Lecture 14: Announcements PDF

Lecture #14, p. 1 Lecture 14: Announcements Today: Brown Ch. 20 Electrochemistry 20.1 Oxidation States and Oxidation–Reduction Reactions 20.2 Balancing Redox Equations 20.3 Voltaic Cells...

Lecture #14, p. 1 Lecture 14: Announcements Today: Brown Ch. 20 Electrochemistry 20.1 Oxidation States and Oxidation–Reduction Reactions 20.2 Balancing Redox Equations 20.3 Voltaic Cells 20.4 Cell Potentials under Standard Conditions 20.5 Free Energy and Redox Reactions When you see someone balancing a redox in their mind Chemistry  Lecture #14, p. 2 Lecture 14: Announcements Problem Set 12: Not chosen for grading Problem Set 13: Posted on Moodle; due tomorrow by 14:00 Problem Set 14: Posted on Moodle with solutions Do not need to hand in! Office Hours: Prof. Norris and Brisby have office hours today Exam: Covers Lectures 1–14, Problem Sets 1–14, and Exercises 1–13 Practice exams posted Allowed one A4 “Cheat Sheet” (both sides), German–English dictionary, and standard scientific calculator... Chemistry  Lecture #14, p. 3 Standard Scientific Calculators digitec.ch TI 30XS, 30X, 30 EcoRs... See exam section of Moodle for “white list” Exam questions? Chemistry  Lecture #14, p. 4 Credits Big Thank You! Chemistry  Lecture #14, p. 5 Review In Lecture 13, we discussed acid–base equilibria Autoionization of water, ion-product constant, !! Proton-transfer reactions, Brønsted-Lowry Acids and Bases Acid–base pairs, conjugate acids and conjugate bases Acid-dissociation constant, !" Base-dissociation constant, !# equilibrium constants: pH, pOH, p!" , p!# , p!! !$ , !% , !&% , !! , !" , and !# Common-ion effect Buffers, Henderson–Hasselbalch Equation Chemistry  Lecture #14, p. 6 Review H, O 1 + H, O 1 ⇌ H. O/ -. + OH - (-.) !! = H. O/ OH - = 1.0×10-01 pH = −log [H. O/ ] - / '! (" )# HA -. + H, O 1 ⇌ A -. + H. O -. !" = [('] p!" = −log [!" ] B -. + H, O 1 ⇌ HB/ -. + OH - -. (2# )(! !# = p!# = −log [!# ] Chemistry  Lecture #14, p. 7 Today: Electrochemistry Lecture 3: Oxidation–reduction reactions (redox) Electron(s) transferred between reactants Important! Basis of batteries, fuel cells, electroplating, corrosion... Quick review: Oxidation: substance loses electrons Reduction: substance gains electrons “LEO the lions says GER” Bookkeeping system? Chemistry  Lecture #14, p. 8 Review: Oxidation Numbers Rules: Atoms in elemental form: 0 Monoatomic ions: ionic charge Nonmetals in ionic/molecular compounds: negative oxidation numbers Oxygen: −2 (except peroxide ion, O,- , , −1) H: +1 (except if bonded to metal, −1) F: −1 (always) Cl, Br, I: −1 (except positive if bonded to oxygen) Sum of oxidation numbers for atoms in compound equals its net charge Ex: Mg ; + 2HCl -. → MgCl, (-.) + H, (?) 0 +1 −1 +2 −1 0 Mg is oxidized; Mg reduces H; Mg is reductant H is reduced; H oxidizes Mg; H is oxidizer Chemistry Recallfromlecture 3 Activityseries Anymetalaboveit onserieswillreactwithacidtoformHagas Mgisabove H seebelowp16  Lecture #14, p. 9 Electrochemistry Study of relationship between electricity and chemical reactions Example ≡ “Nicad” battery disoxidized Ex: Cd ; + NiO, ; + 2H, O(1) → Cd(OH), ; + Ni(OH), ; i isreduced Used to generate power; electrons must be conserved Adds constraint when we balance redox reactions! How? “Method of half-reactions” Divide electrochemical reaction into oxidation and reduction half-reactions Balance each and then combine Procedure? Chemistry  Lecture #14, p. 10 Procedure How to balance redox reactions in acidic aqueous solutions 1. Divide reaction into oxidation half-reaction and reduction half-reaction 2. Balance each half-reaction First, balance elements other than H and O Next, balance O atoms by adding H2O as needed Then, balance H atoms by adding H+ as needed Finally, balance charge by adding e−s as needed 3. Multiply two half-reactions by integers to equate electrons lost and gained 4. Add half-reactions and, if possible, cancel species appearing on both sides 5. Check to make sure atoms and charges are balanced Sounds complicated, but very logical Chemistry  Lecture #14, p. 11 permanganationtoxalation it i thief ua on Example Balance MnO- ,- 1 + C, O1 → Mn ,/ + CO, in acidic H, O Procedure Byq Reduction Oxidation o Half-rxns: MnO" ! → Mn #$ C# O#" ! → CO# +7 −2 +2 +3 −2 +4 −2 o Balance half-rxns: Balance non H, O: MnO" ! → Mn #$ C# O#" ! → 2CO# Balance O: MnO" ! → Mn #$ + 4H# O C# O#" ! → 2CO# fin Byaf Balance H: MnO" $ ! + 8H → Mn #$ + 4H# O C# O#" ! → 2CO# in ingByq Balance charge: MnO" $ " ! + 8H + 5e → Mn #$ + 4H# O C# O#" ! → 2CO# + 2e " o Balance e− gained/lost: 2MnO" $ " ! + 16H + 10e → 2Mn #$ + 8H# O 5C# O#" ! → 10CO# + 10e " o Combine half-rxns: fiii I 2MnO" #" $ ! + 5C# O! + 16H → 2Mn #$ + 10CO# + 8H# O o Check balance: atoms: 2 Mn 20 oxygen charge: 4+ Iiii 10 C 16 H Chemistry Areatomsandchargesbalance Balanced  Lecture #14, p. 12 What If Solution is Basic? How to balance redox reactions in basic aqueous solutions 1. Balance half-reactions as if they occur in acidic solution 2. Count number of H+ in each half-reaction, and add same number of OH− to each side Why? Neutralizing H+ with OH− Equal H+ and OH− on same side form H2O Cancel any H2O appearing on both sides 3. Multiply two half-reactions by integers to equate electrons lost and gained Same as 4. Add half-reactions and, if possible, cancel species appearing on both sides above 5. Check to make sure atoms and charges are balanced Example? Chemistry  Lecture #14, p. 13 wi Asiijusth.irfectia skePetin'equation co cyanate Example Balance CN - + MnO- - 1 → CNO + MnO, (+) in basic H, O Procedure Reduction Oxidation iiiit o Half-rxns: o Balance half-rxns: MnO" ! → MnO# +7 −2 +4 −2 CN" → CNO" +2 −3 +4 −3 −2 Balance non H, O: MnO" ! → MnO# CN" → CNO" Balance O: MnO" ! → MnO# + 2H# O CN" + H# O → CNO" Balance H: MnO" $ ! + 4H → MnO# + 2H# O CN" + H# O → CNO" + 2H$ v Balance charge: MnO" $ " ! + 4H + 3e → MnO# + 2H# O CN" + H# O → CNO" + 2H$ + 2e" 1499g in o Add OH−: MnO" $ " " ! + 4H + 3e + 4OH → MnO# + 2H# O + 4OH " CN" + H# O + 2OH" → CNO" + 2H$ + 2e" + 2OH" Chemistry  Lecture #14, p. 14 Example Balance CN - + MnO- - 1 → CNO + MnO, (+) in basic H, O Procedure Reduction Oxidation o Add OH−: MnO" $ " " ! + 4H + 3e + 4OH → MnO# + 2H# O + 4OH " CN" + H# O + 2OH" → CNO" + 2H$ + 2e" + 2OH" OH o Neutralize H+: MnO" " ! + 4H# O + 3e → MnO# + 2H# O + 4OH " Hay CN" + H# O + 2OH" → CNO" + 2H# O + 2e" neelHaoon sothsides o Cancel H2O: MnO" " ! + 2H# O + 3e → MnO# + 4OH " CN" + 2OH" → CNO" + H# O + 2e" o Balance e− gained/lost: 2MnO" " ! + 4H# O + 6e → 2MnO# + 8OH " 3CN" + 6OH" → 3CNO" + 3H# O + 6e" Initiisi o Combine half-rxns: 2MnO" " " ! + H# O + 3CN → 2MnO# + 3CNO + 2OH " v Chemistry o Check balance: atoms: 2 Mn 3C 3N 2H 9 oxygen charge: 5− Iiii Areatomsandchargesbalanced  Lecture #14, p. 15 Why Do We Care? Voltaic cells: If redox rxn is spontaneous (i.e., ∆. < 0 ), we can extract electrical work Half-rxns above can be physically separated into two “half-cells” Ex: Zn # + Cu,/(()) → Zn,/(()) + Cu(#) Zn is oxidized by Cu Is this redox rxn spontaneous? Chemistry  Lecture #14, p. 16 Recall: Activity Series Any metal is oxidized by metals below it Zn # + Cu,/ (()) → Zn,/ (()) + Cu(#) Spontaneous! We can use this to make battery! Chemistry  Lecture #14, p. 17 Copper–Zinc Battery Cu electrode in 1. CuSO1 Zn electrode in 1. ZnSO1 This is voltaic cell! Solutions in contact through porous glass disc Chemistry  Lecture #14, p. 18 Justtobeclearcommercialalkalinebatteriesarenotbasedoncazn A Salt & Battery Anode − Cathode + Ilipani where where oxidation reduction occurs occurs vowels consonants Chemistry  Lecture #14, p. 19 Battery Components Electrodes: Solid metals connected to external circuit In Cu–Zn example, electrode materials participate in rxn, but not always Zn anode losing mass; Cu cathode gaining mass Electrolyte: Liquid with ions that can conduct electrical current To maintain charge neutrality, ions need to move in electrolyte Salt bridge: Tube containing electrolyte solution allowing ion migration Electrolyte can be gel to stay in tube Anions move toward anode; cations move toward cathode Chemistry  Lecture #14, p. 20 Energy and Batteries If we connect two half-rxns, how do we know if we get a good battery? We discuss batteries in terms of voltage... I obviously w Electric potential ⇒ potential energy difference per unit charge antitise Onevolt 2 1.6 × 10−19 J Energy 1V≡ i.e., each electron carries 1.6 × 10!"# J 1.6 × 10−19 C Charge v cathode + e− At 1 volt, each electron can do 1.6 × 10!"# J of work 1V At 5 volt, each electron carries 5× this energy Expected voltage for a given pair of half reactions? anode − Chemistry Note in a it Fer a.eeipdai eniatn arsino henonw P it Pints  Lecture #14, p. 21 Calculating Cell Potentials ° !!"## ≡ cell voltage at standard conditions IMconcentra forreactan To calculate cell potential for any combination of half cells all half-reactions have been tabulated in terms of reduction Standard reduction potentials ° !345 ≡ potential energy available if reduced i iii Cell potential ⇒ ° !$477 ° = !345 ° cathode − !345 anode Minus sign because oxidation actually occurs at anode Chemistry  Lecture #14, p. 22 Standard Reduction Potentials ° !%"& Voltages always referenced What is our “ground”? Convention: standard H electrode Physical meaning of potentials? Chemistry  Lecture #14, p. 23 Oxidizing and Reducing Agents Chemistry  Lecture #14, p. 24 Determining Cell Potentials Cu-Zn battery From Table 20.1 Chemistry  Lecture #14, p. 25 Standard Hydrogen Electrode SHE Note standard conditions Can operate as anode or cathode Depends on other half-cell Chemistry  Lecture #14, p. 26 Notes ° !!"## To be useful, !$477 ° >0 (by convention) Useful means e−’s flow spontaneously from anode to cathode to do work Electrons have higher potential at anode and flow to cathode ° !$477 is “intensive” quantity—does not depend on size of system ° !%"& Half-cell potentials cannot be measured directly Because we need full cell to measure something But can determine !345° for any half-reaction by measuring with SHE All half-cell potentials listed as reduction potentials But half-rxns are reversible, depending on what half-cell it is connected to More positive !345 ° means greater tendency for reduction Chemistry  Lecture #14, p. 27 Connection to Gibbs Free Energy ° ∆# ° = −& ( !!"## ∆# ° = J/mol of rxn ( = Faraday’s constant = 96,485 C/mol e−’s &= positive unitless number moles of e−’s transferred in balanced cell rxn ° If !!"## > 0, ∆# ° < 0 Spontaneous! Chemistry Simpsons: Season 24, Episode 16 checkunits For moi v Ii 4 FF  Lecture #14, p. 28 Big Picture By combining half-cell rxns, we can design battery Lithium is most easily oxidized; !345 ° = −3.05V Makes sense that lithium is used in batteries! If instead of extracting work, we apply external potential (i.e. voltage)... We can recharge battery by “reversing” chemistry Today just basics Fuel cells, electroplating, corrosion all related topics! Important for engineers Chemistry  Lecture #14, p. 29 What We Learned Reviewed oxidation and reduction, oxidation numbers Electrochemistry = relationship between electricity and chemical rxns Balancing redox reactions: additional constraint of balanced electron flow Method of half-reactions Redox reactions in acidic solutions and in basic solutions Voltaic cells Cell potentials Standard reduction potentials Standard hydrogen electrode Relationship between ∆+ ° and !$477 ° Chemistry  Lecture #14, p. 30 Why Chemistry? ETH engineers must know basic chemistry as part of their education MAVT graduates now do everything! Need chemistry knowledge so you are not left behind! MAVT students were previously told to ignore chemistry Later regretted their poor chemistry knowledge World Needs You! Examples: Face critical crises SAF? Need world-class engineers Hydrogen-fired turbines? Batteries? Your generation must save world Carbon capture? Transportation? Biomedical devices? Chemistry