Chapter 21 electrochemistry
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Chapter 21 Electrochemistry. Electrochemical Processes. Chemical processes can either release energy or absorb energy. The energy can sometimes be in the form of electricity. Electrochemistry has many applications in the home as well as in industry.

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Chapter 21 Electrochemistry

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Chapter 21 electrochemistry

Chapter 21Electrochemistry


Electrochemical processes

Electrochemical Processes

Chemical processes can either

release energy or absorb energy.

The energy can sometimes be in

the form of electricity.

Electrochemistry has many

applications in the home as well as in industry.

Flashlight and automobile batteries are examples of devices used to generate electricity.

Biological systems also use electrochemistry to carry nerve impulses.


Spontaneous redox reaction

Spontaneous Redox Reaction

When a strip of zinc metal is dipped into an aqueous solution of blue copper sulfate, the zinc becomes copper-plated.

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

The net ionic equation involves only zinc and copper.

Electrons are transferred from zinc atoms to copper ions.

This is a redox reaction that occurs spontaneously


Spontaneous redox reaction1

Spontaneous Redox Reaction

As the reactions proceeds, zinc atoms lose electrons as they are oxidized to zinc ions.

The zinc metal slowly dissolves.

0 +2 +2 0

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

At the same time, copper ions in solution gain the electrons lost by the zinc.

They are reduced to copper atoms are deposited as metallic copper.


Spontaneous redox reaction2

Spontaneous Redox Reaction

As the copper ions in solution are gradually replace by zinc ions, the blue color of the solution fades.

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

Reduction: Cu2+(aq) + 2e- Cu(s)


Activity series of metals

Activity Series of Metals

Zinc is higher on the list than copper.

For any two metals in an activity series, the more active metal is the more readily oxidized.


Electric current

Electric Current

When zinc is dipped into a copper sulfate solution, zinc becomes plated with copper.

In contrast, when a copper strip is dipped into a solution of zinc sulfate, the copper does not spontaneously become zinc-plated.

This is because copper metal is not oxidized by zinc ions.

When a zinc strip is dipped into a copper sulfate solution, electrons are transferred from zinc metal to copper ions.

This flow of electrons is an electric current.


Electrochemical process

Electrochemical Process

The zinc-metal–copper-ion system is an example of the conversion of chemical energy into electrical energy.

Electrochemical process – any conversion between chemical energy and electrical energy

All electrochemical processes involve redox reactions.

If a redox reaction is to be used a a source of electrical energy, the two half-reactions must be physically separated.


Electrochemical cell

Electrochemical Cell

The electrons released by zinc must pass through an external circuit to reach the copper ions if useful electrical energy is to be produced.

In this case the system serves as an electrochemical cell.

Also, an electric current can be used to produce a chemical change.

That system, too, serves as an electrochemical cell.

Electrochemical cell – any device that converts chemical energy into electrical energy or vice versa.


Voltaic cells

Voltaic Cells

In 1800, Italian physicist Alessandro Volta build the first electrochemical cell that could be used to generate a dire electric current.

Voltaic cells – are electrochemical cells used to convert chemical energy into electrical energy.

Electrical energy is

produced in a voltaic cell

by spontaneous redox

reactions within the cell.


Constructing a voltaic cell

Constructing a Voltaic Cell

Half cell – one part of a voltaic cell in which either oxidation or reduction occurs.

Typical half-cell consists of a piece of metal immersed in a solution of its ions.

Example: one half-cell is a zinc rod immersed in a solution of zinc sulfate.

Other half-cell is a copper rod immersed in a solution of copper sulfate.

Half cells are connected by a salt bridge which is a tube containing a strong electrolyte, often potassium sulfate.


Constructing a voltaic cell1

Constructing a Voltaic Cell

A porous plate my be used instead of a salt bridge

The porous plate allows ions to pass from on half-cell to the other but prevents the solutions from mixing completely.

A wire carries the electrons in the external circuit from the zinc rod to the copper rod.

A voltmeter or light bulb can be connected in the circuit.

The driving force of such a voltaic cell is the spontaneous redox reaction between zinc metal and copper(II) ions in solution.


Constructing a voltaic cell2

Constructing a Voltaic Cell

The zinc and copper rods in this voltaic cell are the electrodes.

Electrode – a conductor in a circuit that carries electrons to or from a substance other than a metal.

The reaction at the electrode determines whether the electrode is labeled as an anode or a cathode

Anode – the electrode at which oxidation occurs

Cathode – the electrode at which reduction occurs.


Constructing a voltaic cell3

Constructing a Voltaic Cell

Electrons are consumed at the cathode and its labeled the positive electrode. (reduction occurs)

Electrons are produced at the anode and its labeled the negative electrode. (oxidation occurs)

The reaction at the electrode determines whether the electrode is labeled as an anode or a cathode

All parts of the voltaic cell remain balanced in terms of charge at all times.

The moving electrons balance any charge that might build up as oxidation and reduction occur.


Questions

Questions

At which electrode does oxidation take place?

At the anode (negative electrode)

Where does reduction take place?

At the cathode (positive electrode)

What path do the electrons given up by zinc follow?

They go through the wire and the electric light to the copper electrode.

What happen to the electrons at the copper electrode?

They reduce copper ions to copper.


How a voltaic cell works

How a Voltaic Cell Works

  • These steps actually occur at the same time.

  • Zn(s) Zn 2+(aq) + 2e-

  • Electrons are produced at the zinc rod according to the oxidation half reaction

  • The electrons leave the zinc anode and pass through the external circuit to the copper rod.

  • Electrons enter the copper rod and interact with copper ions in solution.

  • Cu 2+ (aq) + 2e- Cu(s)


How a voltaic cell works1

How a Voltaic Cell Works

  • To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge.

  • The overall cell reaction (note the electrons in the overall reaction must cancel.

  • Zn(s) Zn2+(aq) + 2e-

  • Cu2+ (aq) + 2e- Cu(s)

  • Zn(s) + Cu2+ (aq)Zn2+(aq) + Cu(s)


The need for a salt bridge

The Need for a Salt Bridge

As zinc is oxidized at the anode, Zn2+ ions enter the solution and they have no negative ions to balance their charges.

So a positive charge tends to build up around the anode.

At the cathode, Cu2+ ions are reduced to Cu and taken out of the solution leaving behind unbalanced negative ions.

Thus, a negative charge tends to develop around the cathode.


The need for a salt bridge1

The Need for a Salt Bridge

The salt bridge allows negative ions, such as SO42-, to be drawn to the anode to balance the growing positive charge.

Positive ions, such as K+, are drawn from the salt bridge to balance the growing negative charge at the cathode.


Representing electrochemical cells

Representing Electrochemical Cells

You can represent the zinc-copper voltaic cell by suing the following shorthand form.

Zn(s) ZnSO4(aq) CuSO4(aq) Cu(s)

The single vertical lines indicate boundaries of phases that are in contact.

The double vertical lines represent the salt bridge or porous partition that separates the anode compartment from the cathode compartment.

The half-cell that undergoes oxidation (the anode) is written first.


Using voltaic cells as energy sources

Using Voltaic Cells as Energy Sources

The zinc-copper voltaic cell is no longer used commercially.

Current technologies that use electrochemical processes to produce electrical energy include dry cells, lead storage batteries and fuel cells.

Dry Cells – a voltaic cell in which the electrolyte is a paste.

Dry cells used when a compact, portable electrical energy source is required.


Dry cells

Dry Cells

A common type of dry cell is a flashlight battery.

A zinc container is filled with a thick, moist electrolyte paste of manganese (IV) oxide, zinc chloride, ammonium chloride, and water.

A graphite rode is embedded in the paste.

The zinc container is the anode and the graphite rod is the cathode.

The thick paste and its surrounding paper liner prevent the contents of the cell from freely mixing, so a salt bridge is not needed.


Dry cells1

Dry Cells

Oxidation:

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

Reduction:

2MnO2(s) + 2NH4+(aq) + 2e-

Mn2O3(s) + 2NH3(aq) + H2O(l)


Dry cells2

Dry Cells

In an ordinary dry cell, the graphite rod serves only as a conductor and does not undergo reduction.

MnO2 is the species that is actually reduced.

The electrical potential of this cell starts out at 1.5V but decreases steadily during use to about 0.8V.

Dry cells of this type are not rechargeable because the cathode reaction is not reversible.


Alkaline battery

Alkaline Battery

The alkaline battery is an improved dry cell used for the same purposes.

In the alkaline battery, the reactions are similar to those in the common dry cell, but the electrolyte s a basic KOH past.

This change in design eliminates the buildup of ammonia gas and maintains the Zn electrode, which corrodes more slowly under alkaline conditions.


Alkaline battery1

Alkaline Battery


Lead storage batteries

Lead Storage Batteries

People depend on lead storage batteries to start their cars.

Battery is a group of cells

connected together.

A 12-V car battery consists of

six voltaic cells connected

together.

Each cell produces about 2 V

and consists of lead grids.


Lead storage batteries1

Lead Storage Batteries

One set of grids, the anode, is packed with spongy lead.

The other set, the cathode, is packed with lead(IV) oxide.

The electrolyte for both half-cells in a lead storage batter is concentrated sulfuric acid.

Using the same electrolyte for both half-cells allows the cell to operate without a salt bridge or porous separator.


Lead storage batteries2

Lead Storage Batteries

Pb(s) + SO42-(aq) PbSO4(s) + 2e-

PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O(l)

When a lead storage battery discharges, it produces the electrical energy needed to start a car.

The overall spontaneous redox reaction that occurs is the sum of the oxidation and reduction half-reactions.

Pb(s) + PbO2(s) + 2H2SO4 (aq) 2PbSO4(s) + 2H2O(l)

The equation shows that lead sulfate forms during discharge.


Lead storage batteries3

Lead Storage Batteries

Pb(s) + PbO2(s) + 2H2SO4 (aq) 2PbSO4(s) + 2H2O(l)

The sulfate slowly builds up on the plates, and the concentration of sulfuric acid decreases.

2PbSO4(s) + 2H2O(l) Pb(s) + PbO2(s) + 2H2SO4 (aq)

The reverse reaction occurs when a lead storage battery is recharged. This occurs when the car’s generator is working properly.

The reverse reaction is not a spontaneous reaction. A direct current must pass through the cell in a direction opposite that of the current flow during discharge.


Lead storage batteries4

Lead Storage Batteries

In theory, a lead storage battery can be discharged and recharged indefinitely, but in practice its lifespan is limited.

Small amounts of lead sulfate fall from the electrodes and collect on the bottom of the cell.

Eventually, the electrodes lose so much lead sulfate that the recharging process is ineffective or the cell is shorted out.


Lead storage battery

Lead Storage Battery


Fuel cells

Fuel Cells

To overcome the disadvantages associated with lead storage batteries, cells with renewable electrodes have been developed.

Fuel cells – are voltaic cells in which a fuel substance undergoes oxidation and from which electrical energy is continuously obtained.

Fuel cells do not have to be recharged.

They can be designed to emit no air pollutants and to operate more quietly and more const-effectively than a conventional electrical generator


Fuel cells1

Fuel Cells

Simplest fuel cell is the hydrogen-0xygen fuel cell.

Fuel cells – are voltaic cells in which a fuel substance undergoes oxidation and from which electrical energy is continuously obtained.

Fuel cells do not have to be recharged.

They can be designed to emit no air pollutants and to operate more quietly and more const-effectively than a conventional electrical generator

There are three compartment separated from one another by two electrodes made of porous carbon.


Fuel cells2

Fuel Cells

Oxygen (the oxidizer) is fed into the cathode compartment

Hydrogen (the fuel) is fed into the anode compartment.

The gases diffuse slowly through the electrodes.

The electrolyte in the central compartment is a hot, concentrated solution of potassium hydroxide.

Electrons from the oxidation half-reaction at the anode pass through an external circuit to enter the reduction half reaction at the cathode.


Fuel cells3

Fuel Cells


Fuel cells4

Fuel Cells

Oxidation: 2H2(g) + 4OH-(aq) 4H2O(l) + 4e-

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

The overall reaction in the hydrogen-oxygen fuel cell is the oxidation of hydrogen to form water.

2H2(g) + O2(g) 2H2O (l)

Fuel cells were developed for space travel where lightweight, reliable power systems are needed.

Fuel cells different from lead storage batteries in that they are not self-contained.


Fuel cells5

Fuel Cells

Operation depends on a steady flow of fuel and oxygen into the cell (where combustion takes place) and the flow of the combustion product out of the cell.

In the case of the hydrogen fuel cell, the product is pure water.

Both the electricity generated and the water produced are consumed in space flights.

Fuel cells convert 75% of the available energy into electricity.

Conventional electric power plant converts from 35% to 40% of the energy of coal to electricity.


Fuel cells6

Fuel Cells

Other fuels, such as methane (CH4) and ammonia (NH3), can be used in place of hydrogen.

4NH3(g) + 3O2 (g) 2N2(g) + 6H2O (g)

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

Other oxidizers, such as chlorine (Cl2) and ozone (O3)can be used in place of oxygen.


Question

Question

Make a sketch of a tin/lead voltaic cell

Sn SnSO4 PbSO4 Pb

Label the cathode and anode, and indicate the direction of electron flow. Write the equations for the half-reactions.

Sn(s) Sn2+(aq) + 2e-

Pb2+(aq) + 2e- Pb(s)

Tin is the anode, lead is the cathode.

The electrons flow from tin to lead.


Questions1

Questions

What type of reaction occurs during an electrochemical process?

Redox

What is the source of electrical energy produced in a voltaic cell?

Spontaneous redox reactions within the cell

If the relative activities of two metals are known, which metal is more easily oxidized?

The metal with the higher activity.


Homework

Homework

  • Draw detailed pictures of the following:

  • A Voltaic Cell

  • A Lead Storage Battery

  • Fuel Cell

  • Make sure to include all substances involved as well as the half reactions and overall reaction.

  • Make sure you label the anode and cathode and indicate where oxidation and reduction take place.

  • List common uses of each

  • List advantages of disadvantages of each.


Chapter 21 electrochemistry

End of Section 20.1


Electrical potential

Electrical Potential

Electrical potential of a voltaic cell is a measure of the cell’s ability to produce an electric current. (in Volts)

You cannot measure the electrical potential of an isolated half-cell.

When two half-cells are connected to form a voltaic cell, the difference in potential can be measured.

The electrical potential of a cell results from a competition for electrons between two half-cells.

The half-cell that has a greater tendency to acquire electrons will be the one in which reduction occurs.


Electrical potential1

Electrical Potential

Oxidation occurs in the other half-cell as there is a loss of electrons.

Reduction potential – the tendency of a given half-reaction to occur as a reduction.

The half-cell in which reduction occurs has a greater reduction potential than the half-cell in which oxidation occurs.

Cell potential – the difference between the reduction potentials of the two half-cells.


Cell potential

Cell Potential

reduction potential reduction potential

Cell Potential = of half-cell in which - of half-cell in which

reduction occurs oxidation occurs

E0cell = E0red -E0oxid

Standard cell potential (E0cell)– is the measured cell potential when the ion concentration in the half-cells are 1M, any gases are at a pressures of 101 kPa, ant the temperature is 25ºC.

The symbols E0red and E0oxid represent the standard reduction potential for the reduction and oxidation half-cells, respectively.


The lemon battery

The Lemon Battery

A working voltaic cell made

using a lemon and strips of

copper and zinc.

Which is the anode and which

is the cathode?

Zn is anode & Cu is cathode

What process goes on at the anode and cathode?

Oxidation at anode, reduction at cathode.

Which process are electrons lost?

Oxidation

What role does the lemon play in the battery?

Salt bridge


Standard hydrogen electrode

Standard Hydrogen Electrode

Because half-cell potentials cannot be measured directly, scientists have chosen an arbitrary electrode to serve as a reference.

The standard hydrogen electrode is used with other electrodes so the reduction potentials of the other cells can be measured.

The standard reduction potential of the hydrogen electrode has been assigned a value of 0.00 V.


Standard hydrogen electrode1

Standard Hydrogen Electrode

Consists of a platinum electrode

immersed in a solution with a

hydrogen-ion concentration of 1M

Solution is at 25 C and the electrode

is a small square of platinum foil

coated with finely divided platinum,

known as platinum black.

Hydrogen gas at 101 kPa is bubbled

around the platinum electrode.


Standard hydrogen electrode2

Standard Hydrogen Electrode

2H+(aq, 1M) + 2e- H2(g, 101kPa) EOH+ = 0.00V

Double arrows indicate the reaction is reversible.

Standard reduction potential of H+ is the tendency of H+ ions to acquire electrons and be reduced to H2(g)

Whether the half-cell reaction occurs as a reduction or as an oxidation is determined by the reduction potential of the half-cell to which the standard hydrogen electrode is connected.


Standard reduction potentials

Standard Reduction Potentials

A voltaic cell can be made by connecting a standard hydrogen half-cell to a standard zinc half-cell.


Standard reduction potentials1

Standard Reduction Potentials

To determine the overall reaction for this cell, first identify the half-cell in which reduction takes place.

In all electrochemical cells, reduction takes place at the cathode and oxidation takes place at the anode.

A voltmeter gibes a reading of +0.76 V when the zinc electrode is connect to the negative terminal and the hydrogen electrode is connect to the positive terminal.

Zinc is oxidized – anode and Hydrogen ions are reduced – hydrogen electrode is the cathode. .


Standard reduction potentials2

Standard Reduction Potentials

Oxidation: Zn (s) Zn2(aq) + 2e- (at anode)

Reduction: 2H+(aq) + 2e- H2(g) (at cathode)

You can determine the standard reduction potential of a half-cell by using a standard hydrogen electrode and the equation for standard cell potential

E0cell = E0red -E0oxid

E0cell = E0H+ -E0Zn2+

0.76 V = 0.00 V -E0Zn2+

E0Zn2+ =-0.76 V


Standard reduction potentials3

Standard Reduction Potentials

The standard reduction potential for the zinc half-cell is -0.76 V.

The value is negative because the tendency of zinc ions to be reduced to zinc metal in this cell is less than the tendency of hydrogen ions to be reduced to hydrogen gas.

Consequently, the zinc ions are not reduced. Instead, the opposite occurs: Zinc metal is oxidized to zinc ions.


Standard reduction potentials4

Standard Reduction Potentials

For a standard copper half-cell, the measured standard cell potential is +0.34V.

Copper is the cathode and Cu2+ ions are reduced to Cu metal

Hydrogen half-cell is the anode, and H2 gas is oxidized to H+ ions.

E0cell = E0red -E0oxid

E0cell = E0Cu2+ -E0H+

+0.34 V = E0Cu2+ -0.00

E0Cu2+ =+0.34 V


Discussion

Discussion

The two half-cells of a voltaic cell are competing for electrons.

Oxidation or reduction could occur in either cell.

The half-cell with the more positive reduction potential will win the competition and undergo reduction.

The potential produced by the electrochemical cell is the difference in the reduction potentials of the two half-cell reactions.

The quantitative value of any half-cell potential is obtained by measuring it against the standard hydrogen electrode. (Review table 21.2 page 674)


Discussion1

Discussion

Which reactions have the greatest tendency to occur as reductions?

Activity series of metals have the most active metals at the top. Because active metals lose electrons easily, they are most easily oxidized.

Thus, the ions of active metals are least likely to be reduced.

The potential produced by the electrochemical cell is the difference in the reduction potentials of the two half-cell reactions.


Calculating standard cell potentials

Calculating Standard Cell Potentials

To function, a cell must be constructed of two half-cells.

The half-cell reaction having the more positive (or less negative) reduction potential occurs as a reduction in the cell.

You can use the know standard reduction potentials for various half-cells to predict the half-cell in which reduction and oxidation will occur.

If the cell potential for a given redox reaction is + then the reaction is spontaneous as written. If the cell potential is - , then the reaction in nonspontaneous.


Question1

Question

Determine whether the following redox reaction will occur spontaneously.

3Zn2+(aq) + 2Cr(s) 3Zn(s) + 2Cr3+(aq)

Oxidation: Cr(s) Cr3+(aq) + 3e- E0Cr3+ = -0.74V

Reduction:Zn2+(aq) + 2e- Zn(s)E0Zn2+ = -0.76V

E0cell = E0red -E0oxid

E0cell = E0Zn2+ -E0Cr3+

E0cell = -0.76 – (-0.74)

E0cell =-0.02 V (nonspontaneous)


Question2

Question

Is this redox reaction spontaneous as written?

Co2+(aq) + Fe(s) Co(s) + Fe2+(aq)

Oxidation: Fe(s) Fe2+(aq) + 2e- E0Cr3+ = -0.44V

Reduction:Co2+(aq) + 2e- Co(s)E0Zn2+ = -0.28V

E0cell = E0red -E0oxid

E0cell = E0Co2+ -E0Fe2+

E0cell = -0.28 – (-0.44)

E0cell =+0.16 V (spontaneous)


Question3

Question

Determine the cell reaction for a voltaic cell composed of the following half-cells.

Fe3+(aq) + e- Fe2+(aq) E0Fe3+ = +0.77V

Ni2+(aq) + 2e- Ni(s) E0Ni2+ = -0.25V

The half-cell with the more positive reduction potential is the one in which reduction occurs (the cathode)

Oxidation: Ni(s) Ni2+(aq) + 2e-

Reduction: 2Fe3+(aq) + 2e- 2Fe2+(aq)(balance e-)

Ni(s) +2Fe3+(aq)Ni2+(aq) +2Fe2+(aq)


Discussion2

Discussion

In the previous example, we had to multiply the Fe half-cell reaction by a factor of 2 to cancel out the electrons.

Even though, there were two times as many electrons present, the tendency for the electrons to flow is not two times greater.

The tendency, which is measured by the E0 value, remains the same.


Questions2

Questions

A voltaic cell is constructed using the following half reactions.

Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V

Al3+(aq) + 3e- Al(s) E0Al3+ = -1.66V

2Al(s) +3Cu2+(aq) 2Al3+(aq) +3Cu(s)

A voltaic cell is constructed using the following half reactions.

Ag+(aq) + e- Ag(s) E0Ag+ = +0.80V

Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V

Cu(s) +2Ag+(aq) Cu2+(aq) +2Ag(s)


Question4

Question

Calculate the standard cell potential for the Ni/Fe voltaic cell. Half-reactions are as follows;

Fe3+(aq) + e- Fe2+(aq) E0Fe3+ = +0.77V

Ni2+(aq) + 2e- Ni(s) E0Ni2+ = -0.25V

E0cell = E0Fe3+ -E0Ni2+

E0cell = +0.77 V – (-0.25 V)

E0cell =+1.02 V


Question5

Question

A voltaic cell is constructed using the following half-reactions

Al3+(aq) + 3e- Al(s) E0Al3+ = -1.66V

Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V

E0cell = E0Cu2+ -E0Al3+

E0cell = +0.34 V – (-1.66 V)

E0cell =+2.00 V


Question6

Question

A voltaic cell is constructed using the following half-reactions

Ag+(aq) + e- Ag(s) E0Ag+ = +0.80V

Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V

E0cell = E0Ag+ -E0Cu2+

E0cell = +0.80 V – (+0.34 V)

E0cell =+0.46 V


Questions3

Questions

What causes the electrical potential of a cell?

Competition for electrons between two half-cells

What is the electrical potential of a standard hydrogen electrode?

Assigned a value of 0.00 V at 25ºC

How can you find the standard reduction potential of a half-cell?

By connecting it to a standard hydrogen electrode and measuring the cell potential

What cell potential values indicate a spontaneous reaction? A nonspontaneous reaction?

Positive cell potential - spontaneous


Homework1

Homework

Using the reduction potentials from table 21.2, create an electrochemical cell that will operate spontaneously.

Calculate the cell potential

Write the equations for the two half-reactions and the overall cell reaction.

Use the shorthand method to represent the cell


Chapter 21 electrochemistry

End of section 20.2


Electrolytic cells

Electrolytic Cells

An electric current can be used to make a nonspontaneous redox reaction go forward.

Electrolysis – the process in which electrical energy is used to bring about a nonspontaneous chemical change.

Examples of electrolysis are silver-plated dishes and utensils, gold-plated jewelry, and chrome-plated automobile parts.

Electrolytic cell – the apparatus in which electrolysis is carried out is an electrochemical cell used to cause a chemical change through the application of electrical energy.


Electrolytic cells1

Electrolytic Cells

An electrolytic cell uses electrical energy (direct current) to make a nonspontaneous redox reaction proceed to completion.

In both voltaic and electrolytic cells, electrons flow from the anode to the cathode in the external circuit.

For both types of cells, the electrode at which reduction occurs is the cathode.

The key difference between voltaic and electrolytic cells is that in a voltaic cell, the flow of electrons is the result of a spontaneous redox reaction, whereas in an electrolytic cell, electrons are pushed by an outside power source, such as a battery.


Chapter 21 electrochemistry

Voltaic cell – energy is released from a spontaneous redox reaction.

Electrolytic cell - energy is absorbed to drive a non-spontaneous reaction.


Electrolytic cells2

Electrolytic Cells

Electrolytic and voltaic cells also differ in the assignment of charge to the electrodes.

In an electrolytic cell, the cathode is considered to be the negative electrode, because it is connected to the negative electrode of the battery.

The anode in an electrolytic cell is considered to be the positive electrode because it is connected to the positive electrode of the battery.

In a voltaic cell, the anode is the negative electrode and the cathode is the positive electrode.


Electrolytic cells3

Electrolytic Cells

Electrolytic processes are used to separate active metals such as aluminum, magnesium, and sodium from their salts.

The same process is used to recover metals from ores.


Electrolysis of water

Electrolysis of Water

When a current is applied to two electrodes immersed in pure water, nothing happens.

When an electrolyte such as sulfuric acid or potassium nitrate in low concentration is added to the pure water, the solution conducts electricity and electrolysis occurs.

The products of the electrolysis of water are hydrogen gas and oxygen gas.


Electrolysis of water1

Electrolysis of Water

Water is reduced to hydrogen at the cathode

Reduction: 2H2O(l) + 2e- H2(g) + 2OH-(aq)

Water is oxidized at the anode

Oxidation: 2H2O(l) O2(g) + 4H+(aq) + 4e-

The region around the anode turns acidic due to an increase in H+ ions.

The region around the cathode turns basic due to the production of OH- ions.


Electrolysis of water2

Electrolysis of Water

The overall cell reaction

4H2O(l) + 4e-2H2(g) + 4OH-(aq)(x2 to balance)

2H2O(l) O2(g) + 4H+(aq) + 4e-

6H2O(l) 2H2(g) + 4OH-(aq) + O2(g) + 4H+(aq)

The ions produced tend to recombine to form water, so they are not included in the net reaction.

6H2O(l)electrolysis H2(g) + O2(g)


Electrolysis of brine

Electrolysis of Brine

If the electrolyte in an aqueous solution is more easily oxidized or reduced than water, then the products of electrolysis will be substances other than hydrogen and oxygen.

Example is brine (a concentrated aqueous solution of sodium chloride) which produces chlorine gas, hydrogen gas, and sodium hydroxide.

During electrolysis of brine, chloride ions are oxidized to produce chlorine gas at the anode.

Oxidation: 2Cl-(aq) Cl2(g) + 2e- (at anode)


Electrolysis of brine1

Electrolysis of Brine

Water is reduced to produce hydrogen gas at the cathode.

Reduction: 2H2O(l) + 2e- H2(g) + 2OH-(aq) (at cathode)

Sodium ions are not reduced to sodium metal in the process because water molecules are more easily reduced than are sodium ions.

The reduction of water produces hydroxide ions as well as hydrogen gas. Thus the electrolyte in solution becomes sodium hydroxide.


Electrolysis of brine2

Electrolysis of Brine

The overall ionic equation

2H2O(l) + 2e- H2(g) + 2OH-(aq)

2Cl-(aq) Cl2(g) + 2e-

2H2O(l) + 2Cl-(aq) H2(g) + 2OH-(aq) Cl2(g)

The spectator ion Na+ can be included in the equation (as part of NaCl and of NaOH) to show the formation of sodium hydroxide during the electrolytic process

2NaCl (aq) + 2H2O(l) + 2Cl-(aq) H2(g) + 2NaOH(aq) Cl2(g)


Electrolysis in metal processing

Electrolysis in Metal Processing

Electrolytic cells are commonly used in the plating, purifying and refining of metals.

Many of the shiny, metallic objects you see every day, such as chrome-plated fixtures or nickel-plated coins, were manufactured with the help of electrolytic processes.

Electroplating is the deposition of a think layer of metal on an object in an electrolytic cell.

An object may be electroplated to protect the surface of the base metal from corrosion or to make it more attractive.


Electrolysis in metal processing1

Electrolysis in Metal Processing

An object that is to be silver-plated is made the cathode in an electrolytic cell.

The anode is the metallic silver that is to be deposited

The electrolyte is a solution of a silver salt, such as silver cyanide.

When a direct current is applied, silver ions move from the anode to the object to be plated.

Reduction: Ag+(aq) + e- Ag (s) (at cathode)

The net result is that silver transfers from the silver electrode to the object being plated.


Electrolysis in metal processing2

Electrolysis in Metal Processing

Many factors contribute to the quality of the metal coating that forms.

In the plating solution, the concentration of the cations to be reduced must be carefully controlled.

The solution must also contain compounds to control the acidity and to increase the conductivity.

Other compounds may be used to make the metal coating brighter or smoother.


Electrolysis in metal processing3

Electrolysis in Metal Processing

Electroforming – is a process in which an object is reproduced by making a metal mold of it at the cathode of a cell.

A phonograph record can be coated with metal so it will conduct a current.

It is then electroplated with a thick coating of metal. This coating can be stripped off and used as a mold to produce copies of the record.


Electrowinning

Electrowinning

Electrowinning – a process where impure metals can be purified in electrolytic cells.

The cations of molten salts or aqueous solutions are reduced at the cathode to give very pure metals.

A common use is in the extraction of aluminum form its ore, bauxite. (Al2O3)

In a method know as the Hall-Heroult process, purified alumina is dissolved in molten cryolite (Na3AlF6), and heated to above 1000ºC in a carbon line tank.


Electrowinning1

Electrowinning

The carbon lining, connected to a direct current, serves as the cathode. The anode consists of carbon rods dipped into the tank.

At the cathode, Al3+ ions are reduced, forming molten aluminum. At the anode, carbon is oxidized, forming carbon dioxide gas.

2Al2O3(l) + 3C(s) 4Al(l) + 3CO2(g)


Other electrolytic processes

Other Electrolytic Processes

Electrorefining – a piece of impure metal is made the anode of the cell. It is oxidized to the cation and then reduced to the pure metal at the cathode.

Electrorefining technique is used to obtain ultrapure silver, lead and copper.

Other electrolytic processes are centered on the anode rather than the cathode.

Electropolishing - the surface of an object at the anode is dissolved selectively to give it a high polish.

Electromachining - a piece of metal at the anode is partially dissolved until the remaining portion is an exact copy of the object at the cathode.


Questions4

Questions

What is the difference between an electrolytic cell and a voltaic cell?

Voltaic cell uses an electrochemical reaction to produce electrical energy. An electrolytic cell uses electrical energy to bring about a chemical change.

What products form during the electrolysis of water?

H2 (g) and O2 (g)

What chemical changes occur during the electrolysis of brine?

Chloride ions are oxidized to produce chlorine gas and water is reduced to produce hydrogen gas.


Questions5

Questions

What are some application of electrolysis in the field of metallurgy?

Electroplating (deposition of a thin layer of metal on an object), electrorefining (purification of metals ) and electrowinning (extraction of metals)

What is the charge on the anode of an electrolytic cell? Of a voltaic cell?

Electrolytic cell anode (+); voltaic cell anode (-)

Which process, oxidation or reduction, always occurs at the cathode of an electrolytic cell?

Reduction


Questions6

Questions

Can metallic sodium be obtained by electrolyzing brine?

No; the products are chlorine gas, hydrogen gas, and sodium hydroxide.

Sodium is obtained by electrolysis of molten NaCl in the Downs cell, which operates at 801ºC to keep it melted.

The anode and cathode of the cell are separated to prevent recombination of sodium and chlorine.

Reduction of Na+ occurs at a graphite anode. Liquid sodium rises to the top of the molten NaCl and is drawn off.

2NaCl (l) 2Na (l) + Cl2(g)


Chapter 21 electrochemistry

End of Chapter 20


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