Unit 11 Electrochemistry

What is electrochemistry? The study of the relationship between chemical change and electrical work. ◦ Investigated using redox rxns. 1.Electrolytic- use of electricity to cause a chemical reaction Electroplating 2. Galvanic or Voltaic the production of electricity by a chemical reaction. Battery

Oxidation and Reduction Rxns. Type of chemical reaction in which one substance transfers an electron to another substance. ◦ Called Redox reactions

How do redox rxns. work? Oxidation is the loss of electrons ◦ Referred as “reducing agent” ◦ Ionic charge increases; more (+) Reduction is the gain of electrons ◦ Referred as “oxdizing agent” ◦ Ionic charge decreases; more (-) ◦ “LEO” the lion says “GER” ◦ Zn(s) +2H + (aq)Zn +2 (aq) +H 2 (g)

How to keep track of electrons: Oxidation States Evaluating oxidation state (OS) of each atom in a formula helps us… ◦ Determine if electrons are being transferred ◦ Determine which atoms gains or loses electrons OS is generally not the electric charge, instead it’s used to us if the environment is “electron rich” or “electron poor”.

Rules for Assigning OS numbers 1. Total of O.S. = charge shown 2. group 1 metal = + 1 group 2 metal = F = – 1 4. H = O = – 2

Let’s Practice… What is the OS of … ◦ sulfur (S) in SF 6 ◦ Mn in MnO 4 1– ◦ Fe in Fe3+ Now for a challenge… ◦ O in KO 2.

Balancing redox rxns. Reduction and oxidation are linked forever. Yet, it is most helpful and useful to look at them separately. Conisder a rxn. between Zn and Cu +2 oxidation half-rxn: Zn (s) → Zn 2+ (aq) + 2 e– reduction half-rxn: Cu 2+ (aq) + 2 e– → Cu (s) overall redox rxn: Zn (s) + Cu 2+ (aq) → Cu (s) + Zn 2+ (aq)

Balancing redox rxns. in acid solution In a acid solution, you can add H 2 0 and H + to balance the overall rxn. Here are the rules… 1. Split the rxn. into its two ½-reactions, oxidation & reduction 2. Balance each ½-rxn.: ◦ balance the atoms other that O and H ◦ add H 2 O where needed to balance the O ◦ add H 1+ where needed to balance the H ◦ add e – to the more positive side to make it the same overall charge is the same on both sides 3. Multiply each half-rxn by a whole number to make the e – lost = e – gained. 4. Combine the half-rxns and simplify.

Let’s Practice…

Voltaic Cells What would happen if you could control how an electron was transferred? ◦ If you could do that, you have yourself a battery. A.K.A a voltaic cell – developed by Alessandro Volta.

Voltaic Cell

In a spontaneous (voltaic) cell, electrons flow from the anode to the cathode through the wire, allowing us to divert some of this electron current to do useful work. ◦ Anions move thru salt bridge to anode ◦ Cations move thru salt bridge to cathode

How to represent a Voltaic Cell You can represent a voltaic cell using a simple diagram like this one: Zn (s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu (s)

Cell Notation shorthand description of Voltaic cell electrode | electrolyte || electrolyte | electrode oxidation half-cell on left, reduction half- cell on the right single | = phase barrier ◦ if multiple electrolytes in same phase, a comma is used rather than | ◦ often use an inert electrode double line || = salt bridge

Is there a way to predict the way electrons will flow? Absolutely! Some metals are more active than others. For example:

A closer look at batteries…and that stinkin’ bunny that keeps going and going and going and going and going and going

Alkaline Battery The reactions in an alkaline cell like the one in your calculator are: Zn (s) + 2 OH 1– (aq) → Zn(OH) 2 (s) + 2 e – 2 MnO 2 (s) + H 2 O + 2 e – → Mn 2 O 3 (s) + 2 OH 1– (aq) Zn (s) + 2 MnO 2 (s) + 2 H 2 O → Zn(OH) 2 (s) + 2 MnO(OH) (s) Alkaline batteries are primary cells, which cannot be recharged. When most of the reactants have been converted to products, the cell is “dead.”

Hydrogen fuel-cell Fuel cells “burn” fuel to produce electric current instead of heat, and are much more efficient than ordinary combustion. The overall rxn. : 2 H 2 (g) + O 2 (g) → 2 H 2 O(g).

Lead-acid Battery Is a secondary (rechargable) battery. Anode (ox): Pb (s) + SO 4 -2 (aq) PbSO 4 (s) + 2e - Cathode (red): PbO 2(s) + 4H + (aq) +SO4 -2 (aq) + 2e- PbSO 4 (s)+ 2H 2 O (l) Overall: PbO 2 (s) + 2H 2 O (l) PbO 2 (s) + Pb (s) + 2H 2 SO 4 (aq)

Corrosion The corrosion of steel looks like this: anode: 2 Fe (s) → 2 Fe 2+ (aq) + 4 e – E° = v cathode: O 2 (g) + 2 H 2 O (l) + 4 e – → 4 OH 1– (aq) E° = v overall: 2 Fe (s) + O 2 (g) + 2 H 2 O (l) → 2 Fe 2+ (aq) + 4 OH 1– (aq) E° = v

Corrosion Prevention Block moisture or air from contacting iron. Paint or plate with non-reactive Cu Offer up a more active metal like Zn ◦ Galvanized nails – thin coat of Zn; better electron source ◦ Shipping boats, underground pipes, attach pieces of Zn

Exceptions Gold and Platinum are called noble metals (resistant to losing electrons)