Electricity: الكهربائيه

Electrical conductivity: الكهربائيه التوصيليه

Electrical conductivity (C-) or specific conductance is the reciprocal of electrical resistivity, and measures a material's ability to conduct an electric current. Generally, conductivity is equal to the inverse of resistance of material®

C-= 1/R (its unit is Siemens per meter, S/m)

Electrolytic conductivity: االلكتروليتيه التوصيليه

Conductivity (or specific conductance) of an electrolyte solution is a measure of its ability to conduct electricity via ions that are producing from dissociation of electrolyte to yield cations and anions. The SI unit of conductivity is Siemens per meter (S/m).

Conductivity measurements are used routinely in many industrial and environmental applications as a fast, inexpensive and reliable way of measuring the ionic content in a solution.[1] For example, the measurement of product conductivity is a typical way to monitor and continuously trend the performance of water purification systems.

Electrolytic conductivity of ultra-high purity water as a function of temperature.In many cases, conductivity is linked directly to the total dissolved solids (T.D.S.). High quality deionized water has a conductivity of about 5.5 μS/m at 25 °C, typical drinking water in the range of 5–50 mS/m, while sea water about 5 S/m[2] (or 50,000 μS/cm) (i.e., sea water's conductivity is one million times higher than that of deionized water:).

Conductivity is traditionally determined by connecting the electrolyte in a Wheatstone bridge. Dilute solutions follow Kohlrausch's Laws of concentration dependence and additivity

of ionic contributions. Lars Onsager gave a theoretical explanation of Kohlrausch's law by extending Debye–Hückel theory.

How to measure electrolytic conductivity

The electrical conductivity of a solution of an electrolyte is measured by determining the resistance of the solution between two flat or cylindrical electrodes separated by a fixed distance.[3] An alternating voltage is used in order to avoid electrolysis.[citation needed] The resistance is measured by a conductivity meter. Typical frequencies used are in the range 1–3 kHz. The dependence on the frequency is usually small,[4] but may become appreciable at very high frequencies, an effect known as the Debye–Falkenhagen effect.

A wide variety of instrumentation is commercially available.[5] There are two types of cell, the classical type with flat or cylindrical electrodes and a second type based on induction.[6] Many commercial systems offer automatic temperature correction. Tables of reference conductivities are available for many common solutions.

Figure xx: measuring of electrical conductivity of electrolyte

Types of conductors

There are two main types of electrical conductors:

1-Metalic conductors: This conducts electricity via electrons when applying under electrical field such as conducting metals, Cu, Fe, Al, Mn,etc---. This type of conductors work according to ohm law:

E=I.R (E, electrical potential, I, current and R is resistance in ohm).

2-electrolytic conductors: this conduct electricity by ions in the electrolytic solution. Conducting here occurs via cations and anions. Electrolytes can be classified into two types:

2-1: Strong electrolyte: These are excellent conductors and are completely dissociated in solution and show strong conducting such as HCl, HNO3, NaCl etc..,

2-2: Weak electrolyte: These are bad conductors as they are partially dissociated in the solution such as CH3COOH,

The relation between current (I), and electrical potential (E) for metallic conductors and electrolytic conductors can be presented in the following figure:

Dissociation potential l(Eo)

Potential( E)

Electrolytic conductor

E=Eo+I.R

Current(I)

Metallic conductor

(E=I.R)

Figure xx: metallic and electrolytic conductors

From above figure, it can be seen that, metallic conductors are agree completely with ohms law:

E=I.R, Eα I, (metallic conductor),

For electrolytic conductors, don’t agree completely with ohms law and no current flow after applying dissociation potential(Eo). This causes dissociation of electrolyte to anions and cations and these are contributed in electrical conductivity and Ohms law for this case:

(E=Eo+I.R , Ohms law for electrolytic conductors).

Electrolysis: الكهربائي التحليل

Electrolysis is a technique that uses a direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential.

The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The desired products of electrolysis are often in a different physical state from the electrolyte and can be removed by some physical processes. For example, in the electrolysis of brine to produce hydrogen and chlorine, the products are gaseous. These gaseous products bubble from the electrolyte and are collected.

Figure xx: Illustration of an electrolysis apparatus used in a school laboratory.

Generally, in electrolysis and after applying (Eo), electrolyte decompose to its cations and anions. Cations move to negative electrode while anions move to positive electrode.

Basic principles of electrolysis were reported by English Author Michael Faraday who found 1st and 2nd law in electrolysis.

An electrolytic cell is an electrochemical cell

that drives a non-spontaneous redox reaction through the application of electrical energy. They are often used to decompose chemical compounds, in a process called electrolysis—the Greek word lysis means to break up.

Figure xx: samples of electrolytic cells

Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminium and other chemicals. Electroplating (e.g. of copper, silver, nickel or chromium) is done using an electrolytic cell. Electrolysis is a technique that uses a direct electric current (DC).

An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the

electrodes, the ions in the electrolyte are attracted to an electrode with the opposite charge, where charge-transferring (also called faradaic or redox) reactions can take place. Only with an external electrical potential (i.e. voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided can produce a chemical reaction which would not occur spontaneously otherwise.

Galvanic cells compared to electrolytic cells

A galvanic cell can be considered an electrolytic cell acting in reverse. While electrolytic cells convert electrical energy into chemical energy, galvanic cells convert chemical energy into electrical energy. Galvanic cells are often used in batteries.

A galvanic cell, or voltaic cell, named after Luigi Galvani, or Alessandro Volta respectively, is an electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell. It generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane.

Volta was the inventor of the voltaic pile, the first electrical battery. In common usage, the word "battery" has come to include a single galvanic cell, but a battery properly consists of multiple cells.

Figure xx: Galvanic cell with no cation flow

Basic description Schematic of Zn-Cu galvanic cellIn its simplest form, a half-cell consists of a solid metal (called an electrode) that is submerged in a solution; the solution contains cations (+) of the electrode metal and anions (−) to balance the charge of the cations. The full cell consists of two half-cells, usually separated by a semi-permeable membrane or by a salt bridge.

A specific example is the Daniell cell, where a salt bridge is used as separator (see figure). The zinc (Zn) half-cell has a solution of ZnSO4 (zinc sulfate) and the copper (Cu) half-cell has a solution of CuSO4 (copper sulfate).

Let an external electrical conductor connect the copper and zinc electrodes. In the zinc half-cell, zinc from the zinc electrode dissolves into the solution as Zn2+ ions (oxidation), releasing electrons that enter the external conductor. In addition, via the salt bridge zinc ions leave and sulfate ions (SO2−4) enter the zinc half-cell.

In the copper half-cell, the copper ions plate onto the copper electrode (reduction), taking up electrons that leave the external conductor. Since the Cu2+ ions (cations) plate onto the copper electrode, the latter is called the cathode. Correspondingly the zinc electrode is the anode. The electrochemical reaction is:

Zn + Cu2+ → Zn2+ + CuIn addition, electrons flow through the external conductor, which is the primary application of the galvanic cell.

As discussed under #Cell voltage, the emf of the cell is the difference of the half-cell potentials, a measure of the relative ease of dissolution of the two electrodes into the electrolyte. The emf depends on both the electrodes and on the electrolyte, an indication that the emf is chemical in nature.

Figure xx: Schematic of Zn-Cu galvanic cell

Anode and cathode definitions:

depend on charge and discharge[edit]Michael Faraday defined the cathode of a cell as the electrode to which cations (positively charged ions, like silver ions Ag+

) flow within the cell, to be reduced by reacting with electrons (negatively charged) from that electrode. Likewise he defined the anode as the electrode to which anions (negatively charged ions, like chloride ions Cl−

) flow within the cell, to be oxidized by depositing electrons on the electrode.

To an external wire connected to the electrodes of a Galvanic cell (or battery), forming an electric circuit, the cathode is positive and the anode is negative. Thus positive electric current flows from the cathode to the anode through the external circuit in the case of a Galvanic cell.

Consider two voltaic cells of unequal voltage. Mark the positive and negative electrodes of each one as P and N, respectively. Place them in a circuit with P near P and N near N, so the cells will tend to drive current in opposite directions. The cell with the larger voltage is discharged, making it a galvanic cell, so P is the cathode and N is the anode as described above. But, the cell with the smaller voltage charges, making it an electrolytic cell. In the electrolytic cell, negative ions are driven towards P and positive ions towards N. Thus, the P electrode of the electrolytic cell meets the definition of anode while the electrolytic cell is being charged. Similarly, the N electrode of the electrolytic cell is the cathode while the electrolytic cell is being charged.

Cell voltage The standard electrical potential of a cell can be determined by the use of a standard potential table for the two half cells involved. The first step is to identify the two metals reacting in the cell. Then one looks up the standard electrode potential, E0, in volts, for each of the two half reactions. The standard potential for the cell is equal to the more positive E0 value minus the more negative E0 value.

For example, in the figure above the solutions are CuSO4 and ZnSO4. Each solution has a corresponding metal strip in it, and a salt bridge or porous disk connecting the two solutions and allowing SO2−

4 ions to flow freely between the copper and zinc solutions. To calculate the standard potential one looks up copper and zinc's half reactions and finds:

Cu2+ + 2 e−⇌ Cu   E0 = +0.34 VZn2+ + 2 e−⇌ Zn   E0 = −0.76 V

Thus the overall reaction is:

Cu2+ + Zn ⇌ Cu + Zn2+

The standard potential for the reaction is then +0.34 V − (−0.76 V) = 1.10 V. The polarity of the cell is determined as follows. Zinc metal is more strongly reducing than copper metal; equivalently, the standard (reduction) potential for zinc is more negative than that of copper. Thus, zinc metal will lose electrons to copper ions and develop a positive electrical charge. The equilibrium constant, K, for the cell is given by

Faraday's laws of electrolysis

Faraday's laws of electrolysis are quantitative relationships based on the electrochemical researches published by Michael Faraday in 1834. Faraday's law of induction is a basic law of electromagnetism predicting how a magnetic field will interact with an electric circuit to produce an electromotive force (EMF)—a phenomenon called electromagnetic induction. It is the fundamental operating principle of transformers, inductors, and many types of electrical motors, generators and solenoids.[1][2]

The Maxwell–Faraday equation is a generalization of Faraday's law, and is listed as one of Maxwell's equations

First Faraday law of electrolysis:

In 1832, Michael Faraday reported that the quantity of elements separated by passing an electric current through a molten or dissolved salt is proportional to the quantity of electric charge passed through the circuit. This became the basis of the first law of electrolysis:

(mα Q)

(m=eQ)

where; m is the weight of substance liberated at anode or deposited at cathode, e is known as electrochemical equivalent of the metal deposited or of the gas liberated at the electrode Q is the amount passed electrical current .

first Faraday law of electrolysis

Faraday discovered that when the same amount of current is passed through different electrolytes/elements connected in series, the mass of substance liberated/deposited at the electrodes is directly proportional to their equivalent weight.

(mα e), so that, m=e.Q.( Q is a constant)

Second Faraday law of electrolysis:

The weight of matter that is liberated at anode or that deposited at cathode is directly proportional to the amount electricity that is passed through the electrolytic solution:

(mα I. t),

(mα I.t.e),

(m= (I.t.e)/F

Herein, m: mass of the matter, I: current, t: time, e:equivalent weight.

F= Nx e,( N is Avogadro's number, e is the charge of electron= 1.6 x 10-19 column)

Q= I. time,

F= faraday =96500 column

Example:

When a passing of electrical current in a solution of H2SO4 using Pt electrode for one hour, this leads to evolve of 366 mL from a mixture of O2 and H2 gas at STP conditions, calculate amount of the passed current (I)?

Solution:

Equivalent weight € of O2= 16/2=8 gm,

(e of H2= 2/2=1 gm,

Moles of O2= 8/32= 0.25 mole,

Moles of H2= 1/2= 0.5 mole,

Total evolved gas moles= 0.75 moles

1 mole of gas has 22.414 liter,

Volume of evolved gas= 22.414 x 0.75)/1= 16.8 liter

16.8 liter required 96500 column , so that 366 ml(0.366 liter),

Q= (96500 x 0.366)/ 16.8= 2102.321 colunm,

Q= I.t,, I=Q/t= 2102.321/ 3600= 0.584 amber.

Example 2:

Find the time that is needed to paint a plate of copper (25 cm2), with a thickness of 0.1 mm when passing a constant current of 0.5 Am, density of copper 8.96 g/cm3?

Solution:

Volume= 25 x 0.01= 0.25 cm3 ,

Mass= Volume x density= 0.25 x 8.96= 2.24 gm,

According ot 2nd faraday law:

Mass= (I.t.e)/F

For Cu, e= 63.5/2= 31.75 gm,

2.24= (0.5 x t. 31.75)/96500,

Time= 13616 sec= 3.47 minute.

Example3:

When passing of 0.1 Am through a solution of CuSO4 for ten minutes , find how many gram of Cu would precipitate on cathode with number of Cu atoms?

Solution:

Mass = (I.e.t)/F

(e of Cu= 63.5/2= 31.75 gm.)

M= (0.1 x 600 x 31.75)/96500= 0.0195 grram

No of atoms= no of moles x N= (0.0195/63.5)x (6.02 x 1023)

No of atoms= 1.872 x 1020 atom

Electrolytic conductivity: االلكتروليتي التوصيل

Electrolytic conductivity, it is the process of carrying electricity by ions in the electrolytic solution. This occurs after dissociation of the electrolyte under applying potential (E); if we have wire with length of (L) and cross sectional area of (a). then its resistance:

(R αland Rα1/a)

So that, resistance, R, is proportional to the distance( l) between the electrodes and is inversely proportional to the cross-sectional area of the sample,( A)

(RαL/a)

For this equation, proportional constant(r), is the specific resistance:

(R= r.l/a)

In practice the conductivity cell is calibrated by using solutions of known specific resistance, ρ*, so the quantities l and A need

not be known precisely.[8] If the resistance of the calibration solution is R*, a cell-constant, C, is derived.

R ∗ = C × ρ ∗ {\displaystyle R^{*}=C\times \rho ^{*}}

The specific conductance (conductivity), κ (kappa) is the reciprocal of the specific resistance.

Figure xx: xxxxxxxxxxxxxxxxxxxxxxxxxxx

Specific resistance (or resistivity), : النوعيه المقاومه

Specific resistance (r), it is a resistance of silk of 1 cm in its length with a cross sectional area of 1 cm2. It means that (r) is a resistance of silk of volume of 1 cm3 or a solution of 1 mL.

Conductivity: التوصيليه

Conductivity is the inverse of resistance and denoted as (C-),

And the specific conductivity is the inverse of specific resistance(L-), and it is the conductance of 1 mL of the solution.

C-= 1/R, and L-=1/r

C-: is conductance, L- is the specific conductance, its unit (ohm-

1.cm-1)

C- and L- can be connected as follows:

1/C-=1/L- x l/a

L= C- xl/a,

L= 1/Rx l/a

The factor (l/a)= k, it is called cell constant,

Cell constant, k= l/a: l is the distance between two electrodes, and cross sectional area of the electrode.

Equivalent conductance: المكافيء التوصيل

Equivalent conductance (or normality conductance), it is the conductivity of 1 mL of a desired solution that containing 1 gram equivalent of the matter. Mathematically, Equivalent conductance ( Λ ) is equal to (L/C),

(Λ= L/CN (Λ is Equivalent conductance, L specific conductance (ohm-1.cm-1, C is normal concentration).

Λn= L/C= L/C/1000= 1000L/C

Λn= 100x L/C (its unit in eq-1 ohm-1.cm2).

Molar conductivity Λ m: الموالريه التوصيليه

It’s the conductivity of 1 mL of solution that is containing one molecular weight of the dissolved matter.

Λm= L/CM= 1000x L/C (its unit in eq-1.ohm-1.cm2)

There is a direct proportionality between concentration (C) and the conductance: (C- α C and Lα C).this relation not linear for the end it decreases with highly concentrated solutions.

Generally, relation between conductivity and concentration is studied according to the Kohlrausch's Law . .

The electrical conductivity of aqueous solutions is governed by the presence and concentration of ions in solution. Therefore, pure water does not conduct an electrical current well sine the concentrations of hydrogen and hydroxide ions are very small.

(Concentration)1/2

Λo

Electrical conductance, L, Λ

Strong electrolyte or high concentrations

(Concentration)

Electrical conductance, L, Λ

Λo weak electrolyte or diluted concentrations

Solutes whose solutions are conductive are called electrolytes -- a solute is considered a strong electrolyte if it dissociates completely into its constituent ions.

Ex. strong electrolytes --> ionic compounds like NaCl and KI

compounds with highly polar covalent bonds --> HCl in water

Thus, if concentration was higher, that means more ions are present per liter of water -- more ions = more conductivity.

Kohlrausch's LawKohlrausch law states that, “At time infinite dilution, the molar conductivity of an electrolyte can be expressed as the sum of the contributions from its individual ions” i.e., , where, and are the number of cations and anions per formula unit of electrolyte respectively and, and are the molar conductivities of the cation and anion at infinite dilution respectively. The use of above equation in expressing the molar conductivity of an electrolyte is illustrated as,

The molar conductivity of HCl at infinite dilution can be expressed as,

; For HCl, and So, ; Hence,

Xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx

There are two types of electrolytes: strong and weak. Strong electrolytes will usually undergo complete ionization, and therefore they have higher conductivity than weak electrolytes which undergo only partial ionization. For strong electrolytes, such as salts, strong acids and strong bases, the molar conductivity depends only weakly on concentration. On dilution there is a regular increase in the molar conductivity of strong

electrolyte, due to the decrease in solute-solute interaction. Based on experimental data Friedrich Kohlrausch (around the year 1900) proposed the non-linear law for strong electrolytes:

Λ m = Λ m ∘ − K c = α f λ Λ m ∘ {\displaystyle \Lambda _{\mathrm {m} }=\Lambda _{\mathrm {m} }^{\circ }-K{\sqrt {c}}=\alpha f_{\lambda }\Lambda _{\mathrm {m} }^{\circ }}

where

Resistance, R, is proportional to the distance, l, between the electrodes and is inversely proportional to the cross-sectional area of the sample, A (noted S on the Figure above). Writing ρ (rho) for the specific resistance (or resistivity),

R = l A ρ . {\displaystyle R={\frac {l}{A}}\rho .}

In practice the conductivity cell is calibrated by using solutions of known specific resistance, ρ*, so the quantities l and A need not be known precisely.[8] If the resistance of the calibration solution is R*, a cell-constant, C, is derived.

R ∗ = C × ρ ∗ {\displaystyle R^{*}=C\times \rho ^{*}}

The specific conductance (conductivity), κ (kappa) is the reciprocal of the specific resistance.

κ = 1 ρ = C R {\displaystyle \kappa ={\frac {1}{\rho }}={\frac {C}{R}}}

Conductivity is also temperature-dependent. Sometimes the ratio of l and A is called as the cell constant, denoted as G*, and conductance is denoted as G. Then the specific conductance κ (kappa), can be more conveniently written as

Measuring electrolytic conductance of a solution: قياس للمحلول االلكتروليتيه التوصيليه

The electrical conductivity of a solution of an electrolyte is measured by determining the resistance of the solution between two flat or cylindrical electrodes separated by a fixed distance. An alternating voltage is used in order to avoid electrolysis.[citation needed] The resistance is measured by a conductivity meter. Typical frequencies used are in the range 1–3 kHz. The dependence on the frequency is usually small,[4] but may become appreciable at very high frequencies, an effect known as the Debye–Falkenhagen effect. A wide variety of instrumentation is commercially available.[5] There are two types of cell, the classical type with flat or cylindrical electrodes and a second type based on induction.[6] Many commercial systems offer automatic temperature correction. Tables of reference conductivities are available for many common solutions.

Electrolytic conductance of a solution can be measured by using electrical cell of two identical electrode of (Pt) , this type of cell called (conductance cell).

Cell constant, k=l/a

Figure xx: conductance cell

Instrumentally, electrical conductance can be measured using instrument called (conductoormeter). Its work according to the (whinstone bridge), the net passing current in this bridge is equal to zero which can be detected using Galvanometer.

A Wheatstone bridge:

Is an electrical circuit used to measure an unknown electrical resistance by balancing two legs of a bridge circuit, one leg of which includes the unknown component. The primary benefit of the circuit is its ability to provide extremely accurate measurements (in contrast with something like a simple voltage divider).[1] Its operation is similar to the original potentiometer. The Wheatstone bridge was invented by Samuel Hunter Christie in 1833 and improved and popularized by Sir Charles Wheatstone in 1843. One of the Wheatstone bridge's initial uses was for the purpose of soils analysis and comparison.

Figure xx: Wheatstone bridge circuit diagram. The unknown resistance Rx is to be measured; resistances R1, R2 and R3 are known and R2 is adjustable. If the measured voltage VG is 0, then R2/R1 = Rx/R3.

After calculating the value of Rx, electrical conductance can be measured , after knowing the value of cell constant(k). this can be done using a solution of saturated KCl salt (0.1 N):

L= k/R, and k=Lx R

When calculating electrical conductance of solution, conductance of solvent(H2O) must be omitted as:

Lelectrolyte= Lsolution-L solvent(H2O), (L H2O= 0.8 x 10-6 Ω-1 cm-1

Conductivity of water is related to the presence of H+ and OH- ions.

Example:

Electrical conductance, with a distance between its electrodes equal to 5 cm with cross sectional area of 20 cm2 with AgNO3 solution 0.1 N with a resistance of 400 Ω find equivalent conductance for the solution?

Solution:

L= K/R, AND K= Lx R,

(k=l/a, k= 5 cm/20 cm2, k= 0.25 cm-1

(k= L x R, 0.25=Lx 400), L= 6.25 x 10-3 Ω-1 . cm-1

Λn= 1000L/C,

Λn= 1000x6.25 x 10-3 / 0.1= 62.5 eq-1.Ω-1.cm2 .

Ionic mobility: األيونية االنتقالية

Ionic mobility is the ability of charged particles (such as anions ,cations, H+, OH-, Na+, K+etc..)) to move through a medium in response to an electric field that is pulling them. The separation of ions according to their mobility in liquid phase it is called electrophoresis.

When a charged particle in a gas or liquid is acted upon by a uniform electric field, it will be accelerated until it reaches a constant drift velocity according to the formula.

Generally, ionic mobility is the speed of ions in electrolytic solution when applying an electrical field and it depends on the following factors:

1-volume of the ion( radius, r),

2-charge of ion,

3-number of molecules of the solvent that are surrounding to the moving ion.

Ionic mobility changes according to the applied electrical field and with the concentration of the ions in the solution

For example, the mobility of the sodium ion (Na+) in water at 25 °C is 5.19 × 10−8 m2 V−1 s−1.[1] This means that a sodium ion in an electric field of 1 V/m would have an average drift velocity of 5.19 × 10−8 m/s. Such values can be obtained from measurements of ionic conductivity in solution.

Electrical mobility is proportional to the net charge of the particle. This was the basis for Robert Millikan's demonstration that electrical charges occur in discrete units, whose magnitude is the charge of the electron.

Electrical mobility is also inversely proportional to the Stokes radius a {\displaystyle a} of the ion, which is the effective radius of the moving ion including any molecules of water or other solvent which move with it.

Figure 3: Diagram of the separation of charged and neutral analytes (A) according to their respective electrophoretic and electroosmotic flow mobilities

For different ions with the same charge such as Li+, Na+ and K+ the electrical forces are equal, so that the drift speed and the mobility are inversely proportional to the radius a.[2] In fact, conductivity measurements show that ionic mobility increases from Li+ to Cs+ and therefore that Stokes radius decreases from Li+ to Cs+. This is the opposite of the order of ionic radii for crystals, and shows that in solution the smaller ions (Li+) are more extensively hydrated than the larger (Cs+). Ionic mobility(U) can be calculated as follows:

U= (x/t)x (1/E),

Whereas, U is the ionic mobility, x is the distance of movement, t, is the time of transformation and E is the applied potential.

U= x/tE, for K+, UK+= x/tE,

As e/t= speed then (V),

U= V/E and for K+, UK+=V/E,

V= UK+x E, this means that, Vα E while, Uα 1/E.

Ionic conductivity (µ): االيونيه التوصيليهis the summation of conductivity of its ionic species, i.e

µ= µ+ +µ-

For example Ionic conductivity of NaCl:

µNaCl= µNa+ + µCl-

there is a direct relation between ionic conductivity (µ) and ionic mobility(U) as follows:

µ= U . L, (L is specific conductance),

µ= (xA) .L/t.i, (I is the applied current, A is the area of the cell).

Example: when a passing a constant current of 1.6 mA in a solution of NaCl 0.2 M, with CdCl2 at 25 C, ions moved a distance of 10 cm in 3453 second in a cylinder with a volume of 0.1115 cm3, conductivity of NaCl was 2.313 x 10-3 Ω-1.cm-1. Find ionic kinetic for Na+?

Solution:

µNa+= (x.A.L)/(t.i),

µNa+= (10 cm. 0.1115 cm2. 2.313 x 10-3 Ω-1.cm-1)/

(3453 sec. 1.6 x 10-3 Amb),

µNa+= 4.67 x 10-3 volt-1. Sec-1.cm2

Transferring numbers: االنتقال أعداد

Transfer number (t), is a part of current that is transferred by a particular ion and normally it is equal to a unity for anion and cation of a desired electrolyte.

[ t+= (xA/ti). (CF/1000)],

[t-=(xA/ti). (CF/1000)], C is the concentration of ions, t+, t-: transfer number of cations and anions respectively. NOTE: (t++t-=1).

Example: using a solution of KCl(0.1 N), and a solution of LiCl (0.065 N), the density of applied current was 0.005893 Amb in 2130 sec, ionic movement was 5.6 cm and the area of cell was 0.1192 cm2. Find each of tK+, and tCl-?

Solution:

(t= (xACF)/(1000.t.i),

(t+= (5.6x 0.1192x 0.1 x 96500)/(2130x0.005893x1000),

(tNa+= 0.51), t++t-=1,

(tCl-= 1-0.51= 0.49).

Relation between specific conductance(L) and ionic mobility

Specific conductance(L) can be evaluated from ionic mobility of cations and anions as follows:

L= CF(U++ U-), C is the concentration of ions, F is faraday constant, U+ is the ionic mobility of cation and anion in (V-1, sec-

1.cm2)respectively.

Example: A solution of NaCl (0.1N), find its specific and equivalent conductance from ionic motilities of its ions(UNa+= 48.5 x 10—5 V-1.sec-1.cm2, Ucl-= 68.0 x10-5 V-1.sec -1.cm2)?

Solution:

L= FC((UNa++ UCl-),

L= 0.1x 96500(48.6 x10-5 + 68 x10-5)= 11.252 Ω-1.cm-1,

Λn=100L/C= 1000x 11.252/0.1= 112520 equ-1.Ω-1.cm2.

Application of electrical conductivity: التوصيلية تطبيقات , الكهربائية

1-finding of dissociation constant for weak electrolyte:

By plotting equation (5), Λ.C on y-axis against (1/Λ) gives a straight line with slope equal to (ka Λo

2) with interspet equal to (-kaΛo). the following sketch explains that:

Slope= kaΛo2

Intercept= -kaΛo

Λ.C

Figure-xx: finding ka of weak electrolyte using graphical method

2- Determination of ionic product of water:

Water is a weak electrolyte which is dissociated into H3O+ and OH- , water has a conductivity of 5.5 x10-6 equ-1.Ω-1.cm2 at 25 C.

Its equvalent conductance at infinite dilution is equal to the summation of the conductivity of OH- and H+.

Λo= ΛoH+ + Λo OH-,

Equavelent conductance of water is equal to:

Λn= 1000L/S, (S=C which is ionic product of water),

Kw= [H+].[OH-]= [S].[S]=[S2],

Kw of water= 1 x10-14

Example: calculate ionic product of water(kw) at 25 C, if it has specific conductance at 5.5 x 10-5 Ω-1.cm-1, and ionic conductance of H+ and OH-1 are 34.9 x 10-3 eq-1.ohm-1.cm2 and 19.9 x 10-3 eq-1.ohm-1.cm2 respectively?

Solution:

Note: S=C= [OH-]=[H+]

ΛoH2O=Λo

H3O+ +ΛoOH-,

ΛoH2O= 34.9 x 10-3 + 19.9 x 10-3= 54.87 x10-3 eq-1.ohm-1.cm2,

54.87 x10-3 = 1000L/S,= 1000 x 5.5 x 10-6/S ,

S= 1.002 x 10-7 eq-1/L= mol/lite,

Kw =[H+].[OH-]= S.S= S2= 1.004 x 10-14 mol2/liter2.

3-determination of hydrolysis constant of salts, ثابت تعيينلالمالح المائي التحلل

Weak salts that are derived from strong acid and weak base and inverse of that. These salts can hydrolysis with special value of hydrolysis constant (kh).

Herein, α is referring to the degree of dissociation of weak salt (HB) which is derived from weak base (B-) and strong acid (HCl). For this system the total equivalent conductance equal to:

Λsolution= (1-α)ΛBH+ α(ΛH3O+) + α(ΛB) , ----------------------1

Conductivity of weak base is very weak and can be ignored:

Λsolution= ΛBH –αΛBH + α(ΛH3O+) ----------------- 2

[α= (Λsolution- ΛBH)/(ΛBH+ ΛH3O+)],

ΛBH is the conductivity of the salt derived from weak base and strong acid. From equation (3), α fro the salt can be calculated and then dissociation constant (kh) can be calculated from:

[kh= (α2C)/(1-α), C is the concentration of the weak salt.

4-determination of solubility of sparingly soluble salts: تعيين الذوبان الشحيحة لالمالح التحلل ثابت

There are many rarely soluble salts such as PbS, AgCl, and BaSO4. Solubility constant for these salts can be calculated using conductivity measurements. For example for AgCl:

ΛAgCl= 1000L/C -----------1,

LAgCl= Lsolution – LH2O --------------2,

The relation between Λ and solubility (S) is:

Λ= 1000L/S, S=C,

Λ=1000L/C,

For this type of salt [ S=C, and Λ=Λo], Λo is the conductance at infinite dilution. S is the solubility of salt (gm/L, or equ/Liter or mol/Liter).

ΛAgCl= ΛAg+ + ΛCl-= 1000L/S,

S= 1000L/ ΛAg+ + ΛCl-,

S=1000L/ΛAgCl,

Example: specific conductance of AgCl(0.1N)at 25 C equal to 3.4 x 10-6 Ω-1.cm-1. Specific conductance of water (L=1.6 x 10-6 Ω-

1.cm-1)find solubility of this salt ( take Λo for AgCl = 1338.3 eq-

1.ohm-1.cm2).

Solution:

LAgCl= Lsolution –LH2O= 1.81 x 10-6 ohm-1.cm-1,

Λo=1000L/S,

S= 1000L/Λo ,

S= (1000 x 1.81 x 10-6)/(138.3)= 1.3 x10-5 gm.equ./Liter

S= 1.3 x 10-5 mol/L,

kSp= [Ag+][Cl-]= S2= 1.69 x 10-10 mol2/L2.

5- conductmetric titration: التوصيليه التسحيحات

A-titration of strong acid with strong base: for example titration of HCl with NaOH, at the beginning there is an excess of H+ in the solution, by adding OH- as a titrant concentration of H+ decreases which leads to reduce the conductance. After consumption of all H+, there will be an excess of OH- which leads to increase conductance again. From the obtained curve, the end point can be recorded as follows:

Figure xx: conductometric titration of strong acid (HCl) with strong base(NaOH)

b- titration of strong acid with strong base:

Acid in burette and base in the conical flask:

OH-

H+

NaOH/mL

Conductance(C-)

H+

OH Conductance(C-)

Figure xx: conductometric titration of strong base (NaOH) with strong acid(HCl)

C-titration of weak acid(HAC) with strong base(NaOH),

This case can be seen when titrating of a weak acid CH3COOH with strong base NaOH, herein we can see some increament in C- due to presence of H+ and CH3COO- from acetic acid. After that we can see sharp increase in C- due to exist of OH- which come from strong base. This can be shown in the following figure:

Figure xx: conductometric titration of weak acid (CH3COOH) with strong base(NaOH)

NaOH/mL

H+

Conductance(C-)

End point

OH-

D- titration of strong acid(HCl) with weak base(NH3),

Figure xx: conductometric titration of strong acid (HCl) with weak base(NH3)

E- titration of weaak acid(HOAC) with weak base(NH3),

Figure xx: conductometric titration of weak acid (CH3COOH) with weak base(NH3)

NH3/mL

End point

H+ Conductance(C-)

OH-

Cl-

NaOH/mL

End point

Conductance(C-)

OH-

H+

CH3COO-

F-precipitation titration: الترسيبيه التسحيحات

Precipitation titration can be used to follow end point for this type of titration. For example, titration of AgNO3 with KCl.:

AgNO3 (aq) + KCl(aq) AgCl(ppt)+ KNO3(aq)

By adding Cl- from burette as a titrant (KCl), this leads to react Cl- with Ag+ from solution which leads to reduce C-. after reaching end point there will be more free Cl- in solution which leads to increase C- again as shown below:

Figure xx: conductometric precipitation titration of AgNO3 with KCl

Electrochemical cells: الكهروكيميائيه الخاليا

It is a system that is composed of two main components, electrodes, anode and cathode. At the anode, oxidation processes would occur which leads to dissolve the electrode. At

End point

Conductance(C-) Cl-Ag+

the cathode, reduction processes occur which leads to growing up of the electrode.

An electrochemical cell (EC) is a device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions. There are two types of EC:

1- Voltaic cells or galvanic cells

The electrochemical cells which generate an electric current from chemical reactions are called voltaic cells or galvanic cells.

Figure xx: A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.

Figure xx: Galvanic cell with no cation flow

Voltaic cells use a spontaneous chemical reaction to drive an electric current through an external circuit. These cells are important because they are the basis for the batteries that fuel modern society.

2- The other ones are called electrolytic cells

It is also possible to construct a cell that does work on a chemical system by driving an electric current through the system. These cells are called electrolytic cells. Electrolysis is used to drive an oxidation-reduction reaction in a direction in which it does not occur spontaneously.which are used to drive chemical reactions like electrolysis. An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or

other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the electrodes, the ions in the electrolyte are attracted to an electrode with the opposite charge, where charge-transferring (also called faradaic or redox) reactions can take place. Only with an external electrical potential (i.e. voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided can produce a chemical reaction which would not occur spontaneously otherwise.

The Electrolysis of Molten NaCl

An idealized cell for the electrolysis of sodium chloride is shown in the figure below. A source of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride. Because the salt has been heated until it melts, the Na+ ions flow toward the negative electrode and the Cl- ions flow toward the positive electrode

When Na+ ions collide with the negative electrode, the battery carries a large enough potential to force these ions to pick up electrons to form sodium metal.

Negative electrode (cathode): Na+ + e- Na

Cl- ions that collide with the positive electrode are oxidized to Cl2 gas, which bubbles off at this electrode.

Positive electrode (anode): 2 Cl- Cl2 + 2 e-

The net effect of passing an electric current through the molten salt in this cell is to decompose sodium chloride into its elements, sodium metal and chlorine gas.

Electrolysis of NaCl:

Cathode (-): Na+ + e- Na

Anode (+): 2 Cl- Cl2 + 2 e-

The potential required to oxidize Cl- ions to Cl2 is -1.36 volts and the potential needed to reduce Na+ ions to sodium metal is -2.71 volts. The battery used to drive this reaction must therefore have a potential of at least 4.07 volts.

This example explains why the process is called electrolysis. The suffix -lysis comes from the Greek stem meaning to loosen or split up. Electrolysis literally uses an electric current to split a compound into its elements.

electrolysis

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

This example also illustrates the difference between voltaic cells and electrolytic cells. Voltaic cells use the energy given off in a spontaneous reaction to do electrical work. Electrolytic cells use electrical work as source of energy to drive the reaction in the opposite direction.

The dotted vertical line in the center of the above figure represents a diaphragm that keeps the Cl2 gas produced at the anode from coming into contact with the sodium metal generated at the cathode. The function of this diaphragm can

be understood by turning to a more realistic drawing of the commercial Downs cell used to electrolyze sodium chloride shown in the figure below.

Types of electrodes

1-anode:

An anode is an electrode through which the conventional current enters into a polarized electrical device. This contrasts with a cathode, an electrode through which conventional current leaves an electrical device. A common mnemonic is ACID for "anode current into device".[1] The direction of conventional current (the flow of positive charges) in a circuit is opposite to the direction of electron flow, so (negatively charged) electrons flow out the anode into the outside circuit. In a galvanic cell, the anode is the electrode at which the oxidation reaction occurs.

Figure: Diagram of a zinc anode in a galvanic cell. Note how electrons move out of the cell, and the conventional current moves into it in the opposite direction.

2- cathode:

A cathode is the electrode from which a conventional current leaves a polarized electrical device. (This definition can be recalled by using the mnemonic CCD for cathode current departs.) A conventional current describes the direction in which positive electronic charges move. Electrons have a negative electrical charge, so the movement of electrons is opposite to that of the conventional current flow (consequently, the mnemonic cathode current departs also means that electrons flow into the device's cathode

Figure xx: Diagram of a copper cathode in a galvanic cell (e.g., a battery). Positively charged cations move towards the cathode allowing a positive current (i) to flow out of the cathode.

Types of electrodes:

An electrode is an electrical conductor[ metal], used to make contact with a nonmetallic part of a circuit (e.g. a semiconductor, an electrolyte, a vacuum or air). The word was coined by William Whewell at the request of the scientist

Michael Faraday from two Greek words: elektron, meaning amber (from which the word electricity is derived), and hodos, a way. main types of electrodes are:

1-first type electrode: األول النوع اقطاب

Electrode of the first kind is a simple metal electrode immersed in a solution containing its own ion (e.g., silver immersed in a silver nitrate solution). The equilibrium potential of this electrode is a function of the concentration (more correctly of activity) of the cation of the electrode metal in the solution (see Nernst’s electrode potential equation).for gas electrode Pt is used as an auxiliary electrode.

Examples of this type are:

2-Second type electrode: الثاني النوع أقطاب

It’s a metal(Ag) immersed in a rarely soluble its solution (AgCl) which is immersed in completely another soluble solution containing negative ion of the first electrolyte (NaCl).

Zn metal

First type electrode

ZnSO4 solution

Electrodes of the Second Kind An electrode of the first kind involving an M"+/M redox couple will respond to the concentration of another species if that species is in equilibrium with M"+

The main reactions of 2nd type electrode are:

AgCl solution

Second type electrode

Ag metal

NaCl solution

Beside silver/silver chloride electrode, there is another electrode of 2nd type, its calomel electrode, Hg electrode:

It reactions involve:

3-Third type electrode: الثالث النوع : أقطاب

Its a type of electrode of a metal in contact with two sparingly soluble salts with contact with three completely soluble salt.

CaC2O4 solution

Third type electrode

CaCl2 solution

PbC2O4 solution

Pb metal

Ag Ag+ + e

Ag++Cl- AgCl

Ag +Cl- AgCl + e (oxidation reaction

in case of reduction reaction:

AgCl Ag++Cl-

Ag+ +e Ag

AgCl +e Ag(s)+Cl- (reduction reaction of the electrode

Herein , PbC2O4, CaC2O4 are sparingly soluble salts while CaCl2 is completely dissolved salt.

Reactions of this electrode involve:

4- oxidation-reduction electrodes: االكسده اقطابواالختزالOxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

Figure xx: Sodium and fluorine bonding ionically to form sodium fluoride. Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom exothermically. The oppositely charged ions are then attracted to each other. The sodium is oxidized, and the fluorine is reduced.

In general, redox electrodes, require presence of a metal that can exist in two oxidation states such as Fe2+ and Fe3+, Sn2+ and Sn4+ etc---.

Reaction of reversible cells: العكوسه الخاليا تفاعالت

In this type of electrochemical cells, electrons would produce in one half of cell( usually anode). These electrons are consumed at the second half of the same cell (usually cathode). Important type of this cell is galvanic cell.

Herein, Zn electrode representing anode electrode, while Cu electrode representing cathode electrode and their reactions as follow:

Basic description

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

Schematic of Zn-Cu galvanic cell

In its simplest form, a half-cell consists of a solid metal (called an electrode) that is submerged in a solution; the solution contains cations (+) of the electrode metal and anions (−) to balance the charge of the cations. The full cell consists of two half-cells, usually separated by a semi-permeable membrane or by a salt bridge.

A specific example is the Daniell cell, where a salt bridge is used as separator (see figure). The zinc (Zn) half-cell has a solution of ZnSO4 (zinc sulfate) and the copper (Cu) half-cell has a solution of CuSO4 (copper sulfate).

Let an external electrical conductor connect the copper and zinc electrodes. In the zinc half-cell, zinc from the zinc electrode dissolves into the solution as Zn2+ ions (oxidation), releasing electrons that enter the external conductor. In addition, via the salt bridge zinc ions leave and sulfate ions (SO2−

4) enter the zinc half-cell.

In the copper half-cell, the copper ions plate onto the copper electrode (reduction), taking up electrons that leave the

external conductor. Since the Cu2+ ions (cations) plate onto the copper electrode, the latter is called the cathode. Correspondingly the zinc electrode is the anode. The electrochemical reaction is:

Zn + Cu2+ → Zn2+ + Cu

In addition, electrons flow through the external conductor, which is the primary application of the galvanic cell.

As discussed under #Cell voltage, the emf of the cell is the difference of the half-cell potentials, a measure of the relative ease of dissolution of the two electrodes into the electrolyte. The emf depends on both the electrodes and on the electrolyte, an indication that the emf is chemical in nature.

From above equation , potential of cell (E), can be calculated using Nernst equation as follows:

E=Eo- 2.303RT/nF. Log [product]/[reactant],

E, Eo, is the potential and standard potential of the cell, n number of electrons, and F is the Faraday constant.

.

Nernst equation,: نيرنست معادلة

is an equation that is connected between cell potential and the activity or concentration of redox species

At any specific temperature, the Nernst equation derived above can be reduced into a simple form. For example, at the standard condition of 298 K (25°), the Nernst equation becomes for a reaction:

aA+bB=cC+dDgeneral Nernst equation correlates the Gibb's Free Energy G and the EMF of a chemical system known as the galvanic cell. For the reaction

a A + b B = c C + d D and

[C]c [D]d

Q = --------- [A]a [B]b

It has been shown that G = G° + R T ln Q and G = - n FE.

Therefore - n F E = - n F E° + R T ln Q

where R, T, Q and F are the gas constant (8.314 J mol-1 K-1), temperature (in K), reaction quotient, and Faraday constant (96485 C) respectively. Thus, we have

R T [C]c [D]d

E = E° - ----- ln --------- n F [A]a [B]b

This is known as the Nernst equation. The equation allows us to calculate the cell potential of any galvanic cell for any concentrations. Some examples are given in the next section to illustrate its application.

It is interesting to note the relationship between equilibrium and the Gibb's free energy at this point. When a system is at equilibrium, E = 0, and Qeq = K. Therefore, we have,

R T [C]c [D]d

E° = ----- ln ---------, (for equilibrium concentrations) n F [A]a [B]b

Thus, the equilibrium constant and E° are related.

The Nernst Equation at 298 K

At any specific temperature, the Nernst equation derived above can be reduced into a simple form. For example, at the standard condition of 298 K (25°), the Nernst equation becomes

0.0592 V [C]c [D]d

E = E° - --------- log --------- n [A]a [B]b

Please note that log is the logrithm function based 10, and ln, the natural logrithm function.

For the cell

Zn | Zn2+ || H+ | H2 | Pt we have a net chemical reaction of

Zn(s) + 2 H+ = Zn2+ + H2(g) and the standard cell potential E° = 0.763.

If the concentrations of the ions are not 1.0 M, and the H2 pressure is not 1.0 atm, then the cell potential E may be calculated using the Nernst equation:

0.0592 V P(H2) [Zn2+] E = E° - ------- log ------------ n [H+]2

with n = 2 in this case, because the reaction involves 2 electrons. The numerical value is 0.0592 only when T = 298 K. This constant is temperature dependent. Note that the reactivity of the solid Zn is taken as 1. If the H2 pressure is 1 atm, the term P(H2) may also be omitted. The expression for the argument of the log function follows the same rules as those for the expression of equilibrium constants and reaction quotients.

Indeed, the argument for the log function is the expression for the equilibrium constant K, or reaction quotient Q.

When a cell is at equilibrium, E = 0.00 and the expression becomes an equilibrium constant K, which bears the following relationship:

n E° log K = -------- 0.0592where E° is the difference of standard potentials of the half cells involved. A battery containing any voltage is not at equilibrium.

The Nernst equation also indicates that you can build a battery simply by using the same material for both cells, but by using different concentrations. Cells of this type are called concentration cells.

Example 1:

Calculate the EMF of the cell Zn(s) | Zn2+ (0.024 M) || Zn2+ (2.4 M) | Zn(s)

Solution

Zn2+ (2.4 M) + 2 e = Zn ReductionZn = Zn2+ (0.024 M) + 2 e Oxidation--------------------------------------------Zn2+ (2.4 M) = Zn2+ (0.024 M), E° = 0.00 - - Net reaction

Using the Nernst equation:

0.0592 (0.024) E = 0.00 - ------- log -------- 2 (2.4)

= (-0.296)(-2.0) = 0.0592 V

Example 2

Show that the voltage of an electric cell is unaffected by multiplying the reaction equation by a positive number.

SolutionAssume that you have the cell

Mg | Mg2+ || Ag+ | Ag and the reaction is:

Mg + 2 Ag+ = Mg2+ + 2 Ag Using the Nernst equation

0.0592 [Mg2+] E = E° - ------ log -------- 2 [Ag+]2

If you multiply the equation of reaction by 2, you will have 2 Mg + 4 Ag+ = 2 Mg2+ + 4 Ag

Note that there are 4 electrons involved in this equation, and n = 4 in the Nernst equation:

0.0592 [Mg2+]2

E = E° - ------ log -------- 4 [Ag+]4

which can be simplified as

0.0592 [Mg2+] E = E° - ------ log -------- 2 [Ag+]2

Thus, the cell potential E is not affected.

Example 3

The standard cell potential dE° for the reaction Fe + Zn2+ = Zn + Fe2+

is -0.353 V. If a piece of iron is placed in a 1 M Zn2+ solution, what is the equilibrium concentration of Fe2+?

SolutionThe equilibrium constant K may be calculated using

K = 10(n E°)/0.0592

    = 10-11.93

    = 1.2x10-12

    = [Fe2+]/[Zn2+]. Since [Zn2+] = 1 M, it is evident that        [Fe2+] = 1.2E-12 M.

Example 4

From the standard cell potentials, calculate the solubility product for the following reaction:

AgCl = Ag+ + Cl-

SolutionThere are Ag+ and AgCl involved in the reaction, and from the table of standard reduction potentials, you will find:

AgCl + e = Ag + Cl-, E° = 0.2223 V - - - -(1)Since this equation does not contain the species Ag+, you need,

Ag+ + e = Ag, E° = 0.799 V - - - - - - (2)Subtracting (2) from (1) leads to,

AgCl = Ag+ + Cl- . . . E° = - 0.577Let Ksp be the solubility product, and employ the Nernst equation,

log Ksp = (-0.577) / (0.0592) = -9.75Ksp = 10-9.75 = 1.8x10-10

This is the value that you have been using in past tutorials. Now, you know that Ksp is not always measured from its solubility.

Potential of electrode: القطب جهد

Electrode potential, E, in chemistry or electrochemistry, according to a IUPAC definition,[1] is the electromotive force of a cell built of two electrodes: on the left-hand side of the cell diagram is the standard hydrogen electrode (SHE), and on the right-hand side is the electrode in question.The SHE is defined to have a potential of 0 V, so the signed cell potential from the above setup is

Ecell = Eleft (SHE) − Eright = 0 V − Eelectrode = Eelectrode.

SHE is cathode and electrode is anode.

Electrode potential appears at the interface between an electrode and electrolyte due to the transfer of charged species across the interface, specific adsorption of ions at the interface, and specific adsorption/orientation of polar molecules, including those of the solvent.

Electrode potential is the electric potential on an electrode component. In a cell, there is an electrode potential for the cathode and an electrode potential for the anode. The difference between the two electrode potentials equals the cell potential:

[Ecell = Ecathode − Eanode.].

Standard reduction potential tableMain article: Standard electrode potential (data page)

The larger the value of the standard reduction potentials, the easier it is for the element to be reduced (accept electrons); in other words, they are better oxidizing agents. For example, F2 has 2.87 V and Li+ has −3.05 V. F reduces easily and is therefore a good oxidizing agent. In contrast, Li(s) would rather undergo oxidation (hence a good reducing agent). Thus Zn2+ whose standard reduction potential is −0.76 V can be oxidized by any other electrode whose standard reduction potential is greater than −0.76 V (e.g. H+(0 V), Cu2+(0.34 V), F2(2.87 V)) and can be reduced by any electrode with standard reduction potential less than −0.76 V (e.g. H2(−2.23 V), Na+(−2.71 V), Li+(−3.05 V)).

In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, Gibbs free energy ΔG° must be negative, in accordance with the following equation:

ΔG°cell = −nFE°cell

where n is number of moles of electrons per mole of products and F is the Faraday constant, ~96485 C/mol. As such, the following rules apply:

If E°cell > 0, then the process is spontaneous (galvanic cell)If E°cell < 0, then the process is nonspontaneous (electrolytic cell)

Thus in order to have a spontaneous reaction (ΔG° < 0), E°cell must be positive, where:

E°cell = E°cathode − E°anode

where E°anode is the standard potential at the anode and E°cathode is the standard potential at the cathode as given in the table of standard electrode potential.

Standard electrodes: المرجعيه األقطاب1.Standard hydrogen electrode: الهيدروجين قطبالقياسيThe Standard hydrogen electrode (abbreviated SHE), is a redox electrode which forms the basis of the thermodynamic scale of oxidation-reduction potentials. Its absolute electrode potential is estimated to be 4.44 ± 0.02 V at 25 °C, but to form a basis for comparison with all other electrode reactions, hydrogen's standard electrode potential (E0) is declared to be zero volts at all temperatures.[1] Potentials of any other

electrodes are compared with that of the standard hydrogen electrode at the same temperature.

Hydrogen electrode is based on the redox half cell:

2 H+(aq) + 2 e− → H2(g)

This redox reaction occurs at a platinized platinum electrode. The electrode is dipped in an acidic solution and pure hydrogen gas is bubbled through it. The concentration of both the reduced form and oxidised form is maintained at unity. That implies that the pressure of hydrogen gas is 1 bar (100 kPa) and the activity of hydrogen ions in the solution is unity. The activity of hydrogen ions is their effective concentration, which is equal to the formal concentration times the activity coefficient. These unit-less activity coefficients are close to 1.00 for very dilute water solutions, but usually lower for more concentrated solutions. The Nernst equation should be written as:

SHE

SHE is composed of a 1.0 M H+(aq) solution containing a square piece of platinized platinum (connected to a platinum wire where electrons can be exchanged) inside a tube. Platinum is used because it is inert and does not react much with hydrogen .During the reaction, hydrogen gas is then passed through the tube and into the solution causing the reaction:

 2H+(aq) + 2e- <==> H2(g) [ Eo= 0.00 Volt]

2- Saturated Calomel electrode: الكالوميل قطبThe Saturated calomel electrode (SCE) is a reference electrode based on the reaction between elemental mercury and mercury(I) chloride. It has been widely replaced by the silver chloride electrode, however the calomel electrode has a reputation of being more robust. The aqueous phase in contact with the mercury and the mercury(I) chloride (Hg2Cl2, "calomel") is a saturated solution of potassium chloride in water. The electrode is normally linked via a porous frit to the solution in which the other electrode is immersed. This porous

frit is a salt bridge .

SCE potential

The only variable in this equation is the activity (or concentration) of the chloride anion. But since the inner solution is saturated with potassium chloride, this activity is fixed by the solubility of potassium chloride, which is :@ 20 °C. This makes the SCE redox potential +0.248 V at vs. SHE at 20 °C or +0.2444 V vs. SHE at 25 °C,[1] but slightly higher when the chloride solution is less than saturated. For example, a 3.5M KCl electrolyte solution increases the reference potential to +0.238 V vs. SHE at 25 °C, and a 0.1 M solution to +0.3356 V at the same temperature.

Figure xx: SCE apparatus

A calomel electrode is a reference electrode that is based on reactions between mercury (I) chloride (calomel) and elemental mercury. The aqueous phase in contact with both the calomel and the mercury is a saturated solution of water and potassium chloride. The linking of the electrode is through a porous frit to the solution that contains the other electrode. Its structure consists of an outer glass tube that is fitted with a frit at the bottom. This permits electrical contact with the solution on the outside. Another tube is fitted on the inside of the first tube. The bottom of this inner tube has glass wool at the bottom to allow for further connection of electricity between the contents of both tubes.

In calomel electrodes, mercury paste is packed on the inner tube, with mercurous chloride being dispersed in a saturated potassium chloride solution. This can be represented as:

Hg|Hg2Cl2KCl (xM) saturated

The electrode reaction is:

Hg2Cl2 + 2e Hg2Cl2 == 2Hg + 2Cl-

When the electrode is immersed in a solution, there is an electric contact made between the electrolyte and the sample at a certain opening located near the end of the electrode. This forms a conductive bridge between the sample, the reference electrode and the indicating electrode.

The glass body liquid-filled types include the porous ceramic junctions used for routine applications. Cracked bead junctions are used for samples that require slow electrolyte flow, while sleeve junctions require fast electrolyte flow. The polymer-bodied types of electrodes include electrodes such as the liquid-filled with a junction, ceramic junction and the gel-filled with porous polymer junction. The permanent gel-filled type do not require a lot of maintenance and can be applied to numerous routine applications.

The electrode potential is reported with the concentration of potassium chloride, because the electrode potential depends on the potassium chloride concentration. The common reference voltage for a saturated calomel electrode is +0.244 V. Mercurous chloride can disproportion to mercury and mercuric chloride if supplied with a temperature of above 80°C, thus restricting the operation of the calomel electrode below 80°C. Its design and thermal characteristics make the rate of change of progress slow. They can be used in pH measurements and general aqueous electrochemistry.

3-silver-silver chloride electrode:A silver chloride electrode is a type of reference electrode, commonly used in electrochemical measurements. For environmental reasons it has widely replaced the saturated calomel electrode. For example, it is usually the internal reference electrode in pH meters and it is often used as reference in reduction potential measurements. As an example of the latter, the silver chloride electrode is the most commonly used reference electrode for testing cathodic protection corrosion control systems in sea water environments.

Ag metal

The electrode functions as a redox electrode and the equilibrium is between the silver metal (Ag) and its salt—silver chloride (AgCl, also called silver(I) chloride).

The corresponding half-reactions can be presented as follows:

Ag Ag+ + e

Ag++Cl- AgCl

Ag +Cl- AgCl + e (oxidation reaction

in case of reduction reaction:

AgCl Ag++Cl-

Ag+ +e Ag

AgCl +e Ag(s)+Cl- (reduction reaction of the electrode

This reaction is characterized by fast electrode kinetics, meaning that a sufficiently high current can be passed through the electrode with the 100% efficiency of the redox reaction (dissolution of the metal or cathodic deposition of the silver-ions). The reaction has been proven to obey these equations in solutions of pH values between 0 and 13.5.

The Nernst equation below shows the dependence of the potential of the silver-silver(I) chloride electrode on the activity or effective concentration of chloride-ions:

E= Eo- RT/nFLn[Cl-1]

The standard electrode potential E0 against standard hydrogen electrode (SHE) is 0.230V ± 10mV. The potential is however very sensitive to traces of bromide ions which make it more negative. (The more exact standard potential given by an IUPAC review paper is +0.22249 V, with a standard deviation of 0.13 mV at 25 °C

4-Glass electrode: الزجاج قطب

A glass electrode is a type of ion-selective electrode made of a doped glass membrane that is sensitive to a specific ion. The most common application of ion-selective glass electrodes is for the measurement of pH. The pH electrode is an example of a glass electrode that is sensitive to hydrogen ions. Glass electrodes play an important part in the instrumentation for chemical analysis and physico-chemical studies. The voltage of the glass electrode, relative to some reference value, is

sensitive to changes in the activity of certain type of ions. The pH range at constant concentration can be divided into 3 parts:

Scheme of the typical dependence E (Volt) – pH for ion-selective electrode.[citation needed]

Complete realization of general electrode function, where potential depends linearly on pH, realizing an ion-selective electrode for hydronium.

E=Eo- 2.3030RT/Flog[H+],

E=Eo+ 2.303RT/FpHA glass electrode is an electrode that consists typically of a glass tube, sealed at the bottom, with a thin-walled glass bulb containing a solution of constant pH (as a chloride buffer) and a silver-silver chloride reference electrode. This is immersed in an unknown solution, usually

along with a calomel electrode, for determining the pH of this solution.

Figure xx: image of glass electrode(pH) electrode.

Electrical batteries: الكهربائية البطاريات

Electrical Cells and BatteriesIn it's most simple form a battery can be regarded as a pump that provides the energy

to move charge around a circuit.

In order to provide a potential difference, or electro-motive force (EMF) a store of energy

is required. One such method is a battery or cell. The common usages of the term

battery is any device that converts chemical energy into electrical energy. However,

strictly speaking, the term battery is used when several electrical cells are connected

together to provide a source of a potential difference in a circuit. If it is just a single

chemical source then it is called a cell.

History

Figure 1. Galvani's Frog's Leg Experiment

In 1791, Galvani noticed that a circuit created with two different metals, when touched

on the ends of the leg of a dead frog, would cause it twitch. The two metals were

creating an electric current within the frog's leg, causing the muscles to contract.

Early batteries were an improvement of this method transfering chemical energy into

electrical energy.

The first battery was invented in 1793 by Alessandro Volta. Just as the two different

metals touching the wet skin of a frog's leg, caused an electrical current to flow, early

batteries increased the voltage that could be produced by stacking a pile of discs

made from silver and zinc sandwiched between paper soaked in a salt water solution

as shown in Figure 2. In honour of Volta, we use the Volt as the unit of potential

difference and EMF.

Figure 2. Volta's battery of cells.

Battery OperationWhy does this produce electricity? The flow of current can be understood as the flow

of ions from the more reactice metal to the less reactive metal. The ions moving from

one electrode to the other creates an electrical charge which is neutralised by the flow

of electrons across the wire.

Before considering the reaction of two metals, consider what happens when we place

a single metal electrode in an electrolyte. Some of the metal atoms in the electrolyte

go into solution as ions while the remaining electrons create a negative charge on the

metal. The separation of ions and electrons leads to a separation of charge. However,

this build up of charge cannot continue indefinitely because as the negative charge

builds up in the metal it becomes increasingly difficult for positive metal ions to go

into solution. A similar build up in positive charge in the electrolyte also prevents the

build up of charge. This degree of charge build up depends on the metal and

represents the work required to separate electrons from the ions. This is known as the

electroneutrality principle

Similarly, if a copper strip is placed in an aquaous Copper(II)Sulfate solution the

copper will also lose ions. These reactions are often written as Cu | Cu+2 this is the

half-cell reaction.

The tendancy for Zinc to lose ions is greater than that of Copper. When the two cells

are joined together (using a copper wire to connect the electrodes and porous barrier

that allows the ions to pass known as a salt-bridge connect the elecrolytes, the build

up of electrons on the zinc will flow to through the wire onto the copper.

The copper ions in the electrolyte gain electrons and become copper atoms. Thus the

reaction can be written,

Zn | Zn2+ | | Cu2+ | Cu

Figure 3. Danile Cell

To continue the reaction, the charge must be removed. This can be acheived by

coupling a second reaction which uses the electrons in the metal to convert the ions in

the electrolyte into a metal. For a more specific example, consider a zinc electrode in

an electrolyte of Copper(II)Sulphate solution.

The loss of electrons by the Zinc is known as oxidation. Zn(s) → Zn2+ + 2e-.(1)

A wire connecting the Zinc electrode to a Copper electrode, allows the electrons to

flow to the Copper electrode. Copper ions in the copper sulphate solution take up the

electrons and become atoms of copper on the copper electrode. The gaining of

electrons by ions is known as reduction

Cu+2 + 2e- → Cu(s).(2)

The net reactions is then,

Zn(s) + Cu2+ → Cu(s) + Zn2+

When the two electrodes are joined by a wire the charge stored can flow and the

electrons combine. The simplest kinds of battery have two conductors made of

different materials which are partially emersed in a solution which allow the electrons

and ions to flow freely known as an electrolyte.

At the copper electrode (cathode), the acid dissolves the copper metal producing

hydrogen gas, H+. The reaction will continue until the supply of zinc is used up. The

electrons, with their negative charge, are attracted to the copper electrode which

causes a current to flow. One of the problems with this cell is that the current stops

flowing after a short time because the hydrogen bubbles block the current. Cells using

aqueous (containing water) electrolytes are limited in voltage to less than 2 Volts

because the oxygen and hydrogen in water dissociate in the presence of voltages

above this voltage. Lithium batteries (see below) which use non-aqueous electrolytes

do not have this problem and are available in voltages between 2.7 and 3.7 Volts.

However the use of non-aqueous electrolytes results in those cells having a relatively

high internal impedance.

Activity SeriesIn the activity series, a metal will give up electrons to any other metal which is below

it on the activity series. Elements which gain electrons are called negative irons.

Elements which lose electrons are called positive ions. Any more active metal will give

up electrons to a less active metal. This will serve to protect the less active metal from

corrosion. For example, steel ships often have bars of zinc attached to the sides of the

ship. As the steel is corroded by the oxygen of water and air, the zinc will give up

electrons to the steel and protect it from corrosion.

Batteries in CircuitsJust like any other electrical component, individual cells can be placed in series or

parallel. In series their voltage sum to create a battery with a higher voltage but the

current remain the same as in a single cell. In parallel, the batteries have the same

voltage but the current is summed to create a battery with a higher current.

Batteries and Internal Resistance

Figure 4. A real cell with internal resistance.

A real battery has internal resistance, r, which lowers the voltage when the cell is

connect to a load. If you try to send too much current through a battery, the internal

resistance will convert the battery’s own chemical potential energy into thermal

energy. The battery will get warm and electrons will leave the negative electrode with

relatively little energy.

Figure 5. A cell with internal resistance in series with a load, R.

The EMF of the battery is given by E and represents the voltage when the battery is

open circuit. If the battery had no internal resistance, the current flowing in the circuit

would be given by Ohm's law, I =E/R where R is the resistance of the load.

The internal resistance adds to the total resistance of the circuit. If we call the new

total resistance of the cell and load, Req = r + R

Then current is then I= E/Req = E/(r + R).(3)

Since Req is larger than R the current flowing in the circuit is reduced. Therefore, if we

are making a cell we want the internal resistance to be a small as possible.

What is the voltage across the resistor R? Clearly it is not the same as the EMF

because of the internal resistance. We know that the EMF, E = Ir + IR. We also know

that the voltage across the resistor R is IR. Rearranging these two equations we find

the voltage in terms of E, r and R.

V = E - Ir (4).

This equation also shows that if we draw a lot of current from the circuit we see the

voltage reduces. A good example of this is a car starting with its lights switched on.

When the ignitioin key is turned, the electric starter-motor uses a lot of current before

it starts to turn. The current for this is supplied by the car battery, which also powers

the lights. The voltage available for the lights is reduced at this moment and the lights

dim as the car starts. They quickly become bright again once the engine is turning

over because less current is required by the starter motor. (When the car has started

no current is drawn by the starter motor because it disengages from the engine.)

Metal Metal Ion Reactivity

Lithium Li+ Most Reactive

Potassium K+

Calcium Ca2+

Sodium K+

Magnesium Mg2+

Aluminum Al3+

Manganese Mn2+

Zinc Zn2+

Chromium Cr2+, Cr3+

Iron Fe2+, Fe3+

Lead Pb2+

Copper Cu2+

Mercury Hg2+

Silver Ag2+

Gold Au+,Au3+ Least Reactive

Platinum Pt2+

Table 1. Metal Reactivity Series from most reactive to least reactive.

Power and EfficiencyThe power output of the battery is given by PE = I2R(5)

But from equation (3), I = E/Req = E/(r + R), therefore, P = E2R/(r + R)2(6)

Similarly, the power given off as heat by the battery to due to internal resistance is Pr

= I2r = E2r/(r+R)2(7)

The efficiency, η of the circuit is the ratio of the power actually produced by the

battery with its internal resistance to the power supplied by the source, P0. P0 = I E

The power generated is given by I VR, where VR is the voltage across the resistor R.

η = PE/P0 = IVR/IE = R/(r + R)

The smaller the internal resistance, the closer the efficiency, η will be to its maximum

value of 1.

If we plot the power tranferred to a load resistance R against the increasing R along

with the efficiency we find that the maximum power is transferred by the battery

when R = r. This is a very important result and find applications in many electrical

devices is known as Jacobi's theorem

Power transferred and efficiency of a battery against load resistance.

To check this is true we can also differentiate the expression for the power P against R

and set it equal to zero to find the maximum value of R.

dP/dR = -dE2/dR (r + R)-2 + d(R+r)-2/dR (E2R) = 0

E2/(r + R)2 - 2 E2R/(r + R)3 = 0

(r + R)3 - 2R(r + R)2 = 0

(r + R) - 2R = 0

r - R = 0 or r = R

While the power transferred may be at a maximum, the efficiency at this percent is

only 50%. The higher the load resistance, the greater the efficiency. In practise the

exact loading of the circuit is dependent on the application, a good voltmeter has an

extremely high-resistance so that the power transmitted is as small as possible.

Types of BatteryOver the years, progress in battery technology has been rather slow but the need for

small more powerful batteries in the many small electrical items we carry around with

us has driven research into higher power, longer lasting batteries.

1-ZINC CARBON

This is commonly known as the Leclanché Cell and despite being the oldest type of

battery it is still the most commonly used as it is very low-cost. Traditional Zinc

Carbon batteries cannot be reused when their chemical energy has been released

Zinc-Carbon Cell

ALKALINE CEL-LS

Alkaline chemistry is used in common Duracell and Energizer batteries, the electrodes

are zinc and manganese-oxide, with an alkaline electrolyte. Alkaline batteries can be

re-used upto 100 times with the correct type of battery charger. A normal battery

charger must not be used to charge these batteries.

The active materials used are the same as in the Leclanché cell – zinc and manganese

dioxide. However the electrolyte is potassium hydroxide, which is very conductive,

resulting in low internal impedance for the cell. This time the zinc anode does not form

the container; it is in the form of a powder instead, giving a large surface area.

2-SILVER ZINC

Lightweight but expensive. Used in aeronautical applications.Silver Zinc battery is a  type of rechargeable battery with silver oxide , zinc oxide as electrodes and alkaline electrolyte. Silver Zinc batteries have been mainly used in military equipment and various space ships and submarines because of their low self discharge rate, high energy density and reliability.

Working principle of Silver Zinc Battery:The charging and discharging of a Silver Zinc Battery or cell may be represented by a single reversible equation given below:

When the cell is fully charged the positive plate is Monovalent Silver Oxide (Ag2O) and negative plate is Zinc(Zn), During discharging; the Positive Silver Oxide plate is reduced to Silver (Ag) and the Negative Zinc plate is oxidized to Zinc Oxide (Zn). The chemical process is reversed when charging after the cell is discharged.

Rechargable or Secondary CellsRechargeable batteries are rechargeable because the chemical reaction that leads to

the flow of current is reversible by passing a current through the battery. The

animation shows a battery undergoing charging and discharging. When the battery is

charged the current can flow through a resistive load.

.

LEAD-ACID BATTERY

Lead acide batteries are used to provide large amounts of current for a relatively short

time. They consist of plates of lead and lead oxide in a solution of sulphuric acid. Lead

combines with SO4 (sulphate) to create PbSO4 (Lead Sulphate), plus one electron.

Lead dioxide, hydrogen ions and SO4 ions, plus electrons from the lead plate, create

PbSO4 and water on the lead dioxide plate. As the battery discharges, both plates build

up PbSO4 and water builds up in the acid. The characteristic voltage is about 2 volts

per cell, so by combining six cells you get a 12-volt battery.

half-reaction V vs SHE

Pb + SO42- → PbSO4 + 2e- .356

PbO2 + SO42- + 4H+ + 2e- → PbSO4 + 2H2O 1.685

N ICKLE CADMIUM (NI -CAD) BATTERY

Nickle-Cadmium cells are the most common type of re-chargable battery. They have a

high-energy density and a EMF of 1.2 V. They can be recharged more times than other

types of rechargable batteries but unless they are fully discharged before recharging

suffer from a memory effect which reduces their capacity to store charge.

NICKLE-METAL HYDRIDE (NI -MH)

L ITH IUM ION (L I - ION)

Lithium is the most electronegative metal in the electrochemical series. It also has a

low density so it is an atractive material for the anode of batteries. However, lithium is

also very reactive with aqueous electrolytes producing hydrogen gas. Because of this

it took many years to develop a stable electrolyte. Lithium batteries must be sealed

from moisture and air due to the reactivity of lithium. Non-rechargable lithium

batteries have been available since the 1980s with rechargable lithium batteries

becoming widely available around 1995. Lithium-ion batteries can be recharged

between 500 - 1000 cycles. The half-reactions are:

The following reactions take place upon discharge:

Anode: xLi+ + Mn2O4 → LixMn2O4

Cathode: LixC6 → xLi+ + 6C + xe-

Overall: LixMn2O4 + 6C → LixC6 + Mn2O4

LITH IUM POLYMER (L I -POLY)

Lithium polymer batteries use a solid polymer electrolyte.

RUSTING

Corrosion of iron and steel due to rusting is responsible for millions of pounds of

damage each year. Rustiing, is oxidation of the metal to form a metal oxide. Rust does

not firmly adhere to the surface of the metal allowing it oxide further. The oxide

causes damage to the surface of the metal known as pitting which, over time, reduces

thes structural integrity of the metal.

What has this got to do with batteries? Rusting is a chemical process the occurs when

the iron or steel is exposed to moist air, it reacts with the oxygen in the air to create

Iron (III) oxide. We saw earlier how electricity is generated by the process of

oxidization and reduction. The formation of rust can occur at some distance away from

theactual pitting or erosion of iron as illustrated below. This is possible because the

electrons produced via the initial oxidation of iron can be conducted through the metal

and the iron ions can diffuse through the water layer to another point on the metal

surface where oxygen is available. This process results in an electrochemical cell in

which iron serves as the anode, oxygen gas as the cathode, and the aqueous solution

of ions serving as a "salt bridge" as shown below.

Rusting of Iron by water droplet.

Fe → Fe+2 + 2e- and Fe → Fe+3 + e- in the anode

The amount of water complexed with the iron (III) oxide (ferric oxide) varies as

indicated by the letter "X". The amount of water present also determines the color of

rust, which may vary from black to yellow to orange brown. The formation of rust is a

very complex process which is thought to begin with the oxidation of iron to ferrous

(iron "+2") ions.

Fe → Fe+2 + 2 e-

Both water and oxygen are required for the next sequence of reactions. The iron (+2)

ions are further oxidized to form ferric ions (iron "+3") ions.

Fe+2 → Fe+3 + 1 e-

Tthe electrons provided from both oxidation steps are used toreduce oxygen as

shown.

O2 (g) + 2 H2O + 4e- → 4 OH-

The ferric ions then combine with oxygen to form ferric oxide [iron (III) oxide] which is

then hydrated with varying amounts of water. The overall equation for the rust

formation may be written as:

Other metals, such as Aluminium, form an oxide layer when they come into contact

with oxygen from the air but the layer of oxide bonds very strongly to the surface of

the Aluminium preventing further oxidation from occurring. However, Aluminium can

rust in a very short time if a thin layer of mercury is applied to the surface. Mercury

readily combines with aluminium to form a mercury-aluminum amalgam when the two

pure metals come into contact. When the amalgam is exposed to air, the aluminium

oxidizes, leaving behind mercury. The oxide flakes away, exposing more mercury

amalgam, which repeats the process thus a small amount of mercury can rust a large

amount of aluminium over time, by progressively forming amalgam and relinquishing

the aluminium as oxide. For this reason mercury is prohibited on aircraft.

.

Hydrogen fuel cell operation

NUCLEAR BATTERIES

Nuclear batteries may seem like a recipe for disaster given the concern for nuclear

safety, however, they have the ability to produce power for long periods of time.

Nuclear batteries are not new. It may surprise you to know that they have been been

implanted into paitents that suffer from heart arrthymia to power cardiac pacemakers

since the 1973. Advances in the power of lithium batteries led to the phasing out of

nuclear batteries by 1975. The development of nuclear batteries has been re-ignited

with the need for long lasting batteries to power the portable devices such as laptops,

mp3 players and mobile telephones.

Nuclear batteries work by converting the heat produced by a nuclear source and

creating a current using the Seebeck-effect a second type of nuclear cell uses beta-

radiation impinging on a semiconductor junction to create electron hole pair which

migrate to the elelectrode of the junction creating a current. Much in the same way

that a solar cell creates energy. Currently these batteries cannot produce enough

power to run a laptop however they can be used to trickle charge batteries to give

longer lifetime for existing batteries.

Fuel cells: الوقود خالياA fuel cell is an electrochemical cell that converts the chemical energy from a fuel into electricity through an electrochemical reaction of hydrogen fuel with oxygen or another oxidizing agent.[1] Fuel cells are different from batteries in requiring a continuous source of fuel and oxygen (usually from air) to sustain the chemical reaction, whereas in a battery the chemical energy comes from chemicals already present in the battery. Fuel cells can produce

electricity continuously for as long as fuel and oxygen are supplied.

Fuel cells come in many varieties; however, they all work in the same general manner. They are made up of three adjacent segments: the anode, the electrolyte, and the cathode. Two chemical reactions occur at the interfaces of the three different segments. The net result of the two reactions is that fuel is consumed, water or carbon dioxide is created, and an electric current is created, which can be used to power electrical devices, normally referred to as the load.

At the anode a catalyst oxidizes the fuel, usually hydrogen, turning the fuel into a positively charged ion and a negatively charged electron. The electrolyte is a substance specifically designed so ions can pass through it, but the electrons cannot. The freed electrons travel through a wire creating the electric current. The ions travel through the electrolyte to the cathode. Once reaching the cathode, the ions are reunited with the electrons and the two react with a third chemical, usually oxygen, to create water or carbon dioxide.

:

Corrosion: التاكل

Corrosion is a natural process, which converts a refined metal to a more chemically-stable form, such as its oxide, hydroxide, or sulfide. It is the gradual destruction of materials (usually metals) by chemical and/or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to controlling and stopping corrosion.

In the most common use of the word, this means electrochemical oxidation of metal in reaction with an oxidant such as oxygen or sulfates. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion. This type of damage typically produces oxide(s) or salt(s) of the original metal, and results in a distinctive orange colouration. Corrosion can also occur in materials other than metals, such as ceramics or polymers, although in this context, the term "degradation" is more common. Corrosion degrades the useful properties of materials and structures including strength, appearance and permeability to liquids and gases.

Many structural alloys corrode merely from exposure to moisture in air, but the process can be strongly affected by exposure to certain substances. Corrosion can be concentrated locally to form a pit or crack, or it can extend across a wide area more or less uniformly corroding the surface. Because corrosion is a diffusion-controlled process, it occurs on exposed surfaces. As a result, methods to reduce the activity of the exposed surface, such as passivation and chromate conversion, can increase a material's corrosion resistance. However, some corrosion mechanisms are less visible and less predictable.

Types of Corrosion

Back to Top

There are different types of corrosion which depend on the environment surrounding the material, type of material, chemical reaction etc. Some general types of corrosion are described below.

1. Uniform Corrosion

This is also called General corrosion. It is a very common method of corrosion. It deteriorates the whole surface of the metal and makes the surface thin. The damage is done at a constant rate on the entire surface. It can be easily detected by it's appearance. It can be controlled but if it is not, it then destroys the whole metal.

2. Galvanic Corrosion

This type of corrosion occurs with an electrolyte like seawater. Metals have different values of electrical potentials. When they become electrically connected and put in an electrolyte, the more active metal which has a high negative potential becomes the anode. Due to it's high negative potential, it corrodes fast. But the less active metal becomes the cathode. 

The flow of electric current continues till the potentials are equal between both electrodes. So at the joint where the two non similar metals meet, the galvanic corrosion appears. The Galvanic Series shows the list of metals from the most active to the least active (most noble). Thus galvanic corrosion can be controlled by selecting the two metals which are close in series. As platinum is the least active, it is also less active for corrosion.

3. Pitting Corrosion

This occurs because of random attacks on particular parts of the metal's surface. This makes holes which are large in depth. These holes are called "pits". The pit acts as the anode while the undamaged part of the metal is the cathode. It begins with a chemical breakdown in the form of a scratch or spot. The pitting process makes the metal thinner and increases fatigue. For example, it can be very harmful in gas lines.

4. Stress Corrosion Cracking (SCC)

It is a complex form of corrosion which arises due to stress and corrosive environment. This generates brittle and dry cracks in the material. The brittle cracks can inter or Trans granular morphology. The stress is developed in the material due to

bending or stretching of the material. It also affects only at a particular section of material. 

The main reasons for stress corrosion are welding, heating treatments, deformation etc. It is very difficult to detect the cracks or detect stress corrosion because they combine with active path corrosion. The active path corrosion occurs generally along grain or crystallographic boundaries. Stress corrosion is strongly affected by alloycomposition.

5. Corrosion fatigue

This occurs in the presence of a corrosive environment like saltwater. It is a combination of cyclic stress and corrosion. Corrosion fatigue is produced when a metal breaks at a stress level which is lower than its tensile strength. It is strongly affected by the environment in which the metal resides which affects the initiation and growth rate of the cracks. These cracks are too fine to detect easily. So the stress coupons (metal sample) are used to detect the corrosion.

It can be produced by the influence of various types of stress like stresses applied, thermal expansion, thermal contraction, welding, soldering, cleaning, heating treatment, construction process, casting etc. To prevent corrosion fatigue, the designing and construction process of the materials should be done properly, by eliminating any stress and environmental factors and by eliminating crevices. 

6. Intergranular Corrosion

In the granular composition of metals and alloys, grains (small crystals) are present and their surfaces join with each other. This forms the grain boundaries. Thus the grains are separated by grain boundaries. Intergranular corrosion is also known as inter crystalline corrosion. The Intergranular corrosion is developed on or near the grain boundaries of a metal. This can be due to welding, stress, heat treating or improper service etc. The metal can loose its strength due to the Intergranular corrosion.

7. Crevice Corrosion

It is also known as concentration cell corrosion. This is due to the trapping of liquid corrosive between the gaps of the metal. As the electrolyte has aggressive ions like chlorides, the corrosion reaction is started after settling of liquid in gaps. Oxygen is consumed during the reaction. 

Thus an anodic area is developed near the oxygen-depleted zone while the external part of the material acts as a cathode. Crevice corrosion is similar to pitting corrosion.

It’ very difficult to detect crevice corrosion. It can be initiated by materials like gaskets, fasteners, surface deposits, washers, threads, clamp etc.

8. Filiform corrosion

It is a type of concentration cell corrosion. This develops on coated metallic surfaces with a thin organic film. The corrosion generates the defect on the protective coating of metallic surface. The filaments of corrosion product is the cause of degradation of the coating. The filaments look like thin threads. They exist as long branching paths. 

The actively growing filaments do not intersect the inactive filaments. The reflection process takes place when filaments collide with each other. Filiform corrosion is a very specific process because it only affects the surface’s appearance, not the metallic material. 

9. Erosion Corrosion

It is also called flow-assisted corrosion. This is due to the movement of corrosive liquids on metal surface which damages the material. It can be seen in ship propellers which are constantly exposed to sea water or in soft alloys. The damage can be seen as waves or rounded holes etc. It shows the flow of the corrosive liquid. It can be controlled by the use of hard alloys, managing the velocity and flow pattern of the fluid.

10. Fretting Corrosion

It is a form of erosion-corrosion. It shows the combined effect of corrosion and fretting of metal. Due to this corrosion, the material surface starts to disappear. Fretting corrosion exists in the form of dislocations of the surface and deep pits. Oxidation is the main cause of fretting corrosion. It can be controlled by using lubricates, controlling movement etc

Corrosion TheoryBack to Top

1. Water on the metal surface dissolves CO2 and O2 from the air.

2. Fe in contact with dissolved CO2 and O2 undergoes oxidation.

Fe $\rightarrow$ Fe2+ + 2e- - Anode

3. Electrons lost by Fe are taken by H+

H+ + e- $\rightarrow$ H

4H + O2 $\rightarrow$ 2H2O

On multiplying the first equation by 4 and adding to the second,

The dissolved O2 can take electrons directly also.

4. Fe2+ reacts with dissolved O2 and water

Rust (Hydrated ferric oxide).

Corrosion ProtectionBack to Top

Given below are some of the factors that cause corrosion.

Reactivity of metal Presence of impurities Presence of air, moisture, gases like SO2 and CO2

Presence of electrolytes

Two methods are used for the protection of materials from corrosion. 

1. Cathodic protection2. Corrosion inhibitors.

Both methods are based on charge control of the metal surface by measuring the potential of the metal. 

1. Cathodic protection

The principle of this method is to alter the electrode potential of the metallic structure so that they can lie in the immunity region. This is the region where the metal is in the stable state of the element and corrosion reactions are not possible. It is mostly used in steel structures in marine and under ground regions.

Two methods are used to apply the cathodic protection to a metal structure. 

Impressed Current - This method is used for the protection of pipelines and the hulls of ships in sea water. In this method, an electric current is applied to the metal surface by use of DC electrical circuit. The negative and positive terminal of the current source is connected to the metal requiring protection and an auxiliary anode respectively. The flow of electric current charges the structure with electrons and changes the electrode potential in the negative direction. This process continues till it reaches the immunity region. The current flows from anode to cathode. Thus it protects the metal surface from corrosion.

Sacrificial Anode - This is especially used for ships, offshore oil and gas production platform etc. In this technique, the more reactive metal is used to alter the electrode potential and get the immunity region. Zinc is generally used as sacrificial anode. It generates the anodic dissolution current with more negative potential. The cathodic curve intersection is now at a more negative potential which is the immunity region. At this region, the corrosion rate of steel is negligible.

2. Corrosion Inhibitors

According to surface chemistry, the presence of foreign molecules affect the surface reactions.

Corrosion processes are also a type of surface reactions. These can be controlled by foreign compounds which are known as inhibitors.

The inhibitors get adsorbed on the reacting metal surface. It attaches directly to the surface or adsorbs up to one molecular layer of the metal surface. This is a well known method for controlling the corrosion.

The inhibitors can work in different ways; it may block the active sites of corrosion and restrict the rate of anodic or cathodic process, or it may increase the electrode potential etc.

Hexylamine or sodium benzoate are used as inhibitors for anodic reactions. Similarly, oxidising agents like nitrite, chromate, red lead, amines, thio-urea

etc are also used as corrosion inhibitors.

It is the intrinsic property of a metal or a material that they have resistance to corrosion attack in some specific environmental conditions like pressure, temperature and velocity of fluid etc. These materials are thermodynamically unfavorable for corrosion reaction as in the case of graphite, zinc, cadmium. In these metals, corrosion continues but at an extremely slow rate.

Corrosion RateBack to Top

Corrosion Testing

1. Corrosion testing is used to measure corrosion. This is done in corrosion testing laboratories.

2. These are laboratories where experimental testing of materials is done for their verification about corrosion according to various industry standards.

3. Some standards are used for this purpose like ASTM, ISO, NACE, or custom corrosion testing.

4. Corrosion testing also includes DOT test, electrochemical and immersion test and heat transfer.

5. Some standardized methods are also used for testing of glass in different media like acidic, basic or neutral.

6. Testing methods are done in specific conditions of environment. In the ISO method, glass is put in the ionized water up to 60 minutes.

7. This solution is then titrated with HCl solution. Thus the amount of HCl for neutralization is measured.

It is also known as corrosion ratio. It is calculated by taking uniform corrosion over the whole surface. Corrosion rate is measured in terms of mpy (mils per year penetration).