CLASS XI (CLASSIFICATION IN ELEMENTS & PERIODICITY IN PROPERTIES

CLASSIFICATION OF ELEMENTS & PERIODICITY IN PROPERTIES

1. INTRODUCTION
Periodic table may be defined as the table which classifies all the known elements in accordance with their properties in such a way that elements with similar properties are grouped together in the same vertical column and dissimilar elements are separated from one another.

2. HISTORICAL DEVELOPMENT OF THE PERIODIC TABLE

All earlier attempts of the classification of the elements were based upon their atomic weights.

2.1 Dobereiner’s Triads

In 1829, Dobereiner classified certain elements in the groups of three called triads. The three elements in a triad had similar chemical properties. When the elements in a triad were arranged in the order of increasing atomic weights, the atomic weight of the middle element was found to be approximately equal to the arithmetic mean of the other two elements.

2.2 Newland’s Law of Octaves

In 1865, an English chemist, John Alexander Newlands observed that when the lighter elements were arranged in order of their increasing atomic weights, the properties of every eighth element were similar to those of the first one like the eighth note of a musical scale. This generalisation was named as Newlands’s law of octaves.

2.3 Lothar Meyer’s Curve

“Physical properties of elements are periodic functions of their atomic masses.”

According to Lothar Meyer, elements having similar properties occupy the similar positions in atomic volume viz atomic mass curve.

2.4 Mendeleev’s Periodic Law & Table

Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasing atomic weights in such a way that the elements with similar properties occupied the same vertical column or group.

"The physical and chemical properties of the elements are periodic functions of their atomic masses."

2.5 Modern Periodic Law & Table

In 1913, the English physicist, Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. A plot of √v (where √v is frequency of X-rays emitted) against atomic number (Z) gave a straight line and not the plot of √v vs atomic mass. Mendeleev’s Periodic Law was, therefore, accordingly modified. This is known as theModern Periodic Law and can be stated as :

"The physical and chemical properties of the elements are periodic functions of their atomic numbers."

Nomenclature of elements with Atomic Numbers > 100 

The naming of the new elements had been traditionally the privilege of the discoverer and the suggested name was ratified by the IUPAC.

3. PREDICTION OF BLOCK, PERIOD & GROUP

4. PROPERTIES OF AN ELEMENT

4.1 Atomic Radius

We cannot measure the exact size of an isolated atom because its outermost electron have a remote chance of being found quite far from the nucleus. So different types of atomic radius can be used based on the environment of atoms i.e; covalent radius, van der Waals’ radius,metallic radius.

4.1.1 Covalent Radius

The half of the distance between the nuclei of two identical atoms joined by single covalent bond in a molecule is known as covalent radius.

So covalent radius for A-A

If covalent bond is formed between two different elements then

where χA and χB are electronegative of A and B

4.1.2 Van der Waal’s Radius

It is half of the internuclear distance between adjacent atoms of the two neighbouring molecules in the solid state.

4.1.3 Metallic Radius (Crystal radius)

It is one-half of the distance between the nuclei of two adjacent metal atoms in the metallic crystal lattice.

So metallic radius for A-A

4.2 Variation of Atomic Radii in the Periodic Table

(a) Variation along a period

In general, the covalent and van der Waals radii decrease with increase in atomic number as we move from left to right in a period.

4.3 Atomic Radii

(a) Variation along a period

It is because with in the period the outer electrons are in the same valence shell & the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.

(b) Variation along a group

Atomic radius in a group increase as the atomic number increases. It is because with in the group, the principal quantum number (n) increases and the valence electrons are farther from the nucleus.

(c) Ionic Radius

The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion.

In general, the ionic radii of elements exhibit the same trend as the atomic radii. A cation is smaller than its parent atom because it has fewer electrons while its nuclear charge remains the same. The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge. For example, the ionic radius of fluoride ion (F^–) is 136 pm whereas the atomic radius of fluorine is only 64 pm. On the other hand ,the atomic radius of sodium is 186 pm compared to the ionic radius of 95 pm for Na⁺.

(d) Isoelectronic Species

Isoelectronic species are those which have same number of electrons. For example, O^2–, F^–, Na⁺ and Mg²⁺ have the same number of electrons (10). Their radii would be different because of their different nuclear charges. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative charge will have the larger radius. In this case, the net repulsion of the electrons will outweigh the nuclear charge and the ion will expand in size. Order of atomic radii is

General Trend-

4.4 Ionization Energy

The minimum amount of energy required to remove the electron from the outermost orbit of an isolated atom in the gaseous state is known as ionization energy.

IE1, IE2, IE3 and IE4 are successive ionization energies.

Variation of Ionisation Energy in Periodic Table

(a) Variation along a period

In a period, the value of ionisation enthalpy increases from left to right with breaks where the atoms have somewhat stable configurations. The observed trends can be easily explained on the basis of increased nuclear charge and decrease in atomic radii. Both the factors increase the force of attraction towards nucleus and consequently, more and more energy is required to remove the electrons and hence, ionisation enthalpies increase.

(b) Variation along a group

On moving the group, the atomic size increases gradually due to an addition of one new principal energy shell at each succeeding element. On account of this, the force of attraction towards the valence electrons decreases and hence the ionisation enthalpy value decreases.

4.5 Units of I.E./I.P.

It is measured in units of electron volts (eV) per atom or kilo calories per mole (kcal mol–1) or kilo Joules per mole (kJ mol–1). One electron volt is the energy acquired by an electron while moving under a potential difference of one volt.

Important Points

* Ionization energy increases with decreasing the size of an atom or an ion

* Ionization energy increases with decreasing screening effect.

* Ionization energy increases with increasing nuclear charge

* Ionization energy increases if atom having half filled and fully filled orbitals

* The penetrating power of orbitals is in the order s > p > d > f

4.6 Electron Gain Enthalpy

When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the Electron Gain Enthalpy (D_egH). Electron gain enthalpy provides a measure of the ease with which an atom adds an electron to form anion as represented by

X(g) + e^– -> X^–(g)

Depending on the element, the process of adding an electron to the atom can be either endothermic or exothermic. For many elements energy is released when an electron is added to the atom and the electron gain enthalpy is negative. For example, group 17 elements (the halogens) have very high negative electron gain enthalpies because they can attain stable noble gas electronic configurations by picking up an electron. On the other hand, noble gases have large positive electron gain enthalpies because the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration.

Variation of Electron Gain Enthalpy

(a) Variation along a period

Electron gain enthalpy becomes more and more negative from left to right in a period. This is due to decrease in size and increase in nuclear charge as the atomic number increases in a period. Both these factors favour the addition of an extra electron due to higher force of attraction by the nucleus for the incoming electron.

(b) Variation along a group

The electron gain enthalpies, in general, become less negative in going down from top to bottom in a group. This is due to increase in size on moving down a group. This factor is predominant in comparison to other factor, i.e., increase in nuclear charge.

4.7 Electronegativity

The tendency of an atom to attract the shared pair of electrons towards itself is known as its electronegativity. According to Pauling, the electronegativity of F is 4.0 and electronegativity of other elements can be calculated as

(where IP = Ionization potential, EA = Electron affinity)

If IP and EA are taken in electron volt

* Percentage ionic character = 16(χ_A – χ_B) + 3.5(χ_A – χ_B)² where χ_A and χ_B are electronegativities of A and B.

* If the difference in the electronegatives of combining atoms is 1.7, the bond is 50% covalent and 50% ionic.

* If the difference in electronegativities of oxygen and element is very high the oxide shows a basic character.

The periodic trends of elements in the periodic table

4.8 Periodic Trends in Chemical Properties

4.8.1 Periodicity of Valence or Oxidation States

The electrons present in the outermost shell of an atom are called valence electrons and the number of these electrons determine the valence or the valency of the atom. It is because of this reason that the outermost shell is also called the valence shell of the atom and the orbitals present in the valence shell are called valence orbitals. In case of representative elements, the valence of an atom is generally equal to either the number of valence electrons (s- and p-block elements) or equal to eight minus the number

of valence electrons.

In contrast, transition and inner transition elements, exhibit variable valence due to involvement of not only the valence electrons but d- or f-electrons as well. However, their most common valence are 2 and 3.

Let us now discuss periodicity of valence along a period and within a group.

(a) Variation along a period

As we move across a period from left to right, the number of valence electrons increases from 1 to 8. But the valence of elements, w.r.t. H or O first increases from 1 to 4 and then decreases to zero. In the formation of Na2O molecule, oxygen being more electronegative accepts two electrons, one from each of the two sodium atoms and thus shows an oxidation state of –2. On the other hand, sodium with valence shell electronic configuration as 3s1 loses one electron to oxygen and is given an oxidation state of +1. Thus, the oxidation state of an element in a given compound may be defined as the charge acquired by its atom on the basis of electronegativity of the other atoms in the molecule.

(b) Variation within a group

When we move down the gorup, the number of valence electrons remains the same, therefore, all the elements in a group exhibit the same valence. For example, all the elements of group 1 (alkali metals) have valence one while all the elements of group 2 (alkaline earth metals) exhibit a valence

of two. Noble gases present in group 18 are zerovalent, i.e., their valence is zero since these elements are chemically inert.

4.8.2 Anomalous Properties of Second Period Elements

It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups. For example, lithium (of group 1) resembles magnesium (of group 2) and beryllium (of group 2) resembles aluminium (of group 13) and so as. This similarity in properties of elements placed diagonally to each other is called diagonal relationship.

The anomalous behaviour is due to their small size, large charge/radius ratio and high electronegativity of the elements. In addition, the first member of group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals (3s, 3p, 3d). As a consequence of this, the maximum covalency of the first member of each group is 4 (e.g., boron

can only form [BF₄]^–), whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons e.g., aluminium forms [AlF₆]^3–. Furthermore, the first member of p-block elements displays greater ability to form pp-pp multiple bonds to itself (e.g., C =C, CºC, N=N, NºN) and to other second period elements (e.g., C=O, C=N, NºC, N=O) compared to subsequent members of the same groups.

4.8.3 Periodic Trends and Chemical Reactivity

Reactivity of Metals

The reactivity of metals is measured in terms of their tendency to lose electrons from their outermost shell.

In a period

The tendency of an element to lose electrons decreases in going from left to right in a period. So, the reactivity of metals decreases in a period from left to right. For example, the reactivity of third period elements follows the order.

In a group

The tendency to lose electrons increases as we go down a group. So, the reactivity of metals increases down the group. Thus, in group 1, the reactivity follows the order.

Reactivity of Non-Metals

The reactivity of a non-metal is measured in terms of its tendency to gain electrons to form an anion.

In a period

The reactivity of non-metals increases from left to right in a period. During reaction, non-metals tend to form anions. For example, in the second period, the reactivity of non-metals increases in the order.

In a group

The reactivity of non-metals in a group decreases as we go down the group. This is because the tendency to accept electrons decreases down the group. The reactivity of halogens follows the order

The normal oxide formed by the element on extreme left is the most basic (e.g., Na2O) whereas that formed by the element on extreme right is the most acidic (e.g., Cl2O7). Oxides of elements in the centre are amphoteric (e.g., Al2O3, As2O3) or neutral (e.g., CO, NO, N2O). Amphoteric oxides behave as acidic with bases and as basic with acids, whereas neutral oxides have no acidic or basic properties.

4.8.4 Inert Pair Effect

In groups 13-16, as we move down the group, the tendency of s-electrons of the valence shell to participate in bond formation decreases. This means that lower oxidation state becomes more stable.

Reason: As we go down these groups, the increased nuclear charge outweighs the effect of the corresponding increase in atomic size. The s-electrons thus become more tightly held (more penetrating) and hence more reluctant to participate in bond formation. Hence, the lower oxidation state becomes more stable.

5. SUMMARY AND IMPORTANT POINTS TO REMEMBER

1. Mendeleev’s periodic table was based on atomic masses of the elements. When Mendeleev presented the periodic table, only 63 elements were known. He left 29 places in the table for unknown elements.

2. Modern Mendeleev periodic table is based on atomic numbers of the elements. The modern periodic law is : “The physical and chemical properties of the elements are periodic function of their atomic numbers”. The horizontal row in the periodic table is called a period and vertical column is called group. There are seven periods and nine groups in the modern Mendeleev periodic table.

3. The long or extended form of periodic table consists of seven periods and eighteen vertical columns (groups or families). The elements in a period have same number of energy shells, i.e., principal quantum number (n). These are numbered 1 to 7.

In a vertical column (group), the elements have similar valence shell electronic configuration and therefore exhibit similar chemical properties.

4. There are four blocks of elements: s-, p-, d- and f-block depending on the orbital which gets the last electron. The general electronic configuration of these blocks are :

s-block elements occupy IA(1) and IIA(2) groups, i.e., extreme left portion of the periodic table. p-block elementsoccupy IIIA(13), IVA(14), VA(15), VIA(16), VIIA(17) and VIIIA(18) groups, i.e., right portion of the periodic table. d-block elements occupy IIIB(3), IVB(4), VB(5), VIB(6), VIIB(7), VIIB(8, 9 and 10), IB(11) and IIB(12) groups, i.e., central portion of the periodic table. There are four d-block series, i.e., 3d series, 4d series, 5d series and 6d series, each consisting of ten elements, i.e., in all forty d-block elements are present in periodic table. f-block elements are accommodated in two horizontal rows below the main periodic table, each row consists of 14 elements, i.e., 28 f-block elements are present in periodic table. The elements in first row are termed 4f-elements or rare earth or lanthanides while the elements of second row are termed 5f-elements or actinides.

5. The elements are broadly divided into three types :

(i) Metals comprise more than 78% of the known elements. s-block, d-block and f-block elements are metals. The higher members of p-block are also metals.

(ii) Non-metals are less than twenty. (C, N, P, O, S, Se, H, F, Cl, Br, I, He, Ne, Ar, Kr, Xe and Rn are non-metals).

(iii) Elements which lie in the border line between metals and non-metals are called semimetals or metalloids. B, Si, Ge, As, Sb, Te, Po and At are regarded metalloids.

6. IUPAC given a new scheme for assigning a temporary name to the newly discovered elements. The name is derived directly from the atomic number of the elements. However, IUPAC has accepted the following names of the elements from atomci numbers 104 to 110.

7. The recurrence of similar properties of the elements after certain definite intervals when the elements are arranged in order of increasing atomic numbers in the periodic table is termed periodicity. The cause of periodicity is the repetition of similar electronic configuration of the atom in the valence shell after certain definite intervals. These definite intervals are 2, 8, 8, 18, 18 and 32. These are known as magic number. Periodicity is observed in a number of properties which are directly or indirectly linked with electronic configuration.

(i) Effective nuclear charge increases across each period.

(ii) Atomic radii generally decrease across the periods.

(iii) Atomic radii generally increase on moving from top to bottom in the groups.

(iv) Atomic radius is of three types :

(a) Covalent radius : It is half of the distance between the centres of the nuclei of two similar atoms joined by a single covalent bond. This is generally used for non-metals.

(b) Crystal or metallic radius : It is half of the internuclear distance between two nearest atoms in the metallic lattice. It is generally used for metals.

(c) van der Waals’ radius : It is half of the internuclear distance between the nearest atoms belonging to two adjacent molecules in solid state.

(v) Cations are generally smaller than anions.

(vi) Cations are smaller and anions are larger than neutral atoms of the elements.

(vii) Elements of 2nd and 3rd transition series belonging to same vertical columns are similar in size and properties due to lanthanide contraction.

(viii) The first element is each group of the representative elements shows abnormal properties, i.e., differs from other elements of the group because of much smaller size of the atom.

(ix) The ions having same number of electrons but different nuclear charge are called isoelectronic ions.

Examples,

(a) N^3–, O^2–, F^–, Na⁺, Mg²⁺, Al³⁺

(b) P^3–, S^2–, Cl^–, K⁺, Ca²⁺, Sc³⁺

In isoelectronic ions, the size decreases if Z/e increases i.e., greater the nuclear charge, smaller is the size of the ion.

(x) The energy required to remove the most loosely held electron from the gaseous isolated atom is termed ionisation enthalpy.

(xi) Ionisation enthalpy values generally increase across the periods.

(xii) Ionisation enthalpy values generally decrease down the group.

(xiii) Removal of electron from filled and half filled shells requires of higher energy. For example, the ionisation enthalpy of nitrogen is higher than oxygen. Be, Mg and noble gases have high values.

(xiv) Metals have low ionisation enthalpy values while non-metals have high ionisation enthalpy values.

(xv) Successive ionisation enthalpies of an atom have higher values. IE_I < IE_II < IE_III ...

(xvi) The enthalpy change taking place when an electron is added to an isolated gaseous atom of the element is called electron gain enthalpy. The first electron gain enthalpy of most of the elements is negative as energy is released in the process but the values are positive or near zero in case of the atoms having stable configuration such as Be, Mg, N, noble gases, etc.

(xvii) Electron gain enthalpy becomes more negative from left to right in a period and less negative from top to bottom in a group.

(xviii)Successive electron gain enthalpies are always positive.

(xix) The elements with higher ionisation enthalpy have higher negative electron gain enthalpy.

(xx) Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a bond.

(xxi) Electronegativity increases across the periods and decreases down the groups.

(xxii) Metals have low electronegativities and non-metals have high electronegativities.

(xxiii)Metallic character decreases across the periods and increases down the group.

(xxiv)Valence of an element belonging to s- and p- block (except noble gases) is either equal to the number of valence electrons or eight minus number of valence electrons.

(xxv) The reducing nature of the elements decreases across the period while oxidising nature increases.

(xxvi)The basic character of the oxides decreases while the acidic character increases in moving from left to right in a period.

6. SOME IMPORTANT FACTS ABOUT ELEMENTS

(i) Bromine is a non-metal which is liquid at room temperature.

(ii) Mercury is the only metal that is liquid at room temperature.

(iii) Gallium (m.pt. 29.8ºC), caesium (m.pt. 28.5ºC) and francium (m.pt. 27ºC) are metals having low melting points.

(iv) Tungsten (W) has the highest melting point (3380ºC) among metals.

(v) Carbon has the highest melting point (4100ºC) among nonmetals.

(vi) Oxygen is the most abundant element on the earth.

(vii) Aluminium is the most abundant metal.

(viii) Iron is the most abundant transition metal.

(ix) Highest density is shown by osmium (22.57 g cm^–3) or iridium (22.61 g cm^–3).

(x) Lithium is the lightest metal. Its density is 0.54 g cm^–3.

(xi) Silver is the best conductor of electricity.

(xii) Diamond (carbon) is the hardest natural substance.

(xiii) Francium has the highest atomic volume.

(xiv) Boron has the lowest atomic volume.

(xv) The most abundant gas in atmosphere is nitrogen.

(xvi) Fluorine is the most electronegative element.

(xvii) Chlorine has the maximum negative electron gain enthalpy.

(xviii)Helium has the maximum ionisation enthalpy.

(xix) Cesium or francium has the lowest ionisation enthalpy.

(xx) Helium and francium are smallest and largest atoms respectively.

(xxi) H– and I– ions are the smallest and largest anions respectively.

(xxii) H+ and Cs+ ions are the smallest and largest cations respectively.

(xxiii) Cesium is the most electropositive element.

(xxiv) Element kept in water is phosphorus, P₄ (white or yellow).

(xxv) Element kept in kerosene are Na, K, Rb, Cs, etc.

(xxvi) Iodine is the element which sublimes.

(xxvii) Hydrogen is the most abundant element in the universe.

(xxviii) Only ozone is the coloured gas with garlic smell.

(xxix) Metalloids have electronegativity values closer to 2.0.

(xxx) First synthetic (i.e., man-made) element is technetium (At. No. 43).

(xxxi) Most poisonous metal-Plutonium.

(xxxii)Rarest element in earth’s crust-Astatine.

(xxxiii) The elements coming after uranium are called transuranic elements. The elements with Z = 104 – 112, 114 and 116 are called trans-actinides or super heavy elements. All these elements are synthetic, i.e., man-made elements. These are radioactive elements and not found in nature.

(xxxiv) The elements ruthenium (Ru), germanium (Ge), polonium (Po) and americium (Am) were named in honour of the countries named Ruthenia (Russia), Germany, Poland and America, respectively.

(xxxv) The members of the actinide series are radioactive and majority of them are not found in nature.

(xxxvi) The element rutherfordium (Rf, 104) is also called Kurchatovium (Ku) and element dubnium (Db, 105), is also called hahnium.

(xxxvii) Promethium (Pm, 61) a member of lanthanide series is not found in nature. It is a synthetic element.

(xxxviii) Special names are given to the members of these groups in periodic table.

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Atish J. Bain

HOD (Dept. of Chemistry)

St. Thomas High School (10+2)

Dhanbad, Jharkhand

CLASS XII (ELECTROCHEMISTRY)

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ELECTROCHEMISTRY

1. ELECTROCHEMISTRY

Electrochemistry is the study of production of electricity from the energy released during a spontaneous chemical reaction and the use of electrical energy to bring about non-spontaneous chemical transformations. 

2. ELECTROCHEMICAL CELLS

A spontaneous chemical process is the one which can take place on its own and in such a process the Gibb's energy of the system decreases. It is this energy that gets converted to electrical energy. The reverse process is also possible in which we can make non-spontaneous processes occur by supplying external energy in the form of electrical energy. These inter conversations are carried out in equipment called electrochemical cells.

3. TYPES OF ELECTROCHEMICAL CELLS

Electrochemical cells are of two types.

3.1 Galvanic Cells - Converts chemical energy into electrical energy.

3.2 Electrolytic Cells - Converts electrical energy into chemical energy.

4. GALVANIC CELL 

Cell energy is extracted from a spontaneous chemical process or reaction and it is converted electric current.

For example, Daniel Cell is a Galvanic Cell in which Zinc and Copper are used for the redox reaction to take place.

Zn(s) + Cu²⁺(aq) -> Zn²⁺(aq) + Cu(s)

Oxidation Half: Zn(s)-> Zn²⁺(aq) + 2e^-

Reduction Half: Cu²⁺(aq) + 2e^- -> Cu(s)

Zn is the reducing agent and Cu²⁺ acts as oxidising agent. The half cells are also known as electrodes. 

The oxidation half cell is called anode and the reduction half cell is called cathode. Electrons flow from anode to cathode in the external circuit.

Anode is assigned negative polarity and cathode is assigned positive polarity.

In Daniel cell Zn acts as anode and Cu acts as cathode.

5. ELECTROLYTIC CELL

These electrodes are dipped in and electrolytic solutions containing cations and anions. On supplying current the ions move towards electrodes of opposite polarity and simultaneous reduction and oxidation processes takes place.

5.1 Preferential Discharge of Ions

Where there are more than one cation or anion the process of discharge becomes competitive in nature. Discharge of ions requires energy and in case of several ions being present the discharge of that ion will take place first which requires the energy.

6. ELECTRODE POTENTIAL

It may be defined as the tendency of an element, when it is placed in contact with its own ions to either lose or gain electrons and in turn becomes positively or negatively charged.

The electrode potential will be named as oxidation or reduction potential depending upon whether oxidation or reduction has taken place.

6.1 Characteristics

a) Both oxidation and reduction potential are equal in magnitude but opposite in sign.

b) It is not a thermodynamic property, so values of E are not additive.

7. STANDARD ELECTRODE POTENTIAL (E⁰)

It may be defined as the electrode potential of an electrode determined relative of Standard Hydrogen Electrode (SHE) under standard conditions. The standard condition taken are: 

i)        1M concentration of each ion solution.

ii)       A temperature of 298 K.

iii)      1 bar pressure for each gas.

8. ELECTROCHEMICAL SERIES

The half cell potential values are standard values and are represented as the standard reduction potential values as shown in the table below which is also called electrochemical series.

9. CELL POENTIAL OR EMF OF A CELL

The difference between the electrode potentials of two half cells is called cell potential.

It is known as Electromotice force (EMF) of the cell if no current is drawn from the cell.

Ecell = Ecathode + Eanode

Ecell = Ereduction + Eoxidation

Ecell = Eright + Eleft

So for Daniel Cell,

10. CELL DIAGRAM OR REPRESENTATION OF A CELL

The following conventions or notations are applied for writing the cell diagram in accordance with IUPAC recommendations.

The Daniel Cell is represented as follows:

a) Anode half cell is written on the left hand side while cathode half cell is written in the right hand side.

b) A single vertical line (|) separates the metal from aqueous solution of its own ions.

c) A double vertical line (||) represents salt bridge.

d) The molar concentration (C) is placed in brackets () after the formula of the corresponding ion.

e) The value of EMF of the cell is written on the extreme right of the cell.

f) If an inert electrode like platinum is involved in the construction of the cell, it may be written along with the working electrode in brackets say for example, when a Zn anode is connected to a H2 electrode.

11. SALT BRIDGE

Salt bridge is used to maintain the charge balance and to complete the circuit by facilitating the flow of ions through it.

It contains a gel in which inert electrolyte like Na₂SO₄ or KNO₃ etc are mixed with agar agar.

Negative ions flow to the anode and positive ions flow to the cathode through the salt bridge and charge balance is maintained and cell is kept on functioning. 

12. SPONTANEITY OF A REACTION

For a spontaneous cell reaction ΔG should be negative and cell potential should positive.

If we take standard value of cell potential in the above equation we will obtain standard value for ΔG as well.

13. TYPES OF ELECTRODES

13.1 Metal-Metal Ion Electrodes

A metal rod/plate is dipped in an electrolytic solution containing metal ions. There is a potential difference between these two phases and this electrode can act as a cathode or anode both.

13.2 Gas Electrodes 

Electrode gases like H2, Cl2 etc are used with their respective ions. For example, H2 gas is used with a dilute solution of HCl(H+ ions). The metal should be inert so that it does not react with the acid.

The hydrogen electrode is also used as the standard to measure other electrode potentials. Its own potential is set to 0 V as a reference. When it is used as a reference the concentration of dil. HCl is taken as 1 M and the the electrode is called Standard Hydrogen Electrode.

13.3 Metal-Insoluble Salt Electrodes

We use some metals which are sparingly soluble with the metal itself as electrodes. For example, if we use AgCl with Ag there is a potential gap between these two phases which can be identified in the following reaction.

This electrode is made by dipping a silver rod dipped in a solution containing AgCl and Cl- ions.

13.4 Calomel Electrode

Mercury is used with two other phases, one is a calomel paste (Hg2Cl2) and electrolyte containing Cl- ions.

 

This electrode is also used as another standard to measure other potentials. Its standard form is also called Standard Calomel Electrode (SCE).

13.5 Redox Electrode

In these electrodes two different oxidation states of the same metal are used in the same half cell. For example, Fe²⁺ and Fe³⁺ are dissolved in the same container and an inert electrode of platinum is used for the electron transfer. Following reaction can take place.

14. NERNST EQUATION

It relates electrode potential with the concentration of ions.

Thus, the reduction potential increases with the increase in the concentration of ions.

For a general electrochemical reaction of the type.

Nernst equation can be given as

Substituting the value of R (gas constant) and F(Faraday's constant)

15. APPLICATION OF NERNST EQUATION

15.1 Equilibrium Constant from Nernst Equation

For a Daniel cell at equilibrium,

But at equilibrium,

In general,

or,

16. CONCENTRATION CELLS

If the two electrodes of the same metal are dipped separately into two solutions of the same electrolytes having different concentrations and the solutions are connected through salt bridge, such calls are known as concentration cells. For example,

These are of two types,

16.1 Electrode Concentration Cells

16.2 Electrolyte Concentration Cells

17. CASES OF ELECTROLYSIS

17.1 Electrolysis of molten Sodium Chloride

The reactions occurring at the two electrodes may be show as follows

17.2 Electrolysis of an aqueous solution of Sodium Chloride 

Thus, H2 gas is evolved at cathode and Na+ ions remain in solution.

Thus, Cl2 gas is evolved at the anode by over voltage concept while OH- ions remain in the solution.

18. BATTERIES

When Galvanic Cells are connected in series to obtain a higher voltage, the arrangement is called Battery.

18.1 Primary Battery

Primary cells are those which can be used so long that the active minerals are present. Once they get consumed the cell will stop functioning and cannot be re-used. Example, Dry Cell, Leclanche cell and Mercury cell.

18.2 Dry Cell

The standard potential of this cell is 1.5 V and it falls as the cell gets discharged continuously and once used, it cannot be charged.

18.3 Mercury Cell

These are used in small equipment like watches, hearing aids.

The cell potential is approximately 1.35 V and remains constant during its life.

18.4 Secondary Battery

Secondary cells are those which can be recharged again and again for multiple uses.

For example, Lead Storage Battery and Nickel Cadmium Battery.

18.5 Lead Storage Battery

To recharge the cell, it is connected with a cell of higher potential and this cell behaves as an electrolytic cell and the reactions are reversed. Pb(s) and PbO2(s) are regenerated at the respective electrodes. These cells deliver an almost consistent voltage.

19. FUEL CELL

A fuel cell differs from an ordinary battery in the sense that the reactants are not contained inside the cell but are externally supplied from an external reservoir.

Fuel cell is used in space shuttles and in this cell the two gasses are supplied from external storage. 

In this cell carbon rods are used as electrodes with KOH as the electrolyte.

20. CORROSION

It involves a redox reaction and formation of an electrochemical cell on the surface of iron or other metals.

At one location oxidation of iron takes place (anode) and at another location reduction of oxygen to form water takes place (cathode).

First Fe gets oxidised to Fe2+ and then in the presence of oxygen it forms Fe3+ which then reacts with water to form rust which is represented by Fe2O3.xH2O.

Rusting of iron can be avoided by painting it or by coating it with some other metals like Zinc.

The later process is called galvanisation. 

As the tendency of Zn to get oxidised is more than iron it gets oxidised in preference and iron is protected. This method of protecting one metal by the other is also called Cathodic Protection.

21. CONDUCTANCE (G)

It is the reciprocal of resistance and may be defined as the ease with which the electric current flows through a conductor.

SI Unit is Siemen (S).

22. CONDUCTIVITY (κ)

It is the reciprocal of resistivity (ρ).

Hence, conductivity of an electrolytic solution may be defined as the conductance of a solution of 1 cm length with area of cross-section equals to 1 cm^2.

23. FACTORS AFFECTING ELECTROLYTIC CONDUCTANCE

23.1 Electrolyte

An electrolyte is a substance that dissociates in solution to produce ions and hence conducts electricity in dissolved or molten state.

For example, HCl, NaOH, KCl are strong electrolytes and CH3COOH, NH4OH are weak electrolytes.

The conductance of electricity by ions present in the solution is called electrolytic or ionic conductance. The following factors govern the flow of electricity through a solution of electrolyte.

a) Nature of electrolyte or inter-ionic attractions- Lesser the solute-solute interactions, greater will be the freedom of movement of ions and higher will be the conductance.

b) Solvation of ions- Larger the magnitude of solute-solvent interactions, greater is the extent of solvation and lower will be the electrical conductance.

c) Nature of solvent and its viscosity- Larger the solvent-solvent interaction, larger will be viscosity and more will be the resistance offered by the solvent to flow of ions hence lesser will be the electrical conductance.

d) Temperature- As the temperature of electrolytic solution rises solute-solute, solute-solvent and solvent -solvent interaction decreases, this leads to the increase of electrical conductance.

24. MEASUREMENT OF CONDUCTANCE

As we know, κ = 1/R x l/A 

The value of κ could be known if we measure l, A and R. 

The value of resistance of the solution R between two parallel electrodes is determined by using Wheatstone Bridge method.

It consists of two fixed resistances R3 and R4, a variable resistance R1 and the conductivity cell having the unknown resistance R2. The bridge is balanced when no current is passes through the detector. Under these conditions,

25. MOLAR CONDUCTIVITY (Λm)

It may be defined as the conducting power of all the ions produced by dissolving one mole of any electrolyte placed between two large electrodes at one centimetre apart.

Mathematically,

 where, V is the volume of solution in cm^3 containing 1 mole of electrolyte and C is the molar concentration.

26. EQUIVALENT CONDUCTIVITY (Λeq)

It is conducting power of one equivalent of electrolyte placed between two large electrodes at one centimetre apart.

Mathematically,

Where, V is the volume of solution in cm^3 containing 1 equivalent of electrolyte and N is normality. 

27. VARIATION OF CONDUCTIVITY AND MOLAR CONDUCTIVITY WITH DILUTION

Conductivity decreases with decrease in concentration, this is because the number of ions per unit volume that carry the current in the solution decreases on dilution.

Molar conductivity increases with decrease in concentration. This is because the total volume of solution containing one mole of electrolyte also increases.

It has been found that the decrease in κ on dilution of a solution is more than compensated by increase in its volume.

 

28. LIMITING MOLAR CONDUCTIVITY (Λm)

The value of molar conductivity when the concentration approaches zero is known as limiting molar conductivity at infinite dilution.

It is possible to determine the molar conductivity at infinite dilution (Λ⁰_m) in case of strong electrolytes by extrapolation of curve Λ_m vs √C. On contrary, the value of molar conductivity of weak electrolyte at infinite dilution cannot be determined by extrapolation of the curve as the curve becomes almost parallel to y-axis when concentration approaches to zero.

The mathematical representation between Λ_m and Λ⁰_m for strong electrolyte was developed by Debye, Huckel and Onsagar. 

In simplified form the equation can be given as 

Where b is a constant which depend on the nature of the solvent and temperature.

29. KOHLRAUSCH'S LAW

It states that the limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.

In general, if an electrolyte on dissociation given V+ cations and V_ anions then its limiting molar conductivity is given by 

Here, λ⁰₊ and λ⁰_- are the limiting molar conductivities of cations and anions respectively.

30. APPLICATIONS OF KOHLRAUSCH'S LAW

30.1 Calculation of molar conductivities of weak electrolytes at infinite dilution

For example, molar conductivity of acetic acid at infinite dilution can be obtained from the knowledge of molar conductivity at infinite dilution of strong electrolytes like HCl, CH3COONa and NaCl as shown below.

30.2 Determination of Degree of Dissociation of Weak Electrolytes

30.3 Determination of Dissociation Constant (K) of Weak Electrolytes

31. USE OF ΔG IN RELATING EMF VALUES OF HALF CELL REACTIONS

When we have two half cells reactions such that on adding them we obtain another half cell reaction then their EMFs cannot be added directly. 

But in any case thermodynamic functions like ΔG can be added and EMF values can be related through them.

Consider the following three half cell reactions,

  We can easily observe that the third reaction can be obtained by subtracting the first reaction from the second.

But the same relation does not apply on the EMF values. That is E₃≠E₂-E₁. But the ΔG values can be related according to the relations. That is,

Note- We should always remember that EMF values are additive only when two half cell reactions are added to give a complete balanced cell reaction. In any other case we will be using ΔG values to obtain relation between EMF values.

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Atish J. Bain

HOD (Dept. of Chemistry)

St. Thomas High School (10+2)

Dhanbad, Jharkhand