Voltaic Cell

The Voltaic Cell

The Voltaic Cell  There

are two types of electric cells, the electrolytic cell and the voltaic or chemical cell.  In the electrolytic cell, a chemical reaction occurs when electricity passes through it. That is, electrical energy is converted into chemical energy.  In a voltaic cell, however, an electric current is produced when a chemical reaction occurs. That is, chemical energy is converted into electrical energy, as in dry cells or batteries.

A SIMPLE VOLTAIC CELL  If

two different types of metal are immersed into a salt solution, a chemical reaction occurs and a voltaic cell is obtained.  The two metals should have different positions in the Electrochemical Series.  If two metals are immersed in a salt solution, the more reactive of the two becomes the negative terminal (anode) as it donates electrons more readily than the other metal. The positive terminal

 Electrons

donated by the metal move from the negative terminal to positive terminal, resulting in a flow of electric current in the opposite direction.  An example of a voltaic cell is where a zinc plate and a copper plate are immersed in a solution of copper(II) sulphate.

voltmeter

e-

e-

V

e-

eZn

+

Cu

CuSO4 solution

The deflection of the galvanometer needle shows that electric current is passing through. Zinc becomes the negative terminal of the cell when the metal donates electrons. It is the anode. Zn(s) Zn2+(aq) + 2e Copper acts as the positive terminal when it receives electrons and Cu2+ ions are converted into Cu. It is the cathode. Cu2+(aq) + 2e- Cu(s) The overall reaction occuring in the cell is: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Daniell Cell This is another example of a simple voltaic cell. It consists of two metals immersed separately in aqueous solutions containing the ions of each metal. e-

V

-

+ Cu Plate

Figure 1(a)

Zn plate

Salt bridge

CuSO4 solution

ZnSO4 solution

e

V

+

+

-

Zn Porous pot

Cu

ZnSO4 solution

CuSO4 solution Figure 1(b)

In Figure 1(a), the two salt solutions are connected by a salt bridge, containing potassium chloride. In Figure 1(b), a porous pot is used to replace the salt bridge.

The functioning of the porous pot are: To separate copper(II) sulphate solution from zinc sulphate solution (the two electrolytes) To allow the electrons and ions to flow through it, thereby connection the circuit.  As

zinc is more reactive than copper, zinc becomes the negative terminal. It donates electron to become Zn2+ ions. Zn(s) Zn2+(aq) + 2eAt the positive terminal, Cu2+ ions the copper(II) sulphate solution accept 2 electrons to form copper. Cu2+(aq) + 2e- Cu(s)

The overall reaction is: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)  The negative terminal (Zn) contains more electrons than the positive (Cu) because electrons are being released by Zn and used up by Cu. Therefore electrons flow from the negative terminal to the positive terminal.  This process occurs because zinc is more electropositive than copper. It tends to donate electrons while Cu2+ ions tends to receive electrons.

The voltaic cell cell is represented by the following equation: Zn(s) / Zn2+(aq) || Cu2+(aq) / Cu(s) The more electropositive metal and its product are written on the left, while the less electropositive ions and product are written on the right. Other metallic pairs can also be used to produce voltaic cells. For example:  Mg(s) 

/ Mg2+(aq) || Pb2+(aq) / Pb(s)

The magnesium/lead cell has Mg immersed in magnesium nitrate, Mg(NO3)2, solution, while Pb is immersed in plumbum nitrate, Pb(NO3)2, solution. Magnesium, being more electropositive, becomes the negative terminal.

Mg(s)

Mg2+(aq) + 2e-

At the positive terminal, Pb2+ are discharged: Pb2+(aq) + 2e- Pb(s)

TYPES OF VOLTAIC CELLS There are two types of voltaic cells: Primary cells, which cannot be recharge Secondary cells, which can be recharged

Primary Cells

Secondary Cells

PRIMARY CELLS These are chemical cells which cannot be recharged. They include Daniell Cell, dry cell, alkaline cell and mercury cell. These cells cannot be used again once its chemical substances are used up. Primary cells store the chemical substances which react to produce an electric current.

Metal cover

+

Zinc casing

Carbon rod

Mixture of carbon powder and manganese (IV) oxide Ammonium chloride _

Figure 2(a) Dry Cell

Dry Cell 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) + 2eZn2+ ions, which form when Zn donates electrons, dissolve in the electrolyte.

At the positive terminal (carbon) NH+4 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.

Steel cover

Zn powder (-) metal rod Magnesium (IV) oxide (+)

Steel casing

Potassium hydroxide

Figure 2(b) Alkaline cell

Alkaline Cell Sodium hydroxide or potassium paste is used as the electrolyte. The negative terminal is made up of zinc powder. The manganese (IV) oxide mixed with carbon powder serves as the positive terminal. A metal rod in the centre acts as an inactive electrode which receives electrons. Equation: At the negative terminal: Zn (s) + 2OH-(aq) At the positive terminal: 2MnO2(s) + H2O(l) + 2eOverall reaction Zn(s) + 2MnO2(s) + H2O(l)

Zn(OH)2 + 2eMn2O3(s) +2OH-(aq) Zn(OH)2(s) + Mn2O3(s)

Zinc casing

Mercury (II) oxide solution

insulator

Potassium hydroxide (electrolyte)

Zinc powder in hydroxide

Figure 2(c) Mercury cell

Mercury cell The mercury cell is used mostly in calculators, wrist watches, cameras and other devices. The negative terminal is made up of zinc. The positive terminal is made up of mercury (II) oxide, HgO. The electrolyte is a mixture of potassium hydroxide and zinc oxide Equations: At the negative terminal: Zn (s) + 2OH-(aq) At

the positive terminal: HgO (s) + H2O(l) + 2e-

reaction: Zn (s) + HgO (s) + H2O(l)

Zn(OH)2(aq) + 2eHg(l) + 2OH-(aq)

Overall

Zn(OH)2(aq) + Hg (l)

SECONDARY CELLS The lead-acid accumulator and the nickelcadmium cell are examples of secondary cells, which can be recharged. It is recharged after it has discharged all charges from it.

Lead plate coated with PbO2

(+)

(-)

Concentrated sulphuric acid

Figure 2(d) Lead-acid accumulator

Lead-acid Accumulator This secondary cell is also called the car battery as it is used in cars and other vehicles. A reversible chemical reaction takes place in the lead-acid accumulator. The negative terminal is a lead plate that is immersed in a concentrated solution of sulphuric acid, 5M. The positive terminal is a lead plate coated with a layer of brown lead(IV) oxide. The accumulator consists of several such cells which are connected in series.

Reactions during discharge: At

the negative terminal: The lead electrode dissolves to form Pb2+ ions: Pb (s) Pb2+(aq) + 2eAt the positive terminal: PbO2 at the positive terminal receives electrons and reacts with hydrogen ions to form lead ions and water, as in: PbO2(s) + 2H+(aq) + 2e- Pb2+(aq) + 2H2O(l) During

the production of an electric current, the Pb2+ ions witch forms at the terminal react with SO42- ions in sulphuric acid to form a layer of white lead (II) sulphate around the electrodes. Pb2+(aq) + So2-4(aq) PbSO4(s)

When the accumulator is being used to produce electricity, the quantity of acid decreases and more water is formed. This means that the concentration of sulphuric acid decreases as the accumulator is used. Therefore, it should be recharged when the sulphuric acid becomes too dilute for further reaction.

Reaction during recharge of cell The accumulator is recharged by passing through an electric current in the opposite direction, that is, electrolysis is carried out to convert lead(II) sulphate dissolves. Sulphuric acid is formed again. At the negative terminal Pb2+(aq) + 2e- Pb(s) At the positive terminal Pb2+(aq) + 2H2O(l) PbO2(s) + 4H+(aq) +2ePbSO4(s) Pb2+(aq) + SO42-(aq)

Nickel-Cadmium Cell  Cadmium

acts as the negative terminal and nickel (IV) oxide, NiO2 as the positive terminal. The electrolyte here is potassium chloride  The chemical reactions are: Negative terminal: Cd (s) + 2OH-(aq) terminal: NiO2(s) + 2H2O(l) + 2e-

Cd(OH)2(s) +2e-

Positive

reaction Cd (s) +NiO2(s) + 2H2O(l)

Ni(OH)2(s) + Ni(OH)2(s)

Overall

Cd(OH)2(s) + Ni(OH)2(s)

Advantages and Disadvantages of Various Voltaic Cells Cell

Advantages

set up in the laboratory

Disadvantages

Daniell Cell

Easily

Wet

cell – electrolyte easily

Dry Cell

No

spillage Small in size Easily carried about Produces regular current and voltage Obtained in different sizes

Does

Alkaline Cell

Lasts

longer than dry cell (x10) Produces a higher and more regular current

Leakage

Mercury Cell

Small

Very

split Voltage cannot last

in size Produces regular current for a longer period of time Lasts a long time

not last Cannot be recharged Leakage can occur if cell cannot be used anymore

occurs if cell is not used anymore Expensive Cannot be recharge expensive Cannot be recharged Mercury produced is poisonous

Cell

Advantages

Disadvantages

Lead-acid Accumulator

Can

be recharged Produces a high voltage (12 V) for a long period Produce a high current (175 A) suitable for heavy duty

Spillage

Nickelcadmium cell

Can

Expensive

be recharged up to 500

of acid can

occur Big in size Heavy, difficult to be carried about Expensive Loses charge if not used for long

times Lower power density No spillage Transformer needed for Long – lasting (15 - 20 years) recharging cell Smaller than accumulator – portable

Comparison between the Electrolytic Cell and Voltaic Cell Similarities

•Consists of an anode and a cathode / a positive terminal and a negative terminal •Contains an electrolyte •Chemical reactions involves donating or receiving electrons •Positive and negative ions move to the electrodes in the electrolyte •Electrons move from the anode to the cathode

ELeCTrOLyTiC CeLL •Electrical energy → chemical energy •Electric current produces chemical reaction •Negative terminal – cathode •Positive terminal – anode •Electrons flow from anode (+) to cathode (-) •At the cathode (-), cations receive electrons •At the anode (+), anions release electrons •Carbon or different/same metal strips are used as electrodes

Differences

VoLTaiC CeLL •Chemical energy → electrical energy •Chemical reaction produces electric current •Negative terminal – anode •Positive terminal – cathode •Electrons flow from negative terminal (-) to positive terminal (+) •At the negative terminal (-), electrons are released •At the positive terminal (+), electrons are accepted •2 differents metals are used as electrons

The End