There Are Two Types of Voltaic Cell

8/6/2019 There Are Two Types of Voltaic Cell

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CHEMISTRY ASSIGNMENT

Elva Kumalasari XI-IC



There are two types of Voltaic Cell:

- Primary Cell, which cannot be recharge

- Secondary Cell, which can be recharge

Application of voltaic cell in battery (Dry Cell)

Dry Cell (one of secondary cell)

Chemistry is the driving force behind the magics of batteries.

A battery is a package of one or more galvanic cells used for the production and storage of electric

energy by chemical means. A galvanic cell consists of at least two half cells, a reduction cell and an

oxidation cell. Chemical reactions in the two half cells provide the energy for the galvanic cell

operations.

Each half cell consists of an electrode and an electrolyte solution. Usually the solution contains ions

derived from the electrode by oxidation or reduction reaction.

We will make this introduction using a typical setup as depicted here. The picture shows a copper

zinc galvanic cell (battery).

A galvanic cell is also called a voltaic cell. The spontaneous reactions in it provide the electric energy

or current.

Two half cells can be put together to form an electrolytic cell, which is used for electrolysis. In this

case, electric energy is used to force nonsponaneous chemical reactions.

- The positive terminal of a dry cell is a carbon rod, while the negative terminal is the zinc

casing around the cell. - The electrolyte includes a mixture of magnesium (IV) oxide and carbon powder, surrounded

by ammonium chloride powder.

- The chemical reaction which takes place are:

At the negative terminal, Zn:

Zn (s) Zn2+(aq) + 2e-

Zn2+ ions, which form when Zn donates electrons,dissolve in the electrolyte.

- At the positive terminal (carbon) NH4+ ions are discharged. They receive electrons to form

two gases, ammonia and hydrogen.

2NH+4(aq) 2NH3(g) + H2(g)

The hydrogen, which results in this reaction, reacts with manganese (IV) oxide as follows:

2MnO2(s) + H2(g) Mn2O3(g) + H2O(l)

- Overall reaction

Zn (s) + 2MnO2(s) + 2NH+4(aq) Zn2+(aq) + Mn2O3(s) + 2NH3(g) + H2O(l)

- Carbon powder is used to increase the surface area of the carbon electrode and manganese

(IV) oxide reduces the formation of gas bubbles.



Accu (one of primary cell)

Lead Accumulator

This can be recharged by passing a current through it in the reverse direction. The chemical processes that occur

at the electrodes during discharge are reversed by this. Thus the cell recovers its original state, except for some

energy loss. Such cells are called secondary cells or accumulators.

The lead-sulphuric acid cells is a common example. It was inverted by a French physicist, Gaston plate, in 1859.

Electrodes: Alternate parallel plates of lead dioxide (+ve electrode) (oxidised from PbO)

Spongy lead (reduced from PbO) (-ve electrode)

These are kept separate by porous separators made of wood, plastic or glass fibre.

Electrolytes: Dilute sulphuric acid

Container: Glass or bakelite

Here stored chemical energy is converted to electrical energy or current is drawn from the cell.

The hydrogen ions go to the +ve electrode and SO42-

to the -ve electrode. After giving their charges they react

with the electrodes and reduce the active material to lead sulphate.

Therefore, at the -ve electrode

At the +ve electrode,

Both plates (but only half of the active materials) are converted into PbSO4 (whitish). Water is formed thus

lowering the specific gravity of H2SO4 (electrolyte).

The emf of the cell falls and sulphuric acid is consumed.



Discharging Process

Here stored chemical energy is converted to electrical energy or current is drawn from the cell.

The hydrogen ions go to the +ve electrode and SO42-

to the -ve electrode. After giving their charges they react

with the electrodes and reduce the active material to lead sulphate.

Therefore, at the -ve electrode


Both plates (but only half of the active materials) are converted into PbSO4 (whitish). Water is formed thus

lowering the specific gravity of H2SO4 (electrolyte).

The emf of the cell falls and sulphuric acid is consumed.

Recharging Process

Current is passed through the two terminals in the reverse direction to that in which the cell provided current.

That is, the anode is connected to the positive terminal of the d.c. source, and the cathode to the negative

terminal.

The hydrogen ions move to the -ve electrode and sulphate ions to the +ve electrode.

At -ve electrode,


Water is consumed and sulphuric acid is formed thus raising the specific gravity of the electrolyte. In the charging

process, the +ve electrode is coated with dark brown lead peroxide and the -ve electrode with grey spongy lead.

The emf of the cell rises, and the electrical energy supplied is converted into chemical energy which is stored in

the cell.



Application of electrolysis

Production of Chemical

Example: production of sodium chloride, chlorine and hydrogen

Sodium hydroxide, NaOH, also known as lye and caustic soda, is one of the most important

of all industrial chemicals. It is produced at the rate of 25 billion pounds a year in the United

States alone. The major method for producing it is the electrolysis of brine or "salt water," a

solution of common salt, sodium chloride in water. Chlorine and hydrogen gases are

produced as valuable byproducts.

When an electric current is passed through salt water, the negative chloride ions, Cl-,

migrate to the positive anode and lose their electrons to become chlorine gas.

(The chlorine atoms then pair up to form Cl 2molecules.) Meanwhile, sodium ions, Na+, are

drawn to the negative cathode. But they do not pick up electrons to become sodium metal

atoms as they do in molten salt, because in a water solution the water molecules

themselves pick up electrons more easily than sodium ions do.

The hydroxide ions, together with the sodium ions that are already in the solution,

constitute sodium hydroxide, which can be recovered by evaporation.

This so-called chloralkali process is the basis of an industry that has existed for well over a

hundred years. By electricity, it converts cheap salt into valuable chlorine, hydrogen and

sodium hydroxide. Among other uses, the chlorine is used in the purification of water, the

hydrogen is used in the hydrogenation of oils, and the lye is used in making soap and paper.



Purification of metal/

efining metal

Example of purification of metal (purifying copper)

An example of the process described above is the refinement of copper. A common copper

containing ore is chalcocite (Cu2S). This ore is first treated at high temperatures (by a process called,

roasting) or by blowing oxygen through the melted ore. As the oxygen comes in contact with the

ionic complex it causes the reduction of the copper (I) ions in the ore to copper metal. At the same

time sulfur is oxidized to sulfur dioxide as shown in the following equation.

Cu2S + O2 -----> 2 Cu (s) + SO2 (g)

The copper metal that is formed by this reduction process still contains a small amount of impurities such as zinc, iron, silver, and gold. The impure copper is further refined via electrolysis. In this

process, as shown below, the impurities are removed in one of two ways. The more electropositive

zinc and iron are oxidized into their respective ions and enter into solution. The Noble metals, silver

and gold, are not oxidized at the anode, but settle out as metal atoms in a "sludge" as the impure

copper anode dissolves. The copper obtained by this refinement is about 99.5% pure, and the

amount of silver and gold recovered in the process is often sufficient to pay for the cost the

electricity required for the electrolytic process.



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oplating cell are

oth connected to an e

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attery or,

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commonly, a rectifier. The anode is connected to the positive terminal of the s pply, and the cathode

article to e plated) is connected to the negative

terminal.

hen the e

ternal power s pply is

switched on, the metal at the anode isoxidized from

the zero valence state to form cations with a

positive charge. These cations associate with the

anions in the solution. The cations are reduced at

the cathode to deposit in the metallic, zero valence

state. For example, in an acid solution, copper is

oxidized at the anode to u

2+

y losing two

electrons. The u

2+ associates with the anion S 4

2- in

the solution to form copper sulfate. !

t the cathode, the

u

2+ is reduced to metallic copper

y gaining

two electrons. The result is the eff ective transf er of

copper from the anode source to a plate covering

the cathode.

The plating is most commonly a single metallic element, not an alloy. However, some alloys

can e electrodeposited, notablybrass and solder.

" any plating baths include cyanides of other metals

e.g., potassium cyanide) in addition

to cyanides of the metal to be deposited. These free cyanides facilitate anode corrosion, help to

maintain a constant metal ion level and contribute to conductivity. !

dditionally, non-metal

chemicals such as carbonates and phosphates may be added to increase conductivity.



Corrosion

Definition

Corrosion is the disintegration of an engineered material into its constituent atoms due to chemical

reactions with its surroundings. In the most common use of the word, this means electrochemical

oxidation of metals in reaction with an oxidant such as oxygen. Formation of an oxide of iron due to oxidation of the iron atoms in solid solution is a well-known example of electrochemical corrosion,

commonly known as rusting. This type of damage typically produces oxide(s) and/or salt(s) of the

original metal. Corrosion can also refer to other materials than metals, such as ceramics or polymers,

although in this context, the term degradation is more common

Reaction

It has been demonstrated that potential differences within a metal, or between two metals, will

cause chemical reactions at the anode and cathode. Anodic reactions are typified by the dissolution

of iron:

Fe --> Fe+2 + 2e-

Analogous reactions occur in other metals. The electrons migrate through the metal to the cathode

area where they react in any one of several ways.



Some typical cathodic reactions are as follows:

a. Hydrogen ion reduction 2 # + + 2e- ---- > H2 Important in acidic solutions.

b. Reduction of water 2H20 + 2e- ---- > H2+2OH- Occurs normally in natural waters.

c. Oxygen reduction 02 + 4H+ +4e- ---- > 2H20 Occurs in aerated acidic solutions.

d. Oxygen reduction of water O2 + 2H2O + 4e- --- > 40H- Important in natural, aerated waters.

e. Ferric ion reduction Fe+3 + e- --- > Fe+2 Occurs under acidic, turbulent conditions (e.g.

acid cleaning).

f. Sulfate ion reduction 4H2 + S04-2 ---- > S-2 + 4H2O Occurs in the presence of sulfate

reducing bacteria.

g. Metal ion reduction (plating) $ +N + Ne- ----> MO Involves more noble metals in solution. The

most frequent cathodic reactions are a, b, c and d.

Negatively charged ions, such as hydroxyl ions produced at the cathode, migrate to the anode of the

corrosion cell. Positively charged ions will move toward the cathode.

This movement of ions can cause additional reactions at the anode.

Hydroxyl ions will combine with the ferrous cations produced by dissolution of the metal:

Fe+2 + 2OH- ----> Fe(OH)2

The ferrous hydroxide produced has a very low solubility and is quickly precipitated as a white floe at

the metal-water interface. The floe is then rapidly oxidized to ferric hydroxide:

4Fe(OH)2 + O2 + 2H2O ----> 4Fe(OH)3

Dehydrolysis of this product leads to the formation of the corrosion products normally seen on

ferrous surfaces, red dust and hydrated ferric oxide:

2Fe(OH)3 ----> Fe203 + 3H2O Fe(OH)3 ----> FeOOH + H2O



As solid corrosion products are precipitated at the anode, they may cause the precipitation of other

ions from the water. Thus, a corrosion film may show traces of hardness salts, or suspended matter

like mud, sand, silt, clay or microbiological slime.

The structure of the entire surface film, including corrosion products and inclusions, is a major factor

in determining the total amount of corrosion which will take place. If a porous film forms over the

metal, corrosion can continue, because metal ions can penetrate it and reach the solution interface. If, however, a tight, adherent film is formed, ionic diffusion is prevented and the metal will no longer

dissolve.

Most corrosion occurs at the beginning of a metal's service life. Initially, metal dissolution is not

impeded by a film of corrosion products. In time, the film will retard, or halt the corrosion. The

degree to which such a film can impede corrosion is a complex function of the corrosion reactions,

the structure of the deposit and the water velocity



Factors inf luencing corrosion

Accordingly on this basis we list below some of the more important factors, discussing their general

significance with respect to the mechanism of corrosion, and postponing until later chapters the

detailed discussion of others.

Factors Associated Mainly with the Metal

y Effective electrode potential of a metal in a solution

y Overvoltage of hydrogen on the metal

y Chemical and physical homogeneity of the metal surface

y Inherent ability to form an insoluble protective film

Factors Which Vary Mainly with the Environment

y Hydrogen-ion concentration (pH) in the solution

y Influence of oxygen in solution adjacent to the metal

y Specific nature and concentration of other ions in solution

y Rate of flow of the solution in contact with the metal

y

Ability of environment to form a protective deposit on the metal y Temperature

y Cyclic stress (corrosion fatigue)

y Contact between dissimilar metals or other materials as affecting localized corrosion.

Cause of corrosion, many factors affecting corrosion, the type, speed, cause, and seriousness of

metal corrosion. Some of these factors can be controlled and some can not. C limate. The

environmental conditions under which an aircraft is maintained and operated greatly affect

corrosion characteristics. In a predominately marine environment ( with exposure to sea water and

salt air ), moisture-laden air is considerably more detrimental to an aircraft than it would be if all

operations were conducted in a dry climate. Temperature considerations are important because the

speed of electrochemical attack is in creased in a hot, moist climate. Size and Type of Metal. It is

well known fact that some metals will corrode faster that others. It is a less known fact that

variations in size and shape of metal can indirectly affect is corrosion resistance. Thick structural

sections are more susceptible to corrosive attack that thin sections because variations in physical

characteristics are greater. When large pieces are machined or chemically milled after heat

treatment, the thinner areas will have different physical characteristics than the thicker areas.



Preventing corrosion

Sacrificial Protection

'Rusting' can be prevented by connecting iron to a more reactive metal (e.g., zinc or magnesium). This is referred

to as sacrificial protection or sacrificial corrosion, because the more reactive protecting metal is preferentially

oxidized away, leaving the protected metal intact.

Alloying

Iron or steel along with other metals can also be protected by 'alloying' or mixing with other metals (e.g.,

chromium) to make non-rusting alloys. Stainless steel is an example of a non-rusting alloy of iron and carbon.

Brass, an alloy containing copper is another metal alloy which is less expensive and non reactive.

Galvanizing

Coating iron or steel with a thin zinc layer is called 'galvanizing'. This layer is produced by electrolytic deposition.

Dipping the iron/steel object in molten zinc and using it as the negative cathode zinc is coated on it. Zinc

preferentially corrodes or oxidizes to form a zinc oxide layer that does not flake off like iron oxide rust. Also, if thesurface is scratched, the exposed zinc again corrodes before the iron and continues to protect it.

Electroplating

Coating the surface with metals like tin, chromium, nickel etc. by electroplating is also utilized to prevent

corrosion. Steel cans are protected by relatively un-reacted tin and works well as long as the thin tin layer is

complete.

Prevention for iron rusting :

Barrier protectionThe metal surface is not allowed to come in contact with moisture, O 2 and CO2.

i) Coating the metal surface with paint.

ii) Applying oil or grease.

iii) Electroplating with non-corroding metals like Ni, Cr, Al, Sn, Zn.

iv) Coating with alkaline phosphate (anti rust) solution.

Sacrificial protection

Covering the surface with a more electro positive metal than Fe. The more electro positive metal loses electrons

and as long as this coating is present Fe is protected.

Example:

Galvanization - Covering with zinc.

The zinc also forms a protective coating of ZnCO3.Zn(OH)2.



Electrical protection (Cathodic protection)

The iron object is connected to a more active metal either directly or through a wire. Fe acts as the cathode. The

more reactive metal is the anode. It loses electrons and gradually disappears.

Example: Fe can be connected to Mg, Zn or Al which are called the sacrificial anodes. Used for protecting under

ground pipes from rusting.

Elva Kumalasari XI KI

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There Are Two Types of Voltaic Cell

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