Chemistry Form 4 Chapter 6

Chapter 6: Electrochemistry 1.

Electrolyte

 Electrolytes are substances in molten state or aqueous solution that can conduct electricity due to the presence of free moving ions

2.

Non-electrolytes

 Non-electrolytes are substance that cannot conduct electricity either in molten state or aqueous solution.

3.

Electrolysis

 Electrolysis is a process whereby compounds in molten or aqueous state are broken down into their constituent elements by passing electricity through them.

4.

Electrolytic cell

 The electrolytic cell is the set of apparatus needed to conduct electrolysis.  It consists of a battery, an electrolyte and two electrodes.

5.

Electrode

 Electrodes are electrical conductors.  Graphite or platinum is usually used as electrodes because they are inert, they do not react with electrolyte or the products of electrolysis.

6.

Anode

 The electrode which is connected to the positive terminal of an electric source.  Negatively charged ions (anions) in the electrolyte are attracted to the anode.

7.

Cathode

 The electrode which is connected to the negative terminal of the batteries.  Positively charged ions (cations) in the electrolyte are attracted to the cathode.

8.

Electrolysis of

Molten compound:

Molten

 A molten compound consists of one type of cations and one type of anion only.

Compounds

 In solid state, ions do not move freely but are held in fixed positions in a lattice.  In molten electrolyte, the ions move freely.  During electrolysis, the negative ions or anions move to the anode.  The positive ions or cations move to the cathode.  A new substance is then formed at each electrode.  Example: Electrolysis of molten lead (II) bromide, PbBr2.  PbBr2 is an ionic compound. It consist Pb2+ and Br-.  In solid PbBr2, these ions do not move freely but are held in fixed positions in a lattice. When it melts, the ions are free to move.  During the electrolysis of molten PbBr2, Br- are attracted to the anode.  At the anode, Br- undergo discharge whereby each of these ions releases an electron to form a neutral bromine atom.  Two bromine atoms combine to form a bromine gas, Br2 molecule.  Thus, Br2 is released at the anode.  Half equation: 2Br-(l)  Br2(g) + 2e At the cathode, Pb2+ undergo discharges whereby each of the ions accepts two electrons to form a lead atom.  Thus, lead metal is formed at the cathode.  Half equation: Pb2+(l) + 2e-  Pb(s)  Combining the two half equations, we get the overall equation. Pb2+(l) + 2Br-(l)  Pb(s) + Br2(g)

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Example 1.

2. Battery +

Battery

-

+

ammeter A

A +

anode / +ve electrode

2Br-  Br2 + 2e-

anion/ -ve ion

-

+ cathode / -ve electrode

Pb

+

-

2+

-

anode / +ve electrode

-

+ 2e  Pb

cathode / -ve electrode

cation/ +ve ion

anion/ -ve ion

Molten PbBr2

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-

ammeter

cation/ + +ve ion

-

Molten NaCl

2

3.

4.

Battery

Battery +

+

-

ammeter

ammeter

A

A -

+ anode / +ve electrode

anion/ -ve ion

+

cation/ + +ve ion

anion/ -ve ion -

Molten lead (II) oxide

9.

-

anode / +ve electrode

cathode / -ve electrode

-

-

cathode / -ve electrode

+

cation/ +ve ion

Molten zinc iodide, ZnI2

Electrolysis of Aqueous Solutions  An aqueous solution is produced when a solute is dissolved in water.  An aqueous solution of a salt consists of 2 types of cations (cations of the salt and hydrogen ions, H+), and 2 types of anions (anions of the salt and hydroxide ions, OH-).  H+ and OH- are always present together with the ions produced from the dissociation of salts in aqueous solutions.  This is because water dissociates partially to form H+ and OH-. H2O H+ + OH There are three factors that may influence the selective discharge of ions during the electrolysis of an aqueous solution. i. -

Position of ions in the electrochemical series The ions that are lower in the electrochemical series will be selectively discharged

-

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ii.

Concentration of ions in the electrolytes If the concentration of a particular ion is high, the ions is selectively discharged.

iii.

Types of electrodes used in the electrolysis

 The common materials used as electrodes are carbon and platinum because they are inert.  Example: Electrolysis of copper (II) sulphate, CuSO4 solution. CuSO4  Cu2+ + SO42H+ + OH-

H2O

 If carbon is used as the electrodes, OH- ions are discharged at the anode because of the position of OH- ion in the electrochemical series.  If copper is used as the anode, both SO42- ions and OH- ions are not discharge.  Instead the copper anode dissolves by releasing electrons to form copper (II) ions, Cu2+. Hence, the mass of anode decrease.  Copper acts as an active electrode here because it takes part in the chemical reactions during electrolysis. 1.

2. Battery +

Battery

-

+

ammeter

ammeter

A

A -

+

-

+

anode / +ve electrode

anion/ -ve ion

-

anode / +ve electrode

cathode / -ve electrode

cation/ +ve ion

+

-

anion/ -ve ion

Dilute copper (II) chloride CuCl2

cathode / -ve electrode

+

-

cation/ +ve ion

Silver nitrate, AgNO3 solution

3.

4. Battery +

Battery -

+

ammeter

ammeter

A

A +

anode / +ve electrode

anion/ -ve ion

-

-

+

cation/ +ve ion

anion/ -ve ion

Copper (II) sulphate, CuSO4 solution

-

Zinc iodide, ZnI2 solution

5. © MHS 2009

-

anode / +ve electrode

cathode / -ve electrode

+

-

6.

4

cathode / -ve electrode

+

cation/ +ve ion

Battery +

Battery

-

+

ammeter

ammeter

A

A -

+ anode / +ve electrode

anion/ -ve ion

anode / +ve electrode

+

cation/ +ve ion

anion/ -ve ion

Concentrated potassium chloride, KCl

cathode / -ve electrode

cation/ +ve ion

+

-

Concentrated copper (II) bromide, CuBr2

7.

8. Battery

Battery +

+

-

ammeter

ammeter

A

A

anode / +ve electrode

+

+

Ag  Ag + e anion/ -ve ion

anode / +ve electrode

cathode / -ve electrode

cathode / -ve electrode

-

Ag + e  Ag Ag

C

+

-

anion/ -ve ion

cation/ +ve ion

Ni

Electrolysis in Industries

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Ni +

-

NiSO4

Ag2SO4

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-

-

+

-

+

10

-

+ cathode / -ve electrode

-

-

cation/ +ve ion

Extraction of metals -

Reactive metals such as aluminium and magnesium can be extracted from their ores be electrolysis.

-

Aluminium can be extracted from its ore, molten aluminium oxide, Al2O3 using carbon electrodes.

-

In this process, a substance known as cryolite, Na3AlF6 is added to aluminium oxide, Al2O3 to lower its melting point.

Purification of metals -

Pure copper and silver can be obtained through the process of electrolysis.

-

In the purification of copper, the impure copper is made to be the anode while the cathode is a thin layer of pure copper.

Electroplating of metals -

In electroplating of metals, electrolysis is used to coat one metal onto another metal.

-

In the process of electroplating, a more expensive or attractive metal such as silver or gold is coated onto the object to make it look more attractive and more resistant to corrosion.

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Voltaic Cells

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- A simple voltaic cell consists of two different metals immersed into an electrolyte. Chemical energy is converted to electrical energy in the cell. - The more electropositive metal (metal that is higher position in the ECS) will release electron  Negative terminal - The less electropositive metal (metal that is lower position in the ECS) will be positive terminal. - The electron flow form negative terminal to positive terminal. Cation which is lower at ECS will be discharged.

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Daniell Cell -

In a Daniell cell, zinc and copper are used as electrodes. Each electrode is immersed into a different electrolyte. The electrolytes are connected by a salt bridge or a porous pot. The porous pot and salt bridges are: i. to allow the flow of ions so that the circuit is completed ii. to prevent the two aqueous solution from mixing

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The Electrochemical Series (ECS)

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-

The electrochemical series is an arrangement of metals based on the tendency of each metal atom to donate electrons.

-

The electrochemical series can be constructed based on the potential difference between two metals, and the ability of a metal to displace another metal from its salt solution.

-

The electrochemical series is used to determine the terminals and voltage of a cell. It is also used to predict the ability of one metal to displace another metal from its salt solution.

-

The further the two metals are in the ECS, the greater the voltage produced by the cell.

The Advantages Disadvantages of Various Voltaic Cells Cell

Advantages

Disadvantages

Daniell cell

 Easily set up in the laboratory

 Wet cell – electrolyte easily split  Voltage cannot last

Dry cell

 No spillage

 Does not last

 Small in size

 Cannot be recharged

 Easily carried about

 Leakage can occur if cell cannot be used

 Produces regular current and voltage

anymore

 Obtained in different sizes Alkaline cell

 Lasts longer than dry cell (10 x)

 Leakage can occur if cell cannot be used

 Produces a higher and more regular current

anymore  Expensive  Cannot be recharged

Mercury cell

 Small in size

 Very expensive

 Produces regular current for a longer period

 Cannot be recharged  Mercury that is produced is poisonous

of time  Lasts a long time Lead-acid

 Can be recharged

 Spillage of acid can occur

accumulator

 Produces a high voltage (12V) for a long

 Big in size  Heavy, difficult to be carried about

period time  Produces a high current (175A) suitable for

 Expensive  Loses charge if not used for long

a heavy duty Nickel-

 Can be recharged up to 500 times

 Expensive

cadmium

 Dry cell no spillage

 Transformer needed for recharging cell

cell

 Smaller than accumulator - portable

Electrolytic cell © MHS 2009

Voltaic cell

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Battery +

-

ammeter A -

+ anode / +ve electrode

anion/ -ve ion

cathode / -ve electrode

cation/ + +ve ion

-

Similarities  Contains an electrolyte  Consist of an anode and a cathode  Electron move from the anode to the cathode in the external circuit (connecting wires)  Positive ions and negative ions move in the electrolyte  Chemical reactions involve the release or acceptance of electrons Differences Characteristics

Electrolytic Cell

Voltaic Cell

Energy change

 Electrical energy  chemical energy

 Chemical energy  electrical energy

Electric current

 Electric current results in a chemical

 Chemical reaction produces an electric

and reaction

reaction

current

Electrode /

 Cathode: Negative terminal

 Cathode: Positive terminal

Terminal

 Anode : Positive terminal

 Anode : Negative terminal

Flow of electron

 Electron flow from the positive electrode

 Electrons flow from the negative electrode

(anode) to the negative electrode

(anode) to the positive electrode (cathode)

(cathode) Negative terminal

 Cation receives electrons from the cathode (negative terminal)

Positive terminal

terminal

 Anion release electrons to the anode (positive terminal)

Types of electrodes

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 Electron are released at the negative

 Electrons are received by the positive terminal

 Same or two different types of metal, or graphite electrodes

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 Two different types of metal