Chapter 1 - Intoduction CHG8300F

Voltage and Current in Electrochemistry Electrochemistry deals with physically separated oxidation-reduction reactions. Such reactions take place in a cell as two half-reactions. – The force with which electrons travel from the oxidation half-reaction to the reduction half-reaction is measured as voltage . – The rate the electrons are being transferred with is measured as current .

6 State-of-the-art Intel, IBM, and AMD chips have 70 nm copper plated wires connecting a billion transistors! Electroplating

9 Li- metal and Li-ion batteries for Electric Vehicles and Plug-in Hybrid Electric Vehicles (PHEV) Normal hybrid vehicles can run on batteries up to 60 km/h and are charged by the Internal Combustion Engine when it is running. With plug-in hybrid cars, the cars can rely on the extra battery which has higher power and can get the car up to speeds of 100 km/h. This battery is then plugged in to recharge.

12 Corrosion: The opposite of electroplating What we dump into the atmosphere affects corrosion rates

15 Aluminum Production 2Al 2 O 3 (melt) + 3C(s) 4Al(l) + 3CO 2 (g) • Alumina Al 2 O 3 from bauxite (Al 2 O 3 . 2H 2 O) • Electrolyte: Al 2 O 3 dissolved in cryolite Na 3 [AlF 6 ] (10 wt. %) • T= 980 ° C; conductivity = 2.6 W -1 cm -1 • Anode: graphite; Cathode: stainless steel • Anodic reaction: 6(O 2- ) + 3C 3CO 2 + 12e - • Cathodic reaction: 4Al 3+ + 12e - 4Al 12 F Prof. Baranova, uOttawa The Hall-Heroult method of aluminum production:

Salt Crystal Dissolution Ionic solvation through dissolution of a NaCl crystal in water • Ionic Attractive Force Reduction • Solvation (Hydration) Chemical compounds that are dissociated into ions in solid, liquid or dissolved forms are termed electrolytes From the dissolution of monovalent electrolytes each ion carries one unit of elementary charge e 0 of magnitude I .602 x l0 -19 C. The ion oxidation state z is called charge number of the ion. For an electrolyte A n +B n - given the electroneutrality requirement z+ n + = z- n - and is termed electrolyte number

Redox Potentials

Standard electrode potentials / reduction potentials 0 V 0.337 V Cu 2+ + 2e - ® Cu -0.763 V Zn 2+ + 2e - ® Zn SHE More easily reduced / better oxidising agents More easily oxidised / better reducing agents Impossible to measure the potential of a single 1/2 cell, so... To draw up a scale of electrode potentials, one half cell was chosen as a fixed standard to measure other half cells against it. By convention, the primary standard is the STANDARD HYDROGEN ELECTRODE, defined as having an electrode potential of ZERO.

Galvanic Cell - Atomic view Danielle cell, invented in 1836

Transition from Electronic to Ionic Conductivity in an Electrochemical Cell Electrochemical cell for the electrolysis of aqueous CuCI 2 solution. If the ions in an electrolyte solution are subjected to an electric field, E, they will experience a force At the negative electrode: At the positive electrode: Overall Cell Reaction Cu ++ + 2e - à Cu 0

Electrolysis Cells and Galvanic Cells Galvanic cell based on the H 2 /Cl 2 reaction Electrolysis current as a function of the cell Voltage E C .. E D is the decomposition voltage Electrochemical Decomposition of HCI • Cell Voltage D V = E anode – E cathode • Electromotive Force EMF E 0 =lim t→0 D V • Decomposition Voltage E D » E 0

Electrolysis Cells vs Galvanic Cells Cathode: Reduction Reaction (G) Cl 2 + 2e - à 2Cl - (E) Cu 2+ + 2e - à Cu 0 Anode: Oxidation Reaction (E) 2Cl - à Cl 2 + 2e - (G) 2H 2 O + H 2 à 2H 3 O + + 2e - Schematic variation of cell voltage Ec against load current i for (a) a galvanic cell: b) an electrolvsis cell. For a galvanic cell: The power output P of the galvanic cell : For an electrolysis cell:

27 Basic principles and laws in electrochemistry Fuel cell circuit Δ G < 0 Electrolysis circuit Δ G > 0 Polarity of fuel cell and electrolysis cells is different Load Cathode: D z+ + ze - → D Anode: A z- → A + ze - Cathode Anode Electricity source + - Anode: D → D z+ + ze - Cathode: A + ze - → A z- Cathode Anode + -

28 Faraday’s law Faraday’s law — The number of moles of a substance, n, produced or consumed during an electrode process, is proportional to the electric charge passed through the electrode, Q. Assuming that there are no parallel processes, Q = zFn, where z and F are the number of electrons appearing in the electrode reaction equation, and the Faraday constant, respectively. Michael Faraday ( 1791 –1867 ) Electron charge e 0 = 1.6021917 10 -19 C Avogadro’s number A = 6.022169x 10 23 mol -1 e 0 A = F = 96,486.69 C mol -1 = 26.8 A-h zF It M zF MQ m ò = = zF It zF Q n = = (3) (4)

29 Faraday’s Law (continue) For general electrode reaction: n s S+ ze - → n p P Faraday’s law can be formulated as: Or with respect to the consumed substance S: Division of eq. (5) by M p and differentiation with respect to t yields the point rate of production: ò = Idt zF M m p p p n ò = - Idt zF M m s s s n I zF dt dn p p n = I zF M dt dm p p p n = (5) (6) (7)

30 Production Rates and Current Densities Equation (7) shows that the rate of an electrochemical reaction is simply proportional to the current flowing through electrolyzer cell. Experimental method of determining electrochemical reaction rate is to measure the electric current Current related to the electrode area A e is called current density, i[Acm -2 ], [Am -2 ] i zF A I zF dt dn A p e p p e n n = = 1 (8)

31 Faradaic efficiency (or coulometric efficiency) — Relates the moles of product formed in an electrode reaction to the consumed charge. The faradaic efficiency is 1.00 (or 100%) when the moles of product correspond to the consumed charge as required by Faraday’s law. Faradaic Efficiency Prerequisite for the validity of eqs. (5) – (8) is that only the considered reaction takes place at the electrode. But very often several parallel electrochemical reactions occur at the electrode Example cathodic copper deposition: Cu 2+ +2e - → Cu 2H + + 2e - → H 2 parallel formation of hydrogen 1/2O 2 + 2e- + 2H + → H 2 O parallel reduction of dissolved oxygen Individual contribution of each of these parallel reactions is given by their current efficiency F i 1 = F å i (9)

32 Faradaic Efficiency (continue) The current efficiency of an individual reaction may be calculated applying Faraday’s law: Q zF n p p n = F (10) ò = t e idt A Q 0 With Q: (11) A e – electrode surface area [m 2 ]

Faraday ’ s Law for Electrolysis

Coulometry Schematic diagram of a silver coulometer for determination of the total quantity of electricity passed Basic Formulation: • Silver Coulometer: Platinum Crusible (cathode): Ag + + e - ® Ag 0 Silver Rod (anode): Ag 0 ® Ag + + e - Q [C] = m [mg]/1.118 • Gas Combustion Coulometer: Cathode: 2H 2 O + 2e - ® H 2 + 2OH - Anode: 2OH - ® H 2 O + 1/2O 2 + 2e - Overall: H 2 O ® H 2 + 1/2O 2

Electrogravimetry

Electrogravimetry - Examples

37 Daniell cel (1836) Zn (anode) in solution of zinc salt Cu (cathode) in solution of copper salt: Zn + Cu 2+ Zn 2+ + Cu Standard cell potential is given E 0 cell = 0.34 – (-0.76) = 1.1 V Leclanché cell: Zn S /Zn 2+ , NH 4 Cl/MnO 2 S , Mn 2 O 3 S /C Zn is anode, carbon is a cathode E 0 cell = 1.55V Reactions: Mn(IV) to Mn(III) and Zn(0) to Zn(II): Zn + 2MnO 2 + H 2 O à Zn 2+ + Mn 2 O 3 + 2OH - Voltaic cells - Examples

Example of balancing electrochemical reactions 38 Unbalanced reaction: KMnO 4 + Na 2 SO 3 + H 2 O → MnO 2 + Na 2 SO 4 + KOH Reduction : 2 H 2 O + MnO 4 – → MnO 2 + 4 OH – - 3 e – Oxidation : 2 OH – + SO 3 2– → SO 4 2– + H 2 O + 2 e – By multiplying electrons to opposite half reactions: 4 H 2 O + 2 MnO 4 – → 2 MnO 2 + 8 OH – - 6e – 6 OH – + 3 SO 3 2– → 3 SO 4 2– + 3 H 2 O + 6e – Balanced reaction: 2 KMnO 4 + 3 Na 2 SO 3 + H 2 O → 2 MnO 2 + 3 Na 2 SO 4 + 2 KOH

Exercise 1.1 39 If we were to plate 20 g of copper from a copper sulfate bath (CuSO 4 ) with electrical charge passed through Q = 25 A-h. Determine the Faradaic (current) efficiency of the process

Exercise 1.2 40 High purity hydrogen is produced commercially by the electrolysis of water in basic solution (5 M KOH) a)Identify the electrode reactions. Note that in 5M KOH the concentration of hydrogen ions is extremely low, and the reaction: H + (aq.) + 2e - à H 2(g) is not involved a) Calculate the standard cell potential b) Calculate the minimum charge required to produce 1 kg of hydrogen

Electrolysis of water

Electrolysis From Greek electron (amber) and lysis (dissolution)- is a method of using direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction in electrolytic cells. Electrolysis of water is the decomposition of water into O 2 and H 2 due to an electric current being passed through the water. H 2 O H 2 (g) + ½ O 2 (g) In acidic solution (H 2 SO 4 ) half reactions are: 2H + (aq) + 2e - H 2 (g) (cathode) H 2 O (l) 1/2O 2 (g) + 2H + (aq) + 2e - (anode) If we balance for 4e - 2H 2 O (l) 2H 2 (g) + O 2 (g)

Alkaline media 2H 2 O 2H 2 (g) + O 2 (g) Cathode (Reduction) 2H 2 O (l) + 4e - 2H 2 (g) + 4 OH - (aq) Anode (oxidation) 4 OH - (aq) O 2 (g) + 2H 2 O (l) + 4e -

Exercise 1.2 High purity hydrogen is produced commercially by electrolysis of water in basic solution (5M KOH) a) Identify the electrode reactions. b) Calculate the standard cell potential. c) Calculate minimum charge required to produce 1 kg of H 2 .

a) Identify the electrode reactions. b) Calculate the standard cell potential. ࠵? !"## $ = 0.401- (-0.828) = 1.229 V Alkaline media 2H 2 O 2H 2 (g) + O 2 (g) Cathode (Reduction) 2H 2 O (l) + 4e - 2H 2 (g) + 4 OH - (aq) Anode (oxidation) 4 OH - (aq) O 2 (g) + 2H 2 O (l) + 4e -

c) Calculate minimum charge required to produce 1 kg of H 2 .

47 Ohm’s Law and Electrolyte Conductivities Any electrochemical circuit contains ohmic resistors: Wires composed of electronically conducting materials – electrons act as charge carriers Electrolyte representing an ionic conductor – ions are the carriers of electric charge where r is specific resistivity and k the specific conductivity of the resistor material r [ohm-cm] and k [ohm -1 -cm -1 ] Voltage drop caused by a current flowing through a resistor is given by Ohm’s law: U = RI Resistance of a resistor of length l and cross-section A is given by k A A R 1 1 1 × = = r

48 Ion conduction The current flow in electrolyte results from the movement of positive and negative ions; ion motion is caused by potential and concentration gradients Electrolytes : aqueous, molten salts, polymer, nonaqueuse liquid, supercritical fluid, ionic liquid Ion conductivity differs fundamentally from electrical conductivity in metal! Ion conductivity at T = 20 o C of the order of 10 -2 ohm -1 cm -1 for 0.1M salt Electrical conductivity of a typical metal (Fe) is of the order of 10 5 ohm -1 cm -1 For electrolytes, conductivity increases with increasing T For the electrical conductivity of metals and alloys, the temperature coefficient of conductivity is negative and about an order of magnitude lower

49 Ohm’s Law and Electrolyte Conductivities The ohmic potential drop: where l i the equivalent ionic conductivity of ionic species i Specific conductivity k of an electrolyte containing several ionic species i is given as: i k I k A U e 1 1 1 = × × = D W å = i i i i z c l k

50 In 1900 Fredrich Kohlrausch proposed to use equivalent conductance: n + is the number of positive ions into which chemical species dissociates Λ = ν + λ + + ν - λ -

51 Standard electrode potentials

52 Useful conversion factors

Concluding remarks Prof. Elena Baranova, uOttawa ¡ Electrochemistry and electrochemical engineering have wide range of technological application in modern life for making chemicals (electrosynthesis, H 2 , Al, etc. production), energy generation (fuel cell, batteries) or protecting materials (corrosion, coatings) ¡ Novel applications of electrochemistry: nanoparticle preparation by electrodeposition, electrochemical promotion of catalysis, bio- electrochemistry, etc. ¡ A key difference between electrochemical and non-electrochemical processes is our ability to manipulate an additional driving-force variable – electric potential. ¡ Faraday’s law is a basic law in electrochemistry that allows us to find out the quantity of chemical compounds produced or consumed during an electrochemical process.