Oxidation Number OR Oxidation State:

The donation of electrons is called the oxidation and the gain of electrons is called the reduction. Oxidation and reduction can further be explained by a knowledge of “Oxidation number”.

The oxidation state of an atom in its free or ground state is taken as zero. When the atom loses electrons its oxidation state increases and when the atom gains electrons its oxidation state decreases. The term oxidation-number represents the positive or negative character of the atom in a compound.

Oxidation number is defined as the charge an atom appears to have when electrons are assigned in accordance with the following arbitrary rules.

Electrons shared by two like atoms are divided equally between the two atoms. Electrons shared between two unlike atoms are assigned to the more electronegative atom of them.

Conventions Used in Assigning Oxidation Number or Oxidation State:

• The oxidation number of an element in a free atomic state (Na, H, Cl, O, P etc) or in its poly-atomic state (graphite, H₂, Cl₂, O₂ etc) is always zero.
• The oxidation number of hydrogen is always +1 in its compounds.  However, in metal hydrides like NaH, MgH₂  etc. the oxidation number of hydrogen is -1 because metals are more electropositive than hydrogen.
• O.N. of oxygen is always -2 in its compounds.  However, in peroxides like H₂O₂, Na₂O₂, BaO₂ etc. the oxidation number of oxygen is -1. In OF the oxidation number of oxygen is +2 because F is more electronegative than O.
• O.N. of group IA element i.e. Li, Na, K etc is always +1 in their compounds.
• O. N. of group IIA elements i.e. Be, Mg, Ca, Sr and Ba are always +2 in their compounds.
• O. N. of F is always -1 in its compounds because it is most highly electronegative.  Oxidation O. N. of other elements of group VIIA. (17) i.e. Cl, Br and I are also generally –1.
• In an ion, the sum of the oxidation numbers of different atoms is equal to charge over the ion.
• In a complex compound (involving co-ordination by ligands) it is more convenient to use oxidation number of group (ligand) as a whole instead of the oxidation number of individual atoms. For example, in HCN the oxidation number of CN- ion is –1. Here CN-  as a whole is considered and not of individual C or N.
• on the basis of the above standard oxidation numbers, which may be taken as rules, the oxidation, a number of a particular given atom in a compound can be determined.

Valency and Oxidation State:

Valency is a different term than oxidation number though sometimes the valency and the oxidation number of an element are same in a compound. Valency of an element is given by the number of electrons it actually loses or gains or shares during the formation of a compound, Whereas oxidation number is just the apparent charge (not necessarily actual) over the atom when the electrons are counted according to the arbitrary rules given earlier.

In most of the cases, the valency of an element is constant whereas the oxidation state of an element may vary in its different compounds. Valency and oxidation states of carbon in its different compounds give a good example of this. In CH₄, CH₃Cl, CH₂Cl₂, CHCl₃ and CCl₄ the valency of carbon is always four (due to sharing of four electrons) but its oxidation number is – 4, -2, 0, +2 and +4 respectively.

Oxidation-Reduction in Terms of Oxidation Number:

On the basis of oxidation number a reaction involving the increase in oxidation number is called as oxidation while a reaction involving the decrease in oxidation number is called as reduction (Remember increase in O.N. means increase in positive O.N. or decrease in negative O.N., while decrease in O.N. means decrease in positive O.N. or increase in negative O.N.).

For example, in the reaction, 2Mg + O₂ →  2MgO, The O.N. of Mg increase from 0 to +2.while the O.N. of O decreases from 0 to -2.  Thus, magnesium is oxidised while oxygen is reduced.

Science > Chemistry > Redox Reactions > Oxidation Number or Oxidation State

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In this article we shall study about redox reactions, in which both the oxidation and reduction reactions take place simultaneously.

Oxidation:

Old Concept:

• It is a process in which addition of oxygen takes place.

2Mg + O₂  →      2MgO

• It is a process in which addition of electronegative radical takes place.

2FeCl₂ +  Cl₂ →  2FeCl₃

• It is a process in which removal of hydrogen takes place.

H₂S + 2 [Cl] →  S + 2HCl

• It is a process in which removal of electropositive radical takes place.

2KI + H₂O₂ → I₂ + 2KOH

Moden Concept:

According to the electronic concept, a reaction in which loss of electrons from an atom or an ion takes place is called oxidation. Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, the valency of magnesium is increased’ from zero (in the atomic state) to + 2 (in MgO).

i.e.  Mg⁰ →  Mg²⁺ + 2 e^–

In this reaction, magnesium is losing electrons. And hence oxidation of magnesium takes place.

Reduction:

Old Concept:

• It is a process in which addition of hydrogen takes place.

Cl₂ + H₂  →  2HCl

• It is a process in which addition of electropositive radical takes place.

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄ .

• It is a process in which removal of oxygen takes place.

CuO + 2 [H] → Cu + 2H₂O

• It is a process in which removal of electronegative radical takes place.

FeCl₃ + H → FeCl₂ + HCl

Modern Concept:

According to the electronic concept, a reaction in which the gain of electrons by an atom or an ion takes place is called reduction. Consider reaction, 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

In this reaction, the valency of mercury is decreased’ from +2  (in HgCl₂) to +1 (in Hg₂Cl₂).

i.e.   Hg²⁺ +   e^–  →   Hg ⁺

In this reaction mercury is gaining electron. And hence reduction of mercury takes place.

Redox Reaction:

In any of a chemical reaction if one of the reactants is oxidized, other is surely reduced, i.e. oxidation and reduction always take place simultaneously.

Example – 1: 

2Mg + O₂ →  2MgO  

Mg is oxidized to MgO (addition of oxygen, i.e. increase in positive valency of Mg i.e. loss of electrons by Mg), whereas oxygen is reduced to MgO (addition of positive radical, i.e. increase in negative valency of oxygen, i.e. gain of electrons by oxygen)

Example – 2: 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

HgCl₂ is reduced to Hg₂Cl₂ whereas SnCl₂ is oxidised to SnCl₄.

Thus oxidation and reduction take place simultaneously.  Therefore, all such reactions are called as reduction-oxidation reactions or redox reactions.  In all such reactions, one of the reactants loses the electrons (oxidized) while other gains those electrons (reduced)

However, it should be remembered that all the chemical reactions are not redox reactions.  There are several other types of reactions also.

NaCl + AgNO₃ → AgCl + NaNO₃

In such reactions none of’ the reactants is oxidized or reduced; simply the exchange of cation or anion takes place.

Oxidizing Agent (Oxidant):

The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, oxygen is making magnesium to lose electrons and hence in this reaction oxygen is the oxidizing agent.

Example – 2:

Consider reaction,

2K + Cl₂  → 2KCl

In this reaction, chlorine is making potassium to lose an electron and hence in this reaction chlorine is the oxidizing agent.

Characteristics of Oxidizing Agent:

• The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.
• In a reaction, the oxidizing agent oxidizes the other substance but is itself reduced.
• Oxygen, or a substance capable of giving oxygen, is always a good oxidizing agent.
• According to electron concept, an oxidizing agent is that which is capable of de-electronating the other substance.
• An oxidizing agent is an electron acceptor and during the redox reaction, it is electronated.
• Fluorine (F) has a maximum tendency to accept electrons hence it is the strongest oxidizing agent.

Examples of Common oxidizing Agents:

Oxygen (O or O₂), Ozone (O₃), Hydrogen peroxide (H₂O₂), Sulphuric acid (H₂SO₄), Nitric acid (HNO₃), Perchloric acid (HClO₄), Potassium chlorate (KClO₃), Acidified potassium dichromate (K₂ Cr₂ O₇ + H₂SO₄), Acidified potassium permanganate (KMnO₄ + H₂SO₄), Alkaline potassium permanganate (KMnO₄ + KOH)

Reducing Agent (Reductant):

The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, magnesium is making oxygen to gain electrons and hence in this reaction magnesium is reducing agent.

Example – 02:

Consider reaction, 2K + Cl₂  → 2KCl, In this reaction potassium, is making chlorine to gain an electron and hence in this reaction potassium is reducing agent.

Characteristics of Reducing Agent:

• The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.
• In a reaction, the reducing agent reduces the other substance but is itself oxidised.
• Hydrogen, or a substance capable of giving hydrogen, is always a good reducing agent.
• According to electron concept, a reducing agent is that which is capable of electronating the other substance.
• A reducing agent is an electron donor and during the redox reaction, it is de-electronated.
• Sodium (Na) has a maximum tendency to donate electron hence it is the strongest reducing agent.

Examples of Common Reducing Agents:

Hydrogen (H or H₂), Hydrogen iodide (HI), Hydrogen sulphide (H₂S), Lithium aluminium hydride (LiAI H₄), Sodium borohydride (NaB H₄), Sulphur dioxide (SO₂), Carbon (C), Ozone (O₃), Hydrogen peroxide (H₂O₂), Tin & hydrochloric acid (Sn  + HCl), Sodium & alcohol (Na + C₂ H₅OH), Metallic salts (ous) like SnCl₂ , FeSO₄ etc.

Science > Chemistry > Redox Reactions > Introduction to Redox Reactions

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A series of electrodes or half cells arranged in order of their increasing standard oxidation potentials or in the decreasing order of their standard reduction potentials is called an electromotive force series or electrochemical series. Electrochemical series is also known as e.m.f. series

Characteristics Electrochemical Series:

• In this series, all reduction potentials are given on hydrogen scale whose, Eo is taken as zero.
• The standard reduction potential of an element is a measure of the tendency of that element to get reduced.
• The element which has greater reduction potential gets reduced easily.  While the elements with low reduction potential will get easily oxidized
• Elements that lose electrons more easily have lower (negative) reduction potential and those which lose electrons with greater difficulty or instead of losing they accept electrons more easily have a higher (positive) reduction potential.
• In EMF series elements having higher (+ ve), the reduction potential is placed at the top.  While those having lower (-ve) reduction potential are placed at the bottom.  SHE has the middle position in the electrochemical series.
• The substances which are stronger reducing agents than hydrogen are placed below the hydrogen in the series and have negative standard reduction potential. The substances which are weaker reducing agents than hydrogen are placed above the hydrogen in the series and have positive standard reduction potential. Thus as we move down the group strength of reducing agent increases while the strength of the oxidizing agent decreases.
• Metal at the bottom is the most active metal. As we move down in the series activity and electropositivity of metals increase. Nonmetal at the Top is the most active nonmetal. As we move down in the series activity and electronegativity of nonmetal decreases.

Applications of Electrochemical Series:

For Choosing Elements as Oxidising Agents:

The elements which have more electron-accepting tendency are oxidizing agents. Elements at the top of the electrochemical series have higher (+ ve) reduction potential.  Hence they gain an electron from other elements and oxidize them.  So they are good oxidizing agents.

Element (F₂) at the topmost position of electrochemical series which has the highest reduction potential is the strongest oxidizing agent. Oxidizing power decreases from top to bottom in the series.

e.g. The elements like Cu, Ag, Hg, Br₂, Cl₂, etc. are good oxidizing agents. F₂ is the strongest oxidizing agent.

For Choosing Elements as Reducing Agents:

The elements which have more electron losing tendency are reducing agents. The elements at the bottom in the electrochemical series have lower (- ve) reduction potential. Hence they lose electrons readily and supply to other elements and reduce them. So bottom elements in electrochemical series are reducing agents.  Reducing strength goes on increasing from top to bottom in the series.

Element (Li) having the bottom-most position has the lowest reduction potential hence it is the strongest reducing agent.

e.g. The element like Zn, Cd, Ni, K, etc. are good reducing agents.

For Studying displacement reaction:

One metal can be displaced from a salt solution by another metal is known as a redox reaction. Elements having higher reduction potential will gain electrons and that having lower reduction potential will lose electrons. Hence element lower in electrochemical series can displace an element placed higher in electrochemical series from its salt solution.

Example-1:

Zn displaces Cu from CuSO₄, because, zinc is placed lower in electrochemical series and has lower reduction potential while Cu is placed higher in electrochemical series and has higher reduction potential.  Hence zinc can easily displace copper from CuSO₄

Zn + CuSO₄ → ZnSO₄ + Cu    i.e.

Zn + Cu⁺⁺_(aq) → Zn ⁺⁺_(aq) +  Cu

Example-2:

Fe displaces Cu from CuSO₄ because Fe is placed lower in electrochemical series and has lower reduction potential while Cu is placed higher in electrochemical series and has higher reduction potential.  Hence Fe can easily displace copper from CuSO₄.

Fe + CuSO₄ → FeSO₄(aq)+ Cu  i.e.

Fe + Cu⁺⁺_(aq) → Fe⁺⁺_(aq) +  Cu

To predict whether a given metal will displace another, from its salt solution:

A metal lower in the series will displace the metal from its solution which is higher in the series, i.e., the metal having low standard reduction potential will displace the metal from its salt’s solution which has a higher value of standard reduction potential. A metal lower in the series has a greater tendency to provide electrons to the cations of the metal to be precipitated.

Displacement of one nonmetal from its salt solution by another nonmetal:

A nonmetal higher in the series having the high value of standard reduction potential will displace another nonmetal with lower reduction potential i.e., occupying the position below in the series. The nonmetal’s which possess high positive reduction potentials have the tendency to accept electrons readily. These electrons are provided by the ions of the nonmetal having the low value of reduction potential. Thus, Cl2 can displace bromine and iodine from bromides and iodides.

Displacement of hydrogen from dilute acids by metals:

The metal which can provide electrons to H⁺ ions present in dilute acids for reduction evolve hydrogen from dilute acids. The metal having negative values of reduction potential possess the property of losing electron or electrons. Thus, the metals occupying lower positions in the electrochemical series readily liberate hydrogen from dilute acids and on ascending in the series tendency to liberate hydrogen gas from dilute acids decreases.

The metals which are above hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids.

Displacement of hydrogen from water:

Iron and the metals below iron are capable of liberating hydrogen from water. The tendency increases from top to bottom in electrochemical series. Alkali and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.

For Calculation of standard EMF of cell ( E^o_cell):

From the standard electrode potential values, it is easy to calculate EMF of cell. Standard oxidation potential values are given in EMF series. Eo  cell is calculated using formula:

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

e.g. in Daniell cell,

Now, From the series, E^o_Zn  = – 0.763 V ,  E^o_Cu =  + 0.337 V

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

∴  E^o_cell =    E^o_{red (Cu)}    –    E^o_{red (Zn)}

∴  E^o_cell =     0.337  –  ( -0.763)

∴  E^o_cell =     0.337  + 0.763

∴  E^o_cell =     1.10 V

For Checking Spontaneity of Redox Reactions:

If cell is assembled such that one electrode has higher positive oxidation potential and other has lower negative oxidation potential then redox cell reaction will be spontaneous and cell will have positive EMF.  On the contrary if EMF of cell is negative then redox cell reaction will be non spontaneous.

e.g. in Daniell cell,

Now, From the series, E^o_Zn  = – 0.763 V ,  E^o_Cu =  + 0.337 V

E^o_cell =    E^o_{red (cathode)}    –    E^o_{red (anode)}

∴  E^o_cell =    E^o_{red (Cu)}    –    E^o_{red (Zn)}

∴  E^o_cell =     0.337  –  ( -0.763)

∴  E^o_cell =     0.337  + 0.763

∴  E^o_cell =     1.10 V

Since cell has positive EMF, following redox cell reaction is spontaneous.

Zn + Cu⁺⁺_(aq) → Zn ⁺⁺_(aq) +  Cu

Thus higher the positive EMF of the cell, the more is the spontaneity of the redox cell reaction.

For Construction of a Cell:

Various cells can be constructed by combining standard electrodes given in EMF series as per the requirement of e.m.f.

e.g If a cell of e.m.f. 1.1 V is required, then from e.m.f. series we can locate zinc and copper electrode whose combination gives required e.m.f.

For Corrosion Treatment:

When two metals which are in contact with each other are exposed to the atmosphere, the element lower in series will be oxidized. i.e. it is rusted and destroyed.

If there is a scratch on the galvanized sheet of iron, and iron is exposed then zinc is rusted and iron is protected. This is because in e.m.f. series zinc is below the iron. But if there is a scratch on the tin-plated iron, iron gets rusted because in e.m.f. series iron is below tin.

To Find Reactivity of Metals:

As we move down in the electrochemical series reactivity of metal increases. Alkali metals and alkaline metals at the bottom are highly reactive. They can react with cold water and evolve hydrogen. They can dissolve in acid-forming salt.

Metals like Fe, Pb, Sn, Ni, Co which are in little higher in the series do not react with cold water but react with steam and evolve hydrogen. Metals like Cu, Ag, and Au which lie above the hydrogen are less reactive and do not react with water in any form to evolve hydrogen.

To Ascertain Electropositivity of Metals:

Strongly electropositive metals:

Metals having standard reduction potential near about -2.0 volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in nature.

Moderately electropositive metals:

Metals having values of standard reduction potentials between 0.0 and about -2.0 volt are moderately electropositive. Al, Zn, Fe, Ni, Co, etc., belong to this group.

Weakly electropositive metals:

The metals which are above hydrogen and possess positive values of standard reduction potentials are weakly electropositive metals. Cu, Hg, Ag, etc., belong to this group.

To Find Thermal Stability of Metallic Oxides:

The thermal stability of the metal oxide depends on its electropositive nature. As the electropositivity increases from top to bottom, the thermal stability of the oxide also increases from top to bottom.

The oxides of metals having high positive reduction potentials are not stable towards heat. The metals which are above copper form unstable oxides, i.e., these are decomposed on heating.

To Determine the Products of Electrolysis:

In case two or more types of positive and negative ions are present in solution, during electrolysis certain ions are discharged or liberated at the electrodes in preference to others.

In general, in such competition, the ion which is the stronger oxidizing agent (higher value of standard reduction potential) is discharged first at the cathode. e.g. In a mixture of copper and silver ions, silver will be deposited first because the reduction potential of silver is higher than copper.

• The increasing order of deposition of few cations is: K⁺, Ca⁺⁺, Na⁺, Mg⁺⁺+, Al⁺⁺⁺, Zn⁺⁺, Fe⁺⁺, H⁺, Cu⁺⁺, Ag⁺, Au⁺⁺⁺.
• The anion which is a stronger reducing agent (low value of standard reduction potential) is liberated first at the anode.
• The increasing order of discharge of few anions is SO₄^– –, NO₃^–, OH^–, Cl^–, Br^–, I^–
• When an aqueous solution of NaCl containing Na⁺, Cl^–, H^+, and OH- ions is electrolyzed, H+ ions are discharged at cathode and Cl- ions at the anode, i.e., H2 is liberated at cathode and Cl2 at the anode.
• When an aqueous solution of CuS04 containing Cu++, H+ and OH- ions is electrolyzed, Cu⁺⁺ ions are dis­charged at the cathode and OH^– ions at the anode.

Previous Topic: Use of Nernst Equation

Science > Chemistry > Electrochemistry > Electrochemical Series

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In this article, we shall study the use of the Nernst equation to find e.m.f. of cell and electrodes.

Convention Followed While Calculation of Cell Potential (e.m.f.):

• In the symbolic representation of the cell, the right-hand side electrode is the cathode (positive electrode) and the left-hand side is the anode (negative electrode).
• All standard potentials are reduction potentials that are they refer to a reduction reaction.
• The cathode has a higher standard potential than the anode.
• For spontaneous reaction to take place the cell potential should be positive.

Illustrations for Use of Nernst Equation:

When Reactions are given:

Example – 1: 

Cr_(s) + 3Fe³⁺ _(aq) → Cr³⁺_(aq) + 3Fe²⁺ _(aq)

The cell formation is

Cr_(s)| Cr³⁺_(aq)|| Fe²⁺_(aq),Fe ³⁺_(aq)| Pt

The half cell reactions are

Cr_(s) → Cr³⁺(aq)+ 3e^–  (Oxidation)

Fe³⁺_(aq)+ 3e^– → Fe²⁺_(aq)(Reduction)

Hence n = 3

Nernst equation is

Example – 2:  

Al³⁺(aq) + 3e-  → Al_(s) ,

Here n = 3,  Nernst equation is

When Type of Electrode is Given:

Redox Electrode:

Example – 1:  

Pt | Sn²⁺, Sn⁴⁺

The Reduction reaction is

Sn⁴⁺_(aq)+ 2e^– → Sn²⁺_(aq)(Reduction),

Hence n = 2, Nernst equation is

Example – 2:

Pt | Fe²⁺, Fe³⁺

The Reduction reactions are

Fe³⁺_(aq) + 1e^– →  Fe ²⁺_(aq)(Reduction)

Hence n = 1, Nernst equation is

Metal Metal Ion Electrode:

Example – 1: 

Zn_(s)| Zn⁺⁺_(aq)

Reduction reaction for it is

Zn_(s)  →  Zn⁺⁺_(aq)     +   2e^–

Here n = 2,  Nernst equation is

Example – 2:

Al_(s)| Al³⁺_(aq)

The reduction reaction is

Al³⁺(aq) + 3e-  → Al_(s)

Here n = 3,  Nernst equation is

Metal Sparingly Soluble Salt Electrode:

Example – 1:

Cl^– _(aq) | AgCl_(s)| Ag

The Reduction reaction is

AgCl_(s)+ e^– → Cl^– _(aq) + Ag_(s) (Reduction)

Hence n = 1, Nernst equation is

Gas Electrode:

Example – 1:

Cl^– _(aq) |   Cl_2(g), (1 atm)| Pt

The Reduction reaction is

½ Cl_2(g) + e ^–  →   Cl^– _(aq)  (Reduction)

Hence n = 1, Nernst equation is

Important Terms:

Half-Cell:

An electrode in contact with an electrolyte containing its own ions is called a half cell. e.g. In Daniel cell, the zinc rod dipped in zinc sulphate solution is called zinc half cell.

Half-Cell Reaction:

The reaction taking place in a half cell or reaction taking place at each electrode is called half-cell reaction. e.g. In Daniel cell in zinc half cell oxidation takes place. Therefore the half-cell reaction is

Zn_(s)  →  Zn⁺⁺_(aq)     +   2e^–

Cell:

A combination of two half-cells such that oxidation takes place at one half cell and reduction takes place at other half-cell is called the cell. e.g. A Daniel cell is formed by the combination of zinc half cell and copper half cell. Oxidation takes place at zinc half cell and the reduction takes place at the copper half cell.

Single Electrode Potential:

The difference of potential between the electrode and its salt solution around it at equilibrium is called a single electrode potential. Electrode potential depends upon

• Nature of the element/ metal,
• Concentration or activity of ions in solution
• Temperature and
• Pressure in case of gas.

Standard Electrode Potential (E°):

The difference of potential between the electrode and its salt solution around it containing ion concentration at a unit activity at 298 K is called standard electrode potential.

Oxidation  Electrode Potential (E_ox):

The difference of potential between the electrode and its salt solution around it at equilibrium and at constant temperature due to oxidation is called oxidation potential.

Standard Oxidation Potential (E°_ox):

The difference of potential between the electrode and its salt solution around it containing ion concentration at a unit activity at 298 K  due to oxidation is called standard oxidation potential (S.O.P.).

Reduction Electrode Potential (E°_red):

The difference of potential between the electrode and its salt solution around it at equilibrium and at constant temperature due to reduction is called reduction potential.

Standard Reduction Potential (E°_red):

The difference of potential between the electrode and its salt solution around it containing ion concentration at a unit activity at 298 K  due to reduction is called standard reduction potential (S.R.P.).

Standard e.m.f. of Cell:

The algebraic sum of the standard oxidation potential of one electrode (anode) and the standard reduction potential of another electrode (cathode) is called the standard e.m.f. of a cell.

Note:

The oxidation potential of electrode is equal to the reduction potential of the electrode with the opposite sign

Gibb’s Energy Change:

In thermodynamics, the Gibbs free energy is a thermodynamic potential that measures the maximum or reversible work that may be performed by a thermodynamic system at a constant temperature and pressure (isothermal, isobaric).

As the cell reaction in an electrochemical cell progresses, electrons move through a wire connecting the two electrodes until the equilibrium point of the cell reaction is reached, at which point the flow of electrons ceases. Just cell performs the work. In electrochemistry, the maximum amount of electrical work a galvanic cell can do at constant temperature and pressure is Gibb’s free energy.

The amount maximum work a galvanic cell can do is given as

Electrical work = Amount of charge (nF) × Cell potential (E_cell)

Electrical work = n F E_cell

The reversible electrical work done in a galvanic cell reaction is equal to the decrease in its Gibb’s energy

Thus,     Electrical work = – ΔG

∴ – ΔG = n F E_cell

∴ ΔG = –  n F E_cell

The standard Gibb’s energy change is given by

ΔG° = –  nFE°_cell

Gibb’s energy is an extensive property, which depends on the amount of substance. But the electrical potential is an intensive property which does not depend on the amount of substance. Thus E°_cell remains constant. Thus if ΔG° changes there is the corresponding change in the number of electrons. It can be explained as follows

ΔG° = –  n F E°_cell 

If the stoichiometric equation of redox reaction is multiplied by 2, then the standard Gibb’s energy ΔG° gets doubled and the number of electrons ‘n’ also gets doubled.

From (1) and (2) we can see that the e.m.f. of cell in both cases is the same. It shows that electrical potential is an intensive property that does not depend on the amount of substance.

Relation between Standard Cell Potential and Equilibrium Constant:

The Gibb’s free energy of a galvanic cell is given by

G° = –  n F E°_cell

By thermodynamical and equilibrium concept

Previous Topic: Nernst Theory of Electrode Potential

Next Topic: Electrochemical Series and its Applications

Science > Chemistry > Electrochemistry > Use of Nernst Equation

The post Use of Nernst Equation appeared first on The Fact Factor.

In this article, we shall study the Nernst theory of electrode potential, Nernst equation, and its use.

Single Electrode or Half cell or Electrode Couple:

A single electrode or half cell or electrode couple is produced when a metal is dipped in the solution of its own ions.

e.g. Cu | Cu⁺⁺_(aq),   Zn| Zn⁺⁺_(aq) etc

A single vertical line indicates physical contact between the metal and its ions. Sometimes a couple is produced from gas and solution of its ions. In such cases noble metal like platinum is used as a conductor to adsorb the gas.

e.g. Pt| H_2(g) | H⁺_(aq)

Concept of electrode potential (Nernst Theory):

Nernst in 1889 gave his theory of electrode potential. An electrode is a couple of active element, and its ionic solution. When metal is immersed in its salt solution it shows two opposite tendencies called de-electronation (oxidation) and electronation (reduction).

Metals have a tendency to pass into solution as cations and liberate electrons.  This process is oxidation or de – electronation. The tendency of a metal to pass into its salt solution in the form of cations liberating electrons is called the solution pressure of metal (P_s). There is a reverse tendency of cations to deposit on the electrode by taking electrons.  This process is called as electronation or reduction. The tendency of the ions in the solution to be deposited back on the surface of the metal by taking electrons is called an osmotic pressure of ions (P_o).

Nernst suggested the mechanism of the establishment of the difference of potential in a cell.  His theory is based on the theory of electrolytic dissociation and his ideas of solution pressure and formation of the electrical double layer.

De-electronation:

The process in which an atom or ion of an element loses one or more electrons is called de-electronation. De-electronation takes place at the electrode when the solution pressure of metal is greater than its osmotic pressure.

Electrons released in oxidation are accumulated on the electrode and solution, on the other hand, acquires excess positive charge because of the excess of cations. This gives rise to the electrical double layer at the electrode surface. Because of this electrical double layer, a potential difference is set up.  As this potential is due to the oxidation, it is called oxidation potential.

e.g. In Daniel cell solution pressure of zinc is more.  Zinc passes into its ion solution as Zn ⁺⁺ ions and electrons released in oxidation get accumulated on the zinc rod. Thus de-electronation takes place at zinc half cell in a Daniel cell.

Zn_(s)     →      Zn⁺⁺_(aq)     +   2e^–

Thus at zinc electrode, a negative potential is developed which is due to the oxidation.

Electronation:

The process in which an atom or ion of an element gains one or more electrons is called electronation.  Electronation takes place at the electrode when the osmotic pressure of metal is greater.

Due to electronation the electrode loses electrons continuously and acquires a positive charge. And the solution, on the other hand, acquires an excess negative charge.  Thus electrical double layer is set up across the surface of the metal. Because of the electrical double layer potential difference is set up. As this potential is due to reduction it is called a reduction potential.

e.g. In Daniel cell osmotic pressure of copper ions is more. Cu⁺⁺ ions from solution take electrons and deposited on the surface of the metal. Thus due to reduction, there is a removal of electrons from the metal surface. Thus electronation takes place at the copper half-cell in Daniel cell.

Cu⁺⁺_(aq)     +   2e^–    →  Cu_(s)

Thus at the copper electrode, a positive potential is developed due to reduction.

Rate of Electronation or De-electronation:

The rate of de-electronation and electronation differs from metal to metal.  There are three possibilities.

P_s > P_o: When the solution pressure is greater than the osmotic pressure. The rate of de-electronation is greater than the rate of electronation. The electrode undergoes oxidation. Thus negative potential develops on the electrode and it acts as an anode.

P_o > P_s: When the osmotic pressure is greater than the solution pressure. The rate of electronation is greater than the rate of de-electronation. The electrode undergoes a reduction. Thus positive potential develops on the electrode and it acts as a cathode.

P_s = P_o: When the solution pressure is equal to the osmotic pressure. The rate of electronation is equal to the rate of de-electronation. Thus there is no double layer formation and hence no potential is developed on the electrode.  Such an electrode is called the null electrode.

Nernst Equation:

For single electrode potential: Let M be the metal, ‘n’ be the number of electrons involved in the electrode.  Then the reactions are,

M    →   Mⁿ⁺ +   n e^–    (oxidation)   OR

Mⁿ⁺  +    n e^–    → M    (reduction)

According to the Nernst equation at 25° C

Where, E^o = standard oxidation potential

E = oxidation potential

R = 8.314 J K^–1 mol ^–1

n = no. of electrons involved in electrode reaction.

F = Faraday’s constant = 96500 C

T = temperature in K

[Oxidation state] = concentration of Mⁿ⁺ ions in mol dm^-3

[Reduced state] = activity of pure metal = 1

This expression gives variation of electrode potential with respect to electrolyte concentration.

The first part of the equation represents standard state electrochemical conditions and the second term is a correction for non-standard state electrochemical conditions.

Mathematical expression for EMF of a cell:

Let us consider general cell reaction

aA   +   bB  →  cC   +  dD

Let ‘n’ be the number of electrons in cell reaction.

Then according to Nernst equation, EMF of cell is given by

Calculation of Cell Potential Using Nernst Equation:

Consider a cell

Zn_(s)| Zn⁺⁺_(aq)|| H⁺_(aq) (1 M)|H_2(g) (1 atm.) | Pt  +

Oxidation reaction at anode:

Zn_(s)  →  Zn⁺⁺_(aq)     +   2e^–

Reduction reaction at anode: 

2H⁺_(aq) + 2e^–    → H_2(g)

Net cell reaction:   

Zn_(s)  + 2H⁺_(aq) → Zn⁺⁺_(aq)  + H_2(g)

The e.m.f. of a cell at 25 °C by Nernst equation is given by

Where concentrations are in mol dm^-3 and pressure is in atm.

Calculation of Electrode Potential Using Nernst Equation:

Consider an electrode  Zn_(s)| Zn⁺⁺_(aq)

Reduction reaction for it is

Zn_(s)  →  Zn⁺⁺_(aq)     +   2e^–

The electrode potential at 25 °C by Nernst equation is given by

at 25 °C. Where concentrations are in mol dm^-3 and pressure is in atm.

Previous Topic: Types of electrodes

Next Topic: Use of Nernst Equation

Science > Chemistry > Electrochemistry > Nernst Theory

The post Concept of Electrode Potential: Nernst Theory appeared first on The Fact Factor.

In this article, we shall study different types of electrodes, their representation, writing cell reactions, and finding e.m.f. of a cell.

There are four types of electrodes

• Gas electrodes
• Metal–sparingly soluble metal salt electrodes
• Metal – metal ion electrodes
• Redox Electrodes

Gas Electrodes:

A gas electrode consists of a gas (e.g. H2, Cl2, O2) in contact with a solution containing the ions derivable from the gas e.g. H+, Cl-, OH-. The potential of the gas electrode depends upon the concentration of its ions in the solution and the pressure of a gas.

A gas electrode consists of gas, bubbled about inert metal wire (platinized platinum electrode) immersed in a solution containing ions with which gas is irreversible. Platinum is used as conductor and to adsorb the gas. e.g. Standard hydrogen electrode.

Examples of gas Electrodes:

Standard Hydrogen Electrode (SHE):

SHE is represented as,

Pt| H_2(g) (1 atm.)| H⁺_(aq) (1 M)

The half cell reactions are

H_2(g) →   2H⁺_(aq) + 2e^– (oxidation) (L.H.S.)

2H⁺_(aq) + 2e^– → H_2(g)  (reduction) (R.H.S.)

The electrode potential is arbitrarily assigned zero. This electrode is cation electrode.

Chlorine gas electrode:

This electrode is anion electrode.  Chlorine gas electrode is represented as,

Pt| Cl_2(g) (1 atm.)| Cl^–_(aq) (1 M)

The half cell reactions are

2Cl^–_(aq) → Cl_2(g) + 2e^–   (oxidation) (L.H.S.)

Cl_2(g) + 2e^– →  2Cl^–_(aq)   (reduction) (R.H.S.)

Oxygen gas electrode:

Oxygen gas electrode is represented as,

Pt | O_2(g) (1 atm)| OH^– _(aq) (1M)

The half cell reaction is

4OH^– → 2H₂O+ O_2(g) + 4e^–   (oxidation) (L.H.S.)

2H₂O + O_2(g) + 4e^– →  4OH^– (reduction) (R.H.S.)

Metal-Sparingly Soluble Metal Salt Electrode:

Reversible anion electrode is also called as metal- sparingly soluble metal salt electrode. In this electrode a metal, a sparingly soluble salt of the metal in equilibrium with a solution containing the same anion as the sparingly soluble salt. e.g. Calomel electrode.

Metal – Metal Ion Electrodes:

In this case, the metal strip is kept in contact with the solution of a water-soluble salt-containing cation of the same metal.

e.g. Zn_(s) | Zn⁺⁺_(aq)

In the electrochemical cell, the electrode having higher oxidation potential undergoes oxidation and acts as the anode or negative electrode and the electrode having lower oxidation potential undergoes reduction and acts as the cathode or positive electrode.

Examples of metal – metal ions electrodes:

Zn_(s) | Zn⁺⁺_(aq)

Zn_(s) →  Zn⁺⁺_(aq) +   2e^–   (Oxidation)

Zn⁺⁺_(aq) +   2e^–  → Zn_(s) (Reduction)

Cu_(s) | Cu⁺⁺_(aq)

Cu_(s) →  Cu⁺⁺_(aq) +   2e^–   (Oxidation)

Cu⁺⁺_(aq) +   2e^–  → Cu_(s) (Reduction)

Redox Electrode:

In these electrodes, an inert metal like Pt is dipped in a solution containing ions of an active metal in two different oxidation states.

Pt | Fe²⁺, Fe³⁺

Fe²⁺      →     Fe³⁺  e^– (Oxidation)

Fe⁺⁺⁺ +   e^–      →   Fe⁺⁺ (Reduction)

Pt | Sn²⁺, Sn⁴⁺

Sn²⁺  →  Sn⁴⁺  +     2e^– (Oxidation)

Sn⁴⁺ +  2e^–    →   Sn²⁺ (Reduction)

Writing Cell Reaction and Finding E.M.F. of a Cell:

Redox Potential:

The potential developed due to the ability of ions to lose or gain electrons forming a higher or lower stable oxidation state is called redox potential.

The redox potential depends upon the ratio of concentrations of two types of ions.

Pt | Fe²⁺_(aq)(1M),  Fe³⁺ _(aq) (1M)        E²_ox= – 0.771 V

Representation of cells containing standard and reference electrodes:

A cell composed of zinc rod contact with 1 molar zinc ion solution and saturated calomel electrode.

Zn_(s)| Zn²⁺(1M) || KCl_(aq) (saturated) | Hg₂Cl_2(s)|Hg_(l), Pt +

   Cell composed of SHE and saturated calomel electrode

Pt | H_2(g) (1 atm)| H⁺_(aq) (1M) || KCl_(aq)(saturated)|Hg₂Cl_2(s)| Hg_(l) ,Pt  +

Cell Reactions:

Steps to Write Cell Reaction of Galvanic Cell:

• Represent the given galvanic cell with standard convention.
• The electrode on the left side of the representation shows that it is anode and oxidation takes place at this electrode. Write half cell oxidation reaction half-cell reaction for it.
• The electrode on the right side of the representation shows that it is cathode and reduction takes place at this electrode. Write a half-cell reduction reaction half-cell reaction for it.
• Balance above two reactions for electrons for oxidation and reduction reaction.
• Add the two reactions and obtain net (overall)  cell reaction.

Step – 1: Represent the cell conventionally:

Pb_(s) | Pb²⁺_(aq) (1M) || Ag⁺_(aq) (1M)| Ag_(s) +

Step – 2: Write left hand side half-cell reaction: Pb(s) is on the left side of the representation shows that it is anode and oxidation takes place at Pb(s) electrode.

Pb_(s) →  Pb²⁺_(aq) +   2e^– (Oxidation) … (1)

Step – 3: Write right hand side half-cell reaction: Ag_(s) is on right side of the representation shows that it is cathode and reduction takes place at Ag_(s) electrode.

Ag⁺_(aq)   +   e^– → Ag_(s)    (Reduction)  … (2)

Step – 4: Balance the Electrons of above two half cell reactions:

Multiply equation (2) by 2 to balance electrons.

2Ag⁺_(aq)   +   2e^– → 2Ag_(s)    (Reduction)  … (2)

Step – 5: Adding equations (1) and (3) we get overall reaction.

Pb_(s) +  Ag⁺_(aq) →  Pb²⁺_(aq)    +  Ag_(s)

Steps to Find E.M.F. of Galvanic Cell:

• Represent the given galvanic cell with standard convention.
• The electrode on the left side of the representation shows that it is anode and oxidation takes place at this electrode.
• The electrode on the right side of the representation shows that it is cathode and reduction takes place at this electrode.
• Obtain standard oxidation potential values from the electromotive series for the material of cathode and anode.
• Use the following formula for calculation of e.m.f. of a cell.

E^o_Cell =  E^o_(ox/cathode) –   E^o_(ox/anode)

OR

E^o_Cell =  E^o_(ox/cathode)+   E^o_(red/anode)

To find e.m.f. of Daniel Cell :

Step – 1: Represent the cell conventionally

Step – 2: Decide anode and cathode: Pb(s) is on the left side of the representation shows that it is anode and oxidation takes place at Pb(s) electrode. Ag(s) is on the right side of the representation shows that it is cathode and reduction takes place at Ag(s) electrode.

Step – 3: Get values of oxidation potential or reduction potential for electrodes from From electrochemical series

E^o_(ox/Zn) = 0.76 V and EE^o_(ox/Cu) =-0.34 V

Step – 4: calculate e.m.f of cell:

E^o_Cell = E^o_(ox/cathode) –   E_(ox/anode)

E^o_Cell =  E^o_(ox/Zn) –  E^o_(ox/Cu)

E^o_Cell   =      0.76    –  (- 0.34)

E^o_Cell  =  0.76    +   0.34

E^o_Cell   =      1.1 V

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Science > Chemistry > Electrochemistry > Types of Electrodes

The post Types of Electrodes appeared first on The Fact Factor.

The electrode whose potential is arbitrarily fixed or is exactly known at a given constant temperature is known as a reference electrode. Using reference electrodes unknown potential of any other single electrode can be found out e.g. two commonly used reference electrodes are standard hydrogen electrode (SHE) and Calomel electrode.

Use of Reference Electrodes:

A cell is constructed using the given electrode and reference electrode. Using a potentiometer and standard cell (like Weston cell) the e.m.f. of the cell can be measured. By knowing the e.m.f. of the cell and potential of the reference electrode, the potential of the electrode in question can be easily determined.

Standard Hydrogen Electrode (SHE).

SHE is defined as the electrode in which pure and dry hydrogen gas is bubbled at 1 atmospheric pressure and 298 K on a platinized platinum foil through a solution containing H⁺ ions at unit activity.

Construction:

SHE consists of a glass jacket which has a small inlet at the top and many outlets at the bottom.  Inside the glass jacket, there is a glass tube closed at both ends. It has a platinum wire sealed in it. At the lower end of the platinum wire, there is a platinized platinum plate. At the bottom of the glass tube, there is little mercury which is meant for good electrical contact. The glass jacket along with a glass tube is dipped in a vessel containing 1 M HCI solution.

Working:

Pure and dry hydrogen gas is bubbled through HCI solution from the inlet at a constant pressure of 1 atm.  Hydrogen gas is adsorbed on the platinum plate and acts as a hydrogen electrode.  An equilibrium between H2 gas and H+ ion is established across the metal.

Electrode reaction:

The electrode is reversible with respect to hydrogen ions. During working, hydrogen gas from platinum plate changes into hydrogen ions and electrons are set free. These electrons accumulate on the platinum plate.

If the electrode is serving as an anode, then the half-cell reaction is

H_2(g) →   2H⁺_(aq) + 2e^– (oxidation)

The electrons set free remains on the platinum plate and transferred to the other electrode through Pt. wire.  As the process is oxidation, a positive potential is developed.  It is comparatively very small, it is arbitrarily taken as a zero.

If the electrode is serving as a cathode, then the half-cell reaction is

2H⁺_(aq) + 2e^–  →  H_2(g) (reduction)

Representation of electrode:

When acting as an anode,     Pt| H_2(g) (1 atm.)| H⁺_(aq) (1 M)

When acting as cathode,   H⁺_(aq) (1 M) |  H_2(g) (1 atm.) | Pt

Advantages of SHE:

• SHE is used as a reference electrode.  When it is coupled with any other electrode whose potential is to be determined.  The potential of the cell is then measured using a potentiometer.  Since the potential of SHE is zero, the potential of the cell is equal to the potential of another electrode or e.m.f. of the cell itself. Thus when SHE is used, the correction for its own potential is not necessary.
• It can be used over the entire pH range.
• It gives no salt error.
• It consists of a pH scale with voltage measurement.

Difficulties in setting up of SHE:

• It is difficult to obtain 100 % pure and dry hydrogen gas. Even traces of impurities in the hydrogen gas makes the electrode inactive and irreversible.
• It is difficult to maintain exactly 1 atmospheric pressure on hydrogen gas for a longer time.
• It is difficult to maintain the concentration of HCI solution as 1 M because due to the bubbling of hydrogen gas through HCI solution, water is evaporated and hence the concentration of HCI solution may change.
• Since it is made up of glass, it is not so handy.
• Platinum used is rather expensive.
• It is difficult to prepare ideal platinized platinum.

Arrangement to Find Oxidation Potential of another Electrode Using SHE:

Calomel Electrode:

Construction:

The calomel electrode consists of a broad glass tube having sidearm as shown in the figure. The sidearm is used for dipping it any solution used for coupling the calomel electrode. At the bottom of the glass tube, there is pure mercury and a platinum wire is sealed into it at the bottom for electrical connections. The wire runs through a separator glass tube to the top of the tube for electrical contact. Above pure mercury, there is a paste of mercurous chloride (calomel) (Hg2Cl2) in mercury. The rest of the glass vessel and sidearm A is filled with a saturated KCl solution.  KCI solution of 0.1 M or of 1 M can also be used.  Sidearm is plugged with glass wool. The glass tube is closed from the top.

Working:

Since the calomel electrode is reversible, two types of reactions are possible depending upon the nature of another electrode with which it is coupled.

When acting as  negative electrode:

2 Hg_(l) → 2 Hg⁺ + 2 e^–

2 Hg⁺ + 2 Cl – → Hg₂Cl_2(s)

The net oxidation reaction is

2Hg_(l) + 2Cl^–_(sat) → Hg₂Cl_2(s)+  2e^–

Thus oxidation takes place when it is coupled with other electrode having lower oxidation potential.

When acting as positive electrode:

Hg₂Cl_2(s)      →   2Hg_(l) + 2Cl^––

2 Hg⁺ + 2 e^–    → 2 Hg

The net reduction reaction is

Hg₂Cl_2(s)   +  2 e^–   →  2 Hg⁺ + 2Cl^–

Thus reduction takes place when it is coupled with other electrode having greater oxidation potential.

Representation of Electrode:

When acting as anode:   Pt | Hg_(l) | Hg₂Cl_2(s)  | KCl_(sat)

When acting as anode:  KCl_(sat) |  Hg₂Cl_2(s) | Hg_(l) |Pt

Oxidation Potential of Calomel Electrode:

The oxidation potential of the calomel electrode depends upon the concentration of KCl solution used. The negative potentials indicate that when combined with SHE reduction takes place at the calomel electrode.

Advantages of calomel electrode:

• It is easy to set up and easily reproducible.
• It is convenient and easy to transport.
• It is very compact and smaller in size requires little space.
• No separate salt bridge is required as it has already a side tube containing KCl solution.
• Potential does not change appreciably with time and a slight change in temperature.

Disadvantages of Calomel Electrode:

• When half-cell potentials are to be measured, compensation for potential is necessary.
• The calomel electrode cannot be used in the measurement of potentials of the cell where K+ and Cl – ions interfere in the electrochemical reactions of the cell.
• The oxidation potential of the electrode depends on the concentration of KCl. If the concentration of KCl changes, the oxidation potential of electrode changes.

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Science > Chemistry > Electrochemistry > Reference Electrodes

The post Reference Electrodes appeared first on The Fact Factor.

In this article, we shall study the representation of cell using conventions on paper. We shall also discuss the construction and working of the salt bridge.

Electrochemical (Voltaic or Galvanic) cell consists of two half cells. In one half-cell, an oxidation reaction takes place and electrons are generated in the process. While in the second half cell reduction reaction takes place and electrons are absorbed or consumed in this process. The two half-cells are connected to each other internally by porous partition and connected externally by means of connecting wire.

Both the oxidation and reduction reaction takes place simultaneously and separately. The movement of electrons from oxidation half-cell to reduction half-cell through external circuit constitutes the electrical current.

Convention for Representing the Voltaic Cell:

Every time it is not possible to draw neat diagrams of the cell.  So as to avoid this difficulty, following conventions (rules) are used.

• The electrode with greater oxidation potential (negative electrode or anode) is written to the left.  Here oxidation takes place.
• The electrode material is written first, followed by its electrolyte.
• The electrode with lower oxidation potential ( positive electrode or cathode ) is written to the right.  Here reduction takes place.
• The electrolyte is written first followed by the electrode material.
• A single vertical line is drawn between the electrode and its electrolyte which represents direct contact but the separation of the two phases.
• The double vertical line is drawn between two salt solutions which indicate indirect contact by means of a porous pot or a salt bridge.
• Concentration or activity of solutions at two electrodes is written in brackets as (C = 1),    (C = 2) or (a = 1), (a  = 2).
• In the case of gas electrodes, inert metal conductors like platinum are used to establish electrical contact. It should be incorporated in the convention. e.g. a) Hydrogen electrode is represented as

• The representation of the Daniell cell by the above convention is as follows.

More Examples:

A cell consisting of zinc and cadmium electrodes:

As zinc has higher oxidation potential, it acts as -ve electrode.  Cadmium acts as a +ve electrode.  Oxidation takes place at zinc while reduction takes place at Cadmium.

–   Z n   |   Z n⁺⁺   ||   Cd⁺⁺ |   Cd   +

(1 M)       (1 M)

At anode (Zn) :  Zn  → Zn⁺⁺  + 2 e^–   (oxidation)

At Cathode (Cd): Cd⁺⁺  + 2 e^– →   Cd  (reduction)

Net cell reaction :  Zn   + Cd⁺⁺ →  Zn⁺⁺ + Cd     (redox)

Representation of a cell consisting of hydrogen and copper electrodes:

As hydrogen has greater oxidation potential than copper, it acts as -ve electrode (anode) and copper act +ve electrode (cathode).  Oxidation takes place at hydrogen and reduction takes place at copper.

At anode (H₂) :     H_2(g) →  2 H⁺ +   2^– (oxidation)

At cathode (Cu) :  Cu⁺⁺ + 2 e^–  → Cu (reduction)

Net cell reaction :  H_2(g) + Cu⁺⁺  →  2H⁺ + Cu (redox)

Representation of a cell consisting of copper and silver electrodes:

As copper has greater oxidation potential than silver, acts as anode and silver act as the cathode.  Oxidation takes place at copper and reduction takes place at silver.

–    Cu | Cu⁺⁺ ||    Ag⁺  |   Ag⁺

(C =1)   (C = 2)

At anode (Cu):  Cu  →  Cu⁺⁺  + 2 e^–   (oxidation)

At cathode (Ag):  2 Ag⁺   +   2 e^– → 2 Ag     (reduction)

Net cell reaction: Cu + 2 Ag⁺   →  Cu⁺⁺ + 2 Ag (redox)

Representation of a cell from the cell reaction:

1/2 Cl₂+   e^– →  Cl^–

2 OH^–  →  1/2 O₂ +  H₂O   +  2e^–

As chlorine undergoes reduction hence chlorine is a cathode and it must be written to the right.  In the second reaction OH^– ions undergo oxidation hence oxygen electrode is an anode and it must be written to the left.  The cell can be represented as,

Pt, O_2(g) | OH^– _(aq) || Cl^–_(aq) | Cl_2(g) ,Pt

Salt Bridge:

It consists of an inverted U shaped glass tube containing the saturated solution of a strong electrolyte like KCl, KNO3, NH4NO3 immobilized by agar-agar gel with glass wool plugs at the two ends.

Working: The two electrolytic solutions in two half-cells are connected by dipping the arms of the tube in an inverted position in the solutions.

Electrodes and its Types:

Electrodes are metallic or non-metallic rods immersed in the electrolyte. They conduct electric current through them. Carbon and platinum are mainly used electrodes because they are inert and do not get dissolved in the electrolytic solution.

Electrodes are of two types, a) indicator electrode b) Reference Electrode

Indicator Electrode:

The electrode whose potential depends upon the concentration of a particular ion in the solution in which it is dipped and is usually used to find out the concentration of ions in the solution is known as indicator electrode.

It usually consists of metal in the form of wire or rod kept in contact with its salt solution.

e.g. Ag_(s) | Ag⁺ _(aq)     (a = x M)

The electrode system is used for the determination of the concentration of the solution used in that half-cell. All electrodes except the reference electrode are called indicator electrodes.

Reference Electrode:

The electrode whose potential is arbitrarily fixed or is exactly known at a given constant temperature is known as a reference electrode. Using reference electrode unknown potential of any other single electrode can be found out e.g. Two commonly used reference electrodes are Standard hydrogen electrode (SHE) and Calomel electrode.

A cell is constructed using the given electrode and reference electrode. Using a potentiometer and standard cell (like Weston cell) the e.m.f. of the cell can be measured. By knowing the e.m.f. of the cell and potential of the reference electrode, the potential of the electrode in question can be easily determined.

Previous Topic: Secondary Electrochemical Cells

Next Topic: Reference Electrodes

Science > Chemistry > Electrochemistry > Representation of Electrochemical Cell

The post Representation of Electrochemical Cell appeared first on The Fact Factor.

In this article, we shall study the construction and working of lead accumulator, and Ni-Cd cells.

Lead Accumulator:

A lead accumulator is a secondary cell because electrical energy is not generated within the cell itself but it is previously stored in it from the external source. It is reversible cell because cell reactions are reversed if external e.m.f. just greater than the e.m.f. of this cell is applied. Thus in this cell, the net cell reaction can be reversed by applying external opposing e.m.f. greater than the cell e.m.f.

This cell can store electrical energy in form of from external source (charging) and can supply it during discharging.  The energy is stored in the form of chemical energy. So this cell is a storage cell or accumulator or storage battery. Its voltage does not depend on the size of the electrodes or the size of its cell but depends upon the strength of a sulphuric acid solution.

Construction:

A lead accumulator consists of lead plates as the negative electrode. Lead plates impregnated with lead oxide act as the positive electrode. Negative and positive electrodes are arranged in an alternate manner. This assembly of lead plates is dipped in a non-conducting vessel made up of glass or plastic or ebonite and containing 38% H₂SO₄ (sp. gr. 1.215). All positive plates are connected to each other and all negative plates are connected to each other.

Representation of Cell:

Pb_(s)|PbSO_4(s)|38%H₂SO_4(aq)| PbSO4 _(s)| PbO_2(s) |Pb_(s) +

Working: 

Discharging of Cell: 

When cell starts working, oxidation takes place at lead plates and reduction takes place at lead plates with lead oxide.  This is called as discharging of cell.

2 H₂SO_4(aq)  → 4H⁺ +  2 SO₄ ^– –

Reaction at negative electrode (Anode)

Pb  →    Pb⁺⁺ +   2 e^–

Pb⁺⁺ + SO4^– – → PbSO4_(s)

Pb_(s) + SO₄^– –   →   PbSO_4(s) + 2e^–         (oxidation)

Reaction at positive electrode (Cathode)

PbO_2(s) +  4H⁺ + SO4^– –  + 2e^–  → PbSO_4(s)+ 2H₂O_(l)

Net cell  reaction during discharging

Pb_(s) + PbO_2(s) +  2H₂SO_4(aq)  → 2 PbSO_4(s) +  2H₂O_(l)

Thus during discharging sulphuric acid is converted into water and hence specific gravity of sulphuric acid solution falls to 1.17.

Charging of cell: 

When external emf greater than the emf of this cell is applied, exact reverse reactions take place.  At positive electrode oxidation takes place and a negative electrode reduction takes place.  This is called the charging of the cell.

Reaction at negative electrode (Cathode)

PbSO_4(s) + 2e^–  →   Pb_(s) + SO₄^– –   

Reaction at positive electrode (Anode)

PbSO_4(s)+ 2H₂O_(l)      →  PbO_2(s) +  4H⁺ + SO4^– –  + 2e^– 

Net Cell  Reaction During Charging

2 PbSO_4(s) +  2H₂O_(l)

Pb(s) + PbO_2(s) +  2 H₂SO_4(aq)

Pb_(s) + PbO_2(s) +  2H₂SO_4(aq)  → 2 PbSO_4(s) +  2H₂O_(l)

E.M.F. of Cell:

A fully charged accumulator with 38% H₂SO₄ has a voltage about 2.041 V i.e. approx 2 V.

Uses of Lead Accumulator:

• Lead accumulators are used in motor cars and other automotive vehicles and therefore popularly known as car batteries. Here generally six cells are connected in series to give a voltage of 12 V.
• They are used in laboratories as a source of constant DC voltage.
• They are used in telephone and telegraph offices.
• They are used in electric clocks, radio sets, burglar’s alarms etc.
• They are also used as a source of electricity in the initial stages of the rocket.

Maintenance of Lead Accumulator:

• It is kept erect and should not be tilted.
• When the cell is not in use for a longer time, the terminals should be disconnected and acid from the cell should be completely removed and it is to be kept dry.
• The concentration of sulphuric acid should always be maintained to get the optimum value of cell potential.
• The cell should not be allowed to run down much. When in use the e.m.f. of the cell decreases. When the voltage is about 1.8 V, the cell should be recharged.
• The cell should not be overcharged i.e. e.m.f. of the cell should not be more than 2.1 V.
• The acid level should be maintained properly by adding distilled water.
• Positive and negative terminals of the lead accumulator should be cleaned to remove insoluble lead sulphate from time to time for good electrical contact.
• The cell should be charged regularly and never left in the discharged condition.

Nickel Cadmium Cells:

The first Ni–Cd battery was created by Waldemar Jungner of Sweden in 1899. Nickel-Cadmium battery is a rechargeable battery with Nickel Oxide-Hydroxide (cathode) and Cadmium as a cathode and Alkaline Potassium Hydroxide as the electrolyte.

Reaction at the anode (Oxidation):

Cd_(s) + 2OH^– _(aq) →  Cd(OH)_2(s) + 2e^–

Reaction at the cathode (Reduction):

NiO_2(s)+2H₂O_(l) + 2e^– →   Ni(OH)_2(s)  + 2OH^– _(aq)

Net Reaction:

Cd_(s) + 2H₂O_(l) + 2e^–   →  Cd(OH)_2(s)  + Ni(OH)_2(s) 

The reaction products are solid which gets adhered to the respective electrode surface, hence these cells can be recharged. The e.m.f obtained is 1.4 V. These cells have much more life than other dry cells. They are used in electronic watches, calculators and photographic equipment. They have low internal resistance, Can tolerate deep discharge cycle and can also get charged rapidly.  These batteries also have a long lifetime and low maintenance costs.

Nickle-Cadmium batteries are costlier than lead-acid battery also the Cadmium used in the battery causes severe environmental pollution.

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Science > Chemistry > Electrochemistry > Secondary Electrochemical Cells

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Types of Electrochemical Cell:

An electrochemical cell is a device which is used to generate electrical energy at the expense of spontaneous oxidation-reduction reaction. An electrochemical cell is also known as a galvanic cell or a voltaic cell.  In an electrochemical cell, chemical energy is converted into electrical energy. There are two types of voltaic or galvanic cells used to get an electric current of low voltage: a) Primary cells and b)Secondary cells

Primary Cell:

A cell in which electrical energy is generated within the cell itself is called a primary cell. e.g. Dry cell, Daniell cell. Such a cell cannot be reused.

Secondary Cell:

A cell in which electrical energy is not generated within the cell itself but it previously stored in it from an external source is called the secondary cell. e.g. Lead accumulator.  Secondary cells can be reused.

Primary Cells:

Daniel cell:

Principle:

Due to the spontaneous redox reaction, electrical energy is generated within the cell itself.  So Daniel cell is a primary voltaic cell. When an opposing voltage slightly greater than the cell voltage is applied, the cell reaction is reversed so Daniel cell is a reversible cell. Thus Daniel cell is a primary reversible electrochemical cell.

Construction:

Daniel cell consists of two electrodes or two half cells i.e. zinc half-cell and copper half-cell. Daniel cell with porous partition consists of a vessel which is divided into two compartments

In one compartment, a solution of zinc sulphate ZnSO₄ is placed in which a zinc rod is dipped. This rod acts as a negative electrode or anode. This is known as the negative half cell or zinc half cell. The other compartment is filled with copper sulphate CuSO4 solution. In this solution, a copper rod is dipped which acts as a positive electrode cathode. This is known as a positive half cell or copper half cell.

The two half-cells are connected externally by a conducting wire and internally by means of a porous partition or salt bridge or porous pot. The porous partition allows the diffusion of ions through it but prevents excessive mixing of the two solutions.

Representation of the Cell:

Working of the cell:

As long as the circuit is closed, the following reactions take place. Oxidation takes place at zinc half cell and reduction takes place at the copper half cell.  Zinc rod acts as anode i.e.   -ve electrode while copper rod acts as cathode i.e. + ve electrode

Cell reactions:

At anode (Zn) or – ve electrode: 

Zinc atoms from rod enter the solution forming Zn ions ( Zn⁺⁺) leaving behind 2 electrons on the rod.

Zn  → Zn⁺⁺  + 2 e^–   (oxidation)

At cathode (Cu) or + ve electrode: 

Cu⁺⁺ ions from CuSO₄ solution acquire 2 electrons from the copper rod and during this process get deposited on the copper electrode in the form of Cu atoms.

Cd⁺⁺  + 2 e^– →   Cd  (reduction)

Net cell reaction:

Zn   + Cu⁺⁺ →  Zn⁺⁺ + Cu     (redox)

Both the oxidation and reduction reaction take place simultaneously and separately. Due to oxidation, the zinc rod becomes electron-rich and due to the reduction copper rod becomes electron deficient. Thus there is a flow of electrons from electron-rich zinc rod to the electron-deficient copper rod through an external circuit which constitutes the electric current.

The movements of sulphate (SO₄^{– –}) ions from copper sulphate to zinc sulphate through the salt bridge or porous pot completes the Internal circuit and maintain electrical neutrality of the electrolyte.

The size of the zinc rod reduces because it dissolves in ZnSO₄ solution as zinc ions while the size of copper rod increases because of Cu⁺⁺ in the CuSO₄ solution discharge and copper deposits on the rod. The concentration of ZnSO₄ solution increases while the concentration of CuSO₄ solution decreases. The solutions in both the compartments remain electrically neutral.

E. M. F. of Daniel cell:

If the concentrations of zinc sulphate and copper sulphate solutions are  1 M each then the e.m.f. of Daniel cell is about 1.1 volt.

Reversibility of Daniel cell:

If an external opposing e.m.f. slightly greater than 1.1 V is applied then the reverse cell reaction takes place.

Zn⁺⁺ + Cu →  Zn   + Cu⁺⁺   (redox)

Now, reduction takes place at the zinc electrode and oxidation takes place at the copper electrode. Thus Daniel cell is a reversible cell.

Notes:

Explanation of why there is a gradual decrease in e.m.f. of Daniel cell

E^o_Cell = E^o_(ox/cathode) –   E_(ox/anode)

E^o_Cell =  E^o_(ox/Zn) –  E^o_(ox/Cu)

E^o_Cell   =      0.76    –  (- 0.34)

E^o_Cell  =  0.76    +   0.34

E^o_Cell   =      1.1 V

As the concentration of Zn⁺⁺ ions goes on increasing in zinc half cell, the oxidation potential of zinc electrode goes on decreasing. As the concentration of Cu⁺⁺ ions go on decreasing in the copper half cell, the reduction potential of copper electrode goes on decreasing.

Thus the quantity E^o_Cell =  E^o_(ox/Zn) + E^o_(red/Cu) decreases.

Daniel Cell with Porous Pot:

Daniel Cell with Copper Vessel as Cathode:

Daniel Cell with Salt Bridge:

Dry Cell:

The dry cell is actually a modification of the Leclanche cell. A dry cell is a primary cell because electrical energy is generated within the cell itself. It is an irreversible cell as cell reaction doesn’t reverse if higher external emf is applied to it. Since it doesn’t contain any liquid it is called the dry cell.  It is very convenient to carry.

Construction:

A dry cell consists of an outer zinc vessel (can) which works as an anode or negative electrode. In the centre of the zinc vessel, there is a stout graphite rod surrounded by compressed is an aqueous paste of MnO₂, carbon powder and ammonium chloride. This compressed solid is placed at the centre of the vessel.

The gap between zinc can and the compressed mixture is filled up with a jelly paste prepared from zinc chloride, ammonium chloride, and starch. Starch jelly helps to keep ammonium chloride moist for a long time. These Two different pastes are separated by means of a cotton partition.

The top of the can is sealed with sealing wax or resin or plastic to avoid drying of the paste. The whole-cell is covered with insulating safety cover.

Working:

On connecting carbon (+) and zinc (-) terminals through an external circuit, the cell starts working. When cell functions, oxidation takes place at zinc vessel and reduction takes place at graphite.

Reactions at negative electrode (Zn) :

Zn  → Zn⁺⁺  + 2 e^–   (oxidation)

Reactions at positive electrode (C) :

2 MnO₂ +   2 NH₄⁺ +   2e^–  → Mn₂O₃ + 2NH_3(g) + H₂O_(l)   (reduction)

Net Cell Reaction:

The overall reaction is,

Zn + 2 MnO₂ +   2 NH₄⁺    →  Zn⁺⁺ +Mn₂O₃ + 2NH_3(g) + H₂O_(l) 

Thus zinc is consumed continuously during working and holes are formed in the container which results in the leakage of the cell contents.

Side Reaction :

Zn⁺⁺ +  NH₃ →   [Zn(NH₃]⁺⁺

The EMF of Dry Cell: 

The EMF of this cell is 1.5 V.

Reversibility of Cell:

A dry cell is a non-reversible cell.

Uses of the dry cell.

• The dry cell is used in torches, calculators, and toys.
• Dry cells are used in flashlights.
• They are extensively used in many electronic appliances like clocks, transistor radios, tape recorders, calculators, etc.

Disadvantages of a Dry Cell:

• The acidic ammonium chloride corrodes the zinc container and thus the dry cell does not have a long time.
• Zinc is consumed continuously during working and holes are formed in the container which results in the leakage of the cell contents.
• This cell cannot be recharged by applying an external potential. This is because Zn++ ions and NH₃ molecule react each other to form a complex [Zn(NH₃)]⁺⁺ and they are not available for the reverse reaction. Therefore the dry cell is an irreversible cell.

Irreversibility of the dry cell:

During discharging of a dry cell following reaction takes place.

Zn  → Zn⁺⁺  + 2 e^–   (oxidation)

Thus zinc of vessel dissolves in the surrounding electrolyte forming zinc ions. These Zn⁺⁺ ions and NH₃ molecule react each other to form complex {Zn(NH₃)]⁺⁺

Thus if we apply an external e.m.f. zinc ions are not available for the reverse reaction. Hence instead of the reduction of zinc, the reduction of ammonium ion takes place as follows.

2 NH₄⁺  +2  e^–  →   2NH_3(g) +  H₂ (reduction)

Hence due to the absence of reverse reaction. The cell is irreversible.

Button Cell:

Construction:

• Anode can: Usually it contains zinc metal and electrolyte mixture (zinc powder + mercury). This forms the negative terminal of the cell.
• Cathode can: It contains the cathode material/electrolyte mixture (ZnO + KOH). This forms the positive terminal of the cell.
• Separator: It contains porous material with electrolyte.

Working:

On connecting cathode (+) and anode (-) terminals through an external circuit, the cell starts working.

Cell Reactions:

Reactions at anode:

Zn   + 2OH^–  →   ZnO_(s)  + H₂O_(l)   +  2 e^–          (oxidation)

Reactions at cathode:

HgO_(s) + H₂O_(l)  +  2 e^– → Hg_(l) + 2OH^–     (reduction)

Net Cell Reaction:

The overall reaction is,

Zn   + HgO_(s)    →  ZnO_(s) +  Hg_(l)

The EMF of cell:

The EMF of this cell is 1.35 V. Practically this e.m.f. remains constant because no ions are involved whose concentration can change during use.

Reversibility of Cell: 

The button cell is a non-reversible cell.

Uses of the button cell.

• It is small in size hence can be used in digital and electronic instruments like watches, computers, hearing aids, etc.
• The extremely constant voltage over its useful life. Suitable for low drain and intermittent high drain applications.
• It has a long shelf life – up to 3 years.

Disadvantages of the Button Cell:

• It contains mercury, which in certain forms is highly toxic to humans and animals.

Fuel cells:

Fuel cells are galvanic cells in which fuel, in the form of gases, which constitute electrode material are constantly supplied to the cell and electrochemical reaction takes place to produce electricity. The galvanic cells in which the energy of the combustion of fuels is directly converted into electrical energy are called fuel cells.

As it is a direct method to convert chemical energy into electrical energy, the cell has an efficiency of 70%. In another method by burning fossil fuel like coal, coke, petrol, diesel heat is generated, which is used to raise steam, which in turn used to rotate generators to produce electricity. This method has an efficiency of 40 % only.

Construction of Fuel Cell:

It consists of a cylindrical container of alkali-resistant material. It is divided into three compartments. The middle compartment contains a hot concentrated solution of KOH. Pure and dry hydrogen gas enters the anodic compartment from one side and pure dry oxygen enters the cathodic compartment from the other side.

The anode is a porous carbon electrode impregnated with platinum or palladium which acts as a catalyst and provides electrical contact. The cathode is also a porous carbon electrode impregnated with cobalt oxide, platinum or silver which acts as a catalyst and provides electrical contact.

Working of Fuel Cell:

Anode reaction: Hydrogen gas under high pressure diffuses through porous carbon anode

2 H_2(g) + 4OH^–(aq)→ 4H₂O_(l)  + 4e^–

Cathode reaction: The electron generated at anode reach cathode through external circuit and oxygen adsorbed on the surface of cathode get reduced to hydroxyl ions.

O_2(g) + H₂O_(l)  +  4e^– →  4OH^–

Thus the net reaction of the cell is

2 H_2(g) + O_2(g) →   2H₂O_(l)

Thus water is the by-product of the reaction. These cells do not pollute the atmosphere.

E.M.F. of a Cell:

E^o_Cell = E^o_(ox/cathode) –   E_(ox/anode)

E^o_Cell    =      0.83    +   0.40

E^o_Cell    =      1.23 V

Advantages of Fuel Cells:

• The reacting substances are continuously supplied to the electrodes. Hence unlike conventional cells, the fuel cells do not have to be discharged when the chemicals are consumed.
• Thus water is the by-product of the reaction. These cells do not pollute the atmosphere.
• As it is a direct method to convert chemical energy into electrical energy, the cell has an efficiency of 70%.

Drawbacks of Fuel Cells:

• The calculated cell voltage is 1.23 V. Due to the reversibility of the reaction the cell voltage is less than 1.23 V.
• Hydrogen gas is hazardous and its production is expensive.

Uses Fuel Cells:

• These cells have been used in automobiles on an experimental basis.
• These cells are used in space shuttle and space stations to produce electricity and water for consumption for onboard astronauts.
• They may be used in hospitals, hotels, and homes for the production of electricity.

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