Chem. 1B – 11/3 Lecture

Announcements I Exam 2 - Results Average about 65% Broader Distribution than Exam 1 Time was more of an issue Several questions with low % correct were easy conceptual questions Questions with lowest % correct: Score Range # 90 - 104 7 80s 22 70s 20 60s 38 50s 28 <50 21 Version 24% 27% 31% 37% 40% 41% A 23 21 12 9 19 17 B 10 11 18 22

Announcements II Lab Mastering – Some problems recently? Starting Experiment 9 (Wed./Thurs.) Next Week: Quiz 8 (Experiment 9 + 10 Questions + Electrochem Questions) Mastering – Some problems recently? Today’s Lecture Electrochemistry (Ch. 18) Some review + basic concepts today Definitions Standard Half-Cells and Cells Standard Reduction Potential

Chapter 18 Electrochemistry Electrochemical Reactions Balancing Redox Reactions: 6 step method: Assign oxidation states Separate overall reaction into oxidation and reduction reactions Balance each half reaction with respect to mass in order a) mass all elements other than H, O, b) O by adding H2O, c) by adding H+, d) Add OH- to both side if in alkaline sol’n Balance each half reaction for charge by adding electrons Use common multiplier to get equal numbers of electrons for each half-reaction Add each half reaction together to get net reaction without electrons as reactants or products Note: steps 5 and 6 are skipped if stopping at half reactions

Chapter 18 Electrochemistry Electrochemical Reactions Balancing Redox Reactions: Examples (unbalanced): AgNO3(aq) + Zn(s) ↔ Ag(s) + Zn(NO3)2(aq) HClO(aq) + Fe2+(aq) ↔ Cl2(g) + Fe3+(aq) MnO4- (aq) + C2O42-(aq) ↔ Mn2+(aq) + CO2(g)

Chapter 18 Electrochemistry Electrochemical Reactions – Different Forms “Beaker” Reactions Products form along with heat (assuming DH < 0) Little control of reaction Products co-mingled (from reduction and oxidation) Example: nail “rusts” (oxidation of Fe, reduction of O2) Voltaic (Galvanic) Cells Oxidation and reduction reactions may be divided into different parts (half-cells sometimes physically separated through two reaction cells) Two electrodes are also needed Reaction can be “harnessed” through voltage/power production Examples: batteries, pH measuring electrodes

Chapter 18 Electrochemistry Electrochemical Reactions – Different Forms Electrolytic Cell In this type of cell, external electrical energy is used to force unfavorable reactions (e.g. 2H2O(l) ↔ 2H2(g) + O2(g)) to occur Also requires two electrodes – but some differences from electrodes of voltaic cells Examples: Production of Cl2 gas from NaCl(aq), production of H2 gas from water (above), instruments that measure degree of oxidation/reduction at specific voltages (analogous to spectrometers)

Chapter 18 Electrochemistry Voltaic Cells - Description of how example cell works Reaction on anode = oxidation Anode = Zn electrode (as the Eº for Zn2+ is less than for that for Ag+) So, reaction on cathode must be reduction and involve Ag Oxidation produces e-, so anode has (–) charge (galvanic cells only); current runs from cathode to anode Salt bridge allows replenishment of ions as cations migrate to cathode and anions toward anodes GALVANIC CELL voltmeter Ag+ + e- → Ag(s) Zn(s) Ag(s) + – AgNO3(aq) ZnSO4(aq) Zn(s) → Zn2+ + 2e- Salt Bridge

Chapter 18 Electrochemistry Basic Electrical Quantities Current: the flow of electrons (although defined where a positive current has electrons moving backwards) Current units: Amperes (A) with 1 A = 1 C/s and 1 C = 1 Coulomb where 1 electron (elementary charge) has a value of 1.60 x 10-19 C Potential or Voltage: The potential energy associated with the movement of charge (e.g. to electrode of opposite sign) Potential units: Volts (V) = 1 J/C

Chapter 18 Electrochemistry Basic Electrical Quantities – From Voltaic Cells Current: related to the flow of electrons Potential: related to the reaction occurring (more energetic means higher potential) The ability of a metal (or other elements) to reduce can be measured under standard conditions Example: Zn(s) + 2Ag+(aq) ↔ Zn2+(aq) + 2Ag(s) If [Ag+] and [Zn2+] = 1 M, Ecellº = 1.56 V

Chapter 18 Electrochemistry Voltaic Cells Cell notation Example Cell: Zn(s)|Zn2+(aq)||Ag+ (aq)|Ag(s) GALVANIC CELL voltmeter Zn(s) Ag(s) “|” means phase boundary left side for anode (right side for cathode) “||” means salt bridge AgNO3(aq) ZnSO4(aq) Salt Bridge

Chapter 18 Electrochemistry Example Questions Given the following cell, answer the following question: MnO2(s)|Mn2+(aq)||Cr3+(aq)|Cr(s) What compound is used for the anode? What compound is used for the cathode? Write out both half-cell reactions and a net reaction

Chapter 18 Electrochemistry Given the following cell, write the cell notation: GALVANIC CELL voltmeter – reads +0.43 V Pt(s) Ag(s) – + Note: In this case the Pt(s) is an “inert” electrode (provides electrons but doesn’t react AgCl(s) NaCl(aq) FeSO4 (aq), Fe2(SO4)3(aq) Salt Bridge 13

Chapter 18 Electrochemistry Standard Reduction Potential A cell used to determine the standard reduction potential consists of two half cells One half-cell, the anode, is the standard hydrogen electrode (2H+(aq) + 2e- ↔ H2(g)) Eanodeº = 0 (defined) Other is the test cell (compound being reduced when half-cell is coupled to standard hydrogen electrode (oxidation electrode) Both cells under standard conditions (1 M, 1 atm) Ecellº = Ecathodeº The SHE is not actually used much any more (just a reference for relative potential) Pt(s) Ag(s) H2(g) AgNO3(aq) H+(aq) 14

Chapter 18 Electrochemistry Standard Reduction Potential Meaning of Values Half-cells that exhibit positive values have electrodes with compounds that easily reduce (e.g. Ag+(aq), MnO4-, PbO2(s)) Half-cells that exhibit negative values have electrodes that easily oxidize (e.g. alkali metals) What if we have two half-cells (neither SHE), can we find Ecellº? Example: Zn(s)|Zn2+(aq)||Ag+ (aq)|Ag(s) Ecellº = ? Ag+ reduction Eº = +0.80 V Eº = 0 Zn2+ reduction Eº = -0.76 V

Chapter 18 Electrochemistry Example Question An Ag/AgCl electrode is a common reference electrode. What is the standard potential of a cell made up of a Cu2+ solution being reduced to Cu(s) and AgCl(s) being reduced to Ag(s)? E°(Cu2+ + 2e- ↔ Cu(s)) = 0.34 V E°(AgCl(s) + e- ↔ Ag(s) + Cl- (aq)) = 0.22 V What is the balanced reaction and what species must be present at 1 M?