• A device is used to convert the chemical energy produced in a redox reaction in to Electric energy is called a electrochemical cell or simply a chemical cell.
• These are also called as Galvanic cell or simply a Voltaic cell.
ELECTROCHEMICAL CELL
Oxidation -
loss of electron-Gain of oxygen
Reduction
-gain of electron -Loss of oxygen
Redox Reactions.
Redox Reactions.
Loss of oxygen Gain of electrons
• An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped in a
solution of electrolyte. These half-cells are designed to contain the oxidation half-
reaction and reduction half-reaction separately as shown below.
• The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown below.
• Zn (s) ---> Zn+2 (aq) + 2e-
• During the oxidation of zinc, the zinc
electrode will slowly dissolve to produce zinc ions (Zn+2), which enter into the
solution containing Zn+2 (aq) and SO4-2 (aq) ions.
Voltaic Cell: How it works
Voltaic Cell: How it works
• The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below.
• Cu+2 (aq) + 2e- ---> Cu (s)
• When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface of the solid copper electrode.
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How a Voltaic Cell Works
The two half-reactions can be summed to show the overall reaction.
• Note that the electrons must cancel.
Zn(s) → Zn2+(aq) + 2e– Cu2+(aq) + 2e– → Cu(s)
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
The half-cell that undergoes
The half-cell that undergoes oxidation oxidation (the anode) is written first, to the left of the double vertical lines.(the anode) is written first, to the left of the double vertical lines.
Representing Electrochemical Cells
• The single vertical lines indicate boundaries of phases that are in contact.
• The double vertical lines represent the salt bridge or porous partition that separates the
anode compartment from the cathode compartment.
You can represent the zinc-copper voltaic cell by using the following shorthand form.
Zn(s) ZnSO
4(aq) CuSO
4(aq) Cu(s)
Salt Bridge Salt Bridge
• The reaction in each half-cell does not occur unless the two half cells are connected to each other.
• It is an inverted U-tube contaning aan electrolyte e.g KCL,KNO3 etc it act as bridge by connecting two half cells
---Helps in
• To completing the electric circuit
• To prevent mixing of soln of two half cell.
• To help maintain electric neutrality
Summary of Voltaic Cells (Galvanic Summary of Voltaic Cells (Galvanic
Cell) Cell)
• A device that spontaneously produces electricity by redox
• Uses chemical substances that will participate in a spontaneous redox reaction.
– The reduction half-reaction (SOA) will be above the oxidation half-reaction (SRA) in the activity series to ensure a spontaneous reaction.
• Composed of two half-cells; which each consist of a metal rod or strip immersed in a solution of its own ions or an inert electrolyte.
– Electrodes: solid conductors connecting the cell to an external circuit – Anode: electrode where oxidation occurs (-)
– Cathode: electrode where reduction occurs (+)
– The electrons flow from the anode to the cathode (“a before c”) through an electrical circuit rather than passing directly from one substance to another
– A porous boundary separates the two electrolytes while still allowing ions to flow to maintain cell neutrality
• Often the porous boundary is a salt bridge, containing an inert aqueous electrolyte (such as Na2SO4(aq) or KNO3(aq)),
• Or you can use a porous cup containing one electrolyte which sits in a container of a second electrolyte.
Electrolytic Cell Electrolytic Cell
• It is a device to convert Electric Energy into Chemical Energy
• An electrolytic cell is an electrochemical cell in which the energy from an applied voltage is used to drive an otherwise
nonspontaneous reaction. Such a cell
could be produced by applying a reverse voltage to a voltaic cell like the Daniell cell.
Difference b/w Electrochemical Difference b/w Electrochemical
cell and Eletrolytic Cell cell and Eletrolytic Cell
Electrochemical cell Chemical enrgy
converted into Electric energy
It is based on redox rxn which is spontaneous
i.e rxn occurs its own
Eletrolytic Cell
• Electric enrgy converted into Chemical energy
• It is based on redox rxn which is non-
spontaneous
Electrode Potential Electrode Potential
• It is the tendency of an electrode in half cell to lose or gain electrons when it is in contact with solution of its own ion
• Zn (s) ---> Zn+2 (aq) + 2e- Oxidation half rxn
• Cu+2 (aq) + 2e- ---> Cu (s) Reduction half rxn
Electomotive force (emf) Electomotive force (emf)
• The two half-cells are also connected externally.
In this arrangement, electrons provided by the oxidation reaction are forced to travel via an
external circuit to the site of the reduction reaction.
The fact that the reaction occurs spontaneously once these half cells are connected indicates that there is a difference in potential energy. This
difference in potential energy is called an
electomotive force (emf) and is measured in
terms of volts. The zinc/copper cell has an emf of about 1.1 volts under standard conditions.
Nernst Equation Nernst Equation
• Electrochemistry deals with cell potential as well as energy of chemical reactions.
The energy of a chemical system drives the charges to move, and the driving force give rise to the cell potential of a system called galvanic cell. The energy aspect is also related to the chemical equilibrium. All these relationships are tied together in the concept of Nearnst equation.
• Walther H. Nernst (1864-1941) received the Nobel prize in 1920 "in recognition of his work in thermochemistry". His
contribution to chemical thermodynamics led to the well known equation correlating chemical energy and the electric potential of a galvanic cell or battery.
• The general Nernst equation correlates the Gibb's Free Energy ΔG and the EMF of a chemical system known as the galvanic cell.
• Ecell = E0cell - (RT/nF)lnQ
– Here, Ecell = Electrode potential
– E0cell = Standard potential of a particular reaction – F=Faraday constant, R=Universal Gas constant,
T=Absolute temperature
– n= number of electron involved – Q= Amount of current flow
Nernst Equation Nernst Equation
• Remember that
G = G + RT ln Q
• This means
−nFE = −nFE + RT ln Q
• Dividing both sides by −nF, we get the Nernst equation:
or, using base-10 logarithms, E = E − RT
nF ln Q E = E − 2.303 RT
nF log Q
Nernst Equation Nernst Equation
At room temperature (298 K), and R = 8.314 J/mol K
F = 96,485 J/V-mol
The final form of the Nernst Equation becomes
E = E − 0.0592
n log Q
2.303 RT
F = 0.0592 V
• Primary cell
• A primary cell is a special type of electrochemical cell in which the reaction cannot be reversed, and the identities of the anode and cathode are therefore fixed.
The anode is always the negative electrode. The cell can be discharged but not recharged.
• Secondary cell
• A secondary cell, is one in which the chemical reactions are reversible.
• When the cell is being charged, the anode becomes the positive (+) and the cathode the negative (−) electrode. This is also the case in an electrolytic cell.
• When the cell is being discharged, it behaves like a primary cell, with the anode as the negative and the cathode as the positive electrode.
• For example, Rechargeable battery
Reference Electrodes Reference Electrodes
• A reference electrode is an electrode which has a stable and well-known electrode potential.
• There are many ways reference electrodes are used. The simplest is when the reference
electrode is used as a half cell to build an
electrochemical cell. This allows the potential of the other half cell to be determined.
• Common reference electrodes are
• Standard hydrogen electrode (SHE)
• Normal hydrogen electrode (NHE)
• Saturated calomel electrode (SCE)
Potentiometric titration Potentiometric titration
• Potentiometric titration is a technique similar to direct titration of a redox reaction. No indicator is used. To do this, two electrodes are used, an indicator electrode and a reference electrode.
• The indicator electrode forms an electrochemical half cell with the interested ions in the test solution. The reference electrode forms the other half cell, holding a consistent electrical potential.
Buffer solution Buffer solution
• A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
• Its pH changes very little when a small amount of strong acid or base is added to it and thus it is used to prevent changes in the pH of a solution.
• Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical application.
• One example of a buffer solution found in nature is blood.
• Henderson–Hasselbarch equation
• In chemistry, the Henderson–Hasselbalch equation is useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base reactions.
• The equation is given by:
• Here, [HA] is the molar concentration of the undissociated weak acid, [A ] is the molar ⁻ concentration of this acid's conjugate base
• And pKa= -log ka where ka is the acid dissociation constant
Concentration Cells Concentration Cells
• Electrolyte concentration cell
– Where the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells.
• So, a concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.
• For example an electrochemical cell, where two copper electrodes are submerged in two
copper(II) sulfate solutions, whose
concentrations are 0.05 M and 2.0 M, connected through a salt bridge.
• This type of cell will generate a potential that can be predicted by the Nernst equation. Both
electrodes undergo the same chemistry
(although the reaction proceeds in reverse at the cathode)
• Cu2+(aq) + 2 e– → Cu(s) (at anode)
Concentration Cells (II) Concentration Cells (II)
• Electrode concentration cells
– the electrodes themselves have
different compositions. This may be due to.
• Different fugacity of gases involved in
electrode reactions (e.g., The H+ (aq)/H2 (g) electrode).
• Different compositions of metal amalgams in electrode materials.
Applications of Electrochemistry Applications of Electrochemistry
• Measurement of activities and activity coefficients.
• Electrochemical series.
• Equilibrium constants and
thermodynamic functions of cell reactions
• the coating of objects with metals or metal oxides through electrodeposition
Cont..
Cont..
• The generation of chemical energy
through photosynthesis is inherently an electrochemical process
• production of metals like aluminum and titanium from their ores
• Certain diabetes blood sugar meters measure the amount of glucose in the blood through its redox potential
Cont..
Cont..
• Electrochemistry has also important
applications in the food industry, like the
– assessment of food/package interactions, – the analysis of milk composition,
– the characterization and the determination of the freezing end-point of ice-cream mixes, – the determination of free acidity in olive oil
