Progress paradigm ‐ 1
1 Define Electrochemistry.
2 Define Electrochemical Cells.
3 What is a galvanic cell?
4 Explain redox reaction in a galvanic cell.
5 Explain Daniell cell,.
Progress paradigm ‐ 2
6. Fill in the blanks ‐
i. In a Daniel cell The reduction half reaction occurs on the .............. electrode while
ii. the oxidation half reaction occurs on the ..............electrode. These two portions of the cell are also called half-cells or ..............couples. can construct innumerable number of galvanic cells on the pattern.
iii. The electrolytes of the two half-cells are connected
internally through a ............... According to IUPAC convention, standard reduction potentials are now called ...............
iv. In a galvanic cell, the half-cell in which oxidation takes place is called anode
and it has a ..............potential with respect to the solution. The potential difference between the two electrodes of a galvanic
cell is called the cell potential and is measured in ..............
v. Convention is that we keep the .............. on the left and the .............. on the right while representing the galvanic cell.
vi . A galvanic cell is generally represented by putting a vertical line between .............. and .............. solution and putting a double vertical line between
the two..............connected by a salt bridge.
7 Define electrode potential.
8 Define electromotive force (emf) of the cell
9 What is a Standard Hydrogen Electrode?
10 The pressure of the hydrogen gas present in this half cell equals ...............
11 Platinum is used in the Standard Hydrogen Electrode due to the following reasons.............. A standard hydrogen electrode has five components: ..............
12 What is the use of a standard hydrogen electrode?
13 Why is a platinum electrode used?
14 Which foil is used in standard hydrogen electrode?
15 Explain The Core Process of Redox Reactions.
(Eyref 17) Nernst Equation
(Eyref 18) Nernst equation for a Nickel- Silver Cell
(Eyref 19) Nernst Equation for a General Electrochemical reaction
(Eyref 20) Nernst Equation - Applications
(Eyref 21) Nernst Equation - Limitations
Progress paradigm ‐ 3
16. Explain Nernst equation.
17. Derive Nernst equation.
18. Write Applications of the Nernst Equation
19 Aqueous copper sulphate solution and aqueous silver nitrate solution are electrolysed by 1 ampere current for 10 minutes in separate electrolytic cells.
Will the mass of copper and silver deposited on the cathode be same or different? Explain your answer.
20. Depict the galvanic cell in which the cell reaction is
Progress paradigm ‐ 4
23 Can E0cell or ΔrG0 for a cell reaction ever be equal to zero?
24 Under what conditions is E0cell = 0 and ΔrG0 = 0 ?
25 What does the negative sign in the expression E0Zn2+ /Zn = – 0.76 V means?
25. Explain Faraday’s First law of electrolysis
26 Consider a cell given below
Cu |Cu²+ || Cl- | Cl₂, Pt
Write the reactions that occur at anode and cathode.
27. Explain Faraday’s Second law of electrolysis.
28. Depict the galvanic cell in which the cell reaction is : Cu + 2Ag+→ 2Ag + Cu2+.
29. Give a few Application of electrolysis.
30. Write the Nernst equation for the cell reaction in the Daniell cell. How will the Ecell be affected when concentration of Zn2+ ions is increased ?
31. What does the negative sign in the expressionEoZn2+/Zn = -0.76 V mean.
32. Under what condition is ECell = 0 and ΔrG = 0?
Electrochemistry
SVTs
MCQs
Electrochemistry is the branch of chemistry which deals with the relationship between electrical energy and chemical energy and inter-conversion of one form into another.
Electrochemical Cells and types
An electrochemical cell consists of two metallic electrodes dipped in electrolytic solutions. The cells are of two types:
(a) Electrolytic cells
(b) Galvanic cells
3. A galvanic cell consists of two half cells. Each half cell contains an electrolytic solution and a metallic electrode.The electrode at which- oxidation takes place is called an anode and the electrode at which reduction takes place is called the cathode. The half-cells are separated from each other by means of a porous pot or a salt bridge.
If an external opposite potential is applied in the galvanic cell (Fig a) and increased slowly, we find that the reaction continues to take place till the opposing voltage reaches the value 1.1 V (Fig.) when, the reaction stops altogether and no current flows through the cell. Any further increase in the external potential again starts the reaction but in the opposite direction (Fig.). It now functions as an electrolytic cell, a device for using electrical energy to carry non-spontaneous chemical reactions.
a. When Eext < 1.1 V
(i) Electrons flow from Zn rod to Cu rod hence current flows from Cu to Zn.
(ii) Zn dissolves at anode and copper deposits at cathode.
b. When Eext = 1.1 V
(i) No flow of electrons or current.
(ii) No chemical reaction.
When Eext >1.1 V
(i) Electrons flow from Cu to Zn and current flows from Zn to Cu.
(ii) Zinc is deposited at the zinc electrode and copper dissolves at copper electrode.
Galvanic Cells
A galvanic cell is an electrochemical cell that converts the chemical energy of a spontaneous redox reaction into electrical energy. In this device the Gibbs energy of the spontaneous redox reaction is converted into electrical work which
may be used for running a motor or other electrical gadgets like heater,
fan, geyser, etc.
Daniell cell discussed earlier is one such cell in which the following redox reaction occurs.
Zn(s) + Cu2+ (aq) →Zn2+ (aq) + Cu(s)
This reaction is a combination of two half reactions whose addition
gives the overall cell reaction:
(i) Cu2++ 2e- → Cu(s) (reduction half reaction)
(ii) Zn(s) →Zn2++ 2e- (oxidation half reaction)
These reactions occur in two different portions of the Daniell cell.
i. The reduction half reaction occurs on the copper electrode while
ii. the oxidation half reaction occurs on the zinc electrode.
These two portions of the cell are also called half-cells or redox couples.
The copper electrode may be called the reduction half cell and the zinc electrode, the oxidation half-cell.
We can construct innumerable number of galvanic cells on the pattern
of Daniell cell by taking combinations of different half-cells.
Each half cell consists of a metallic electrode dipped into an electrolyte. The two half-cells are connected by a metallic wire through a voltmeter and a
switch externally.
The electrolytes of the two half-cells are connected
internally through a salt bridge.
Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
do not require a salt bridge.
At each electrode-electrolyte interface there is a tendency of metal
ions from the solution to deposit on the metal electrode trying to make
it positively charged.
At the same time, metal atoms of the electrode
have a tendency to go into the solution as ions and leave behind the
electrons at the electrode trying to make it negatively charged.
At equilibrium, there is a separation of charges and depending on the
tendencies of the two opposing reactions, the electrode may be positively
or negatively charged with respect to the solution. A potential difference
develops between the electrode and the electrolyte which is called
electrode potential.
When the concentrations of all the species involved
in a half-cell is unity then the electrode potential is known as standard
electrode potential.
According to IUPAC convention, standard reduction potentials are now called standard electrode potentials.
In a
galvanic cell, the half-cell in which oxidation takes place is called anode
and it has a negative potential with respect to the solution.
The other half-cell in which reduction takes place is called cathode and it has a
positive potential with respect to the solution.
Thus, there exists a potential difference between the two electrodes and as soon as the switch is in the on position the electrons flow from negative electrode to positive electrode. The direction of current flow is opposite to that of electron flow.
The potential difference between the two electrodes of a galvanic
cell is called the cell potential and is measured in volts.
The cell potential is the difference between the electrode potentials (reduction potentials) of the cathode and anode. It is called the cell electromotive force (emf) of the cell when no current is drawn through the cell.
It is now an accepted convention that we keep the anode on the left and the cathode on the right while representing the galvanic cell.
A galvanic cell is generally represented by putting a vertical line between metal and electrolyte solution and putting a double vertical line between
the two electrolytes connected by a salt bridge.
Under this convention
the emf of the cell is positive and is given by the potential of the half cell
on the right hand side minus the potential of the half-cell on the
left hand side i.e.,
Ecell = Eright – Eleft
This is illustrated by the following example:
Cell reaction:
Cu(s) + 2Ag+ (aq) → Cu2+ (aq) + 2 Ag(s)
Half-cell reactions:
Cathode (reduction): 2Ag+(aq) + 2e- → 2Ag(s)
Anode (oxidation): Cu(s) → Cu2+(aq) + 2e-
It can be seen that the sum of overall reaction
in the cell and that silver electrode acts as a cathode and copper
electrode acts as an anode. The cell can be represented as:
Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s)
and we have Ecell= Eright – Eleft = EAg+|Ag – ECu2+|Cu
⇉
In a Daniell cell electrons flow from zinc electrode to copper electrode through an external circuit, while metal ions form one half cell to the other through the salt bridge.
Here current flows from copper electrode to zinc electrode that is cathode to anode via an external circuit.
Daniell cell is a reversible cell while a voltaic cell may be reversible or irreversible.
Measurement of Electrode Potential
What is a Standard Hydrogen Electrode?
The Standard Hydrogen Electrode is often abbreviated to SHE, and its standard electrode potential is declared to be 0 at a temperature of 298K. This is because it acts as a reference for comparison with any other electrode.
The half cell reaction of SHE can be written as follows:
2H+ (aq) + 2e– → H2 (g)
The reaction given above generally takes place on a platinum electrode. The pressure of the hydrogen gas present in this half cell equals 1 bar.
Platinum is used in the Standard Hydrogen Electrode due to the following reasons:
Nernst equation-
The Nernst equation provides the relation between the cell potential of an electrochemical cell, the standard cell potential, temperature, and the reaction quotient.
With the help of the Nernst equation the cell potentials of electrochemical cells can be determined even under non-standard conditions.
The Nernst equation is often used to calculate the cell potential of an electrochemical cell at any given temperature, pressure, and reactant concentration. The equation was introduced by a German chemist, Walther Hermann Nernst.
What is the reaction quotient Q?
The reaction quotient (Q) measures the relative amounts of products and reactants present during a reaction at a particular point in time. The reaction quotient aids in figuring out which direction a reaction is likely to proceed, given either the pressures or the concentrations of the reactants and the products.
Derivation Fig.
Nernst Equation Applications
a. Single electrode reduction or oxidation potential at any conditions
b. Standard electrode potentials
c. Comparing the relative ability as a reductive or oxidative agent
d. Finding the feasibility of combining such single electrodes to produce an electric potential
e. Emf of an electrochemical cell
f. Unknown ionic concentrations
g. The pH of solutions and solubility of sparingly soluble salts can be measured with the help of the Nernst equation.
Limitations of Nernst Equation
1. Because the activity of an ion in a very dilute solution approaches infinity, it can be defined in terms of ion concentration. But even so, in very high concentration solutions, the ion concentration does not equal the ion activity. In order to use the Nernst equation in such cases, experimental measurements must be conducted to obtain the true activity of the ion.
2. It cannot be used to measure cell potential when a current flows through the electrode. This is because the current flow affects the ions’ activity on the surface of the electrode. Additional factors such as resistive loss and overpotential must be considered when a current flows through the electrode.
Measurement of the Conductivity of Ionic Solutions-
Accurate measurement of an unknown resistance can be performed on a Wheatstone bridge.
However, for measuring the resistance
of an ionic solution there are two problems.
a. Firstly, passing direct current (DC) changes the composition of the solution. Secondly, a solution cannot be connected to the bridge like a metallic wire or other solid conductor.
This difficulty is resolved by using an alternating current (AC) source
of power.
b. The second problem is solved by using a specially designed vessel called conductivity cell. It is available in several designs.
Wheatstone bridge -
Wheatstone bridge, also known as the resistance bridge, calculates the unknown resistance by balancing two legs of the bridge circuit. One leg includes the component of unknown resistance.
The Wheatstone Bridge Circuit comprises two known resistors, one unknown resistor and one variable resistor connected in the form of a bridge. This bridge is very reliable as it gives accurate measurements.
Construction of Wheatstone Bridge
A Wheatstone bridge circuit consists of four arms, of which two arms consist of known resistances while the other two arms consist of an unknown resistance and a variable resistance. The circuit also consists of a galvanometer and an electromotive force source.
The Wheatstone bridge is
fed by an oscillator O (a source of a.c. power in the
audio frequency range 550 to 5000 cycles per
second).
P is a suitable detector (a headphone or
other electronic device) and the bridge is balanced
when no current passes through the detector.
Wheatstone Bridge Application
The Wheatstone bridge is used for the precise measurement of low resistance.
Wheatstone bridge and an operational amplifier are used to measure physical parameters such as temperature, light, and strain.
Quantities such as impedance, inductance, and capacitance can be measured using variations on the Wheatstone bridge.
Limitations of Wheatstone Bridge
For low resistance measurement, the resistance of the leads and contacts becomes significant and introduces an error.
For high resistance measurement, the measurement presented by the bridge is so large that the galvanometer is insensitive to imbalance.
The other drawback is the resistance change due to the current’s heating effect through the resistance. Excessive current may even cause a permanent change in the value of resistance.
What is Wheatstone Bridge?
Wheatstone bridge, also known as the resistance bridge, calculates the unknown resistance by balancing two legs of the bridge circuit. One leg includes the component of unknown resistance.
What is the Wheatstone Bridge Principle?
The Wheatstone bridge works on the principle of null deflection, i.e. the ratio of their resistances are equal and no current flows through the circuit.
When is the Wheatstone Bridge balanced?
A Wheatstone bridge is said to be in a balanced condition when no current flows through the galvanometer. This condition can be achieved by adjusting the known resistance and variable resistance.
When is the Wheatstone Bridge said to be unbalanced?
Under normal conditions, the bridge is in an unbalanced condition where current flows through the galvanometer.
What are the limitations of Wheatstone Bridge?
For low resistance measurement, the resistance of the leads and contacts becomes significant and introduces an error.
Variation of Conductivity and Molar Conductivity with Concentration
Both conductivity and molar conductivity change with the
concentration of the electrolyte. Conductivity always decreases with
decrease in concentration both, for weak and strong electrolytes.
This can be explained by the fact that the number of ions per unit
volume that carry the current in a solution decreases on dilution.
The conductivity of a solution at any given concentration is the
conductance of one unit volume of solution kept between two
platinum electrodes with unit area of cross section and at a distance
of unit length. This is clear from the equation:
G = k A/l = kappa when A and l are unity as m or cm.
Molar conductivity of a solution at a given concentration is the
conductance of the volume V of solution containing one mole of
electrolyte kept between two electrodes with area of cross section A and
distance of unit length. Therefore,
When concentration
approaches zero, the molar conductivity
is known as limiting molar
conductivity.
What is Kohlrausch’s Law?
Kohlrausch’s law states that the equivalent conductivity of an electrolyte at infinite dilution is equal to the sum of the conductances of the anions and cations.
The molar conductivity of a solution at a given concentration is the conductance of the volume of solution containing one mole of electrolyte kept between two electrodes with the unit area of cross-section and distance of unit length. The molar conductivity of a solution increases with the decrease in concentration. This increase in molar conductivity is because of the increase in the total volume containing one mole of the electrolyte. When the concentration of the electrolyte approaches zero, the molar conductivity is known as limiting molar conductivity.
Kohlrausch observed certain regularities while comparing the values of limiting molar conductivities of some strong electrolytes. On the basis of his observations, Kohlrausch proposed “limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anions and cations of the electrolyte”. This law is popularly known as Kohlrausch law of independent migration of ions. For example, limiting molar conductivity, Λ of sodium chloride can be determined with the knowledge of limiting molar conductivities of sodium ion and chloride ion.
Who discovered the law of independent migration of ions? This law was discovered by Friedrich Kohlrausch after observing experimental data on conductivities of various electrolytes.
What is kohlrausch law and its applications?
It’s used to determine an electrolyte’s dissociation constant. It’s used to find out what a weak electrolyte’s limiting molar conductivity is. This law can also be used to determine the degrees of dissociation of weak electrolytes.
What is Kohlrausch law of independent migration?
According to Kohlrausch’s law of independent ion movement, the limiting molar conductivity of an electrolyte can be described as the sum individual contributions of its cations and anions. Since fewer ions are present for conduction, a solution’s conductivity decreases with dilution.
Why do we need Kohlrausch law?
The Kohlrausch law can be used to calculate the limiting molar conductivities of any electrolyte. At larger concentrations, weak electrolytes have lower molar conductivities and a lower degree of dissociation.
What is infinite dilution in electrochemistry?
Infinite dilution is a state of dilution in which the concentration does not change when more solvent is added. In chemistry, the concept of infinite dilution is used to investigate how compounds dissolve in different solvents.
Electrolytic Cells and Electrolysis
In an electrolytic cell external source of voltage is used to bring about a chemical reaction.
One of the simplest electrolytic cell consists of two copper strips dipping in an aqueous solution of copper sulphate.
Quantitative Aspects of Electrolysis
Michael Faraday was the first scientist who described the quantitative
aspects of electrolysis.
Faraday’s Laws of Electrolysis
After his extensive investigations on electrolysis of solutions and melts
of electrolytes, Faraday published his results during 1833-34 in the
form of the following well known Faraday’s two laws of electrolysis:
(i) First Law:
The amount of chemical reaction which occurs at any
electrode during electrolysis by a current is proportional to the
quantity of electricity passed through the electrolyte (solution or
melt).
(ii) Second Law:
The amounts of different substances liberated by the
same quantity of electricity passing through the electrolytic solution
are proportional to their chemical equivalent weights (Atomic Mass
of Metal ÷ Number of electrons required to reduce the cation).
There were no constant current sources available during Faraday’s times.
The general practice was to put a coulometer (a standard electrolytic cell) for determining the quantity of electricity passed from the amount of metal (generally silver or copper) deposited or consumed.
Coulometers are now obsolete and we now have constant current (I) sources available and the quantity of electricity Q, passed is given by
Q = It
Q is in coulombs when I is in ampere and t is in second
The amount of electricity (or charge) required for oxidation or reduction depends on the stoichiometry of the electrode reaction.
Charge on one electron is equal to 1.6021 x 10-19 C.
Charge on one mole of electrons is equal to NAx1.6021x 10-19 C
= 6.02 x 1023 mol-1 x 1.6021 x 10-19 = 96487 C mol-1
For approximate calculations we use 1F = 96500 C mol-1.
Products ofElectrolysis
Products of electrolysis depend on the nature of material being
electrolysed and the type of electrodes being used. If the electrode is
inert (e.g., platinum or gold), it does not participate in the chemical
reaction and acts only as source or sink for electrons.
On the other
hand, if the electrode is reactive, it participates in the electrode reaction.
Thus, the products of electrolysis may be different for reactive and inert electrodes.The products of electrolysis depend on the different oxidising
and reducing species present in the electrolytic cell and their standard
electrode potentials.
Moreover, some of the electrochemical processes
although feasible, are so slow kinetically that at lower voltages these do
not seem to take place and extra potential (called overpotential) has to
be applied, which makes such process more difficult to occur.
Batteries
A Battery is a device consisting of one or more electrical cells that convert chemical energy into electrical energy. Every battery is basically a galvanic cell where redox reactions take place between two electrodes which act as the source of the chemical energy.
Battery types
Primary Cell / Primary battery
Secondary Cell / Secondary battery
Based on the application of the battery, they can be classified again. They are:
Household Batteries
These are the types of batteries which are more likely to be known to the common man. They find uses in a wide range of household appliances (such as torches, clocks, and cameras). These batteries can be further classified into two subcategories:
Rechargeable batteries Nickel
Examples: Cadmium batteries, Lithium-Ion
Non-rechargeable batteries
Examples: Silver oxide, Alkaline & carbon zinc
Industrial Batteries
These batteries are built to serve heavy-duty requirements. Some of their applications include railroad, backup power and more for big companies. Some examples are:
Nickel Iron
Wet Nickel Cadmium (NiCd)
Vehicle Batteries
These are more user-friendly and a less complicated version of the industrial batteries. They are specifically designed to power cars, motorcycles, boats & other vehicles. An important example of a vehicle battery is the Lead-acid battery.
Primary Cell
In the primary batteries, the reaction occurs only once and after use
over a period of time battery becomes dead and cannot be reused
again.
The most familiar example of this type is the dry
cell (known as Leclanche cell after its discoverer) which is
used commonly in our transistors and clocks.
The cell
consists of a zinc container that also acts as anode and
the cathode is a carbon (graphite) rod surrounded by
powdered manganese dioxide and carbon (Fig.3.8). The
space between the electrodes is filled by a moist paste of
ammonium chloride (NH4Cl) and zinc chloride (ZnCl2).
Structure of Cell
The structure of the zinc-carbon dry cell is shown in the figure. It consists of the anode terminal as zinc or in general graphite rod. The carbon forms the cathode terminal. It may be observed that in older versions of dry cell the zinc was used as cathode and graphite was used as anode terminal. The selection of the elements is fundamentally based on its chemical configuration of the outermost orbit of the elements.
If it has more number electrons in the outermost orbit, then it can act as a donor, and hence forms the cathode. Similarly, if the outermost orbit has fewer electrons, it can easily accept and hence forms the anode. The electrolyte placed in between acts as a catalyst for the chemical reactions. In general, we use ammonium chloride jelly as the electrolyte. In the figure shown, the electrolyte used is a mixture of zinc and chloride. Also, sodium chloride is also used as an electrolyte. A mixture of manganese dioxide and carbon is surrounded around the anode rod.
The whole configuration is placed in a metal tube. The jelly is prevented from drying up by using a pitch at the top of the cell. A carbon washer is placed at the bottom. The purpose of this washer is to prevent the zinc anode rod from coming in contact with the container.
This is also called a spacer as shown in the diagram. The zinc can is also surrounded by paper insulation for insulation purposes. For large batteries, other insulating materials such as mica, etc. are also used. The positive terminal of the ell is formed at the top. The negative terminal of the cell is formed at the base.
Working of Dry Cell
A dry cell fundamentally works on chemical reactions. Due to the reactions that take place between the electrolyte and the electrodes, the electrons flow from one electrode to the other. Substances such as acids dissolve in water to form ionized particles. The ionized particle is of two types. The positive ions are called cations and the negative ions are called anions. The acids which are dissolved in water are called electrolytes.
In the above-mentioned diagram, the zinc chloride forms as the electrolyte. Similarly ammonium chloride jelly also forms as an electrolyte. The metal rods immersed in electrolytes form electrodes. Based on the chemical characteristics of the metal rods, we have a positive electrode as anode and a negative electrode as the cathode.
The electrodes attract the oppositely charged ions to their side. For example, the cathode attracts the anions and the anode attracts the cations. In this process the electrons flow from one direction to the other, hence we get a flow of charges. This is called current.
Chemical Reactions
The reactions taking place in the cell is shown below. First is the oxidation reaction.
In this, the zinc cathode is oxidized to positively charged zinc ions releasing two ions. These electrons are collected by the anode. Then comes the reduction reaction.
The reduction reaction at the anode is shown above. This reaction produces an electric current. It releases oxide ions with magnesium oxide. This reaction forms when magnesium is combined with the electrolyte.
The other two reactions represent an acid-base reaction and precipitation reaction taking place in the dry cell. In the acid-base reaction, NH is combined with OH to produce NH3 along with water. The outcomes are NH3 and water base.
Difference Between a Dry Cell and Wet Cell
The main difference between the dry cell and the wet cell is the form of electrolyte. As discussed before, in a dry cell, the electrolyte such as ammonium chloride is dry in nature. Such dry cells are more common and used in toys, radios, etc. But in a wet cell, the electrolyte is in the liquid state.
Liquid electrolytes such as sulfuric acid, which is a dangerous corrosive liquid is used. Due to the nature of such liquids, the wet cell is more explosive in nature and needs to be handled with care. The best advantage of such wet cells are they can b easily recharged and used for numerous applications. Such batteries find common usage in aviation, utilities, energy storage, and cell phone towers.
Dry Cell Functions
The dry cell function based on the chemical reactions between the electrode and the electrolytes. When the electrodes are placed in the electrolytes, it attracts the oppositely charged ions towards themselves. This causes the flow of charges, and hence current is produced.
Advantages
The advantages of the dry cell include the following.
The dry cell has numerous advantages such as
It is small in size.
It can come in a variety of voltage levels.
It is handy and has numerous applications.
It is the only source of DC voltage.
It can be used along with power electronic circuits to regulate the output voltage
It is rechargeable.
Disadvantages
The disadvantages of the dry cell include the following.
It must be handled with care
It is explosive
Large rating batteries are very heavy
Applications
The applications of the dry cell include the following.
Toys
Aviation
Cell phones
Radio
Calculator
Watches
Hearing Aids
Hence we have seen the operation, classification, and applications of the dry cells. One interesting point to be noted is the battery works only when the electrodes are physically in touch with each other. There must exist a conducting medium between the two electrodes. The question is can water be used as a conducting medium between the electrodes of the dry cell? In that case, what will happen if this cell is dipped in water?
Mercury cell
(Fig.)
suitable for low current devices
like hearing aids, watches, etc. consists of zinc – mercury
amalgam as anode and a paste of HgO and carbon as the
cathode. The electrolyte is a paste of KOH and ZnO. The
electrode reactions for the cell are given below:
Types of Mercury Cells
There are two varieties of mercury cells. One is a zinc-mercuric oxide cell, and the other is a cadmium-mercuric oxide cell.
Advantage of Mercury Cell over Dry Cell
1. Long shelf time of up to 10 years
2. High capacity per size
3. The constant voltage output of 1.35V
4. Inexpensive to produce mercury cells with known technology
Limitations
Environment issues and economic issues – Inhalation of mercury vapour is harmful to the human body, including organs like the kidney, nervous system, digestive system, eye, skin and immunity systems. Even a small amount is very toxic to the human body. Dangerous to the development of children during pregnancy and early childhood.
Why does the cell potential of mercury remain constant?
Since mercury ion doesn’t involve any ions whose concentration changes, it has a constant potential of 1.35V.
Is mercury cell rechargeable?
No, the mercury cell is a primary cell, which is not rechargeable.
What is the oxidising agent in a mercury cell?
HgO serves as an oxidising agent, and the reaction takes place at the cathode.
How much mercury will be present in a mercury cell battery?
Button cell batteries can contain 5 mg of mercury in a single unit.
What are the cathode and the anode in a mercury cell?
A mercury cell consists of a zinc anode and mercuric oxide cathode.
Is a mercury cell a primary cell or a secondary cell?
A mercury cell is a primary cell. It is non-rechargeable and non-usable.
What is the electrolyte used in a mercury cell?
Potassium or sodium hydroxide is used as an electrolyte in this cell.
What are the two types of mercury cells?
Cadmium-mercury cells and zinc-mercury cells are the two types of mercury cells.
What is the voltage output of a mercury cell?
1.35V is the voltage output of a mercury cell.
Secondary Batteries
A secondary cell after use can be recharged by passing current
through it in the opposite direction so that it can be used again. A
good secondary cell can undergo a large number of discharging
and charging cycles. The most important secondary cell is the lead
storage battery (Fig. ) commonly used in automobiles and
invertors. It consists of a lead anode and a grid of lead packed with
lead dioxide (PbO2 ) as cathode. A 38% solution of sulphuric acid
is used as an electrolyte.
Another important secondary
cell is the nickel-cadmium cell
(Fig.) which has longer life
than the lead storage cell but
more expensive to manufacture.
We shall not go into details of
working of the cell and the
electrode reactions during
charging and discharging.
The overall reaction during
discharge is:
Fuel cell
A fuel cell is a device that converts chemical potential energy (energy stored in molecular bonds) into electrical energy. A fuel cell is a lot like a battery. It has two electrodes where the reactions take place and an electrolyte which carries the charged particles from one electrode to the other.
Fuel cells require a continuous input of fuel and an oxidizing agent (generally oxygen) in order to sustain the reactions that generate the electricity. Therefore, these cells can constantly generate electricity until the supply of fuel and oxygen is cut off.
Despite being invented in the year 1838, fuel cells began commercial use only a century later when they were used by NASA to power space capsules and satellites. Today, these devices are used as the primary or secondary source of power for many facilities including industries, commercial buildings, and residential buildings.
A fuel cell is similar to electrochemical cells, which consists of a cathode, an anode, and an electrolyte. In these cells, the electrolyte enables the movement of the protons.
Working of Fuel Cell
The reaction between hydrogen and oxygen can be used to generate electricity via a fuel cell. Such a cell was used in the Apollo space programme and it served two different purposes – It was used as a fuel source as well as a source of drinking water (the water vapour produced from the cell, when condensed, was fit for human consumption).
Types of Fuel Cells
Despite working similarly, there exist many varieties of fuel cells. Some of these types of fuel cells are discussed in this subsection.
The Polymer Electrolyte Membrane (PEM) Fuel Cell
These cells are also known as proton exchange membrane fuel cells (or PEMFCs).
The temperature range that these cells operate in is between 50oC to 100oC
The electrolyte used in PEMFCs is a polymer which has the ability to conduct protons.
A typical PEM fuel cell consists of bipolar plates, a catalyst, electrodes, and the polymer membrane.
Despite having eco-friendly applications in transportation, PEMFCs can also be used for the stationary and portable generation of power.
Phosphoric Acid Fuel Cell
These fuel cells involve the use of phosphoric acid as an electrolyte in order to channel the H+
The working temperatures of these cells lie in the range of 150oC – 200oC
Electrons are forced to travel to the cathode via an external circuit because of the non-conductive nature of phosphoric acid.
Due to the acidic nature of the electrolyte, the components of these cells tend to corrode or oxidize over time.
Solid Acid Fuel Cell
A solid acid material is used as the electrolyte in these fuel cells.
The molecular structures of these solid acids are ordered at low temperatures.
At higher temperatures, a phase transition can occur which leads to a huge increase in conductivity.
Examples of solid acids include CsHSO4 and CsH2PO4 (cesium hydrogen sulfate and cesium dihydrogen phosphate respectively)
Alkaline Fuel Cell
This was the fuel cell which was used as the primary source of electricity in the Apollo space program.
In these cells, an aqueous alkaline solution is used to saturate a porous matrix, which is in turn used to separate the electrodes.
The operating temperatures of these cells are quite low (approximately 90oC).
These cells are highly efficient. They also produce heat and water along with electricity.
Solid Oxide Fuel Cell
These cells involve the use of a solid oxide or a ceramic electrolyte (such as yttria-stabilized zirconia).
These fuel cells are highly efficient and have a relatively low cost (theoretical efficiency can even approach 85%).
The operating temperatures of these cells are very high (lower limit of 600oC, standard operating temperatures lie between 800 and 1000oC).
Solid oxide fuel cells are limited to stationary applications due to their high operating temperatures.
Molten Carbonate Fuel Cell
The electrolyte used in these cells is lithium potassium carbonate salt. This salt becomes liquid at high temperatures, enabling the movement of carbonate ions.
Similar to SOFCs, these fuel cells also have a relatively high operating temperature of 650oC
The anode and the cathode of this cell are vulnerable to corrosion due to the high operating temperature and the presence of the carbonate electrolyte.
These cells can be powered by carbon-based fuels such as natural gas and biogas.
Applications of fuel cell
Fuel cell technology has a wide range of applications. Currently, heavy research is being conducted in order to manufacture a cost-efficient automobile which is powered by a fuel cell. A few applications of this technology are listed below.
Fuel cell electric vehicles, or FCEVs, use clean fuels and are therefore more eco-friendly than internal combustion engine-based vehicles.
They have been used to power many space expeditions including the Appolo space program.
Generally, the byproducts produced from these cells are heat and water.
The portability of some fuel cells is extremely useful in some military applications.
These electrochemical cells can also be used to power several electronic devices.
Fuel cells are also used as primary or backup sources of electricity in many remote areas.
Frequently Asked Questions – FAQs
1. What is a fuel cell?
A fuel cell is an electrochemical cell that generates electrical energy from fuel via an electrochemical reaction. It offers high efficiency and zero emissions.
2. How does a fuel cell differ from conventional methods of energy generation?
A fuel cell is different from the conventional methods of energy generation because, in a fuel cell, chemical energy is directly converted into electrical energy without intermediate conversion into mechanical power.
3. Why is fuel cell better than the conventional methods of energy generation?
A fuel cell is preferred over conventional methods of energy generation because, in a fuel cell, zero combustion takes place. Thus, carbon dioxide is not produced.
4. What are the benefits of a fuel cell?
Fuel cells provide clean energy and emit no pollution. Moreover, it also offers high efficiency and zero emissions. No carbon dioxide is produced while generating chemical energy from a fuel cell.
5. Which electrolyte is used in molten carbonate fuel cells?
Lithium potassium carbonate salt is used as an electrolyte in molten carbonate fuel cells.
Thus, the different types of fuel cells and the working of an alkaline fuel cell are briefly discussed in this article along with some applications of these electrochemical cells.
The working of this fuel cell involved the passing of hydrogen and oxygen into a concentrated solution of sodium hydroxide via carbon electrodes. The cell reaction can be written as follows:
Corrosion
Corrosion: Corrosion is the process of slowly eating up metals by gas and water vapours present in the atmosphere due to the formation of certain compounds like oxide, sulphides, carbonate, etc.
Rust:
Corrosion of iron is known as rusting.
When Iron comes in contact with oxygen in presence of moisture(Water), a reddish-brown coating is formed on the surface of Iron which is called rust.
Prevention of Rusting
Barrier Protection - In this process, the metal is not allowed to come in contact with air, moisture, Oxygen and Carbon dioxide.
The methods used are
i) The metal surface is coated with paint.
ii) A film of oil or grease on the surface of iron tools and machinery.
iii) Coating the iron surface with non-corroding metals like Nickel, Chromium, Aluminium, etc., by the process of electroplating.
iv) Coating the iron surface with Phosphate or other chemicals gives protection from air and moisture.
Sacrificial protection: It involves covering the iron surface with a layer of metal that is more active i.e., electropositive than Iron. So, the more active metal will lose electrons, not Iron. When zinc is coated on the iron surface the process is called galvanization.
Using anti-rust solutions: Alkaline phosphate and alkaline chromate solutions due to their alkaline nature prevent the availability of H+ ions.
Phosphates provide a thick film of insoluble Iron phosphate which protects the Iron objects from rusting.
Corrosion is one of the most common phenomena that we observe in our daily lives. You must have noticed that some objects made of iron are covered with an orange or reddish-brown coloured layer at some point in time. The formation of this layer is the result of a chemical process known as rusting, which is a form of corrosion.
Cossorion, in general, is a process through which refined metals are converted into more stable compounds such as metal oxides, metal sulfides, or metal hydroxides. Likewise, the rusting of iron involves the formation of iron oxides via the action of atmospheric moisture and oxygen. If we look at the science behind corrosion, then we can say that it is a spontaneous/irreversible process wherein the metals turn into a much more stable chemical compound like oxides, sulphides, hydroxides, etc. We will delve deeper into the concept of corrosion and understand its different factors, including its meaning, types, prevention and more in this lesson.
Corrosion Definition
What is Corrosion? It is basically defined as a natural process that causes the transformation of pure metals into undesirable substances when they react with substances like water or air. This reaction causes damage and disintegration of the metal, starting from the portion of the metal exposed to the environment and spreading to the entire bulk of the metal.
Corrosion is usually an undesirable phenomenon since it negatively affects the desirable properties of the metal. For example, iron is known to have good tensile strength and rigidity (especially alloyed with a few other elements). However, when subjected to rusting, iron objects become brittle, flaky, and structurally unsound. On the other hand, corrosion is a diffusion-controlled process, and it mostly occurs on exposed surfaces. Therefore, in some cases, attempts are made to reduce the activity of the exposed surface and increase a material’s corrosion resistance. Processes such as passivation and chromate conversion are used for this purpose. However, some corrosion mechanisms are not always visible, and they are even less predictable.
On the other hand, corrosion can be classified as an electrochemical process since it usually involves redox reactions between the metal and certain atmospheric agents such as water, oxygen, sulphur dioxide, etc.
Do All Metals Corrode?
Metals placed higher in the reactivity series, such as iron, zinc, etc., get corroded very easily, and metals placed lower in the reactivity series, like gold, platinum and palladium, do not corrode. The explanation lies in the fact that corrosion involves the oxidation of metals. As we go down, the reactivity series tendency to get oxidised is very low (oxidation potentials are very low).
Interestingly, aluminium doesn’t corrode, unlike other metals, even though it is reactive. This is because aluminium is covered by a layer of aluminium oxide already. This layer of aluminium oxide protects it from further corrosion.
Factors Affecting Corrosion
1. Exposure of the metals to air containing gases like CO2, SO2, SO3 etc.
2. Exposure of metals to moisture, especially salt water (which increases the rate of corrosion).
3. Presence of impurities like salt (For example, NaCl).
4. Temperature: An increase in temperature increases corrosion.
5. Nature of the first layer of oxide formed: Some oxides like Al2O3 form an insoluble protecting layer that can prevent further corrosion. Others, like rust, easily crumble and expose the rest of the metal.
6. Presence of acid in the atmosphere: Acids can easily accelerate the process of corrosion.
Types of Corrosion
Some of the corrosion types include the following:
(i) Crevice Corrosion
Whenever there is a difference in ionic concentration between any two local areas of a metal, a localised form of corrosion known as crevice corrosion can occur. For instance, this form of corrosion mostly occurs in confined spaces (crevices). Examples of areas where crevice corrosion can occur are gaskets, the undersurface of washers, and bolt heads. All grades of aluminium alloys and stainless steels also undergo crevice corrosion. This is mainly because of the formation of a differential aeration cell that leads to the formation of corrosion inside the crevices.
(ii) Stress Corrosion Cracking
Stress corrosion cracking can be abbreviated to ‘SCC’ and refers to the cracking of the metal as a result of the corrosive environment and the tensile stress placed on the metal. It often occurs at high temperatures.
For example, stress corrosion cracking of austenitic stainless steel in chloride solution.
(iii) Intergranular Corrosion
Intergranular corrosion occurs due to the presence of impurities in the grain boundaries that separate the grain formed during the solidification of the metal alloy. It can also occur via the depletion or enrichment of the alloy at these grain boundaries.
For example, Aluminum-base alloys are affected by IGC.
(iv) Galvanic Corrosion
When there exists an electric contact between two metals that are electrochemically dissimilar and are in an electrolytic environment, galvanic corrosion can arise. It refers to the degradation of one of these metals at a joint or at a junction. A good example of this type of corrosion would be the degradation that occurs when copper, in a salt-water environment, comes in contact with steel.
For example, when aluminium and carbon steel are connected and immersed in seawater, aluminium corrodes faster, and steel is protected.
(iv) Pitting Corrosion
Pitting Corrosion is very unpredictable and, therefore, is difficult to detect. It is considered one of the most dangerous types of corrosion. It occurs at a local point and proceeds with the formation of a corrosion cell surrounded by the normal metallic surface. Once this ‘pit’ is formed, it continues to grow and can take various shapes. The pit slowly penetrates metal from the surface in a vertical direction, eventually leading to structural failure if left unchecked.
For example, consider a droplet of water on a steel surface, pitting will initiate at the centre of the water droplet (anodic site).
(v) Uniform Corrosion
This is considered the most common form of corrosion wherein an attack on the surface of the metal is executed by the atmosphere. The extent of the corrosion is easily discernible. This type of corrosion has a relatively low impact on the performance of the material.
For example, a piece of zinc and steel immersed in diluted sulphuric acid would usually dissolve over its entire surface at a constant rate.
(vi) Hydrogen Grooving
This is a corrosion of the piping by grooves that are formed due to the interaction of a corrosive agent, corroded pipe constituents, and hydrogen gas bubbles. The bubbles usually remove the protective coating once it comes in contact with the material.
(vii) Metal Dusting
Metal dusting is a damaging form of corrosion that occurs when vulnerable materials are exposed to certain environments with high carbon activities, including synthesis gas. The corrosion results in the break-up of bulk metal to metal powder. Corrosion occurs as a graphite layer is deposited on the surface of the metals from carbon monoxide (CO) in the vapour phase. This graphite layer then goes on to form meta-stable M3C species (where M is a metal) that usually move away from the metal surface. In some cases, no M3C species may be observed. This means that the metal atoms have been directly transferred into the graphite layer.
(viii) Microbial Corrosion
Microbial corrosion, which is also known as microbiologically influenced corrosion (MIC), is a type of corrosion that is caused by microorganisms. The most common one is chemoautotrophs. Both metallic and non-metallic materials, either in the presence or absence of oxygen, can be affected by this corrosion.
(viii) High-temperature Corrosion
High-temperature corrosion, as the name suggests, is a type of corrosion of materials (mostly metals) due to heating. Chemical deterioration of metal can occur due to a hot atmosphere that contains gases such as oxygen, sulphur, or other compounds. These compounds are capable of oxidising the materials (metals in this case) easily. For example, materials used in car engines have to resist sustained periods at high temperatures, during which they can be affected by an atmosphere containing corrosive products of combustion.
Corrosion Examples, Reactions and Effects
Here are some typical examples of corrosion, as seen mostly in metals.
1. Copper Corrosion
When copper metal is exposed to the environment, it reacts with the oxygen in the atmosphere to form copper (I) oxide, which is red in colour.
2Cu(s) + ½ O2(g) → Cu2O(s)
Cu2O further gets oxidised to form CuO, which is black in colour.
Cu2O(s) + ½ O2(g) → 2CuO(s)
This CuO reacts with CO2, SO3 and H2O (present in the atmosphere to form Cu2(OH)2(s) (Malachite), which is blue in colour and Cu4SO4(OH)6(s) (Brochantite), which is green in colour.
This is why we observe copper turning bluish-green in colour.
A typical example of this is the colour of the Statue of Liberty, which has the copper coating on it turning blue-green in colour.
2. Silver Tarnishing
Silver reacts with sulphur and sulphur compounds in the air, giving silver sulphide (Ag2S), which is black in colour. Exposed silver forms Ag2S as it reacts with the H2S(g) in the atmosphere, which is present due to certain industrial processes.
2Ag(s) + H2S(g) → Ag2S(s) + H+2+(g)
3. Corrosion of Iron (Rusting)
Rusting of iron, which is the most commonly seen example, happens when iron comes in contact with air or water. The reaction could be seen as a typical electrochemical cell reaction. Consider the diagram given below.
Here, metal iron loses electrons and gets converted to Fe{aq}2+ (this could be considered as the anode position). The electrons lost will move to the other side, where they combine with H+ ions. H+ ions are released either by H2O or by H2CO3 present in the atmosphere (this could be considered as the cathode position).
�2�⇌�++��−
�2��3⇌2�++��32
The Hydrogen, thus formed by the reaction of H+ and electrons, react with oxygen to form H2O.
Anode reaction
2Fe(s) → 2Fe2+ + 4e– ;
���2+/��0=–0.44 �
Cathode reaction
�2(�)+4�+(��)+4�−⟶2�2�(�)���+/�2/�2/�=1.23�
Overall reaction
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)Eocell = 1.67V
The Fe2+ ions formed at the anode react with oxygen in the atmosphere, thereby getting oxidised to Fe3+ and forming Fe2O3, which comes out in the hydrated form as Fe2O3.xH2O
Fe2+ + 3O2 → 2Fe2O3
Fe2O3 + xH2O → Fe2O3. xH2O (rust)
Other examples include,
• Corrosion of Zinc when it reacts with oxygen and HCl to form white-coloured ZnCl2.
• Corrosion of Tin to form black-coloured Na2[Sn(OH)2].
Effects
Corrosion can have a varying degree of effect on a lot of things. As such, it mainly causes waste of natural resources. Additionally, it can further cause hazardous situations such as building structures becoming weak and unstable, accidents caused by corroded parts as well as other unwanted failures, such as cracked pipelines, bridge collapsing, transport vehicle crashes or other catastrophes. It is, therefore, important to check and prevent corrosion at all costs.
Prevention of Corrosion
Preventing corrosion is of utmost importance in order to avoid huge losses. The majority of the structures that we see and use are made out of metals. This includes bridges, automobiles, machinery, household goods like window grills, doors, railway lines, etc. While this is a concerning issue, several treatments are used to slow or prevent corrosion damage to metallic objects. This is especially done to those materials that are frequently exposed to the weather, saltwater, acids, or other hostile environments. Some of the popular methods to prevent corrosion include,
• Electroplating
• Galvanization
• Anodization
• Passivation
• Biofilm Coatings
• Anti-Corrosion Protective Coatings
• Painting and Greasing
• Use of Corrosion Inhibitor or Drying Agents
• Periodic Cleaning of Metal Surface
•
• Frequently Asked Questions on Corrosion
Q1
What is corrosion?
Corrosion is a natural phenomenon of conversion of metal into a more stable chemical oxide.
Q2
Do all metals corrode?
No, all metals do not corrode. The metals which are higher in the reactivity series corrode easily as their oxidation potential is high. For example, iron.
Q3
What affects corrosion?
Metals on exposure to air or moisture (salt water) increase corrosion.
Q4
What is the effect of temperature on corrosion?
Corrosion increases with an increase in temperature.
Q5
What is meant by high-temperature corrosion?
Corrosion of metals due to heating is called high-temperature corrosion.
SVTs
SVTs
031203.004
SVT031203.001
SVT031203.003
1. What is kohlrausch law and its applications?
2. Why do we need Kohlrausch law?
3. What is infinite dilution in electrochemistry?
4. Give Faraday’s first and second Law of Electrolysis.
5. What is a Battery?
6. What are Primary and secondary cells.
7. What are types of Mercury Cells. Give examples.
8. Is mercury cell rechargeable? Give reasons for your answers.
MCQs
MCQs
True or false,
Also give reasons for the answers
1. An electrochemical cell works only if emf is negative.
2. In an electrolytic cell reduction occurs at cathode.
3. Electrolytic conductance generally decreases with rise in temperature.
4. Greater reduction potential represents greater reducing power of the substance.
5. In Zn-Cu cell, copper acts as a cathode while in Cu-Ag cell, copper acts as anode.
6. Metallic conductance decreases with increase in temperature.
7. Both E⁰ cell and ΔG⁰ for the cell reaction are intensive properties.
8. Ferrous sulphate solution can be safely stored in a copper vessel.
9. Reduction potential and oxidation potential for a half cell reaction are numerically equal but of opposite sign.
10. In a Daniel cell, the electrons flow from anode to cathode in the external circuit and from cathode to anode through the salt bridge to complete the circuit.
11. Out of HCI and NaCl, The molar conductivity of HCI is higher.
12. CuSO4 solution can be stirred with a silver spoon but AgNO3, solution cannot be stirred with a copper spoon.
Which of the following is not a good conductor ? a. Cu metal b. NaCl(aq.) c. NaCl(molten) d. NaCl (s)
D
When 0.5 Faraday of electricity is passed through NaCl solution, the amount of chlorine liberated is : (a) 71.0 g (b) 142.0 g (c) 35.5 g (d) 17.75 g,
D
The amount of silver (at. mass = 108) deposited from a solution of silver nitrate when a current of 965 coulombs was passed is : (a) 10.8 g(b) 0.108 g (c) 1.08 g (d) 1.08 x 103 g.
C
The SI units of molar conductivity are : (a) Sm2 mol-l (b) Sm-1 mol-l c. Sm-2 mol (d) S m3 mol-l.
A
Fill in the blanks, also justify your answer.
1. When 0.5 Faraday of electricity is passed through NaCl
solution, the amount of chlorine liberated is..........
2. The amount of silver (at. mass = 108) deposited from a solution of silver nitrate when a current of 965 coulombs was passed is ............
E1. Represent the cell in which the following reaction takes place
Mg(s) + 2Ag+ (0.0001M) Mg2+(0.130M) + 2Ag(s)
Calculate its E(cell) if Eo(cell) = 3.17 V.
E2. Calculate the equilibrium constant of the reaction:
Cu(s) + 2Ag
Eo(cell) = 0.46 V
E3. The standard electrode potential for Daniell cell is 1.1V.
Calculate
the standard Gibbs energy for the reaction:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
E4. Resistance of a conductivity cell filled with
0.1 mol L–1 KCl solution is
100 W.
If the resistance of the same cell when filled with
0.02 mol L–1 KCl solution is 520 Ω ,
calculate the conductivity and molar conductivity of 0.02 mol L–1KCl solution.
The conductivity of 0.1 mol L–1KCl
solution is 1.29 S/m.
E5. The electrical resistance of a column of
0.05 mol L–1NaOH solution of
diameter 1 cm and length 50 cm is
5.55 × 103 ohm. Calculate its
resistivity, conductivity and molar conductivity.
E6. The molar conductivity of KCl solutions at different concentrations at
298 K are given below: Table NCERT ……..(0312031a)
E7. Calculate Λ⁰m for CaCl2 and MgSO4 from the data given in Table ……(0312031a)
E8. Λ⁰m for NaCl, HCl and NaAc are 126.4, 425.9 and 91.0 S cm2 mol-1 respectively.
Calculate Λ⁰m for HAc.
E9. The conductivity of 0.001028 mol L–1 acetic acid is 4.95 × 10-5 S cm-1.
Calculate its dissociation constant if Λ⁰m for acetic acid is 390.5 S cm2mol-1.
E10. A solution of CuSO4 is electrolysed for 10 minutes with a current of
1.5 amperes. What is the mass of copper deposited at the cathode?
(0312031a)
N1. How would you determine the standard electrode potential of the system
Mg2+|Mg?
N2. Can you store copper sulphate solutions in a zinc pot?
N3. Consult the table of standard electrode potentials and suggest three
substances that can oxidise ferrous ions under suitable conditions
N4. Calculate the potential of hydrogen electrode in contact with a solution
whose pH is 10.
N5. Calculate the emf of the cell in which the following reaction takes place:
Ni(s) + 2Ag+ (0.002 M) Ni2+ (0.160 M) + 2Ag(s)
Given that Eo(cell) = 1.05 V
N6. The cell in which the following reactions occurs:
2Fe3+(aq) +2 I- Fe2+(aq) + I2(s)
has Eo(cell) = 0.236 V at 298 K.
Calculate the standard Gibbs energy and the equilibrium constant of the cell reaction.
N7. Why does the conductivity of a solution decrease with dilution?
N8. Suggest a way to determine the Λ⁰m value of water.
N9. The molar conductivity of 0.025 mol L-1methanoic acid is 46.1 S cm2 mol-1.
Calculate its degree of dissociation and dissociation constant.
Given λ⁰(H+) = 349.6 S cm2 mol-1and λ⁰ (HCOO-1) = 54.6 S cm2 mol-1.
N10. If a current of 0.5 ampere flows through a metallic wire for 2 hours,
then how many electrons would flow through the wire?
N11. Suggest a list of metals that are extracted electrolytically.
N12. Consider the reaction:
Cr2O72-+ 14H+ + 6e- 2Cr3++ 7H2O
What is the quantity of electricity in coulombs
needed to reduce 1 mol of CrO72-?
N13. Write the chemistry of recharging the lead storage battery,
highlighting all the materials that are involved during recharging.
N14. Suggest two materials other than hydrogen that
can be used as fuels in fuel cells.
N15. Explain how rusting of iron is envisaged as setting up of an electrochemical cell.
B1. Arrange the following metals in the order in which they displace each other
from the solution of their salts.
Al, Cu, Fe, Mg and Zn.
B2. Given the standard electrode potentials,
K+/K = –2.93V, Ag+/Ag = 0.80V,
Hg2+/Hg = 0.79V
Mg2+/Mg = –2.37 V, Cr3+/Cr = – 0.74V
Arrange these metals in their increasing order of reducing power.
B.3 Depict the galvanic cell in which the reaction
Zn(s)+2Ag+(aq) ®Zn2+(aq)+2Ag(s) takes place. Further show:
(i) Which of the electrode is negatively charged?
(ii) The carriers of the current in the cell.
(iii) Individual reaction at each electrode.
B4. Calculate the standard cell potentials of galvanic cell in which the following
reactions take place:
(i) 2Cr(s) + 3Cd2+(aq) ® 2Cr3+(aq) + 3Cd
(ii) Fe2+(aq) + Ag+(aq) ® Fe3+(aq) + Ag(s)
Calculate the D
r
Go and equilibrium constant of the reactions.
B5. Write the Nernst equation and emf of the following cells at 298 K:
(i) Mg(s)|Mg2+(0.001M)||Cu2+(0.0001 M)|Cu(s)
(ii) Fe(s)|Fe2+(0.001M)||H+(1M)|H2(g)(1bar)| Pt(s)
(iii) Sn(s)|Sn2+(0.050 M)||H+(0.020 M)|H2(g) (1 bar)|Pt(s)
(iv) Pt(s)|Br–(0.010 M)|Br2(l )||H+(0.030 M)| H2(g) (1 bar)|Pt(s).
B6. In the button cells widely used in watches and other devices the following
reaction takes place:
Zn(s) + Ag2O(s) + H2O(l ) ® Zn2+(aq) + 2Ag(s) + 2OH–(aq)
Determine D
r
G
o and E
o
for the reaction.
B7. Define conductivity and molar conductivity for the solution of an electrolyte.
Discuss their variation with concentration.
B8. The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 S cm–1. Calculate
its molar conductivity.
B9. The resistance of a conductivity cell containing 0.001M KCl solution at 298
K is 1500 W. What is the cell constant if conductivity of 0.001M KCl solution
at 298 K is 0.146 × 10–3 S cm–1.
B10. The conductivity of sodium chloride at 298 K has been determined at different
concentrations and the results are given below:
Concentration/M 0.001 0.010 0.020 0.050 0.100
102 × k/S m–1 1.237 11.85 23.15 55.53 106.74
Calculate Λ
m for all concentrations and draw a plot between Λ
m and c½.
Find the value of 0
m
.
B11. Conductivity of 0.00241 M acetic acid is 7.896 × 10–5 S cm–1. Calculate its
molar conductivity. If 0
m
for acetic acid is 390.5 S cm2 mol–1, what is its
dissociation constant?
B12. How much charge is required for the following reductions:
(i) 1 mol of Al3+ to Al?
(ii) 1 mol of Cu2+ to Cu?
(iii) 1 mol of MnO4
– to Mn2+?
B13. How much electricity in terms of Faraday is required to produce
(i) 20.0 g of Ca from molten CaCl2?
(ii) 40.0 g of Al from molten Al2O3?
B14. How much electricity is required in coulomb for the oxidation of
(i) 1 mol of H2O to O2?
(ii) 1 mol of FeO to Fe2O3?
B15. A solution of Ni(NO3)2 is electrolysed between platinum electrodes using a
current of 5 amperes for 20 minutes. What mass of Ni is deposited at the
cathode?
B16. Three electrolytic cells A,B,C containing solutions of ZnSO4, AgNO3 and CuSO4,
respectively are connected in series. A steady current of 1.5 amperes was
passed through them until 1.45 g of silver deposited at the cathode of cell B.
How long did the current flow? What mass of copper and zinc were deposited?
B17. Using the standard electrode potentials given in Table 3.1, predict if the
reaction between the following is feasible:
(i) Fe3+(aq) and I–(aq)
(ii) Ag+ (aq) and Cu(s)
(iii) Fe3+ (aq) and Br– (aq)
(iv) Ag(s) and Fe
3+ (aq)
(v) Br2 (aq) and Fe2+ (aq).
B18. Predict the products of electrolysis in each of the following:
(i) An aqueous solution of AgNO3 with silver electrodes.
(ii) An aqueous solution of AgNO3 with platinum electrodes.
(iii) A dilute solution of H2SO4 with platinum electrodes.
(iv) An aqueous solution of CuCl2 with platinum electrodes.