Corrosion Corrosion of Iron Since E redFe 2

Corrosion • • • Corrosion of Iron Since E red(Fe 2+) < E red(O 2) iron can be oxidized by oxygen. Cathode: O 2(g) + 4 H+(aq) + 4 e- 2 H 2 O(l). Anode: Fe(s) Fe 2+(aq) + 2 e-. Dissolved oxygen in water usually causes the oxidation of iron. Fe 2+ initially formed can be further oxidized to Fe 3+ which forms rust, Fe 2 O 3. x. H 2 O(s).

• Oxidation occurs at the site with the greatest concentration of O 2. Preventing Corrosion of Iron • Corrosion can be prevented by coating the iron with paint or another metal. • Galvanized iron is coated with a thin layer of zinc.

• Zinc protects the iron since Zn is the anode and Fe the cathode: Zn 2+(aq) +2 e- Zn(s), E red = -0. 76 V Fe 2+(aq) + 2 e- Fe(s), E red = -0. 44 V • With the above standard reduction potentials, Zn is easier to oxidize than Fe.

Preventing Corrosion of Iron • To protect underground pipelines, a sacrificial anode is added. • The water pipe is turned into the cathode and an active metal is used as the anode. • Often, Mg is used as the sacrificial anode: Mg 2+(aq) +2 e- Mg(s), E red = -2. 37 V Fe 2+(aq) + 2 e- Fe(s), E red = -0. 44 V

Electrolysis of Aqueous Solutions • Nonspontaneous reactions require an external current in order to force the reaction to proceed. • Electrolysis reactions are nonspontaneous. • In voltaic and electrolytic cells: – reduction occurs at the cathode, and – oxidation occurs at the anode. – In electrolytic cells, electrons are forced to flow from the anode to cathode.

Electrolysis – In electrolytic cells the anode is negative and the cathode is positive. (In galvanic cells the anode is positive and the cathode is negative. )

• • Example, decomposition of molten Na. Cl. Cathode: 2 Na+(l) + 2 e- 2 Na(l) Anode: 2 Cl-(l) Cl 2(g) + 2 e-. Industrially, electrolysis is used to produce metals like Al.

Electroplating • Active electrodes: electrodes that take part in electrolysis. • Example: electrolytic plating.

• Consider an active Ni electrode and another metallic electrode placed in an aqueous solution of Ni. SO 4: • Anode: Ni(s) Ni 2+(aq) + 2 e • Cathode: Ni 2+(aq) + 2 e- Ni(s). • Ni plates on the inert electrode. • Electroplating is important in protecting objects from corrosion.

• • • Quantitative Aspects of Electrolysis We want to know how much material we obtain with electrolysis. 1 Ampere is 1 Coulomb per second (A= C/s) 1 mole of electrons = 96, 485 C = 1 Faraday Use balanced half-equation to equate moles of substance to moles of electrons Molar mass (in g) = 1 mole of substance

Examples 1. A car bumper is to be electroplated with Cr from a solution of Cr 3+. What mass of Cr will be applied to the bumper if a current of 0. 50 amperes is allowed to run through the solution for 4. 20 hours? 2. What volume of H 2 gas (at STP) will be produced from the SHE after 2. 56 minutes at a current of 0. 98 amperes? 3. What volume of F 2 gas, at 25°C and 1. 00 atm, is produced when molten KF is electrolyzed by a current of 10. 0 A for 2. 00 hours? What mass of potassium metal is produced? At which electrode does each reaction occur?