fe 3+ (aq) + Fe(s) 3Fe 2+ (aq) E cell = 1.21 V Fe(s) Fe 2+ (aq) + 2e- E = +0.44 V 69

PRACTICE Completely describe the galvanic cell based on the following half-reactions under standard conditions. MnO 4 - + 8 H + +5e - Mn +2 + 4H 2 O; Eº=1.51 V Fe +3 +3e - Fe(s) Eº=0.036V 70

DETERMINING IF REDOX REACTION IS SPONTANEOUS + E cell ; spontaneous reaction E cell = 0; equilibrium More positive E cell ; stronger oxidizing agent or more likely to be reduced - E cell ; nonspontaneous reaction 71

CELL NOTATION 1. Anode 2. Salt Bridge 3. Cathode Anode Salt Bridge Cathode : symbol is used whenever there is a different phase 72

Cell Notation Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M More detail.. Zn (s) Zn 2+ (1 M) Cu 2+ (1 M) Cu (s) anode cathode Zn (s) Zn +2 (aq, 1M) K(NO 3 ) (saturated) Cu +2 (aq, 1M) Cu(s) anode Salt bridge cathode 19.2 73

WRITING GALVANIC CELL Zn Zn 2+ + 2e- E = +0.76 V Cu 2+ + 2e - Cu E = 0.34 V Cu 2+ + Zn Cu + Zn +2 E cell = 1.10 V Notice, we reverse the potential for the anode. E cell = E cathode - E anode 74

WRITING GALVANIC CELL Shorthand Notation Zn Zn 2+ Cu 2+ Cu Anode Cathode Salt bridge 75

CELL POTENTIAL MEASUREMENTS We can measure the magnitude of the EMF causing electron (i.e., current) flow by measuring the voltage. e - Anode Cathode 76

HALF-CELL POTENTIALS What we seek is a way to predict what the voltage will be between two 1/2 cells without having to measure every possible combination. To accomplish this, what we need to is to know what the inherent potential (data booklet) for each 1/2 cell is. The above statement requires that we have a reference to use in comparing 1/2 cells. That reference is the standard hydrogen electrode (SHE) 77

STANDARD REDUCTION POTENTIALS Most reference tables for electrochemistry are written as reductions in a half-reaction format, with the most negative reduction on the top and the most positive on the bottom. By definition, all half reductions are compared to the hydrogen half reaction that has the standard value of 0.00 V under standard conditions. Standard state: 25 o C, 1 atm 2H + + 2 e- H 2 ξ 0 cell = 0.00 All other reduction potentials are based on this zero point. 78

STANDARD ELECTRODE POTENTIALS Standard reduction potential (E 0 ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. Reduction Reaction 2e - 2H (1 M) H 2 (1 atm) E 0 0 V Standard hydrogen electrode (SHE) 79 19.3

Zn (s) + 2 H + (aq) -> H 2 (g) + Zn 2+ (aq) K(NO 3 ) Zn(s) Zn +2 KNO 3 H + (aq) H 2 (g) Pt 80

HALF-CELL POTENTIALS With our zero we can then measure the voltages of other 1/2 cells. In our example, Zn/Zn +2 is the anode: oxidation Zn Zn 2+ + 2e - 2H + + 2e - H 2 E Zn/Zn2+ = 0.76 V E SHE = 0 V Zn + 2H + Zn 2+ + H 2 E cell = E SHE + E Zn/Zn+2 = 0.76 V 0 81

PREDICT TITLE HERE 82

CORROSION METALS BY ELECTROCHEMICAL PROCESS 83

RUST FORMATION 84

CATHODIC PROTECTION Iron Nail with Zinc 85

TESTING KNOWLEDGE 86