ELECTROCHEMISTRY 9.1 and 9.2 To play the movies and simulations included, view the presentation in Slide Show Mode.

Electron Transfer Reactions Electron transfer reactions are oxidation-reduction or redox reactions. Eg: rusting and corrosion; all types of batteries, alkaline, NiCad, car; metabolism

Terminology for Redox Reactions OXIDATION—loss of electron(s) by a species; REDUCTION—gain of electron(s OXIDIZING AGENT—electron acceptor; species is reduced. REDUCING AGENT—electron donor; species is oxidized.

You can’t have one… without the other! Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation LEO the lion says GER! GER!

Another way to remember OIL RIG

Typical redox reactions Magnesium burns in air to produce a bright light 2Mg(s) + O2 (g) 2MgO(s) Mg is oxidized by O2 ;O2 is called the oxidizing agent. 2Mg 2Mg2+ + 4e- Each Mg loses 2 electrons (4 e- lost) O2 is reduced by Mg; Mg is called the reducing agent. O2 + 4e- 2O2- Each O in O2 gains 2 electrons (4 e- gained) IF ONE STUBSTANCE IS OXIDIZED, ANOTHER IN THE SAME REACTION MUST BE REDUCED.

Oxidation States A way of keeping track of the electrons. need the rules for assigning (memorize). The oxidation state of elements in their standard states is zero. Cu(s), Na(s), O2(g) Oxidation state for monoatomic ions are the same as their charge. Cu2+, I-, N3- Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide. H2O, NO2 (ox. # -2) H2O2, Na2O2 (ox. # -1) In compounds with nonmetals hydrogen is assigned the oxidation state +1. H2SO4, HCl, NaHCO3

In its compounds fluorine is always –1. SF6 The sum of the oxidation states must be zero in compounds. PbCl4 NaBrO3 +4 + 4(-1) = 0 +1 + +5 + 3(-2) = 0 7. The sum of the oxidation states must be equal the charge of the ion. NO3- Cr2O7 2- +5 + 3(-2) = -1 2(+6) + 7(-2) = -2 8. Group IA elements are always +1 CsF, NaCl Hydrogen is always +1 with the exception of hydrides. (H is -1). LiH 10.Group IIA elements are always +2. CaF2 , BaCl2

Oxidation States Assign the oxidation states to each element in the following. CO2 NO3- H2SO4 Fe2O3 Na2Cr2O7 C = +4 O = -2 O = -2 N = +5 S = +6 O = -2 H = +1 Fe = +3 O = -2 Cr = +6 O = -2 Na = +1

Common oxidation numbers Li +1 Na Cs Rb K Fr Ba +2 Be Mg Sr Ca Ra B +3 C +4 -2 -4 N +5 +4 +3 +2 +1 -3 O -1 -2 F -1 Ne H +1 He Al +3 Si +4 -4 P +5 +3 -3 S +6 +4 +2 -2 Cl +7 +5 +3 +1 -1 Ar Sc 3+ Ti +4 +3 +2 V +5 +4 +3 +2 Cr +6 +3 +2 Mn +7 +6 +4 +3 +2 Fe +3 +2 Co +3 +2 Ni +2 Cu +2 +1 Zn +2 Ga +3 Ge +4 -4 As 5+ 3+ 3- Se 6+ 4+ 2- Br +5 +1 -1 Kr +4 +2 Y +3 Zr +4 Nb +5 +4 Mo +6 +4 +3 Tc +7 +6 +4 Ru +8 +6 +4 +3 Rh +4 +3 +2 Pd +4 +2 Ag +1 Cd +2 In +3 Sn +4 +2 Sb +5 +3 -3 Te +6 +4 -2 I +7 +5 +1 -1 Xe +6 +4 +2 Lu +3 Hf +4 Ta +5 W +6 +4 Re +7 +6 +4 Os +8 +6 Ir +4 +3 Pt +4 +2 Au +3 +1 Hg +2 +1 Tl +3 +1 Pb +4 +2 Bi +5 +3 Po +2 At -1 Rn Lr +3 Common oxidation numbers

Oxidation numbers and the periodic table Some observed trends in compounds. Metals have positive oxidation numbers. Transition metals typically have more than one oxidation number. Nonmetals and semimetals have both positive and negative oxidation numbers. No element exists in a compound with an oxidation number greater than +8. The most negative oxidation numbers equals 8 - the group number

LEO says GER Loss of electrons; oxidation / Gain of electrons; reduction) +4 +3 +2 +1 0 -1 -2 -3 -4 Decreased oxidation number REDUCTION OXIDATION increased oxidation number

Agents Oxidizing agents Reducing agents gets reduced gains electrons. More negative oxidation state. Reducing agents gets oxidized. Loses electrons. More positive oxidation state.

Half-Reactions All redox reactions can be thought of as happening in two halves. One produces electrons - Oxidation half. The other requires electrons - Reduction half.

Steps to Balancing Redox Equations In aqueous solutions the key is the number of electrons produced must be the same as those required. For reactions in acidic solution an 8 step procedure. For reactions in basic solutions, one more step is required.

Acidic Solution Write separate half reactions For each half reaction balance all reactants except H and O Balance O using H2O Balance H using H+ Balance charge using e- Multiply equations to make electrons equal Add equations and cancel identical species Check that charges and elements are balanced.

BrO3- + I- Br- + I2 in an acidic solution Write separate half reactions BrO3- Br- I- I2 For each half reaction balance all reactants except H and O BrO3- Br- 2 I- I2 done Balance O using H2O BrO3- Br- + 3 H2O 2 I- I2 Balance H using H+ BrO3- + 6H+ Br- + 3 H2O 2 I- I2

BrO3- + 6H+ + 6e- Br- + 3 H2O 2 I- I2 + 2e- Balance charge using e- BrO3- + 6H+ + 6e- Br- + 3 H2O 2 I- I2 + 2e- Multiply equations to make electrons equal BrO3- + 6H+ + 6e- Br- + 3 H2O 3[2 I- I2 + 2e- ] Add equations and cancel identical species BrO3- + 6H+ + 6I- Br- + 3 I2 + 3 H2O 8. Check that charges and elements are balanced.

Practice The following reactions occur in acidic solutions. Balance them. a) MnO4- + H2O2 ® Mn+2 + O2 b) I- + NO3- ® I2 + NO(g) c) Cr2O72- + Fe2+ ® Cr3+ + Fe3+ d) Mn+2 + BiO3- ® Bi+3 + MnO4-

Basic Solution Do everything you would with acid, but add one more step. Add enough OH- to both sides to neutralize the H+

Al + NO3- Al+3 + NH3 in a basic solution Write separate half reactions Al Al3+ NO3- NH3 For each half reaction balance all reactants except H and O done done Balance O using H2O Al Al3+ NO3- NH3 + 3H2O done Balance H using H+ Al Al3+ NO3- + 9 H+ NH3 + 3H2O

Balance charge using e- Al Al3+ + 3e- NO3- + 9H+ + 8e- NH3 + 3 H2O Add OH- to both sides to balance H+ . Create H2O. Al Al3+ + 3e- NO3- + 9 H+ + 8e- + 9OH- NH3 + 3H2O + 9 OH- Al Al3+ + 3e- NO3- + 9H2O + 8e- NH3 + 3H2O + 9OH- 7. Multiply equations to make electrons equal 8[Al Al3+ + 3e- ] 3 (NO3- + 9H2O + 8e- NH3+ 3H2O + 9OH-) 8. Add equations and cancel identical species 8Al + 3NO3- + 18H2O 8Al3+ + 3 NH3 + 27OH- 9. Check that charges and elements are balanced.

Practice The following reactions occur in basic solutions. Balance them. Cr(OH)3 + OCl- + ® CrO4-2 + Cl- MnO4- + Fe+2 ® Mn+2 + Fe+3 Fe2+ + H2O2 ® Fe3+ + OH - d) S2O32- + OCl- ® SO42- + Cl-