Warmup 2. Balance this redox reaction in acid: + Fe 2+ + Cr 2 O 7 2-  Fe 3+ + Cr 3+ Ni e¯  Ni E° = V Zn e¯  Zn E° = V 1. Write the notation for a zinc-nickel cell, and calculate the voltage given the reduction potentials above. a)Ni 2+ | Ni || Zn | Zn V b)Zn | Ni 2+ || Zn 2+ | Ni V c) Zn 2+ | Zn || Ni 2+ | Ni V d) Zn | Zn 2+ || Ni 2+ | Ni V Done in a second!

Electrochemistry Day 2 And then a lab!!!!

Balancing Redox Equations (in acid or base sol’n) Write the unbalanced equation for the reaction ion ionic form. Separate equation into two half-reactions. Use oxidation numbers if necessary. The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2- in acid solution? Fe 2+ + Cr 2 O 7 2- Fe 3+ + Cr 3+ Oxidation: Cr 2 O 7 2- Cr Reduction: Fe 2+ Fe Balance the atoms other than O and H in each half-reaction. Cr 2 O Cr 3+

For reactions in acid, add H 2 O to balance O atoms and H + to balance H atoms. For bases, add OH - Cr 2 O Cr H 2 O 14H + + Cr 2 O Cr H 2 O Add electrons to one side of each half-reaction to balance the charges on the half-reaction. Fe 2+ Fe e - 6e H + + Cr 2 O Cr H 2 O If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients. 6Fe 2+ 6Fe e - 6e H + + Cr 2 O Cr H 2 O

Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. 6e H + + Cr 2 O Cr H 2 O 6Fe 2+ 6Fe e -Oxidation: Reduction: 14H + + Cr 2 O Fe 2+ 6Fe Cr H 2 O

Calculate the standard reduction potential of the half- reaction Pt 2+ (aq) + 2e -  Pt(s) Ag + (aq) + e -  Ag(s) E° = Ag(s)  2e - + 2Ag + (aq) E° = e - + Pt 2+ (aq)  Pt(s) E° = ? 2Ag(s) + Pt 2+ (aq)  Pt(s) + 2Ag + (aq)E T ° = 0.38 V a)-1.18 V b) 1.18 Vc) V d) 2.00 Ve) 0.40 V Is the reaction between silver and platinum ions spontaneous or nonspontaneous? Explain.

Prelaboratory 1. Electrons flow from the substance that is which electrons, to the substance that is, which electrons. 2. Write the balanced chemical equation for the decomposition of water into its elements. 3. Assign oxidation states to the hydrogen and oxygen atoms in each substance in the above chemical equation. Then, label the atom that is oxidized and the atom that is reduced. 4. Balance the following oxidation and reduction half-reactions for the decomposition of water. H 2 O  O 2 + H + + e - H 2 O + e -  H 2 + OH - 5. When pH is 6.0 or lower, bromothymol blue, an acid-base indicator, is _________. When pH is 7.6 or higher, it turns __________. oxidized loses reduced gains 2H 2 O  2H 2 + O H 2 O  O 2 + 4H + + 4e - 4H 2 O + 4e -  H 2 + 4OH -

Lab time!