Metal Ions/Metal Systems
If a metal plate or rod is placed in a solution of its own salt, two processes - oxidation and reduction, take place simultaneously. The atoms from the metal's surface, lose electrons and go into solution as ions, as the electrons lost are left on the surface of the metal plate, thereby making it electrons surplus and negatively charged.
M(s) ----> M+(aq) + e-
Conversely, the metallic ions in the salt solution gain electrons from the surface of the metal plate and get deposited as metallic atoms. The electrons gained from the surface of the electrode renders it electron deficient and positively charged.
M+(aq) + e- ----> M(s)
Depending on the nature of the metal, a particular process will predominate. Hence, a potential difference, known as the electrode potential for the metal ions/metal system [M+(aq)/M(s)] is established between the electrode and the electrolyte.
For instance, if a zinc electrode is placed in a solution of zinc tetraoxosulphate (VI) solution, ZnSO4, the zinc atoms from the electrode, spontaneously lose electrons and go into solution as zinc ions, Zn2+. Hence, the zinc electrode becomes negatively charged on the surface due to excess electrons, while the solution becomes positively charged due to the presence of excess Zn2+ in it. A potential difference, known as the electrode potential of zinc, is therefore set up between the zinc electrode and the solution.
As this continues, the zinc rod keeps diminishing in size until all the metallic atoms are converted to ions.
Zn(s) ----> Zn2+(aq) + 2e-
electrode solution
(decreases in size)
Similarly, if a copper electrode is dipped into a solution of copper (II) tetraoxosulphate (VI), CuSO4. The copper (II) ions, Cu2+, from the solution pick up electrons from the copper rod and get deposited as metallic copper atoms on the electrode. Consequently, the copper rod becomes positively charged on the surface due to a deficit of electrons, while the solution becomes negatively charged because of the excess tetraoxosulphate (VI) ions in solution, SO4--.
Also, as this continues, the electrode keeps increasing in size until all the metallic ions in the solution have been converted to metallic atoms.
Cu2+(aq) + 2e- ----> Cu(s)
solution electrode
(increases in size)
The two systems discussed above are known as half cells. If the two electrodes are connected to a galvanometer, by means of a wire, it would be observed that current flows from the zinc electrode to the copper electrode. The zinc electrode serves as the negative terminal or the anode, while the copper electrode serves the positive terminal or the cathode. This is because zinc, which is more electropositive than copper, will always get oxidized by losing electrons in the presence of copper, while the latter will always gain electrons and get reduced in the presence of zinc.
Zn(s) ----> Zn2+(aq) + 2e- (oxidation half-equation)
Cu2+(aq) + 2e- ----> Cu(s) (reduction half-equation)
Zn(s) + Cu2+(aq) ----> Zn2+(aq) + Cu(s) (overall equation)
The set-up consisting of two half-cells with different electrode potentials is discussed further under electrochemical cell.
Standard Electrode Potential
The electrode potential of a given metal ions/metal system depends on three factors, namely - the overall energy change, the concentration of ions in solution and the temperature. Due to these, the standard electrode potential is always used to compare the electrode potential of two or more metal ions/metal systems.
The standard electrode potential of a metal, E°, is the value (based on the hydrogen scale) recorded by a voltmeter, when its electrode potential is measured at standard conditions by connecting the metal electrode placed in 1M concentration of its salt solution, to a hydrogen electrode at 25°C (298K) and 1atm (760mmHg/101325Nm^-2), with the hydrogen electrode having an arbitrary value of 0.00V at all temperatures.
Depending on the nature and position of the metal in the electrochemical series, electrons will either flow from it to the hydrogen electrode or vice-versa. If the electrode potential is negative, it indicates that electrons flow from the metal electrode (anode) to the hydrogen electrode (cathode); and if positive, electrons flow from the hydrogen electrode (anode) to the metal electrode (cathode).
Therefore, the electrode potential of the metal ions/metal system can be calculated as:
E°cell = E°cathode - E°anode .......(i), or
E°cell = E°reduction - E°oxidation .......(ii)
where E°cell is the voltmeter reading or the electromotive force (e.m.f.) of the cell.
For instance, if a Zn2+(aq)/Zn(s) half-cell is used, electrons will flow from the zinc electrode to the hydrogen electrode, because zinc is higher than hydrogen in the electrochemical series, and as such, will be oxidized by losing electrons in the presence of hydrogen. The voltmeter records a value of 0.76V, which therefore gives the electrode potential of the zinc as -0.76V. As stated earlier, the negative sign indicates that the zinc electrode served as the anode and was oxidized, while the hydrogen electrode served as the cathode and was reduced. The two processes can be observed by the decrease in size of the anode and the liberation of hydrogen gas at the cathode according to the reactions:
Zn(s) ----> Zn2+(aq) + 2e- (anodic half-reaction/oxidation half-equation)
2H+(aq) + 2e- ----> H2(g) (cathodic half-reaction/reduction half-equation)
Zn(s) + 2H+(aq) ----> Zn2+(aq) + H2(g) (overall reaction)
This can also be confirmed from the calculation of the emf of the cell, by using equation (i),
E°cell = E°cathode - E°anode
0.76 = 0.00 - E°Zn
E°Zn = -0.76V
Or, by using equation (ii),
E° = E°reduction - E°oxidation
0.76 = 0.00 - E°Zn
E°Zn = -0.76V
On the other hand, when a Cu2+(aq)/Cu(s) half-cell is used, electrons flow from the hydrogen electrode to the copper electrode because copper is lower than hydrogen in the electrochemical series. Therefore, the copper ions in the solution are reduced to metallic copper atoms in the presence of hydrogen. The voltmeter's reading would be seen to be 0.34V, from which the electrode potential of copper is obtained as +0.34V, which shows that the copper electrode served as the cathode, while the hydrogen electrode served as the anode.
H2(g) ----> 2H+(aq) + 2e- (anodic half-reaction/oxidation half-equation)
Cu2+(aq) + 2e- ----> Cu(s) (cathodic half-reaction/reduction half-equation)
H2(g) + Cu2+(aq) ----> 2H+(aq) + Cu(s) (overall reaction )
This can be confirmed by calculating the emf of the cell, by using equation (i),
E° = E°cathode - E°anode
0.34 = E°Cu - 0.00
E°Cu = +0.34V
Or, by using equation (ii),
E° = E°reduction - E°oxidation
0.34 = E°Cu - 0.00
E°Cu = +0.34V
Electrochemical Cell
The study of electrode potential is a branch of electrochemistry that deals with the generation of electricity using chemical reactions. It is applied in the electrochemical cell, which works based on the difference in the standard electrode potentials of the electrodes in two half-cells.
An electrochemical cell is a device used to convert chemical energy into electrical energy. It generates electricity from chemical reactions. A group of electrochemical cells is known as a battery.
Salt-Bridge & Electrochemical Cells
Let us consider an electrochemical cell made up of the two half-cells - the zinc ions/zinc metal system and the copper ions/copper metal system.
In the Zn(s)/Zn2+(aq) half-cell, where oxidation takes place, the net electrical charge of the electrolyte is positive (Zn2+ > SO4-- ), while the net charge of the electrolyte is negative (SO4-- > Cu2+) in the Cu2+(aq)/Cu(s) half-cell, where reduction takes place.
Therefore, to take care of the imbalance in electrical neutrality between the two half-cells, a salt-bridge is usually used to connect both electrolytes. (Please refer to your textbooks for the diagram).
It is made up of a porous material, e.g. filter paper, saturated with a salt solution, say sodium chloride [NaCl(aq)]. As the electrolyte in the oxidation half-cell becomes increasingly positively charged, the presence of the excess Zn2+ ions are counterbalanced by the Cl- ions, which go into solution from the salt-bridge, when one end of the bridge is dipped into the electrolyte.
Also, the net negative charge of the electrolyte in the reduction half-cell, due to the excess SO4--, is neutralized by the introduction of Na+ ions from the salt-bridge, when its other end is dipped into the electrolyte. As some of the cations and anions from the salt-bridge enter the respective electrolytes, some of the excess cations and anions from the electrolytes move out of their respective solutions via the salt-bridge, thereby, establishing the electrical neutrality, and completing the electrical circuit.
Types of Electrochemical Cells
There are two types of electrochemical cells, namely primary cells and secondary cells.
Primary cells are non-rechargeable. When once the chemicals in the cells are used up, they cannot be regenerated and the electrochemical cells will stop discharging electric current. They will, therefore, need to be changed. Most primary cells are dry cells. Examples of primary cells are Daniell Cell, Leclanché Cell etc.
Leclanché Cell
The Leclanché cell is the commonly used battery in our torches, transistor radios, remote controllers and toys. An example is the Duracell battery.
A typical Leclanché cell is cylindrical in shape and is made up of a solution or paste of ammonium chloride as the electrolyte and carbon rod (graphite) as the cathode. Both are enclosed within a thin zinc sheet, which serves as the anode. In between the cathode and the electrolyte is a paste of manganese (IV) oxide, placed inside a thin film of muslin bag. (Please refer to your textbooks for the diagram).
At the anode, the zinc atoms lose two electrons each to dissolve in the electrolyte as zinc ions
Zn(s) ----> Zn2+(aq) + 2e-(oxidation half-reaction)
The electrons lost by the zinc atoms pass through an external circuit to do some work such as lighting-up a torch bulb, making an electronic bell or device work before reaching the cathode.
At the cathode, the ammonium ions accept the electrons to become ammonia and hydrogen gas. The hydrogen gas is removed by the manganese (IV) oxide, which prevents it from adhering to the cathode.
2NH4+(aq) + 2e- ----> 2NH3(g) + H2(g) (reduction half-reaction)
Zn(s) + 2NH4+(aq) ----> Zn2+(aq) + 2NH3(g) + H2(g) (overall reaction)
The cell stops working immediately the zinc anode or ammonium chloride electrolyte is used up; hence, the need for its replacement.
Secondary Cells are those that store electrical energy, when connected to an alternating current (a.c.) power supply, and discharge the stored energy as a direct current (d.c.), when made to work. That is why they are also known as storage cells. Unlike primary cells, they are rechargeable, because the chemicals in the cells can be regenerated whenever they are used up, through the process of electrolysis. They are made up of wet and dry cells. The wet cells have liquid electrolytes, while the dry cells have electrolytes in molten or pasty form. Examples of secondary cells include the lead accumulator, nickel or lithium-ion accumulator etc.
Lead Accumulator
In the lead accumulator, its anode is made of lead, while the cathode is made of lead (IV) oxide, PbO2; the electrolyte is made up of concentrated tetraoxosulphate (VI) acid, H2SO4. (Please refer to your textbooks for the diagram).
Being a secondary cell, the lead accumulator can be discharged and recharged. An example is the car battery.
Discharging
At the anode, the lead atoms undergo oxidation to form lead (II) ions, Pb2+, by losing two electrons each.
Pb(s) ----> Pb2+(aq) + 2e- (oxidation half-reaction)
Just like the primary cell, the electrons move through an external circuit and perform some work, such as powering the engine of a vehicle, or a mobile phone before arriving at the cathode.
At the cathode, the PbO2 undergo reduction to form Pb2+, by accepting two electrons each in the presence of the hydrogen ions, H+, from the electrolyte.
PbO2(s) + 4H+(aq) + 2e- ----> Pb2+(aq) + 2H2O(l) (reduction half-reaction)
The Pb2+ ions produced at both electrodes combine with the tetraoxosulphate (VI) ions, SO4--, from the electrolyte to deposit lead (II) tetraoxosulphate (VI), PbSO4, which eventually covers the electrodes.
Pb2+(aq) + SO4--(aq) ----> PbSO4(s)
The PbSO4 and the water molecules produced in the half-reactions decrease the density of the acid until it falls below 1.16gdm^-3, due to the removal of the H+ and SO4-- ions from the electrolyte. Similarly, the e.m.f. of the cell drops to 1.8V.
Consequently, the cell stops discharging current because the chemicals have been used up, and need to be regenerated. (At this point, we say the battery has 'run' down and needs to recharged.)
Recharging
Here, the lead accumulator functions as an electrolytic cell. When connected to a power supply, the anode in the discharged cell becomes the cathode, while the cathode becomes the anode. This is because the anode and cathode in the discharged cell are connected to the cathode and anode of the power source respectively. (Refer to your textbooks for diagrams of the lead accumulator during discharging and recharging).
At the anode, the PbSO4 dissociates to form Pb2+ and SO4-- ions. These Pb2+ ions then combine with the water molecules to form PbO2 and H+ ions
PbSO4(s) ----> Pb2+(aq) + SO4--(aq)
Pb2+(aq) + 2H2O(l) ----> PbO2(s) + 4H+(aq) + 2e- (oxidation half-reaction)
At the cathode, the PbSO4 dissociates in the same way as it does at the anode. However, the Pb2+ ions produced undergo reduction by accepting two electrons each and are deposited on the cathode as lead metal.
PbSO4(s) ----> Pb2+(aq) + SO4--(aq)
Pb2+(aq) + 2e- ----> Pb(s) (reduction half-reaction)
The 4 moles H+ ions (produced in the oxidation half-reaction) combine with the 2 moles of SO4-- ions (produced from the dissociation of PbSO4 at both electrodes) to regenerate the tetraoxosulphate (VI) acid, H2SO4.
4H+(aq) + 2SO4--(aq) ----> 2H2SO4(aq)
2H+(aq) + SO4--(aq) ----> H2SO4(aq)
Consequently, the density of the acid in the electrochemical cell returns to the optimal level of 1.25gcm^-3 and its e.m.f. increases back to 2.2V, after recharging.
Similarities between Electrolytic and Electrochemical Cells
1. Both cells are made up of the anode, the cathode and the electrolyte.
2. In both cells, oxidation occurs at the anode, while reduction takes place at the cathode.
3. In both cells, electrons flow from the anode to the cathode.
Differences between Electrolytic and Electrochemical Cells
1. In an electrolytic cell, the anode is the positive terminal, while the cathode is the negative terminal; whereas, in an electrochemical cell, the anode is the negative terminal, and the cathode is the positive terminal.
2. An electrolytic cell is used to convert electrical energy into chemical energy, while an electrochemical cell uses chemical energy to generate electrical energy.
3. In an electrolytic cell, no salt-bridge is required, whereas, a salt-bridge is needed in an electrochemical cell.
4. In an electrolytic cell, there is only one electrolyte, while an electrochemical cell can contain more than one electrolyte.
5. In an electrolytic cell, the two electrodes are in the same compartment, while in an electrochemical cell, the electrodes may be in the same or different compartments.
6. An electrolytic cell is known as a voltameter, while an electrochemical cell is called a galvanic cell.
Do These:
Question 1
a) What do you understand as the electrode potential of a metal?
b) Describe the structure of the Leclanché cell and give the chemical reactions that occur in it.
Question 2
The reactions that occur inside a car battery are said to be redox. Explain this statement using suitable equations to support your points.
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WhatsApp: 07034776117
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If a metal plate or rod is placed in a solution of its own salt, two processes - oxidation and reduction, take place simultaneously. The atoms from the metal's surface, lose electrons and go into solution as ions, as the electrons lost are left on the surface of the metal plate, thereby making it electrons surplus and negatively charged.
M(s) ----> M+(aq) + e-
Conversely, the metallic ions in the salt solution gain electrons from the surface of the metal plate and get deposited as metallic atoms. The electrons gained from the surface of the electrode renders it electron deficient and positively charged.
M+(aq) + e- ----> M(s)
Depending on the nature of the metal, a particular process will predominate. Hence, a potential difference, known as the electrode potential for the metal ions/metal system [M+(aq)/M(s)] is established between the electrode and the electrolyte.
For instance, if a zinc electrode is placed in a solution of zinc tetraoxosulphate (VI) solution, ZnSO4, the zinc atoms from the electrode, spontaneously lose electrons and go into solution as zinc ions, Zn2+. Hence, the zinc electrode becomes negatively charged on the surface due to excess electrons, while the solution becomes positively charged due to the presence of excess Zn2+ in it. A potential difference, known as the electrode potential of zinc, is therefore set up between the zinc electrode and the solution.
As this continues, the zinc rod keeps diminishing in size until all the metallic atoms are converted to ions.
Zn(s) ----> Zn2+(aq) + 2e-
electrode solution
(decreases in size)
Similarly, if a copper electrode is dipped into a solution of copper (II) tetraoxosulphate (VI), CuSO4. The copper (II) ions, Cu2+, from the solution pick up electrons from the copper rod and get deposited as metallic copper atoms on the electrode. Consequently, the copper rod becomes positively charged on the surface due to a deficit of electrons, while the solution becomes negatively charged because of the excess tetraoxosulphate (VI) ions in solution, SO4--.
Also, as this continues, the electrode keeps increasing in size until all the metallic ions in the solution have been converted to metallic atoms.
Cu2+(aq) + 2e- ----> Cu(s)
solution electrode
(increases in size)
The two systems discussed above are known as half cells. If the two electrodes are connected to a galvanometer, by means of a wire, it would be observed that current flows from the zinc electrode to the copper electrode. The zinc electrode serves as the negative terminal or the anode, while the copper electrode serves the positive terminal or the cathode. This is because zinc, which is more electropositive than copper, will always get oxidized by losing electrons in the presence of copper, while the latter will always gain electrons and get reduced in the presence of zinc.
Zn(s) ----> Zn2+(aq) + 2e- (oxidation half-equation)
Cu2+(aq) + 2e- ----> Cu(s) (reduction half-equation)
Zn(s) + Cu2+(aq) ----> Zn2+(aq) + Cu(s) (overall equation)
The set-up consisting of two half-cells with different electrode potentials is discussed further under electrochemical cell.
Standard Electrode Potential
The electrode potential of a given metal ions/metal system depends on three factors, namely - the overall energy change, the concentration of ions in solution and the temperature. Due to these, the standard electrode potential is always used to compare the electrode potential of two or more metal ions/metal systems.
The standard electrode potential of a metal, E°, is the value (based on the hydrogen scale) recorded by a voltmeter, when its electrode potential is measured at standard conditions by connecting the metal electrode placed in 1M concentration of its salt solution, to a hydrogen electrode at 25°C (298K) and 1atm (760mmHg/101325Nm^-2), with the hydrogen electrode having an arbitrary value of 0.00V at all temperatures.
Depending on the nature and position of the metal in the electrochemical series, electrons will either flow from it to the hydrogen electrode or vice-versa. If the electrode potential is negative, it indicates that electrons flow from the metal electrode (anode) to the hydrogen electrode (cathode); and if positive, electrons flow from the hydrogen electrode (anode) to the metal electrode (cathode).
Therefore, the electrode potential of the metal ions/metal system can be calculated as:
E°cell = E°cathode - E°anode .......(i), or
E°cell = E°reduction - E°oxidation .......(ii)
where E°cell is the voltmeter reading or the electromotive force (e.m.f.) of the cell.
For instance, if a Zn2+(aq)/Zn(s) half-cell is used, electrons will flow from the zinc electrode to the hydrogen electrode, because zinc is higher than hydrogen in the electrochemical series, and as such, will be oxidized by losing electrons in the presence of hydrogen. The voltmeter records a value of 0.76V, which therefore gives the electrode potential of the zinc as -0.76V. As stated earlier, the negative sign indicates that the zinc electrode served as the anode and was oxidized, while the hydrogen electrode served as the cathode and was reduced. The two processes can be observed by the decrease in size of the anode and the liberation of hydrogen gas at the cathode according to the reactions:
Zn(s) ----> Zn2+(aq) + 2e- (anodic half-reaction/oxidation half-equation)
2H+(aq) + 2e- ----> H2(g) (cathodic half-reaction/reduction half-equation)
Zn(s) + 2H+(aq) ----> Zn2+(aq) + H2(g) (overall reaction)
This can also be confirmed from the calculation of the emf of the cell, by using equation (i),
E°cell = E°cathode - E°anode
0.76 = 0.00 - E°Zn
E°Zn = -0.76V
Or, by using equation (ii),
E° = E°reduction - E°oxidation
0.76 = 0.00 - E°Zn
E°Zn = -0.76V
On the other hand, when a Cu2+(aq)/Cu(s) half-cell is used, electrons flow from the hydrogen electrode to the copper electrode because copper is lower than hydrogen in the electrochemical series. Therefore, the copper ions in the solution are reduced to metallic copper atoms in the presence of hydrogen. The voltmeter's reading would be seen to be 0.34V, from which the electrode potential of copper is obtained as +0.34V, which shows that the copper electrode served as the cathode, while the hydrogen electrode served as the anode.
H2(g) ----> 2H+(aq) + 2e- (anodic half-reaction/oxidation half-equation)
Cu2+(aq) + 2e- ----> Cu(s) (cathodic half-reaction/reduction half-equation)
H2(g) + Cu2+(aq) ----> 2H+(aq) + Cu(s) (overall reaction )
This can be confirmed by calculating the emf of the cell, by using equation (i),
E° = E°cathode - E°anode
0.34 = E°Cu - 0.00
E°Cu = +0.34V
Or, by using equation (ii),
E° = E°reduction - E°oxidation
0.34 = E°Cu - 0.00
E°Cu = +0.34V
Electrochemical Cell
The study of electrode potential is a branch of electrochemistry that deals with the generation of electricity using chemical reactions. It is applied in the electrochemical cell, which works based on the difference in the standard electrode potentials of the electrodes in two half-cells.
An electrochemical cell is a device used to convert chemical energy into electrical energy. It generates electricity from chemical reactions. A group of electrochemical cells is known as a battery.
Salt-Bridge & Electrochemical Cells
Let us consider an electrochemical cell made up of the two half-cells - the zinc ions/zinc metal system and the copper ions/copper metal system.
In the Zn(s)/Zn2+(aq) half-cell, where oxidation takes place, the net electrical charge of the electrolyte is positive (Zn2+ > SO4-- ), while the net charge of the electrolyte is negative (SO4-- > Cu2+) in the Cu2+(aq)/Cu(s) half-cell, where reduction takes place.
Therefore, to take care of the imbalance in electrical neutrality between the two half-cells, a salt-bridge is usually used to connect both electrolytes. (Please refer to your textbooks for the diagram).
It is made up of a porous material, e.g. filter paper, saturated with a salt solution, say sodium chloride [NaCl(aq)]. As the electrolyte in the oxidation half-cell becomes increasingly positively charged, the presence of the excess Zn2+ ions are counterbalanced by the Cl- ions, which go into solution from the salt-bridge, when one end of the bridge is dipped into the electrolyte.
Also, the net negative charge of the electrolyte in the reduction half-cell, due to the excess SO4--, is neutralized by the introduction of Na+ ions from the salt-bridge, when its other end is dipped into the electrolyte. As some of the cations and anions from the salt-bridge enter the respective electrolytes, some of the excess cations and anions from the electrolytes move out of their respective solutions via the salt-bridge, thereby, establishing the electrical neutrality, and completing the electrical circuit.
Types of Electrochemical Cells
There are two types of electrochemical cells, namely primary cells and secondary cells.
Primary cells are non-rechargeable. When once the chemicals in the cells are used up, they cannot be regenerated and the electrochemical cells will stop discharging electric current. They will, therefore, need to be changed. Most primary cells are dry cells. Examples of primary cells are Daniell Cell, Leclanché Cell etc.
Leclanché Cell
The Leclanché cell is the commonly used battery in our torches, transistor radios, remote controllers and toys. An example is the Duracell battery.
A typical Leclanché cell is cylindrical in shape and is made up of a solution or paste of ammonium chloride as the electrolyte and carbon rod (graphite) as the cathode. Both are enclosed within a thin zinc sheet, which serves as the anode. In between the cathode and the electrolyte is a paste of manganese (IV) oxide, placed inside a thin film of muslin bag. (Please refer to your textbooks for the diagram).
At the anode, the zinc atoms lose two electrons each to dissolve in the electrolyte as zinc ions
Zn(s) ----> Zn2+(aq) + 2e-(oxidation half-reaction)
The electrons lost by the zinc atoms pass through an external circuit to do some work such as lighting-up a torch bulb, making an electronic bell or device work before reaching the cathode.
At the cathode, the ammonium ions accept the electrons to become ammonia and hydrogen gas. The hydrogen gas is removed by the manganese (IV) oxide, which prevents it from adhering to the cathode.
2NH4+(aq) + 2e- ----> 2NH3(g) + H2(g) (reduction half-reaction)
Zn(s) + 2NH4+(aq) ----> Zn2+(aq) + 2NH3(g) + H2(g) (overall reaction)
The cell stops working immediately the zinc anode or ammonium chloride electrolyte is used up; hence, the need for its replacement.
Secondary Cells are those that store electrical energy, when connected to an alternating current (a.c.) power supply, and discharge the stored energy as a direct current (d.c.), when made to work. That is why they are also known as storage cells. Unlike primary cells, they are rechargeable, because the chemicals in the cells can be regenerated whenever they are used up, through the process of electrolysis. They are made up of wet and dry cells. The wet cells have liquid electrolytes, while the dry cells have electrolytes in molten or pasty form. Examples of secondary cells include the lead accumulator, nickel or lithium-ion accumulator etc.
Lead Accumulator
In the lead accumulator, its anode is made of lead, while the cathode is made of lead (IV) oxide, PbO2; the electrolyte is made up of concentrated tetraoxosulphate (VI) acid, H2SO4. (Please refer to your textbooks for the diagram).
Being a secondary cell, the lead accumulator can be discharged and recharged. An example is the car battery.
Discharging
At the anode, the lead atoms undergo oxidation to form lead (II) ions, Pb2+, by losing two electrons each.
Pb(s) ----> Pb2+(aq) + 2e- (oxidation half-reaction)
Just like the primary cell, the electrons move through an external circuit and perform some work, such as powering the engine of a vehicle, or a mobile phone before arriving at the cathode.
At the cathode, the PbO2 undergo reduction to form Pb2+, by accepting two electrons each in the presence of the hydrogen ions, H+, from the electrolyte.
PbO2(s) + 4H+(aq) + 2e- ----> Pb2+(aq) + 2H2O(l) (reduction half-reaction)
The Pb2+ ions produced at both electrodes combine with the tetraoxosulphate (VI) ions, SO4--, from the electrolyte to deposit lead (II) tetraoxosulphate (VI), PbSO4, which eventually covers the electrodes.
Pb2+(aq) + SO4--(aq) ----> PbSO4(s)
The PbSO4 and the water molecules produced in the half-reactions decrease the density of the acid until it falls below 1.16gdm^-3, due to the removal of the H+ and SO4-- ions from the electrolyte. Similarly, the e.m.f. of the cell drops to 1.8V.
Consequently, the cell stops discharging current because the chemicals have been used up, and need to be regenerated. (At this point, we say the battery has 'run' down and needs to recharged.)
Recharging
Here, the lead accumulator functions as an electrolytic cell. When connected to a power supply, the anode in the discharged cell becomes the cathode, while the cathode becomes the anode. This is because the anode and cathode in the discharged cell are connected to the cathode and anode of the power source respectively. (Refer to your textbooks for diagrams of the lead accumulator during discharging and recharging).
At the anode, the PbSO4 dissociates to form Pb2+ and SO4-- ions. These Pb2+ ions then combine with the water molecules to form PbO2 and H+ ions
PbSO4(s) ----> Pb2+(aq) + SO4--(aq)
Pb2+(aq) + 2H2O(l) ----> PbO2(s) + 4H+(aq) + 2e- (oxidation half-reaction)
At the cathode, the PbSO4 dissociates in the same way as it does at the anode. However, the Pb2+ ions produced undergo reduction by accepting two electrons each and are deposited on the cathode as lead metal.
PbSO4(s) ----> Pb2+(aq) + SO4--(aq)
Pb2+(aq) + 2e- ----> Pb(s) (reduction half-reaction)
The 4 moles H+ ions (produced in the oxidation half-reaction) combine with the 2 moles of SO4-- ions (produced from the dissociation of PbSO4 at both electrodes) to regenerate the tetraoxosulphate (VI) acid, H2SO4.
4H+(aq) + 2SO4--(aq) ----> 2H2SO4(aq)
2H+(aq) + SO4--(aq) ----> H2SO4(aq)
Consequently, the density of the acid in the electrochemical cell returns to the optimal level of 1.25gcm^-3 and its e.m.f. increases back to 2.2V, after recharging.
Similarities between Electrolytic and Electrochemical Cells
1. Both cells are made up of the anode, the cathode and the electrolyte.
2. In both cells, oxidation occurs at the anode, while reduction takes place at the cathode.
3. In both cells, electrons flow from the anode to the cathode.
Differences between Electrolytic and Electrochemical Cells
1. In an electrolytic cell, the anode is the positive terminal, while the cathode is the negative terminal; whereas, in an electrochemical cell, the anode is the negative terminal, and the cathode is the positive terminal.
2. An electrolytic cell is used to convert electrical energy into chemical energy, while an electrochemical cell uses chemical energy to generate electrical energy.
3. In an electrolytic cell, no salt-bridge is required, whereas, a salt-bridge is needed in an electrochemical cell.
4. In an electrolytic cell, there is only one electrolyte, while an electrochemical cell can contain more than one electrolyte.
5. In an electrolytic cell, the two electrodes are in the same compartment, while in an electrochemical cell, the electrodes may be in the same or different compartments.
6. An electrolytic cell is known as a voltameter, while an electrochemical cell is called a galvanic cell.
Do These:
Question 1
a) What do you understand as the electrode potential of a metal?
b) Describe the structure of the Leclanché cell and give the chemical reactions that occur in it.
Question 2
The reactions that occur inside a car battery are said to be redox. Explain this statement using suitable equations to support your points.
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