Ch. 13 Fundamentals of Electrochemistry

Size: px

Start display at page:

Download "Ch. 13 Fundamentals of Electrochemistry"

Andrew Armstrong
6 years ago
Views:

1 Ch. 13 Fundamentals of Electrochemistry Basic Concepts of electrochemistry redox reaction : reactions with electron transfer oxidized : loses electrons reduced : gains electrons Fe 3+ + V 2+ Fe 2+ + V 3+ oxidizing agent (oxidant) : takes electrons from others and reduces reducing agent (reductant) : oxidizes itself 1. Electric charge (q) q : [coulomb] or [C] = x C for a single electron for 1 mole of e q = n F x C x 6.02x10 23 /mol 9.649x10 4 C/mol Faraday constant [F] Basic Concepts of electrochemistry 13.2 ex) If g of Fe 2+ was reduced in the reaction above, how many coulombs of charge must have been transferred from V 2+ to Fe 3+? 9.649x10 3 C

2 13-1. Basic Concepts of electrochemistry Electric Current : quantity of charge per second [], ampere <a charge of one coulomb per second> ex) Suppose that electrons are forced into a platinum wire immersed in a solution containing Sn 4+, which is reduced to Sn 2+ at a constant rate of 4.24 mmol/h. How much current flows into the solution? C/s = Basic Concepts of electrochemistry voltage, work, free energy difference in electrical potential (E) between two points = potential difference [V], volts work : energy in joules [J] J = E q one joule of energy = one coulomb of charge is moved between points whose potential differ by one volt. Ex) How much work is required to move 2.36 mmol of electrons through a potential difference of 1.05 V? 2.39x10 2 J

3 13-1. Basic Concepts of electrochemistry 13.5 The free energy change : G work done on surrounding : - G G = -E q = - n F E 4) Ohm's law current, I, flowing through a circuit I E R [] Basic Concepts of electrochemistry ) power : p work done per unit time, J/s or watt work p s E I Eq s E q s

13-1. Basic Concepts of electrochemistry 13.7 Type of Electrochemical cells Galvanic Cells : cell generates electric energy from where? chemical reaction.

4 13-1. Basic Concepts of electrochemistry 13.7 Type of Electrochemical cells Galvanic Cells : cell generates electric energy from where? chemical reaction. = spontaneous reaction Electrolytic Cells : requires an external source of E.E. one reagent must be oxidized, the other reduced galvanic cell in ction 13.8 reduction : 2gCl (s) + 2 e - 2 g(s) + 2 Cl - (aq) oxidation : Cd(s) Cd 2+ (aq) + 2 e - Net Cd(s) +2gCl (s) Cd 2+ (aq) + 2 Cl - (aq) + g(s) G = -150 KJ/mol of Cd

13-2. galvanic cell in ction KCl contained gel agar with KCl 2. Salt Bridge 13.

bridge net : Cd(s) + 2 g + (aq) Cd 2+ + 2 g(s) Cd(s) Cd(NO 3 ) 2 (aq) gno 3 (aq) g (s) 13-3.

10 measured voltage potential difference between g and Cd electrodes - Positive voltage when e- flows from anode

5 13-2. galvanic cell in ction KCl contained gel agar with KCl 2. Salt Bridge 13.9 cathode: 2 g + (aq) + 2 e - 2 g(s) 3) Line notation anode : Cd (s) Cd 2+ (aq) + 2 e - phase boundary salt bridge net : Cd(s) + 2 g + (aq) Cd g(s) Cd(s) Cd(NO 3 ) 2 (aq) gno 3 (aq) g (s) Standard Potentials measured voltage potential difference between g and Cd electrodes - Positive voltage when e- flows from anode into neg terminal. - Opposite connection : negative - standard reduction potential (E o ) : activities of all species are unity - left half cell :standard hydrogen electrode(she) if P H2(g) is 1 atm, E H 2, H 0, definition By convention, it is written as SHE g + g(s)

6 13-3. Standard Potentials Standard Reduction Potentials std. Reduction potential = potential difference between the std potential of the reduction of interest and the potential of SHE positive: electrons move from SHE to g + negative: electrons moves reverse direction Nernst eq : net driving force ) Nernst eq. for a half reaction, a + n e - b B E E 0 RT nf log b B a E 0 : std.reduction potential R: gas constant J/mol-K n: number of moles It shows two terms. 1. driving force under std. condition = E 0 2. concentration dependence (activity term) concentration term in Nernst Eq. a B reaction quotient Q pure solids, pure liquids, solvent not included a t 25 0 C Nernst equation simplifies to 59.16mV change per 10 fold change E E log n b B a

7 13-4. Nernst eq : net driving force ) Nernst equation for a complete reaction voltage difference (E) for the two electrodes. E = E + - E - E + : potential of electrode at positive input terminal of potentiometer = E cathode Procedure: 1. write half reactions with S.R.P. 2. write Nernst equation for half reaction 3. The more positive E 0 reduction predominantly happen 4. calculate net cell voltage E = E + -E - 5. write balanced reaction by subtracting the half reaction Nernst eq : net driving force ex) Find the voltage of the cell in Figure 14-5 if the right half cell contains 0.50 M gno 3 (aq) and the left half cell contains M CdCl 2 (aq). Write the net cell reaction and state whether it is spontaneous in the forward or reverse direction. right: 2 g + (aq) + 2 e - 2 g(s) 0.799V left: Cd 2+ (aq) + 2 e - Cd (s) V net : Cd(s) + 2 g + (aq) Cd g(s) E + = 0.781V, E - =-0.461V, E = E + -E - = 1.242V

8 13-4. Nernst eq : net driving force * dvice for finding relevant half-cell reactions Pb(s) PbF 2 (s) F - (aq) Cu +2 (aq) Cu(s) anode cathode oxidation reduction right : Cu e Cu (s) left: PbF 2 (s) + 2 e 2 F - + Pb (s) in fact, 2 F - + Pb (s) PbF 2 (s) + 2 e E 0 and the equilibrium constant E 0 and the equilibrium constant : galvanic cell produces electricity until the cell reaches eq. t eq. E=0 right electrode : a + n e - c C E 0 + left electrode : d D + n e - b B E 0 - E E E ( E 0 E ) log n c C a log c d C D E log{ } a b n B Q : reaction quotient at eq. E=0 as Q=K E 0 = logk n (at 25 0 C) K = 10 0 ne b B d D

9 13-6. Cells as chemical probes right half cell : gcl(s) left half cell : CH 3 COOH g + (aq) + Cl - (aq) CH 3 COO - (aq) + H + (aq) they come to equilibrium without other half-cell --- chemical reaction Cells as chemical probes with cell assembled, gcl(s) + e - g (s) + Cl - (aq, 0.10M) E +0 = 0.222V in left : the only element -- hydrogen gas bubbles out 2 H + (aq,?m) + 2 e - H 2 (g, 1.0atm) E -0 = 0 E = E + -E = ( log[cl - ] P ) - (0- log H 2 ) [ H ] in order to have same electrons 2 2 [ H ] [ Cl ] = log = V P H 2 P H2 we know [Cl - ], adjusted = 1. Thus, [H + ] = 1.8 x 10-4 M we can calculate K for this case (K a ) K a = = cell probe to measure the unknown H + conc.

13-6. Cells as chemical probes 13.19 ex) The cell in the following figure measures the formation constant (K f ) of Hg(EDT) 2-. The solution in the right hand compartment contains 0.

10 13-6. Cells as chemical probes ex) The cell in the following figure measures the formation constant (K f ) of Hg(EDT) 2-. The solution in the right hand compartment contains mmol of Hg 2+ and 2.00 mmol of EDT in a volume of L buffered to ph If the voltage is 0.331V, find the value of K f for Hg(EDT) Selected Problem No. 14-4, -5, -10, -17, -35, -38,

Review: Balancing Redox Reactions. Review: Balancing Redox Reactions

Review: Balancing Redox Reactions Determine which species is oxidized and which species is reduced Oxidation corresponds to an increase in the oxidation number of an element Reduction corresponds to a