UNIT 3 CLASSIFICATION OF ELEMENTS

Syllabus

• Significance of classification

• Brief history of the development of periodic table

• Modern periodic law

• Present form of periodic table

• Periodic trends in properties of elements – atomic radii, ionic radii, inert gas radii, ionic radii,ionisation enthalpy, electron gain enthalpy, electronegativity, valence

NEED FOR CLASSIFICATION OF ELEMENTS

Before the beginning of eighteenth century, only few elements were known. Therefore , it was quite easy to study and remember the properties of these elements and their compounds individually. But gradually this number went on increasing and as many 114 elements, either natural or synthetic , have been discovered. It thus become rather difficult to study individually the chemistry of these elements and millions of their compounds. This necessity led to the classification of various elements into groups.

Although every element is different from every other element, yet some elements have certain similarities. Based upon these similarities, the scientists after numerous attempts were ultimately successful in arranging the various elements into groups or chemical families in such a way that similar elements were grouped together and dissimilar elements were separated from one another. This arrangement of elements is called classification of elements and this led to the formulation of a periodic table. Thus, 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.

HISTORICAL DEVELOPMENT OF THE PERIODIC TABLE

In 1808 , John Dalton proposed his Atomic Theory. According to this theory, atoms of one element can be distinguished from the other on the basis of their atomic weights. Thus, all earlier attempts on the classification of the elements were based on their atomic weights.

1. Dobereiner’s Triads : The first attempt towards the classification of elements was made by Johann Dobereiner, a German chemist in 1829. He pointed out that there were sets of three elements (triads) which showed similar chemical properties. He also noted the fact that the atomic weight of the central element of the Triad was approximately the mean of the atomic weights of the other two members and the properties of the middle element were in between those of the other two members

Some examples of Dobereiner’s Triads are :

Element Atomic weight Element Atomic weight Element Atomic weight

Li 7 Ca 40 Cl 35.5

Na 23 Sr 88 Br 80

K 39 Ba 137 I 127

The major drawback of Dobereiner’s classification was that the concept of triads could be applied to limited number of elements.

Chancourtois telluric helix

In 1862 Chancourtois a French geologist who arranged the elements in order of increasing atomic weights. He plotted the elements on a cylinder, which he called the telluric helix. The paper, which used geological terms and was published without the diagram, was completely ignored until the work of Mendeleev.

Newlands Law of Octaves

In 1865 , an English chemist, John Newlands observed that :

When the lighter elements are arranged in order of their increasing atomic weights, the properties of every eighth element is similar to the first one like the eighth note of a musical scale. This generalization was named as Newlands law of octaves.

Sl. No 1 2 3 4 5 6 7

Element Li Be B C N O F

Sl. No 8 9 10 11 12 13 14

Element Na Mg Al Si P S Cl

Sl. No 15 16

Element K Ca

Thus sodium the eighth element is similar to lithium and so is the next eighth element potassium. The same is true of beryllium, magnesium and calcium ; boron and aluminium and so on. This generalisation is also discarded since it could not be applied to elements having atomic weights greater than that of calcium , i.e., 40 amu. Furthermore, with discovery of noble gases, the properties of eighth element are no longer similar to the first one.

Lothar – Meyer arrangement

In 1869 , Lothar Meyer , a German chemist, studied the physical properties such as atomic volume , melting point and boiling points of various elements. He plotted the physical properties such as atomic volume, melting point and boiling points against atomic weight and obtained periodically repeated pattern.

A plot of atomic volume against atomic weight is shown.In the curve :

(i) The most strongly electropositive alkali metals (Li, Na, K, Rb and Cs) occupy the peaks on the curve.

(ii) The less strongly electropositive alkaline earth metals (Be, Mg, Ca, Sr and Ba) occupy the descending positions on the curve.

(iii) The most electronegaive elements i.e., halogens (F, Cl, Br and I) occupy ascending positions of the curve.

Lothar Meyer’s atomic volume curves

On the basis of these observations, Lother Meyer proposed that the physical properties of the elements are a periodic function of their atomic weights. He arranged the then known elements in the tabular form in the order of their increasing atomic weights. However, his work was not published until after the work of Dimitri Mendeleev, the scientist who is generally credited the development of the Modern Periodic Table.

MENDELEEV’S PERIODIC LAW

In 1869 Dmitri Mendeleev, a Russian chemist , made a remarkable contribution towards the classification of elements. On the basis of the chemical properties of the elements, he proposed that chemical properties of the elements are a periodic function of their atomic weights.

Later on, Mendeleev came to know about the work of Lother Meyer. He then integrated the two statements in the form of a law called Mendeleev – Lothar Meyer Periodic Law or simply Mendeleev Periodic Law. It states that :

The physical and chemical properties of elements are a periodic function of their atomic weights, i.e., when the elements are arranged in order of their increasing atomic weights, elements with similar properties are repeated after certain regular intervals.

MENDELEEV’S PERIODIC TABLE

A tabular arrangement of the elements based on the periodic law is a periodic table. Mendeleev’s 1871 table is reproduced in TABLE. In this table the elements are arranged in 12 horizontal rows and 8 vertical columns or groups. To bring similar elements appropriate groups, Mendeleev left blank spaces for elements undiscovered at the time. He also made some assumptions about the correct values of atomic weights

The elements within a group of Mendeleev’s table have similar properties, and these properties change gradually from the top to bottom in the group. As an example, the alkali metal (Group I ) have low melting points that decrease in the order.

Li(174C)>Na(97.8C) > K(63.7C) > Rb(38.9C) > Cs(28.5C)

The alkali metals also have high atomic volumes (low densities) and are extremely reactive toward water, producing hydrogen gas.

At the top of the Mendeleev’s table are listed the formula of the chlorides, hydrides and oxides of the elements (R) in each group. Mendeleev was able to correlate these formula with the group numerals, e.g., sodium chloride has the formula NaCl ; the arsenic – hydrogen compound arsine, AsH3 and an oxide of molybdenum, MoO3.

With the discovery of more and more elements, it was realised that Mendeleev’ Periodic Table should be revised and the new elements should be included in it at appropriate places. Hence in order to accommodate the new elements and to remove some of the defects of the original table, the original table proposed by Mendeleev was modified.

Characteristics of Mendeleev’s Periodic Table

Inert gases were not known when the Mendeleev periodic table was formulated. With the discovery of inert gases, a new group called zero group was added to Mendeleev’s original table. The modified version of Mendeleev’s table is given below.

It is evident that modified Mendeleev’s periodic table consists of :

(i) Nine vertical columns called groups. These are designated as I, II, III, IV, V, VI, VII , VIII and zero. Except for the groups VIII and zero, each group is further divided into sub-groups designated as A and B. The elements which lie on the left hand side of each group constitute sub-group A while those placed on the right hand side form sub-group B. This sub-division is made on the basis of the difference in their properties. Group VIII contains nine elements in three sets each containing three elements. Group zero consists of inert gases.

(ii) Seven horizontal rows called periods : These are

numbered from 1 to 7. First period contains only two elements and is thus the shortest period. Second and third period contain 8 elements each. These are called short periods. Fourth and fifth periods contain 18 elements each and are called long periods. Sixth period contains 32 elements and hence is called longest period. Seventh period is however, incomplete and contains 24 elements.

Significance of Mendeleev's Periodic Table

Some important contributions of Mendeleev’s periodic Table are :

i) Systematic study of elements : It has made the study of 105 elements quite convenient. Knowing the properties of one element in the group, the properties of other elements in the group can be guessed. Thus it becomes very useful in studying and remembering the properties of a large number of elements.

ii) Prediction of new elements : At the time of Mendeleev only 56 elements were known. While arranging these elements , he left certain vacant places . These gaps represented the undiscovered elements. Mendeleev predicted the properties of these undiscovered elements on the basis of their positions. For example, he predicted the properties of scandium, gallium and germanium which were discovered later. The observed properties of these elements were found to be similar to those predicted by Mendeleev.

iii) Correction of atomic masses : The Mendellev’s periodic table helped in correcting the atomic masses of some elements based on their positions in the table. For example, atomic mass of beryllium was corrected from 13.5 to 9. With the help of this table , atomic masses of indium, gold, platinum etc., were corrected.

iv) Valence of Elements : Mendellev‘s classification helped in understanding the valence of elements. The valence of the elements is given by the group number.

Defects in Mendellev's periodic Table

In spite of its useful role in the study of chemistry , Mendeleev’s Table possessed many drawbacks. Some of these drawbacks are :

i) Position of hydrogen : Hydrogen is placed in Group IA. However, it actually resembles the elements of Group-IA (alkali metals ) as well as the elements of Group VII-A (halogens ). Thus, the position of hydrogen in the periodic table is not clear.

ii) Position of isotopes : On the basis of atomic weight , various isotopes of the same elements should be assigned different places in the periodic table. Mendeleev could not provide separate places for isotopes.

iii) Position of lanthanides and actinides: Fourteen elements following Lanthanum( known as lanthanides or rare earths) and the fourteen elements following Actinium (known as actinides or transuranic elements ) have not been provided separate and proper places in the Mendellev’s table, rather they have been placed in two rows at the bottom of the table.

iv) Disimilar elements placed together : Noble metals like Cu, Ag and Au are placed along with chemically dissimilar alkali metals in Group I. Similarly, Mn possessing very few similarities with halogens have been placed in VII group.

v) Similar elements separated : In Mendeleev’s periodic table, certain chemically similar elements such as copper and mercury ; gold and platinum have been placed in different groups.

vi) Anomalous pairs : In the Mendeleev’s Table based on atomic weight, the positions of certain pairs , e.g. Argon( at. wt = 39.94) and potassium ( at. wt = 39.1) : Cobalt( at wt =58.93 ) and nickel ( at wt = 58.71 ) ; Tellurium at wt = 127.60) and iodine (atomic weight = 126.90 ) would be reversed. In other words, certain pairs of elements are misfit in the periodic table, if atomic weight is the basis of classification.

ATOMIC NUMBER AND MODERN PERIODIC LAW

In 1913, Moseley, an English physicist measured the frequencies of X-rays emitted by a metal when bombarded with high speed electrons. He discovered that the square root of frequency () of the prominent X-rays emitted by a metal was proportional to the atomic number and not to the atomic weight of the atom of the element.

 = a ( Z b)

where ‘a’ is the proportionality constant and ‘b’ is a constant for all the lines in a given series of X-rays.

Therefore he concluded that atomic number was a better fundamental property of an element than its atomic weight. He therefore suggested that atomic number (Z) instead of atomic weight should be the basis of classification of the elements. This forms the basis of the modern periodic law.

Thus, Modern periodic law states that physical and chemical properties of the elements are periodic function of their atomic numbers, i.e., if elements are arranged in the order of their increasing atomic numbers, the elements with similar properties are repeated after certain regular intervals.

Theoretical justification of the Modern Periodic Law

An atom consists of a nucleus (which contains protons and neutrons) surrounded by electrons. Atomic mass is a nuclear property and depends upon the number of protons and neutrons in the nucleus, whereas atomic number implies the number of electrons in the extranuclear part or the number of protons in the nucleus. Now, the chemical properties of the elements depend, among other things, upon the interaction between atom and the reagent. Since nucleus is deep seated in an atom and is also shielded by electrons in the extranuclear part so atomic mass has little effect on the chemical properties of elements. Electrons, however, are exposed to the environments and hence can interact with the reagent. As a result, the physical and chemical properties of elements depend upon their atomic numbers rather than atomic masses. Further it is understandable that the physical and chemical properties of elements could be different depending upon the number of electrons and their electronic configuration in any atom.

CAUSE OF PERIODICITY

According to the Modern Periodic Law, the properties of the elements are repeated after certain regular intervals when these elements are arranged in order of their increasing atomic numbers. Further an examination of the electronic configuration of the various elements clearly indicates that with a gradual increase in atomic number there occurs a repetition of similar outer electronic configurations after certain regular intervals. By correlating these two observations, we can say that, the cause of periodicity in properties is the repetition of similar outer electronic configurations at certain regular intervals.

For example, all the elements of Group I A ( or 1) i.e., alkali metals have the similar outer electronic configuration , i.e., ns1 where n refers to the number of the outer most principal shell. These electronic configurations are given in the Table.

Element Atomic number Electronic configuration

Li 3 1s22s1

Na 11 1s22s22p63s1

K 19 1s22s22p63s23p64s1

Rb 37 1s22s22p63s23p63d104s24p65s1

Cs 55 1s22s22p63s23p63d104s24p64d105s25p66s1

Fr 87 1s22s22p63s23p63d104s24p64d104f105s25p6 5d106s26p67s1

Thus it is because of similarity in electronic configuration that all the elements have similar properties.

In a similar manner all the halogens , i.e., elements of group VII ( or 17) have similar outer electronic configurations i.e., ns2np5 and as such possesses similar properties. The electronic configurations of halogens are given in the following Table.

Element Atomic number Electronic configuration

F 9 1s22s22p5

Cl 17 1s22s22p63s23p5

Br 35 1s22s22p63s23p64s24p5

I 53 1s22s22p63s23p63d104s24p64d105s25p5

At 85 1s22s22p63s23p63d104s24p64d104f145s2 5p6 5d106s26p5

Again all the elements of group 0 ( or 18) , i.e., inert gases have similar outer electronic configuration, i.e., ns2np6 and also have similar properties.

It follows from above two Tables that there is a repetition in the electronic configurations of alkali metals and halogens only after certain regular intervals. These regular intervals being 2, 8, 8, 18 , 18 and 32. These numbers are also sometimes called magic numbers.

THE PERIODIC TABLE

A table in which elements are arranged in the order of increasing atomic number in the manner that the elements with similar properties fall in the same vertical column is known as Periodic Table. The elements which fall in the same column and resemble one another in their properties are said to belong to the same group or family of elements. Thus, elements belonging to one particular group or family resemble one another quite closely but differ from the elements belonging to the other groups quite appreciably.

The Long form of the Periodic Table – Bohr’s Table

The objective of the periodic table is to organise and systematise the chemistry of the elements. A broad study of characteristics of various groups gives us a fairly good insight into the general behaviour and important properties of the most of the elements.

The long form of the periodic table helps us to understand the reason why certain elements resemble one another and why they differ from other elements in their properties. The arrangement of elements in this table is also in keeping with their electronic structures.

The long form of the periodic table is also called Bohr’s periodic table, since it is based on Bohr’s scheme of classification of the elements into four types depending on the number of incomplete shells of electrons in the atom. This table was proposed by Rang(1893) and then modified by Werner(1905) and extended by Bury(1921).

Another significant feature of the long form of the periodic table is that it helps us to understand the cause of periodicity of the properties and the reason why similar properties reccur at certain regular intervals ie., after 2, 8, 8, 18, 18, and 32 elements. The numbers 2, 8, 18, and 32 are sometimes referred as magic numbers.

The long form periodic table is constructed by arranging elements in the order of increasing atomic number. The vertical rows constitute Groups while the horizontal rows constitute Periods.

Periods

A horizontal row of a periodic table is called period. A period consists of a series of elements having the same valence shell. There are seven periods in all , which are numbered as 1, 2, 3, 4, 5, 6 and 7. There is a close connection between electronic configurations of the elements and the long form of the periodic table. The principal quantum number ‘n’ defines the main energy level of the electron. Each period of the periodic table begins with the filling of new energy shell. Each period of the periodic table represents the principal quantum number of the valence shell of elements present in it. The number of elements in each period is equal to the number of electrons which can be accommodated in the orbitals belonging to that electron shell.

The first period corresponds to the filling of electrons in the first energy shell ( i.e., n =1). The energy level has only one orbital ( i.e., 1s ) and therefore it can accommodate two electrons. This means that there can be only two elements in this period. They are hydrogen and helium.

The second period starts with the electrons beginning to enter the second energy shell ( n = 2 ) . Since there are only four orbitals ( one 2s and three 2p-orbitals) to be filled which can accomodate eight electrons. Thus the second period has eight elements in it. They are Li, Be, B, C, N, O , F and Ne.

The third period begins with the electrons entering the third energy shell ( n = 3 ) . Out of the nine orbitals of this energy level ( one s, three p and five d ) , the five 3d orbitals have higher energy than 4s orbitals. As such only four orbitals ( one 3s and three 3p ) corresponding to n = 3 are filled before the fourth energy level begins to be formed . Hence there are only eight elements in the third period.

The fourth period corresponds to n = 4. It starts the filling of 4s orbitals. However, after the 4s but before the 4p orbitals , there are five 3d orbitals also to be filled. Thus, in all , nine orbitals ( one 4s, five 3d and three 4p ) have to be filled and as such there are eighteen elements in the fourth period. The filling of 3d-orbitals starts from Sc ( Z = 21). The elements from Sc ( Z = 21 ) to Zn ( Z = 30 ) are called 3d-transition series.

The fifth period beginning with 5s-orbital ( n = 5 ) is similar to fourth period. There are nine orbitals ( one 5s, five 4d and three 5p ) to be filled and therefore , there are eighteen elements in the fifth period.

The sixth period starts with the filling of 6s-orbital( n = 6 ) . There are sixteen orbitals ( one 6s, seven 4f, five 5d and three 6p ) in which there are thirty two elements in the sixth period. The filling of elements takes place before the next energy level starts. As such there are thirty two elements in the sixth period. The filling up of 4f orbitals begins with Cerium ( Z = 58 ) and ends at Lutetium ( Z = 71 ). It constitutes the first f-inner transition series which is called Lanthanide or Lanthanoid Series.

The seventh period begins with 7s-orbital( n = 7) . It would also have contained 32 elements corresponding to the filling of 7s, 5f, 6d and 7p orbitals. At present there are only 29 elements in it. The filling up of 5f-orbitals begins with thorium ( Z = 90 ) and ends up with Lawrencium ( Z = 103). It constitutes second f-transition series which is called Actinide or Actinoid Series. It mostly includes man-made radioactive elements. In order to avoid undue expansion of the periodic table, the 4f and 5f-transition elements have been placed in separate panels at the bottom.

The relationship between number of elements in a period and electron filling of orbitals are summed up below.

Number of elements in different periods

Pereiod Principal valence shell (n) Orbitals being filled up Electrons to be accomodated Number of elements

First n = 1 1s 2 2

Second n = 2 2s, 2p 2+6 =8 8

Third n = 3 3s,3p 2+6=8 8

Fourth n = 4 4s,3d,

4p 2+10+6 18

Fifth n = 5 5s,4d,

5p 2+10+6 18

Sixth n = 6 6s,4f,

5d,6p 2+14

+10+6 32

Seventh n = 7 7s,5f,

6d,7p 2+14

+10+6 32*

*At present 7th period contains only 24 elements.

The periods 2 and 3 contain 8 elements each and are called Short Periods. There are 18 elements in 4th and 5th periods and are called long Periods. Sixth period containing 32 elements is called the longest Period.

GROUPS

A vertical column of the periodic table is called a Group. A group consists of a series of elements having similar configuration of the outer energy shell . There are eighteen vertical columns in the long form of the periodic table. According to the recommendations of the International Union of Pure and Applied Chemistry (IUPAC) , these groups are numbered from 1 to 18. The elements belonging to the same group are said to constitute a family.

Naming elements with large atomic numbers greater than 100

It has been historical practice to allow the discoverer of an element to assign the element’s name. In recent times this has led to some controversy, because elements with very high atomic numbers are so unstable that only minute quantities of them (sometimes only one or two atoms) are prepared before scientists claim credit for their discovery. This has led on occasion to questions of reliability for the data and whether the sought-after element has in fact been made. For example, both Americans and Soviet scientists claim credit for discovering element 104. The Americans named it Rutherfordium and Soviets named it Kurchatovium.

Because of this problem, the international Union of Pure and Applied Chemistry (IUPAC) has made the official recommendation that until a new element’s discovery has been proved, a systematic nomenclature be applied. The rules of this system are as follows:

1. All elements will end in the letters -ium.

2. The name will be constructed from the following numerical roots :

Digit Name Abbreviation

0 nil n

1 un u

2 bi b

3 tri t

4 quad q

5 pent p

6 hex h

7 sept s

8 oct o

9 enn e

3. The symbol will consist of three letters derived from the numerical roots above.

The following illustrates how this works for element 104.

The name for the element is unnilquadium and its symbol is Unq.

Nomenclature of Elements with atomic number above 100

Atomic

Number Name Symbol IUPAC official name IUPAC symbol

101 Unnilunium Unu Mendelevium Md

102 Unnilbium Unb Nobelium No

103 Unniltrium Unt Lawrencium Lr

104 Unnilquadium Unq Rutherfordium Rf

105 Unnilpentium Unp Dubnium Db

106 Unnilhexium Unh Seaborgium Sg

107 Unnilseptium Uns Bohrium Bh

108 Unniloctium Uno Hassnium Hs

109 Unnilennium Une Meitnerium Mt

110 Ununnilium Uun Darsmstadtium Ds

111 Unununnium Uuu Rontgenium Rg

112 Ununbium Uub * *

113 Ununtrium Uut +

114 Ununquadium Uuq * *

115 Ununpentium Uup +

116 Ununhexium Uuh +

117 Ununseptium Uus +

118 Ununoctium Uuo +

* Official IUPAC name yet to be announced

+ Elements yet to be discovered.

General Characteristics of Groups

1. Number of valence electrons : On moving down the group the number of valence electron does not change, i.e., remains the same.

2. Valency : The valencies of all elements of the same group are the same. The valence of an element with respect to oxygen is equal to its group number.

3. Properties of elements : All the elements of a given group possess very similar physical and chemical properties. There is a regular gradation in their properties when we move from top to bottom in a group. For example:

(a) The alkali metals (Group 1) resemble each other and their base forming tendency increases from Li to Cs.

(b) The reactivity of halogens (Group 17) decreases as we pass from F to I.

4. Size of atoms : Size of atoms increases on descending a group. For example, in group 1, atomic size increases from Li to Cs. Thus :

Li < Na < K < Rb < Cs 5. Metallic character : The metallic character of the elements increases in moving from top to bottom in a group. This is particularly apparent in Groups 14, 15 and 16 which begin with non-metals ( viz., C, N and O respectively) and end with metals (namely Pb, Bi and Po respectively). For example in Group 15 nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids and bismuth is a typical metal. Thus the metallic character of these elements increases from nitrogen to bismuth as shown below: Elements of Group 15 N, P As, Sb Bi Non-metals metalloids metal Metallic character : metallic character increasing  It is because of a gradual increase of metallic character of the elements from top to bottom that the oxides of the elements become more and more basic in the same direction. For example : Oxides of N2O3, P2O5 As2O3,Sb2O3 Bi2O3 Group 15 acidic amphoteric basic Basic character : Basic character increasing  6. Number of electron shells : In going down a group the number of electron shells increases by one at each step and becomes equal to the number of the period to which the element belongs as shown below for elements of Group 1. Elements Electronic confn. Numberof shells Li 2,1 2 Na 2,8,1 3 K 2,8,8,1 4 Rb 2,8,18,8,1 5 Cs 2,8,18,18,8,1 6 Fr 2,8,18,32,18,8,1 7 General Characteristics of Periods 1. Number of valency electrons : Number of valence electrons increases from 1 to 8 when we proceed from left to right in a period. 2. Valency : The valence of the elements with respect to hydrogen in each short period increases from 1 to 4 and then falls to one while the same with respect to oxygen increases from 1 to 7 as shown below for 2nd and 3rd period. Hydrides of second period elements . Valency with respect to hydrogen is shown in brackets. Li Be B C N O F LiH (1) BeH2 (2) BH3 (3) CH4 (4) NH3 (3) H2O (2) HF (1) Elements of 3rd period . Oxides of elements. Valency with reference to oxygen is shown in brackets. Na Mg Al Si P S Cl Na2O (1) MgO (2) Al2O3 (3) SiO2 (4) P2O5 (5) SO3 (6) Cl2O7 (7) 3. Size of atoms : Size of atoms decreases from left to right in a period. Thus alkali metals have the largest size while halogens have the smallest size. 4. Properties of elements : The properties of elements of a given period differ considerably but elements in adjacent periods show marked similarity beteen them. For example, when we consider the elements of 2nd and 3rd periods, we find that Na resembles Li, Mg resembles Be, Al resembles B, Si resembles C, P resembles N, S resembles O, Cl resembles F and Ar resembles Ne. 5. Metallic character : On moving from left to right in a period the metallic character of the elements decreases. For example, in 3rd period, Na, Mg and Al are metals, while Si, P, S and Cl are non-metals as shown below: Elements of 3rd period : Na, Mg, Al Si,P,S,Cl Metals Non-metals Metallic character : Metallic character decreasing  It is becase of the gradual decrease of the metallic character from left to right that the oxides of elements become less and less basic in the same direction. For example : Basic character : Basic character decreasing  6. Number of shells : In going from left to right in a period the number of electron shells remains the same and the number of a period corresponds to the number of the shells found in the elements of that period, e.g. all the elements of 2nd period have the electrons only in the two shells. 7. Diagonal relationship : The earlier members of the 2nd period of the periodic table viz., Li, Be and B are said to be diagonal neighbours of the elements viz. Mg , Al and Si lying on the right hand side in the 3rd period. Thus, Li which is diagonally placed to Mg in the periodic table is the diagonal neighbour of Mg. Similarly , Be and B are diagonal neighbours of Al and Si respectively. Each diagonal pair forms a sort of bridge between the periods to which the elements of the diagonal pair belong and hence these elements are bridge elements. The diagonal pairs viz. Li – Mg, Be –Al and B – Si are shown below : It has been noted that both the members of a given diagonal pair show many similar properties with each other. Thus Li resembles Mg, Be resembles Al and B resembles Si. This similarity in properties existing between two members of a diagonal pair is known as diagonal relationship. Merits of Long Form of Periodic Table Over Mendeleev’s Periodic Table The long form of the periodic table has a number of merits over Mendeleev’s periodic table in the following respects: (i) The classification of elements based on a more fundamental property viz. Atomic number. (ii) It relates the position of an element to its electronic configuration. Each group contains elements with similar electronic configuration and hence similar properties. (iii) It explains the similarities and variations in the properties of elements in terms of their electronic configurations and brings out clearly the trends in chemical properties across the long periods. (iv) The noble gases having completely filled electron shells have been placed at the end of each period. Such location of noble gases represents a logical completion of each period. (v) In this form of periodic table, the elements of two sub-groups have been placed separately and thus dissimilar elements do not fall together. (vi) It provides a clear demarcation of different types of elements, e.g. active metals, transition metals, non-metals, metalloids, noble gases, lanthanides and actinides. (vii) It is easier to remember, understand and reproduce. Defects of Long Form of Periodic Table Although the long form of the periodic table is superior to Mendeleev’s periodic table in many respects, it retains some of the defects as such. For example : (i) The problem of position of hydrogen still remains unsolved. (ii) Like the Mendeleev’s table, it fails to accommodate the lanthanides and actinides in the main body of the table. (iii) The arrangement is unable to reflect the electronic configuration of many elements. Problems 01. What would be the systematic name and symbol for element114 ? 02. What would be the IUPAC name and symbol for the element with atomic number 120 ? Division of elements into s, p, d and f blocks The aufbau principle and the electronic configuration configuration of atoms provide a theoretical foundation for the periodic classification. The elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behaviour. This similarity arises because these elements have the number and same distribution of electrons in their outermost orbitals. We can classify the elements into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. This is illustrated in Fig. We note two exceptions to this categorisation. Strictly , helium belongs to the s-block but its positioning the p-block along with other group 18 elements is justified because it has a completely filled valence shell (1s2) and as a result , exhibits properties characteristic of other noble gases. The other exception is hydrogen. It has a lone s-electron and hence can be placed in Group 1 (alkali metal). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a Group 17 (halogen) element. Because it is a special case , we shall place hydrogen separately at the top of the Periodic Table as shown in Fig. 1. The s-Block Elements The elements of Group 1 (alkali metals) and Group 2(alkaline earth metals) which have ns1 and ns2 outermost electronic configuration belong to the s-Block Elrments. • They are all reactive metals with low ionisation enthalpies. • The lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals). • The metallic character and reactivity increase as we go down the group. • Because of high reactivity , they are never found pure in nature. • The compounds of s-block elements , with the exception of those of lithium and beryllium are predominantly ionic. 2. The p-Block Elements The p-Block elements comprise those belonging to Group 13 to 18 and these together with s-Block Elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns2np1 to ns2p6 in each period. At the end of each period is a noble gas element with a closed valence shell of the ns2np6 configuration. All the orbitals in the valence shell of noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity. Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16). These two groups of elements have high negative electron gain enthalpies and readily add one or two electrons respectively to attain stable noble gas configurations. The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group. 3. The d-Block Elements (Transition Elements) These are the elements of Group 3 to 12 in the centre of the periodic table. These are characterized by the filling of orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n  1) d1 – 10 ns0-2. They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts. However, Zn, Cd and Hg which have electronic configuration , (n  1) d10 ns2 do not show most of the properties of transition elements. In this way , transition metals form a bridge between the chemically active metals of s-block elements and the less active elements of Groups 13 and 14 and thus take their familiar name ‘transition elements’. 4. The f-Block Elements(Inner transition elements) The two rows of the elements at the bottom of the Periodic Table , called Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids (Th ( Z = 90) – Lr(Z = 103) are characterised by the outer electronic configuration (n2) f114 ( n1)d01 ns2. The last electron added to each element is filled in the f-orbital. These two series of the elements are hence called the inner transition elements (f-Block Elements). They are all metals. Within each series, the properties of these elements are quite similar. The chemistry of the early actinoids is more complicated than corresponding lanthanoids due to large number of oxidation states possible for these actinoid elements. Actinoid elements are radio active. Many of the actinoid elements have been made only in nanogram quantities or even less by nuclear reactions and their chemistry is not fully studied. The elements after uranium are called Transuranium Elements. Metal, Non-metals and Metalloids In addition to displaying the classification of elements into s- , p- and d- and f-blocks , elements can be divided into metals and non-metals. Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. Metals are usually solids at room temperature [mercury is an exception ; gallim and caesium have low boiling points(303 K and 302 K) respectively] . Metals usually have high melting and boiling points. They are good conductors of heat and electricity. They are malleable(can be flattened into thin sheets by hammering) and ductile(can be drawn into wires). In contrast non-metals are located at the top right hand side of the Periodic Table. Non-metals are usually solids or gases at room temperature with low melting and boiling points. They are poor conductors of heat and electricity. Most non-metallic solids are brittle and are neither malleable nor ductile. The elements become more metallic as we go down a group ; the non-metallic character increases as one goes from left to right across the Periodic Table. The change from metallic to non-metallic character is not abrupt as shown as shown by thick zig-zag line in the following Figure ; the elements (e.g. germanium, silicon, arsenic, antimony and tellurium ) bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metals and non-metals. These elements are called Semi Metals or Metalloids. Problems 3. Write the electronic configurations of elements with atomic numbers 3,10, 16 and 17 and 24. Which of these is noble gas, chalcogen, alkali metal, halogen and transition element ? Write the name of the element and block to which the element belongs . 4. (a) How would you justify the presence of 18 elements in 5th period of the Periodic of the periodic table. (b) The elements Z = 117 and 120 have yet been discovered. In which family/group would you place these elements and also give the electronic configuration in each case. 5. Arrange the following elements in the increasing order of metallic character : Si, Be, Mg, Na, P. 6. Write the electronic configurations of the elements given below and also predict their period , group and block. A ( Z = 16 ) ; B = ( Z = 37 ); C ( Z = 30) 7. Elements A, B ,C and D have atomic numbers 12, 19, 29 and 36. On the basis of electronic configurations write to which group of the periodic table the element belong. Predict the blocks to which these elements can be classified ? Which of these are representative elements ? 8. The elements Z = 107 and Z = 109 have been made recently; element Z = 108 has not yet been prepared. Indicate the group in which you will place the above elements. 9. Find the number of period and group and block in which the element with atomic number 24 is present. 10. The electronic configurations of some elements are given below. Group these elements which you expect should have similar chemical properties. (i) 1s22s22p63s1 (ii) 1s22s22p6 (iii) 1s22s22p63s23p5 (iv) 1s22s22p63s23p64s1 (v) 1s22s22p63s23p6 (vi) 1s22s22p63s23p63d104s24p5 Procedure to predict the Number of Period. Group and Block of an Element In order to predict the period, group and block of an element proceed as follows : (i) Write the electronic configuration of the element. (ii) Number of the period is the highest value of shell which is present in the electronic configuration of element concerned. (iii) The name of the block is that sub-shell in which the last electron of the element is filled. (iv) Group number of the element can be predicted from the eectronic configuration as follows : Thus, For s-block elements, group number is equal to number of valence electrons (n s) For p-block elements, group number is equal to 10 + number of valence electrons ( n s and n p) For d-block elements , the group number is equal to the number of electrons in ( n  1 ) d and n s sub-shells. PERIODIC PROPERTIES OF ELEMENTS Most the properties of the elements such as valence , atomic size, ionisation energy and electron affinity are directly related to the electronic configurations of the atoms. These properties undergo periodic variation with the change in atomic number within a period or a group. These properties indirectly control the physical properties such as melting point , boiling point, density etc. 1. Valency or oxidation states The valence is the most characteristic property of the elements and can be understood in terms of their electronic configurations. The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and / or equal to eight minus the number of outermost electrons as shown below. Nowadays the term oxidation state is frequently used for valence. Consider two oxygen containing compounds OF2 and Na2O. The order of electronegativity of the three elements involved in these compounds is F > O > Na. Each of the atoms of fluorine, with outer electronic configuration 2s22p5, shares one electron with oxygen in the OF2 molecule. Being highest electronegative element, fluorine is given oxidation state 1. Since there are two fluorine atoms in this molecule, oxygen with outer electronic configuration 2s22p4 shares two electrons with fluorine atoms thereby exhibits an oxidation state +2. In Na2O, oxygen being more electronegative accepts two electrons , one from each of the two sodium atoms and , thus shows oxidation state of 2. On the other hand sodium with electronic configuration 3s1 loses one electron to oxygen and is given oxidation state +1. Thus oxidation state of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.

Some periodic trends observed in the valence of elements (hydrides and oxides) are shown in the Table.

Group 1 12 13 14 15 16 17 18

Number of valence electrons

1

2

3

4

5

6

7

8

Valence 1 2 3 4 3,5 2,6 1,7 0,8

Periodic Trends in valence of Elements as shown by the formula of their compounds

Group 1 12 13 14 15 16 17

Formula of the hydride LiH

NaH

KH

CaH2 B2H6

AlH3 CH4

SiH4

GeH4

SnH4 NH3

PH3

AsH3

SbH3 H2O

H2S

H2Se

H2Te HF

HCl

HBr

HI

Formula of oxide Li2O

Na2O

K2O MgO

CaO

SrO

BaO B2O3

Al2O3

Ga2O3

In2O3 CO2

SiO2

GeO2

SnO2

PbO2 N2O3,

N2O5

P4O6, P4O10

As2O3, As2O5

SO3

SeO3

TeO3

Cl2O7

2. Atomic volume : Atomic volume is defined as the volume in centimetre cube occupied by one gram atom of the element in the solid state and hence is commonly called gram atomic volume. It is obtained by dividing the atomic weight of the element by its density ; i.e.,

Atomic volume is the volume in c.c occupied by 6.023 x 1023 atoms(Avogadro’s number) of an element. Thus the atomic volumes of the elements should be the volumes occupied by the same number of atoms. However, this conclusion is not very accurate due to certain variable factors, which in turn, depend on the manner in which atoms are packed in various elements. It is a common practice to use the term gram atomic volume for volume per atom for assignement of sizes of atoms.

Lothar Meyer studied the atomic volumes of a large number of elements, He plotted the values of atomic volumes of various elements against their atomic numbers and obtained a curve called atomic volume curve. The curve is shown in Fig.

Fig. Atomic volume versus atomic number curve

The atomic volume curve consists of many sharp peaks and broad minima and reflects remarkable periodicity in the atomic volumes of the elements. On the basis of atomic volume curve, the following generalizations can be made.

(i) Each peak is occupied by the first element of the period i.e., alkali metals Li, Na, K, etc.

ii) The less reactive transition metals, e.g., Fe, Mn, Co, Ni etc. occur at the minima.

iii) The electronegative elements occupy the ascending slope of the curves, whereas electropositive elements are present on the descending slopes of the curve.

iv) The elements showing high melting points (C, Si, Fe etc) occupy the descending slopes and minima of the curve.

v) The elements of a particular in the periodic table are placed identically on the atomic volume curve.

Variation of Atomic volume

In a Group : Atomic volume increases more or less regularly in going down a group (Table).

Atomic volumes ( in c.c) of s and p-block elements.

IA IIA IIIA IVA VA VIA VIIA Zero

H 13 He 32

Li 13 Be 5 B 5 C 5 N 14 O 11 F 15 Ne 17

Na 24 Mg14 Al 10 Si 12 P 17 S 16 Cl 19 Ar 24

K 46 Ca 26 Ga 12 Ge 13 As 16 Se 34 Br 23 Kr 33

Rb 56 Sr 34 In 16 Sn 16 Sb 18 Te 21 I 26 Xe 43

Cs 71 Ba 38 Ti 17 Pb 18 Bi 21 Po - At - Rn 50

The increase in atomic volume in going down in a group is due to increase in number of shells. The larger the number of shells , the bigger is the volume of the atom.

(b) In a period : In going from left to right in a period it decreases at first for some elements, becomes minimum in the middle and then increases(Table). The variation of atomic volume in going from left to right in a period is influenced by the following two factors :

(i) Nuclear charge : Nuclear charge increases by one, as we move from left to right in the period. The increased nuclear charge attracts each electron more strongly towards the nucleus, resulting in a decrease in volume of the atom.

(ii) Number of valence electrons : Towards the close of a period, due to increase in the number of valence electrons, the volume of the atom increases, so that it may accommodate all the electrons.

These two factors, one causing an increase while the other a decrease, combine to result that in a period atomic volume decreases at first for some elements, becomes minimum in the middle and then increases.

3. ATOMIC SIZE

The atomic size is very important property of the atoms because it is related to many other chemical and physical properties. In dealing with atomic size the atom is assumed to be a sphere and its radius determines the size. In general the atomic radius is defined as the distance from the centre of the nucleus of the atom to the outermost shell of electrons. However, it is not possible to find precisely the radius of the atoms because of the following reasons.

i) The atom is too small to be isolated.

ii) Wave mechanical model of the atom does not allow us to have its well-defined boundary because probability of finding the electrons is never zero even at a very large distance from the nucleus.

iii) The probability distribution of an atom is affected by other atoms present in its neighbourhood.

iv) Size of an atom also changes from one bonding state to another

The approximate radii of the atoms can be determined by measuring the distance between the atoms in a covalent molecule by X-ray diffraction and other spectroscopic techniques. The radius of the atom is referred to as covalent radius and can be defined as one-half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond. For homonuclear molecule:

rcovalent = ½ [Inter-nuclear distance between two bonded atoms]

For example as shown in Fig. the inter-nuclear distance between two hydrogen atoms is 74pm.Therefore covalent radius of hydrogen atom is equal to 37pm.

Similarly, the atomic radii of chlorine and bromine are 99 pm and 114 pm respectively.

It may be noted that atomic radii may be given in different operational names depending upon the bonding state. These are covalent radius, van der waal’s radius and metallic radius. Covalent radius is always shorter than the other two radii because overlapping of orbitals involved in covalent bond formation decreases the inter-nuclear distance.

VAN DER WAAL'S RADII

The van der Waal’s radius may be defined as one half of the distance between the nuclei of two adjacent identical atoms belonging to two neighbouring molecules of an element in the solid state. It is essentially the distance between two non-bonded atoms of two adjacent molecules and is related to the effective packing size of the atom when the element is in the solid state. The name van der Waal’s radius is used because the forces existing between the molecules are the van der Waal’s forces of attraction. The van der Waal’s radii are obtained from X-ray studies of various elements in the solid state. For example, the inter-nuclear distance between two chlorine atoms of nearest neighbouring molecules is 360 pm.

The van der Waal’s radius is therefore 180 pm. The van der Waal’s and covalent radii of some common elements are given in TABLE.

The van der Waal’s and Covalent Radii of Some Elements

Element H N O F Cl Br I

van der wall’s radius ( pm ) 120

150

140

135

180

195

215

Covalent radius

( pm) 37

75 73 72 99 114 133

The covalent radius is smaller than van der Waal’s radius. The reason is that in the formation of covalent bond, the atoms have to come closer to each other. This also explains why covalent bonds are much stronger than van der Waal’s forces.

It may be noted that the noble gases , ordinarily do not form any covalent bonds. In crystals of noble gases, therefore no chemical forces are operating between the atoms. The van der Waal’s forces are the only forces in these cases. Hence , the atomic radii of noble gas atoms are only the van der Waal’s radii which represents the distance of closest approach of the two adjacent atoms of the noble gas considered in the solid state.

Metallic Radius

It is half the distance between the centres of nuclei of two adjacent atoms in the metallic crystal. Metallic radius of the element is greater than the covalent radius. This is due to the reason that a metallic bond is weaker than a covalent bond and hence the inter-nuclear distance between two atoms in a metallic lattice is larger than the inter-nuclear distance between the atoms held by a covalent bond.

For example, metallic and covalent radii for potassium are 231 pm and 203 pm respectively.

Periodic Trends in Covalent Radii

Generally speaking , the covalent radii decrease in moving from left to right in any given period and increase in moving from top to bottom in any given group. Explanation for these variations is given below.

a) Variation of Covalent radii in a period

The variation of covalent radii in a period can be explained on the basis of electronic configurations of the elements. In order to do so the covalent radii of elements of the second period beginning with Lithium and ending with Fluorine are shown in the TABLE.

Covalent radii of elements of second period

Element Atomic number Nuclear charge Electronic configuration covalent

radius (pm)

Li 3 +3 1s2, 2s1 152

Be 4 +4 1s2, 2s2 111

B 5 +5 1s2, 2s2, 2p1 88

C 6 +6 1s2, 2s2, 2p2 77

N 7 +7 1s2, 2s2, 2p3 70

O 8 +8 1s2, 2s2, 2p4 74

F 9 +9 1s2, 2s2, 2p5 72

It may be noted that the nuclear charge in these elements increases from +3 to +9 . Each of these elements contain two electrons in the K-shell and some( ranging from 1 to 7 ) in the L-shell. As the charge on the nucleus increases, step by step, from +3 to +9 on moving from left to right, the K-electrons are attracted more and more strongly towards the nucleus.

Hence, the K-electrons tend to lie closer and closer to the nucleus. In other words there is more and more contraction of the K-shell in moving from left to right across the period.

The L-electrons are also attracted towards the nucleus but the K-electrons lying in between , tend to ‘screen’ them from the nucleus. In the case of Lithium, for example, the nuclear charge is +3 and of the three electrons , two are in the K-shell and one in the L-shell. The electron in the L-shell is attracted towards the nucleus, but not by a free charge of +3 but by a charge of +3 screened by two electrons of the K-shell which lie in between. The net effect is about the same as that of the attractive force of +1 charge. Similarly, in the case of beryllium atom, the two L-electrons are attracted by a nucleus of +4 charge and are screened by the two electrons present in the K-shell. It may be noted that the electrons present in the same shell (L-shell) do not screen each other from the nucleus. The screening effect is thus only due to the K-electrons.

In spite of the screening effect of the intervening electrons , the fact remains that the L-electrons experience greater and greater attraction towards the nucleus as we move from left to right across the period. Thus, like the K-shell, the L-shell also becomes smaller and smaller. Thus, the covalent radius decreases on moving from left to right across a period. This trend is shown along the third and subsequent periods as well.

In the case of noble gases, the atomic radii are only van der Waal’s radii which are naturally larger than the covalent radii of other elements.

Variation of atomic radius with atomic number

across the second period

b) Variation of Covalent Radii in a group

In the order to understand the change in the size of the atoms in the same group, the values of atomic radii for alkali metals shown in the TABLE is considered.

Alkali metals Atomic number Atomic radius (pm)

Li 3 152

Na 11 186

K 19 231

Rb 37 244

Cs 55 262

It has been seen that there is increase in covalent radius in moving from top to bottom. This may be explained as follows. In moving down a group, the number of principal shell increases and therefore the size of the atom increases. This effect is partly annealed by with drawing in the electron shells on account of the increasing nuclear charge. Hence the radius of the atom increases in moving from top to bottom in a group.

Variation of atomic radius with atomic number

for alkali metals and halogens

3. Ionic Radius

Ions are formed when the neutral atoms loose or gain electrons. A positive ion is formed by the loss of one or more electrons by the neutral atoms whereas a negative ion is formed by the gain of one or more electrons by the atom. The term ionic radii refers to the size of the ions in the ionic crystals. Ionic radius may be defined as the effective distance from the nucleus of the ion to the point up to which it has influence in the ionic bond.

The equilibrium distance between the nuclei of the two adjacent ions can be determined by X-ray analysis of ionic crystals. Assuming ions to be spheres , the inter-nuclear distance can be taken as the sum of the ionic radii of the adjacent ions. Knowing the ionic radius of one of the ion, the radius of the other can be calculated.

The study of ionic radii leads to two important generalisations :

i) The size of the cation is smaller as compared with that of the parent atom.

ii) The size of the anion is larger than in comparison with that of the parent atom.

Radii of Cations and Anions

A positive ion , i.e., a cation results from the loss of one or more electrons from the outer shell called valence shell of an atom. This causes in many cases the removal of the whole of the outer shell of electrons. For example, in Lithium there is only one electron in the outer shell (L) . This electron is removed in the change :

The L-shell (2s) therefore disappears. In magnesium, there are two electrons in the outer (M) shell. Both these electrons are removed in the change.

Mg  Mg2+ + 2e

1s2, 2s2, 2p6, 3s2 1s2, 2s2, 2p6

130 pm 65 pm

The M-shell therefore disappears in this case also.

Thus in the formation of positive ion, the outer shell of electrons is generally removed completely. The cation therefore is much smaller than the corresponding atom. At the same time, the elimination of one or more orbital electrons from an atom, the number of electrons decreases while the nuclear charge remains the same. The nuclear charge, therefore now acts on lesser number of electrons. In other words , the effective nuclear charge per electron increases and electrons are therefore pulled in more towards the nucleus than before. This effect also tends to decrease the radius of the cation.

The decrease in radius , when an atom changes into a cation is shown below in the case of sodium, lead, aluminium, iron and manganese.

Atomic radius of Na = 154 pm

Ionic radius of Na+ = 95 pm

Atomic radius of Pb = 147 pm

Ionic radius of Pb2+ = 120 pm

Atomic radius of Al = 118 pm

Ionic radius of Al3+ = 50 pm

Atomic radius of Mn = 126 pm

Ionic radius of Mn2+ = 80 pm

Ionic radius of Mn4+ = 46 pm

Atomic radius of Fe = 126 pm

Ionic radis of Fe 2+ = 76 pm

Ionic radius of Fe3+ = 64 pm

It is evident that the radius of a cation is invariably smaller than that of the corresponding atom. The size decreases further with the loss of subsequent electrons when cations of higher valencies are formed.

During the change of the atom into negative ion, i.e., an anion, one or more electrons are added to the valence shell of the atom. As a result, the same nuclear charge acts on a large number of electrons. In other words, effective nuclear charge per electron is reduced and therefore, the electrons are held less tightly by the nucleus and the electron cloud expands. Thus, the radius of a negative ion is invariably larger than that of the corresponding atom. This is shown below in the case of chlorine, bromine, iodine, oxygen, sulphur and nitrogen.

Atomic radius (pm) Ionic radius(pm)

Cl 99 Cl 181

Br 114 Br 195

I 133 I 216

O 73 O2 140

S 102 S2 184

N 75 N3 171

Variation of ionic radii in isoelectronic ions

Isoelectronic ions have the same number of electrons but they differ in charge on their nuclei. The examples are furnished by ions given the TABLE.

Radii of iso-electronic ions

Ion N3 O2 F Na+ Mg2+

Number of electrons 10 10 10 10 10

Charge on the nucleus +7 +8 +9 +11 +12

Radius ( pm ) 171 140 136 95 60

The effect of nuclear charge on the ionic size is best illustrated by considering the radii of these iso-electronic species. All of them contain the same number of electrons. They differ only in the charge on the nucleus. It is evident that as the nuclear charge increases, the electrons are held more tightly by the nucleus with the result that the ionic radius decreases.

Periodic Trends in Ionic Radii

Generally speaking the ionic radii increase in moving from top to bottom in any given group and decrease in moving along a period. The ionic radii of some common ions are given in TABLE.

Ionic radii (pm) of some common ions

1 2 3 16 17

Li+

68 Be2+

30 B3+

20

Na+

98 Mg2+

65 Al3+

45 O2

145 F

133

K+

133 Ca2+

94 Ga3+

60 S2

190 Cl

181

Rb+

148 Sr2+

110 In3+

81 Se2

202 Br

196

Cs+

167 Ba2+

129 Tl3+

91 Te2

222 I

219

a. Variation in a Group

It is evident from the data given in the above TABLE that ionic radii increase with increase in atomic number with in a Group. Since the atomic size increases as we move down the group, therefore the ionic radius also increases as we move down the Group.

It can also be seen that within a Group, there is a rapid increase in ionic radius as we move from one element to another amongst the first few elements but the increase is not so rapid when we do so amongst the last members. For example, the increase in ionic radius is quite rapid as we move from Li+ to Na+ and from Na+ to K+ but the increase is not so rapid as we move from K+ to Rb+ and from Rb+ to Cs+. A possible explanation for this behaviour can be that between K+ and Rb+ come the elements of the first transition series. In these elements, the last occupied shell is same, their increasing nuclear charge tends to cause contraction of the atoms and the ions. Thus, the ions which follow any of the transition series would be smaller than if only eight elements had separated them from the lighter members of the family.

b. Variation in a Period

Within a period , the radii of the cations of elements of Group 1 - 3 decreases with increase in atomic number and so do the radii of anions of elements of Group 16 and 17 . Since atomic size decreases as we move along a period, the ionic radii also decreases correspondingly.

4. Ionisation Enthalpy

The ionisation energy of an element is defined as the amount of energy required to remove an electron from an isolated gaseous atom of that element resulting in the formation of a positive ion. This gives a direct measure of the ease with which an atom can change into a cation as represented below:

M ( g ) + IE  M+ (g) + e

The energy required to bring about the above change is called Ionisation Energy . Evidently , the smaller the value of ionisation energy, the easier it is for the neutral atom to change into a positive ion.

Ionisation energies are generally measured by spectroscopic methods. Another method is to have the vapours of the element in a discharge tube and connect it to a source of current. The voltage applied is gradually increased. At certain voltage there will be a sudden rise in the current passing through the tube. The energy corresponding to this voltage is known as the First Ionisation Energy ( IE1). The sudden rise in the current is due to the liberation of an electron from each neutral atom producing an ion M+.

If the applied voltage is increased further, there may again be a stage when the current shows a sudden rise. This is due to the elimination of another electron from each positively charged ion (M+) produced earlier. The energy corresponding to this stage is known as the Second Ionisation Energy (IE2 ) . At this stage doubly charged ions M2+ are produced. If the potential is increased even beyond this point, there may again be sudden and sharp rises in the current at certain points. These points corresponds to loss of three or more electrons and give the third (IE3) or higher ionisation energies of the element. These may be represented as follows:

Here IE1, IE2 and IE3 stand for first, second and third ionisation energies respectively.

The ionisation energy is measured in electron volt (eV ) as well as in Joules or kilojoules ( kJ ).

The values presented in electron volts give ionisation energy per atom while those expressed in kilojoules represent ionisation energy per mole( i.e., per Avogadro’s number of atoms ) of the element. The ionisation energy of hydrogen is 13.595 eV per atom of hydrogen and 13.595 x 1.602 x 1022 x 6.02 x 1023 = 1312 kJ per mol of hydrogen Let us consider the values of IE1 and IE2 for Lithium and sodium given below.

The second ionisation energies are very much higher than the first ionisation energies.

If one electron has taken out of the atom, it becomes increasingly difficult to remove the second and subsequent electrons from the resulting positively charged ion on account of electrostatic attraction. This is due to the fact that after the removal of an electron, the number of electrons decreases while the nuclear charge remains the same. Consequently the remaining electrons are held more tightly by the nucleus and it becomes difficult to remove the second electron. Therefore, the second ionisation energy is greater than the first ionisation energy and similarly the third ionisation energy is greater than the second ionisation energy and so on. Thus,

I E3  I E2  I E1

Factors determining Ionisation Energies

The magnitude of ionisation energy depends upon the following factors.

1. Atomic size : The larger the atomic size, the smaller is the ionisation energy. The reason for this is that as the size of the atom increases , the outer electrons lie farther away from the nucleus. Hence according to Coulom’s law, the attractive pull of the nucleus on the outer electrons decreases and it becomes easier to knock out an electron from the outer shell of the atom.

2. Nuclear Charge : The force of attraction between the nucleus and the outermost electron increases with increase in nuclear charge. Thus , greater nuclear charge, greater is the energy needed to pull out an electron from the atom. Hence ionisation energy increases with increase in nuclear charge.

3. Number of electrons in the inner shell : The larger the number of electrons in the inner shells, the smaller the ionisation energy. The electrons in the inner shell act as a screen or shield between the nucleus and electrons in the outermost shell. This is known as screening effect or shielding effect. The larger the number of electrons in the inner shells, the greater is the screening effect. Consequently, the electrons in the valence shell experience less attraction from the nucleus. Hence ionisation energy should be low.

4. Penetration Effect :The s-electrons in their motion around the nucleus remain closer to the nucleus than do p, d, or f-electrons of the same principal energy shell. In other words , s-electrons penetrate more towards the nucleus than do the p, d, or f-electrons of the same principal energy shell. In other words, s-electrons penetrate more towards the nucleus than p-electrons and the penetrating power of the electrons in the given principal shell varies as :

s  p  d  f .

Thus, the s-electron experience more attraction from the nucleus than the p, d or f-electrons of the same principal energy shell . It follows , therefore , that ionisation energy for pulling out an s-electron is maximum and it decreases in pulling out a p-electron, a d-electron or an f-electron of the same principal energy shell in the order of their mention. Thus, the ionisation energy in pulling out an electron from a given energy level decreases in the order

s  p  d  f orbitals.

5. Electronic configuration

Certain electronic configurations are more stable than others. For example , if an atom has fully filled or half-filled orbitals, its ionisation energy is higher than expected normally from its position in the periodic Table. For example, Beryllium and Nitrogen in the second period and magnesium and phosphorus in the third Period, have slightly higher ionisation energies than would be expected. This is due to the extra stability of the fully completed s-orbitals in Beryllium and Magnesium and that of the half-filled p-orbitals in nitrogen and Phosphorus. Further , it is seen that He, Ne, Ar, Kr and Xe have highest ionisation energies in their respective periods. This shows that an atom with s2P6 arrangement ( the so called octet arrangement) in the outer shell is highly stable. In this connection it also be noted that Li+ and Na+ ions which also have s2p6configuration require very high energies viz., 7297.2 and 4563.5 kJmol1 to permit pulling out of an extra electron from them to give Li2+ and Na2+ respectively.

Variation of ionisation enthalpy in the periodic table

The variation of ionisation enthalpies for elements with atomic numbers 1-60 is shown graphically in Figure.

Variation first ionisation enthalpies for elements

with Z = 1 – 60.

The graph consists of several maxima and minima. The maxima are found at noble gases which possesses stable electronic configurations. Due to very high ionisation enthalpies, these gases are almost inert and show extremely low chemical reactivity. The minima occur at alkali metals. Due to very low ionisation enthalpies, these elements get easily ionised and are very highly reactive.

Ionisation enthalpies of the elements show a periodic trend. In general they increase across a period and decrease in going down a group.

Variation in a group

In general, the ionisation energy decreases in going from top to bottom in a Group. On moving down the group :

i) The nuclear charge increases.

ii) There is a gradual increase in atomic size due to increase in the number of principal energy shell.

iii) There is increase in number of inner electrons which shield the valence electrons from the nucleus.

The overall effect of increase in atomic size and shielding is much more than the overall effect of increase in nuclear charge. Consequently, the outer most electron is less and less tightly held to the nucleus as we move down the group. Hence ionisation energy decreases as we go down from top to bottom in any group. The decreasing trend in ionisation enthalpies in Group 1 is shown in the following Figure.

First ionization enthalpies of alkali metals as a function of atomic number

There are some exceptions, however. For example, the first ionisation energy of Indium is less that of Thalium in Group 13 and first ionisation energy of tin is less than that of lead in Group 14.

b. Variation along a Period

In general, the ionisation enthalpy increases with increase in atomic number across a period.

The gradual increase in the values of ionisation enthalpy on moving from left to right in a period is due to the combined effect of the following factors.

(i) The nuclear charge increases regularly with increase in atomic number.

(ii) Number of shells remain the same and the differentiating electrons enter into the same shell.

(iii) Size of atom decreases.

Due to the combined effect of the above factors, the attractive forces between the nucleus and valence electron increase as one moves from left to right in a period. Thus, more and more energy is required to remove electron from the atom and consequently, the ionisation enthalpy increases with increase in atomic number in a period. The variation in the values of ionisation enthalpies of elements across the second period is shown in Fig.

Variation of Ionisation energies with

Atomic Numbers.

This figure shows that , in general, the values of ionisation enthalpies increases on going from Li to Ne. However, there are certain exceptions also. For example, boron and oxygen have lower ionisation enthalpies as compared to Be and N respectively. These discrepencies can be explained as follows.

(i) The electronic configuration of Be and B are :

Be ( Z = 4 ) : 1s22s2

B ( Z = 5 ) : 1s22s22p1

Although boron possesses a higher nuclear charge, yet its ionisation energy is less as compared to that of beryllium. This may be attributed to the following reasons.

(a) 2s electrons are more penetrating than 2p electrons.

(b) 2s-electrons in beryllium are not shielded by the inner electrons as efficiently as 2p electron in boron.

(c) Beryllium possesses a stable electronic configuration as all the orbitals are completely filled in it.

Due to these factors , beryllium possesses much higher ionisation enthalpy as compared to the value it should have otherwise. The value of ionisation enthalpy of boron is in accordance to the regular trend but due to a sudden increase in value for beryllium, the ionisation enthalpy of boron appears to be low.

(ii) Similarly, a decrease in ionisation enthalpy in going from N to O is due to the presence of half-filled shell in nitrogen which makes its configuration relatively more stable and gives a boost to the ionisation enthalpy of nitrogen.

Due to higher ionisation enthalpy of nitrogen, the ionisation enthalpy of oxygen appears to get lowered.

(iii) The exceptionally very high value of ionisation enthalpy of neon is due to its stable electronic configuration.

Problems

11. Arrange with explanation , the following elements in the order of ionisation enthalpy. X( Z = 4), Y (Z = 5) , D (Z = 6) and E ( Z = 11).

12. The first (IE1) and second(IE2) ionisation enthalpies ( kJ/mol) of few elements are given

Element IE1 IE2

I 2372 5251

II 520 7300

III 900 1760

IV 1680 3380

Which of the the above elements is likely to be :

(i) a reactive metal (ii) a reactive non-metal

(iii) a noble gas

(iv) a metal A that form stable binary halide with fomula

AX2 ( X = halogen)

13. Consider the ground state electronic configurations given

below :

X = 1s22s22P5 ; Y = 1s22s22P4

R = 1s22s22P6 : Z = 1s22s1

Q = 1s22s22P63s1

Pick out the correct answers ?

i) Which of the above configuration is associated with highest and lowest ionisation energy ?

ii) Arrange these configurations in the order of the increasing ionisation energies.

14. Among the elements Li, K, Ca S and Kr which one has the lowest first ionisation energy ? Which has the highest first ionisation energy ?

15. Among the elements of the second period from Li to Ne pick out the element :

i) with highest first ionisation energy.

ii) with highest electronegativity.

iii) with largest atomic radius.

iv) that is most reactive non-metal.

v) that is most reactive metal

16. Which of the following pairs of elements would you expect to have lower first ionisation enthalpy? Explain.

a. Cl or F b) Cl or S c) K or Ar d. Kr or Xe

17. Which of the following species will have the largest and smallest size ?

Mg, Mg2+, Al, Al3+

18. The first ionisation enthalpy (iH) values of the third period elements , Na, Mg and Si are respectively 496, 737 and 786 kJ mol1. Predict whether the first egH value of Al will be close to 575 or 760 kJ mol1 ? Justify your answer.

5. Electron Gain Enthalpy(egH)

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 Electron Gain Enthalpy (egH). Electron gain enthalpy provides a measure of ease with which an atom adds an electron to form anion as represented by equation :

X(g) + e X ; H = egH ………. (1)

Depending upon 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 level leading to a very unstable electronic configuration. It may be noted that the electron gain enthalpies have large negative values towards the upward right of the periodic table preceeding the noble gases.

The variation in electron gain enthalpies of elements is less systematic than ionisation enthalpies.

Electron Gain Enthalpies (kJ/mol) of Some Main Group Elements

H

 73 He

+48

Li

 60 Be

+66 B

83 C

122 N

+31 O

141 F

200 Ne

+116

Na

 53 Mg Al

50 Si

119 P

74 S

200 Cl

349 Ar

+96

K

48 Ca Ga

36 Ge

116 As

77 Se

195 Br

325 Kr

+96

Rb

47 Sr In

29 Sn

120 Sb

101 Te

190 I

295 Xe

+77

Cs

46 Ba Tℓ

30 Pb

101 Bi

110 Po

174 At

270 Rn

+68