18.5 Harnessing the Power of Voltaic...

1 April 12 Harnessing the Power of Voltaic Cells / Corrosion

18.5 Harnessing the Power of Voltaic Cells

Batteries and Corrosion

Dr. Fred Omega Garces Chemistry 201 Miramar College


Commercial Voltaic Cells Voltaic Cells are convenient energy sources Batteries is a self-contained group of voltaic cells arranged in series. Advantage: Portable Disadvantage: Very Expensive (0.80¢ / Kwatt-hr) Need cells in series to provide power

The Processes occurring during the discharge and recharge of a lead-acid battery. When the lead-acid battery is discharging (top) it behaves like a voltaic cell: the anode is negative (electrode-1) and the cathode is positive (electrode-2). When it is recharging (bottom), it behaves like an electrolytic cell; the anode is positive (electrode-2) and the cathode is negative (electrode-1).


Dry Cell or LeClanche Cell Dry Cells Invented in the 1860’s the common dry cell or LeClanche cell, has become a familiar household item. An active zinc anode in the form of a can house a mixture of MnO2 and an acidic electrolytic paste, consisting of NH4Cl, ZnCl2, H2O and starch powdered graphite improves conductivity. The inactive cathode is a graphite rod.

Anode (oxidation) Zn(s) g Zn2+

(aq) = 2e- Cathode (reduction). The cathodic half-reaction is complex and even today, is still being studied. MnO2(s) is reduced to Mn2O3(s) through a series of steps that may involve the presence of Mn2+ and an acid-base reaction between NH4

+ and OH- :

2MnO2 (s) + 2NH4+(aq) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O (l)

The ammonia, some of which may be gaseous, forms a complex ion with Zn2+, which crystallize in contact Cl- ion:

Zn2+(aq) + 2NH3

(aq) + 2Cl-(aq) g Zn(NH3)2Cl2(s)

Overall Cell reaction: 2MnO2 (s) + 2NH4Cl(aq) + Zn(s) g Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s) Ecell = 1.5 V

Uses: common household items, such as portable radios, toys, flashlights, Advantage; Inexpensive, safe, available in many sizes Disadvantages: At high current drain, NH3(g) builds up causing drop in voltage, short shelf life because zinc anode reacts with the acidic NH4

+ ions.


Dry Cell or LeClanche Cell Invented by George Leclanche, a French Chemist.

Acid version: Zinc inner case that acts as the anode and a carbon rod in contact with a moist paste of solid MnO2 , solid NH4Cl, and carbon that acts as the cathode. As battery wear down, Conc. of Zn+2 and NH3 (aq) increases thereby decreasing the voltage. Half reactions: E°Cell = 1.5 V Anode: Zn(s) g Zn+2

(aq) + 2e-

Cathode: 2NH4+(aq) + MnO2(s) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O(l)

Advantage: Inexpensive, safe, many sizes Disadvantage: High current drain, NH3(g) build up, short shelf life

Anode Cathode

Steel cover

Insulating washer

Wax seal Sand cushion

Carbon rod (cathode_

NH4Cl ZnCl2, MnO2 paste

Porous separator

Zinc can (anode)

Wrapper


Alkaline Battery Alkaline Battery The alkaline battery is an improved dry cell. The half-reactions are similar, but the electrolyte is a basic KOH paste, which eliminates the buildup of gases and maintains the Zn electrode.

Anode (oxidation) Zn(s) + 2OH- (aq) g ZnO(s) + H2O (l) + 2e- Cathode (reduction). 2MnO2 (s) + 2H2O (l) + 2e- g Mn(OH)2(s) + 2OH-

(aq) Overall Cell reaction: 2MnO2 (s) + H2O (l) + Zn(s) g ZnO(s) + 2MnO(OH)(s) Ecell = 1.5V

Uses: Same as for dry cell. Advantages: No voltage drop and longer shell life than dry cell because of alkaline electrolyte; sale. Disadvantages; More expensive than common dry cell.


Alkaline Battery Leclanche Battery: Alkaline Version In alkaline version; solid NH4Cl is replaced with KOH or NaOH. This makes cell last longer mainly because the zinc anode corrodes less rapidly under basic conditions versus acidic conditions. Half reactions: E°Cell = 1.5 V Anode: Zn(s) + 2OH-

(aq) g ZnO(s) + H2O(l) + 2e-

Cathode: MnO2 (s) + H2O(l) + 2e- g MnO3 (s) + 2OH-(aq)

Nernst equation: E = E° - [(0.592/n)log Q], Q is constant !!

Advantage: No voltage drop, longer shelf life. Disadvantage: More expensive


Mercury Button Battery Mercury and Silver batteries are similar. Like the alkaline dry cell, both of these batteries use zinc in a basic medium as the anode. The solid reactants are each compressed with KOH, and moist paper acts as a salt bridge. Half reactions: E°Cell = 1.6 V Anode: Zn(s) + 2OH-

(aq) g ZnO(s) + H2O(l) + 2e-

Cathode (Hg): HgO (s) + 2H2O(l) + 2e- g Hg(s) + 2OH-(aq)

Cathode (Ag): Ag2O (s) + H2O(l) + 2e- g 2Ag(s) + 2OH-(aq)

Advantage: Small, large potential, silver is nontoxic. Disadvantage: Mercury is toxic, silver is expensive.


Nickel-Cadmium Battery Battery for the Technological Age Rechargeable, lightweight “Ni-cad” are used for variety of cordless appliances. Main advantage is that the oxidizing and reducing agent can be regenerated easily when recharged. These produce constant potential.

Half reactions: E°Cell = 1.3 V Anode: Cd(s) + 2OH-

(aq) g Cd(OH)2 (s) + 2e- Cathode: 2NiO(OH) (s) + 2H2O(l) + 2e- g Ni(OH)2 (s) + 2 OH-

(aq)

Nickel-Metal Hydride (Ni-MH)Battery The nickel-metal hydride battery has started to replace the nickel-cadmium battery. The anode half-reaction oxidizes the hydrogen adsorbed in metal alloys (designated M; e.g., LaNi5) in a basic (KOH) electrolyte, while nickel(III) as NiO(OH) is reduced at the cathode.

Anode (oxidation); MH(s) + OH- (aq) g M(s) + H2O(l) + e-

Cathode (reduction): NiO(OH)(s) + H2O (l) + e- g Ni(OH)2 (s) + OH-(aq)

Overall (cell) reaction: MH(s) + NiO(OH) (s) g M (s) + Ni(OH)2 (s): E cell = 1.4 V

The cell reaction is reversed during recharging. Uses: Cordless razors, photo flash units, and hand tools. Advantage: Lightweight: high power, avoids the toxicity of cadmium in NiCad battery. Disadvantages: Discharge upon storage.


Lithium Ion Battery Lithium Ion Battery The name refers to a family of rechargeable battery in which lithium ion smove from the anode to the cathode during discharge and back to the anode upon recharging.

Advantage: • Lighter than other types of rechargeable batteries of the same size. • The electrodes of a lithium-ion battery are made of lightweight lithium and carbon. • Lithium is also a highly reactive element, meaning that a lot of energy can be stored in its atomic bonds • They hold their charge. • Losses only about 5 percent of its charge per month, compared to a 20 percent loss per month for NiMH batteries. • They have no memory effect. • Lithium-ion batteries can handle hundreds of charge/discharge cycles.

Disadvantage: • Start degrading upon leaving factory. •Last only two or three years from the date of manufacture whether use them or not. • Extremely sensitive to high temperatures. Heat causes lithium-ion battery packs to degrade faster than they normally. • If you completely discharge a lithium-ion battery, it is ruined. • Must have an on-board computer to manage the battery. Makes more expensive. • There is a small chance that, if a lithium-ion battery pack fails, it will burst into flame.


Lithium Ion Battery Lithium Ion Battery: Work in Progress Half reactions: E°Cell = 3.6 - 3.7 V Anode: LiXXO2 g Li1-xXXO2 + Li+ + xe-

Cathode: C + xLi+ + xe- g LixC Where XX = combination of Cobalt and Manganese.

Advantage: No voltage drop, longer shelf life. Disadvantage: • Start degrading upon leaving factory. •Last only two or three years from the date of manufacture whether use them or not. • Extremely sensitive to high temperatures. Heat causes lithium-ion battery packs to degrade faster than they normally. • If you completely discharge a lithium-ion battery, it is ruined. • Must have an on-board computer to manage the battery. Makes more expensive. • There is a small chance that, if a lithium-ion battery pack fails, it will burst into flame.


Lead Storage Battery Lead-Acid Battery. A typical 12-V lead-acid battery has six cells connected in series, each of which delivers about 2 V. Each cell contains two lead grids packed with the electrode material: the anode is spongy Pb, and the cathode is powered PbO2. The grids are immersed in an electrolyte solution of 4.5 M H2SO4. Fiberglass sheets between the grids prevents shorting by accidental physical contact. When the cell discharges, it generates electrical energy as a voltaic cell.

Half reactions: E°Cell = 2.0 V Anode: Pb(s) + SO4

2- g PbSO4 (s) +2 e- E° = 0.356 Cathode (PbO2): PbO2 (s) + SO4

2- + 4H+ + 2e- g PbSO4 (s) + 2 H2O E° = 1.685V

Net: PbO2 (s) + Pb(s) + 2H2SO4 g PbSO4 (s) + 2 H2O E°Cell = 2.0 V

Note hat both half-reaction produce Pb2+ ion, one through oxidation of Pb, the other through reduction of PbO2. At both electrodes, the Pb2+ react with SO4

2- to form PbSO4(s)


Fuel Cells; Batteries Fuel Cell can be made as an electrochemical device for converting chemical energy into electricity.

In contrast to storage battery, fuel cell does not need to involve a reversible reaction since the reactant are supplied to the cell as needed from an external source. This technology has been used in the Gemini, Apollo and Space Shuttle program.

Half reactions: E°Cell = 0.9 V Anode: 2H2 (g) + 4OH-

(aq) g 4H2O(l) + 4e- Cathode: O2 (g) + 2H2O(l) + 4e- g 4OH-

(aq)

Advantage: Clean, portable and product is water. Efficient (75%) contrast to 20-25% car, 35-40% from coal electrical plant

Disadvantage: Cannot store electrical energy, needs continuous flow of reactant, Electrodes are short lived and expensive.


Corrosion Not all spontaneous redox reaction are beneficial. Natural redox process that oxidizes metal to their oxides and sulfides runs billions of dollars annually. Rust for example is not the direct product from reaction between iron and oxygen but arises through a complex electrochemical process.

Rust: Fe2O3 • X H2O Anode: Fe(s) g Fe+2 + 2e- E° = 0.44 V Cathode: O2 (g) + 4H+ + 4e- g 2H2O (l) E° = 1.23 V Net: Fe+2 will further oxidized to Fe2O3 • X H2O


Conditions for Corrosion Conditions for Iron Oxidation: Iron will oxidize in acidic medium

SO2 g H2SO4 g H+ + HSO4+

Anions improve conductivity for oxidation. Cl- from seawater or NaCl (snow melting) enhances rusting

Conditions for Prevention: Iron will not rust in dry air; moisture must be present Iron will not rust in air-free water; oxygen must be present Iron rusts most rapidly in ionic solution and low pH (high H+) The loss of iron and deposit of rust occur at different placm on object Iron rust faster in contact with a less active metal (Cu) Iron rust slower in contact with a more active metal (Zn)


Iron Corrosion; Chemistry

Iron will not rust in dry air; moisture must be present. Iron will not rust in air-free water; oxygen must be present Iron rusts most rapidly in ionic solutions and at low pH (High H+)

Most common and economically destructive form of corrosion is the rusting of iron. Rust is not a direct product of the reaction between iron and oxygen but arises through complex electrochemical process. The features of a voltaic cell can help explain this process.

The loss of iron and the depositing of rust often occur at different places on the same object. Iron rust faster in contact with a less active metal (such as Cu) and more slowly in contact with a more active metal (such as Zn).


Corrosion Prevention

Date post:	04-Sep-2019
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