Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010)

Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010) Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010) Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010) Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010) Preview Fundamentals of Inorganic Chemistry For Competitive Exams by Ananya Ganguly (2010)

Chandigarh • Delhi • Chennai

Ananya Ganguly

Fundamentals of

InorganIc chemIsty

attempt to render any type of professional advice or analysis, nor is it to be treated as such While much care has been taken to ensure the veracity and currency of the information presented within, neither the publisher, nor its authors bear any responsibility for any damage arising from inadvertent omissions, negligence or inaccuracies (typographical

or factual) that may have found their way into this book

Copyright © 2011 Dorling Kindersley (India) Pvt Ltd

Licensees of Pearson Education in South Asia

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Preface v

Chapter 12 Transition Elements (d-block) and Their Compounds 12.1—12.57

Contents

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In recent years, the question format of non-medical examinations like IIT-JEE and AIEEE has been restructured with greater emphasis on the theoretical and conceptual intricacies and the application of the underlying basic concepts and principles.

The poor performance of students in these examinations is partly due to non-availability of

a comprehensive text book which lays adequate stress on the authentic logical theoretical concept building, numericals and related problems.

The present book grew out of my experience of classroom teaching through lectures, notes and assignments Effort has been made to maintain lucid style and simplicity of expression Apart from this, the book has numerous carefully selected examples and solved illustrations which include almost all the previous years’ questions asked in IIT-JEE and AIEEE examinations.

Selected questions of different formats keeping in mind the recent pattern of examinations have been listed at the end of each chapter Solutions to almost all advanced questions have been incorporated.

I would like to thank Showick Thorpe and Sanjay Sharma for their constant encouragement during the publication of this book I extend my thanks to Nitkiran Bedi for her contribution in bringing out this book in record time I am also thankful to the staff members of my academy for their sincere help.

I will appreciate comments, suggestions and criticism from the readers and will incorporate them in the subsequent editions.

Ananya Ganguly

Preface

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Periodic Properties and Chemical bonding

Appearance and properties Solids, some with high melting

points; lustrous, malleable and ductile Gases, or solids with low melting points Conduction of heat and electricity Very good Poor

Compounds Ionic compounds with non-metals;

alloys with other metals Ionic compounds with metals; covalent compounds with other non-metals

Chemical nature Reducing agents Oxidizing agents

However, it was widely believed that there had to be an underlying reason for the patterns

1 Prout’s Hypothesis

One of the first suggestions was due to Prout Prout’s hypothesis was that all elements were made from a whole ber of hydrogen atoms (Be careful here: an atom’ in Prout’s time was a very different thing to our understanding of the word.) According to him the atomic masses of the elements should be a whole number of times that of hydrogen, i.e., they should be integers Unfortunately, from Prout’s point of view, the results of experiments showed that the atomic masses of many elements were not integers

Triad of atoms Mean of first and last element

Therefore, this hypothesis was not acceptable for all elements

3 The Telluric Helix

It was in 1862, that a periodic classification of the elements was developed that approached the idea we have today

At that time A.E de Chancourtois, a professor of Geology at the Ecole des Mines in Paris presented an account of his telluric helix in which he indicated the relative properties of elements and their atomic weights

He used a vertical cylinder with 16 equidistant lines on its surface, the lines lying parallel to the axes Then he drew a helix at 45 degree to the axis and arranged the elements on the spiral in the order of their increasing atomic weights In this manner, elements that differed from each other in atomic weight by 16 or multiples of 16 fell very nearly on the same vertical line In addition to the 16 vertical lines, de Chancourtois felt that other connecting lines could be drawn, and that all elements lying on such lines were related in some manner His arrangement resulted in the proposal by de Chancourtois that the properties of the elements are the properties of numbers

4 Newland’s Rule of Octave

A major attempt at making a link was made by new lands in 1864 (Table 1.2) He grouped elements into sets of eight and claimed that every eighth element in the pattern was chemically similar Newlands’ law of octaves was largely ignored, or at best treated with mild amusement

The first thorough attempt at relating chemical properties to atomic masses was made by the Russian

Table 1.2 Examples of Newlands’ Octaves*

The author claims the discovery of a law according to which the elements analogous in their properties exhibit peculiar relationships, similar to those subsisting in music between a note and its octave Professor G F Foster humorously enquired of Mr Newlands whether he had ever examined the elements according to the order of their initial letters?

Newlands was not at all happy about the credit that went to Mendeleeff over the discovery of the periodic law In 1884 Newlands wrote: Having been the first to publish the existence of the periodic law more than nineteen years ago, I feel, under existing circumstances, com- pelled to assert my priority in this matter As a matter of simple justice, and in the interest of all true workers in science, both theoretical and practical, it is right that the originator of any proposal or discovery should have the credit of his labour.

5 Lothar Meyer’s Volume Curves

The graphs of atomic volumes against atomic weights are known as Lothar Meyer volume curves

(i) Lothar Meyer plotted a graph between atomic weight and atomic volume (i.e., atomic weight in solid state/

density)

(ii) Elements with similar properties occupied the similar positions on the graph

(iii) Strong electropositive elements of IA except Li, all others Na, K, Rb, Cs etc., occupied the top position on the

graph

8070605040302010

Fig 1.1 Atomic volume versus atomic number curve

(iv) IIA group elements Be, Mg, Ca, Sr, Ba etc., occupied the positions on the ascending part of the graph

(v) Inert gases except He occupied the positions on the descending part of the graph

(vi) Halogens also occupied the descending part of the graph

n Transition elements have very small volumes and therefore these are present at the bottoms of the curve metals have highest atomic volumes

6.(a) Mendeleef’s Periodic Table

n Periodic table is based on atomic weight

n In the periodic table, the horizontal lines are called periods and the vertical lines are called groups

n The periodic table consists of seven periods and nine groups (The earlier periodic table had only 8 groups The noble gases were added later in the zero group because these were not discovered when Mendeleef put forward his periodic table

n All the groups (except VIII and Zero groups) are divided into subgroups A and B

n 2, 8, 18 and 32 are called magic numbers

(b) Merits of Mendeleef’s Periodic Table

n Classification of elements then known, was done for the first time and the elements having similar properties were kept in the same group

n It encouraged research and led to discovery of newer elements

n Mendeleef had even predicted the properties of many elements not discovered at that time This helped in the covery of these elements For example, Mendeleef predicted the properties of the following elements:

(a) ka-boron—This was later called scandium (Sc)

(b) Eka-aluminium—This was later called gallium (Ga)

(c) Eka-silicon—This was later called germanium (Ge)

n Atomic weight of elements were corrected Atomic weight of Be was calculated to be 3 × 4.5 = 13.5 by considering its valency 3 Mendeleef calculated it 2 × 4.5 = 9

(c) Defects of Mendeleef’s Periodic Table

(1) Position of Hydrogen: Hydrogen resembles alkali metals and halogens in its properties Hence, its position was

not sure

(2) Position of Isotopes: The isotopes have different atomic weights and the periodic table is based on atomic

weights Therefore, isotopes should get different places in the periodic table on the basis of atomic weights

(3) The periodic table is not fully based on increasing order of atomic weights

(4) It is not proper to place together the elements having differing properties, such as coinage metals (Cu, Ag and Au)

with alkali metals; Zn, Cd and Hg with alkaline earth metals and metal like Mn with halogens Similarly, Pt and

Au having similar properties have been placed in different groups

(5) There is no indication whether lanthanides and actinides are associated with group IIIA or group IIIB.

(6) Position of Isobars: These elements have different groups when mass remains same.

(7) Lot of stress was given to valency of elements

MODeRN PeRIODIC Law aND MODeRN PeRIODIC TabLe

velocity electrons is proportional to the nuclear charge of the atom This can be represented by the following expression

= a (Z–b) where Z is nuclear charge on the atom and a and b are constants

n The nuclear charge on an atom is equal to the atomic number

n According to modern periodic law, “The properties of elements are the periodic functions of their atomic numbers”

Modern Periodic Table

n On the basis of the modern periodic law, a scientist named Bohr proposed a long form of periodic table that was

pre-pared by Rang and Warner

n In the periodic table, the horizontal lines are periods and the vertical lines are groups

n The periodic table has a total of seven periods and 18 groups But according to CAS system, the number of groups is 16,

because the eighth group has been divided into three groups

n There are two elements in the first period eight elements in each of the second and third periods, eighteen elements in each of the fourth and fifth period thirty two elements in the sixth period and only nineteen elements till now in the seventh period Total 105 elements have been discovered so far

n The first period is very short period, second and third are short periods fourth and fifth are long periods sixth is very long period, while the seventh is incomplete period

n The lanthanides (Elements from atomic numbers 58 to 71) and actinides (elements from atomic numbers 90 to 103) are included in the sixth and seventh periods through these have been kept outside the periodic table

n Period—The details about the seven periods are as follows

(i) Each period starts with an alkali metal whose outermost electronic configuration is ns1

(ii) Each period ends with a noble gas of outermost electronic configuration ns2np6 except He The electronic ration of He is 1s2

configu-(iii) The number of elements in a period is equal to the number of necessary electrons to acquire ns2np6 configuration

in the outermost shell of first element (alkali metal) of the period First period contains two elements

(v) The number of elements in each period may be determined by the number of electrons in a stable configuration as under

Periods Stable electronic configuration Number of electrons From To

Sixth 6s 2 4f 14 5d 10 6p 6 32(including lanthanoids) Cs (55) Rn (86)

Seventh (including Actinoids) Fr (87) Ha (105)

electronic basis for the Periodic Classification

With a better understanding of the part that the electron plays great role in the properties of the elements, a corresponding understanding of the periodic system came about

On the basis of electronic configuration, the elements may be divided into four groups:

s-block elements

n The elements of the periodic table in which the last electron enters in s-orbital, are called s-block elements.

n s-orbital can accommodate a maximum of two electrons

n Their general formulae are ns1 and ns2 respectively, where, n = (1 to 7)

n I A group elements are known as alkali metals because they react with water to form alkali II A group elements are known as alkaline earth metals because their oxides react with water to form alkali and these are found in the soil or earth

n The total number of s block elements are 14

n Fr87 and Ra88 are radioactive elements while H and He are gaseous elements

n Cs and Fr are liquid elements belonging to s-block.

p-block elements

n The elements of the periodic table in which the last electron gets filled up in the p-orbital, called p-block elements.

n A p-orbital can accommodate a maximum of six electrons Therefore, p-block elements are divided into six groups

which are III A, IV A, V A, VI A, VII A and zero groups

n The general formulae of p block elements are ns2 p1–6 (where, n = 2 to 6)

n The zero group elements having general formula ns2p6 are inert, because their energy levels are fully filled

n The total number of p block elements in the periodic table is 30 (excluding He)

n There are nine gaseous elements (Ne, Ar, Kr, Xe, Rn, F2, Cl2, O2 and N2) belonging to p-block Gallium (Ga) and

bro-mine (Br) are liquids

n The step-like thick lines drawn in the periodic table in the p-block divides elements into metals nonmetals and

metal-loids

d-block elements

n The elements of the periodic table in which the last electron gets filled up in the d orbital, called d block elements.

n The d block elements are placed in groups named IIIB, IV B, V B, VI B, VII B, VIII, I B and II B.

n In d block elements, the electron gets filled up in the d orbital of the penultimate shell That is why, these elements are

known as transition elements

n Though the total number of d block elements is 33 in the periodic table but there are only 30 transition elements Because only those elements are transition in which d orbital is partially filled.

n The general formula of these elements is (n–1)s2, p6, d1–10 ns1–2 where n = 4 to 7

n All of these elements are metals

n Out of all the d block elements mercury is the only liquid element.

f-block elements

n The elements of the periodic table in which the last electron gets filled up in the f orbital, called f block elements.

n The f block elements are from atomic number 58 to 71 and from 90 to 103

n The lanthanides occur in nature in low abundance and therefore, these are called rare earth elements

n There are 28 f block elements in the periodic table

n The elements from atomic number 58 to 71 are called lanthanides because they come after lanthanum (57) The ments from 90 to 103 are called actinides because they come after actinium (89)

ele-n All the actinide elements are radioactive

n All the elements after atomic number 92 (i.e., U92) are transuranic elements

n The general formula of these elements is (n–2)s2 p6 d10 f(1–14) (n–1)s2 p6 d0–1 ns2 where n = 6 and 7

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Fig 1.2 Long form of the periodic table of the elements with their atomic numbers and ground state outer

electronic configurations The groups are numbered 1-18 in accordance with the 1984 IUPAC recommendations

This notation replaces the old numbering scheme of IA-VIIA, VIII, IB-VIIB and 0 for the elements

Nomenclature of elements with atomic Numbers > 100

The naming of the new elements had been traditionally privilege of the discoverer (or discovers) and the suggested name was ratified by the IUPAC In recent years, this has led to some controversy The new elements with very high atomic

numbers are so unstable that only minute quantities, sometimes only a few atoms of them are obtained Their synthesis and characterisation, therefore, require highly sophisticated costly equipment and laboratory Such work is carried out with competitive spirit only in some laboratories in the world Scientists, before collecting the reliable data on the new element,

at times get tempted to claim for its discovery For example, both American and Soviet scientists claimed credit for covering element 104 The Americans named it Rutherford turn whereas Soviets named it Kurchatovium To avoid such problems, the IUPAC has made recommendation that until a new element’s discovery is proved, and its name is officially recognized, a systematic nomenclature be derived directly from the atomic number of the element using the numerical roots for O and numbers 1-9 These are shown in table The roots are put together in order of digits which make up the atomic number and “ium” is added at the end The IUPAC names of the elements with Z above 100 are shown in the table

Name nil un bi tri quad pent hex sept oct enn

Abbreviation n u b t q p h s o e

Table 1.3 Nomenclature of Elements with Atomic Number Above 100

Atomic number Name Symbol IUPAC official name IUPAC symbol

104 Unnilquadium Unq Rutherfordium Rf

*Official IUPAC name yet to be announced + Elements yet to be discovered

Thus, the new element first gets a temporary name, with symbol consisting of three letters Later permanent name and symbol are given by a vote of IUPAC preventatives from each country The permanent name might reflect the country (or state of country) in which the element was discovered or pay tribute to a notable scientist As of now, elements with atomic numbers up to 112, 114 and have been discovered Elements with atomic numbers 113, 115, 117, and 118 are not yet known

Worked-out Example

1 What would be the IUPAC name and symbol for the element with atomic number 120?

Solution: From Table, the roots for 1, 2 and O are un, bi and nil, respectively Hence, the symbol and the name

respectively are Ubn and unbinilium

The types of elements

Using electronic configuration as the criterion, we ordinarily recognize four general type of elements; the inert gas ments, the representative elements, the transition elements, and the inner transition elements The classification of the ele-ments into these groups is dependent on the extent to which the s, p, d and f orbitals are filled

(i) Inert Gases

(a) s and p-orbitals of the outer most shell of these elements are completely filled The outermost electronic

configu-ration is ns2np6

(b) Helium is also inert gas but its electronic configuration is 1s2

(ii) Representative or Normal Elements

(a) Outermost shell of these elements is incomplete The number of electrons in the outermost shell is less than

eight

(b) Inner shells are complete

(c) s- and p-block elements except inert gases are called normal or representative elements

(iii) Transition Elements

(a) Last two shells of these elements namely outermost and penultimate shells are incomplete

(b) The last shell contains one or two electrons and the penultimate shell may contain more than eight up to eighteen

electrons

(c) The outermost electronic configuration is similar to d-block elements, i.e., (n–1)d1–10 ns1-2

(d) According to latest definition of transition elements those elements which have partly filled d-orbitals in neutral

state or in any stable oxidation state are called transition elements According to this definition, Zn, Cd and Hg

in neutral as well as in stable +2 oxidation state.

(e) Because of the extra stability which is associated with empty, half-filled, and filled subshells, there are some

apparent anomalies in electronic arrangements in the transition series This empirical rule is illustrated by the chromium and copper configuration in the first d series of elements:

Sc Ti V Cr Mn Fe Co Ni Cu Zn

(iv) Inner Transition Elements

(a) In these elements last three shells i.e., last, penultimate and prepenultimate shells are incomplete

(b) These are related to IIIB i.e., group 3

(c) The last shell contains two electrons Penultimate shell may contain eight or nine electrons and pre-penultimate

shell contains more than 18 up to32 electrons

(d) Their outermost electronic configuration is similar to f-block element i.e., (n–2)f1–14(n–1)s2 (n–1)p6 (n–1)d0–1ns2

(v) Typical Elements: The elements of second and third periods are called typical elements as they summarize the

properties of their respective groups

Group I II III IV V VI VII

3 rd period Na Mg Al Si P S Cl

(vi) Elements of the seventh period after atomic number 93 (i.e., actinides) are synthetic elements and are called

transu-ranium elements

(i) Bridge elements: The typical elements of third period are also called bridge elements as the division between

two subgroups A and B starts from these elements In second group, Mg acts as a bridge element The ties of bridge element are somewhat mixed properties of the elements of two subgroups as magnesium shows similarities with alkaline earth metals (IIA) on one hand and with zinc metals (IIB) on the other

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(vii) Diagonal relationship: The first three members of second period (Li, Be and B) not only show similarities with the

members of their own groups but show similarities with the elements diagonally placed in the higher groups This resemblance is termed diagonal relationship

Lithium shows similarities with magnesium, beryllium with aluminium and boron with silicon

(viii) The elements belonging to the same subgroup ex hibit either same properties or a regular gradation.

(ix) Coinage metals (Cu, Ag, Au) and Alkali metals: There appears hardly any resemblance between the members of

these two groups except that all show monovalency

Table 1.4 Some Properties are Given Below to Justify this Statement

IA (alkali metals) IB (coinage metals)

(i) Soft metals (i) Hard metals (ii) Low densities (ii) High densities (iii) Low melting points (iii) High melting points (iv) Highly reactive (iv) Less reactive, noble metals (v) Never found free in nature (v) Found mainly free in nature (vi) Decompose water readily (vi) Do not decompose water

The Screening effect or Shielding effect

In a multielectron atom, the electrons of the valency shell (outermost shell) are attracted towards the nucleus and also these electrons are repelled by the electrons present in the inner shells On account of this, the actual force of attraction between the

nucleus and the valency electrons is somewhat decreased by the repulsive forces acting in opposite direction This decrease in

the force of attraction exerted by the nucleus on the valency electrons due to the presence of electrons in the inner shells,

is called screening effect or shielding effect The magnitude of the screening effect depends upon the number of inner

elec-trons, i.e., higher the number of inner elecelec-trons, greater shall be the value of screening effect The screening effect constant is

represented by the symbol σ The magnitude of ‘σ’ is determined by the Slater’s rules The contribution of inner electrons to

the magnitude of ‘σ’ is calculated in the following ways:

For ns or np orbital electrons

(i) Write the electronic configuration of the element in the following order and group them as,

(Is), (2s,2p), (3s,3p), (3d), (4s, 4p), (4d,4f), (5s, 5p), (5d, 5f), (6s, 6p), etc

(ii) Electrons to the right of the (ns, np) group are not effective in shielding the ns or np electrons and contribute nothing to σ (iii) All other electrons in the (ns, np) group contribute to the extent of 0.35 each to the screening constant (except for Is for which the value is 0.30)

(iv) All the electrons in the (n – l) th shell contribute 0.85 each to the screening constant

(v) All the electrons in the (n – 2) th shell or lower contribute 1.0 each to the screening constant

For d- or f-electron, rules (i) to (iii) remain the same but rules (iv) and (v) get replaced by the rule (vi).

(vi) All the electrons in the groups lying left to (nd, nf) group contribute 1.0 each to the screening effect

Example 1 Calculation of screening constants of alkali metals for valency electrons.

Example 2 Calculation of screening constants of mem bers of second period for valency electrons.

Example 3 Calculation of screening constant in zinc (a) for a 4s-electron (b) for 3d electron.

(a) The electronic configuration of zinc (30) is,

(1s)2 (2s2p)8 (3s3p)8 (3d)1° (4s)2

σ = 10 × 1.0 + 18 × 0.85 + 1 × 0.35 = 25.65

(b) For 3d-electron,

σ = 18 × l.0 + 9 × 0.35 = 21.15

It is clear from the examples (1) and (2) that the mag nitude of screening constant in the case of s- and p-block

ele-ments increases in a period as well as in a group as the atomic number increases.

From the example (3), it is evident that the electrons in different orbitals are differently affected by the same nuclear

charge depending upon their proximity to the nucleus.

effective atomic Number

Due to screening effect the valency electron experiences less attraction towards nucleus This brings decrease in the nuclear charge (Z) actually present on the nucleus The reduced nuclear charge is termed effective nuclear charge and is repre-sented by Z* It is related to actual nuclear charge (Z) by the following formula:

Z* = (Z – σ) where a is screening constant

It is observed that magnitude of effective nuclear charge or effective atomic number increases in a period when we

move from left to right.

(iv) The inter nuclear distance between the two atoms can be measured by × – ray diffraction or spectroscopic studies.

(v) Covalent radius: One half of the distance between the nuclei (internuclear distance) of two covalently bonded atoms

in a homodiatomic molecule is called the covalent radius of that atom The covalent bond must be single covalent bond The covalent radius (rA) of atom A in a molecule A2 may be given as:

A–A A

dr2

DA–B = rA + rB – 0.09 ∆x Where ∆x is the difference of electronegativities of the atoms A and B

According to Pauling – If the electronegativities of the two atoms A and B are xA and xB respectively then,

Metallic radius > Covalent radius

For example – Metallic radius and covalent radius of potassium are 2.3 Å and 2.03Å respectively

Van der waal’s radius or collision radius

The molecules of non metal atoms are generally gases On cooling, the gaseous state changes to solid state

In the solid state, the non metallic elements usually exist as aggregations of molecules are held together by Van der

Waal forces One half of the distance between the nuclei of two adjacent atoms belonging to two neighbouring

mol-ecules of a compound in the solid state is called Van der Waal’s radius

It may also be defined as half of the inter nuclear distance of two non bonded neighbouring atoms of two adjacent molecules

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Van der waal’s radius > metallic radius> covalent radius

The Van der Waal’s radius and Covalent radius of Chlorine atom are 1.80Å and 0.99Å respectively

Ionic Radius

A neutral atom changes to a cation by the loss of one or more electrons and to an anion by the gain of one or more electrons The number of charge on cation and anion is equal to the number of electrons lost or gained respectively The ionic radii of the ions present in an ionic crystal may be calculated from the internuclear distance between the two ions

It is defined as the distance between the nucleus and outermost shell of an ion or it is the distance between the nucleus and the point where the nucleus exerts its influence on the electron cloud

(i) Radius of a cation: A cation is formed by removing one or more electrons from a neutral atom Hence, the ionic

radii of a cation is always less than the atomic radii of that very atom This is due to the fact that nuclear charge

in the case of a cation is acting on lesser number of electrons and pulls them closer

(a) The effective nuclear charge increases For example, in Na atom, 11 electrons are attracted by 11 protons and in

Na+,10 electrons are attracted by 11 protons Thus in the formation of cation, number of electrons decreases and nuclear charge remains the same

(b) Generally, the formation of cation results in the removal of the whole outer shell

(c) Interelectronic repulsion decreases The interelectronic repulsion in Na is among 11e– and in Na+ among 10e–

(ii) Radius of an anion: In the case of negative ion or an anion, as electron or electrons are added to the neutral atom,

the nuclear charge acts on more electrons so that each electron is held less tightly and the electron cloud expands

Thus, the size of an anion is more than its parent atom.

(a) The effective nuclear charge decrease in the formation of anion Thus the electrostatic force of attraction between

the nucleus and the outer electrons decreases and the size of the anion increases

(b) Interelectronic repulsion increases

(iii) Isoelectronic series: A series of atoms, ions and molecules in which each species contains same number of electrons

but different nuclear charge is called isoelectronic series

(a) Number of electrons is same

(b) Number of protons is increasing

(c) So the effective nuclear charge is increasing and atomic size is decreasing In an isoelectronic series atomic size

decreases with the increase of charge

Some of the examples of isoelectronic series are as under:

H+ and Cs+ are the smallest and largest cations respectively

H– and I– are the smallest and largest anions respectively

Periodicity in atomic radius and ionic radius

1 For normal elements

(a) In a period from left to right effective nuclear charge increases because the next electron fills in the same shell

So the atomic size decreases For example, the covalent radii of second period elements in Å are as follows:

(b) In a group moving from top to bottom the number of shells increases So the atomic size increases Although the

effective nuclear charge increases but its effect is negligible in comparison to the effect of increasing number of shells For example the covalent radii of IA group elements in Å are as follows:

2 For inert gases

The atomic radius of inert gas (zero group) is shown largest in a period because of its Van der Waal’s radius which is

generally larger than the covalent radius The Van der Waal’s radius of inert gases also increases in moving from top

to bottom in a group

3 For transition elements

There are three series of transition elements:

3d – Sc (21) to Zn (30)

4d – Y (39) to Cd (48)

5d – La (57), Hf (72) to Hg (80)

(a) From left to right in a period

(i) The atomic size decreases due to the increase in effective nuclear charge

(ii) In transition elements, electrons are filled in the (n–1)d orbitals These (n–1)d electrons screen the ns

elec-trons from the nucleus So the force of attraction between the ns elecelec-trons and the nucleus decreases.

This effect of (n–1)d electrons over ns electrons is called shielding effect or screening effect The atomic

size increases due to shielding effect and balance the decrease in size due to increase in nuclear charge to about 80%

(iii) Thus moving from left to right in a period, there is a very small decrease in size and it may be considered

that size almost remains the same

(iv) In the first transition series, the atomic size slightly decreases from Sc to Mn because effect of effective

nuclear charge is stronger than the shielding effect The atomic size from the Fe to Ni almost remains the same because both the effects balance each other The atomic size from Cu to Zn slightly increases because shielding effect is more than effective nuclear charge due to d10 structure of Cu and Zn The atomic radii of the elements of 3d transition series are as under

4 Inner transition elements

As we move along the lanthanide series, there is a decrease in atomic as well as ionic radius The decrease in size is

regular in ions but not so regular in atoms This is called lanthanide contraction The atomic radii in Å are as under:

1.88 1.82 1.83 1.82 1.81 1.80 2.04 1.80

1.78 1.77 1.76 1.75 1.94 1.73

There are two peaks one at Eu (63) and other at Yb (70) This is due to the difference in metallic bonding Except

Eu and Yb other lanthanides contribute three elctrons in metallic bond formation These two atoms contribute two

electrons in the bond formation leaving behind half-filled and completely filled 4f-orbitals respectively.

Cause of lanthanide contraction

In lanthanides the additional electron enters into (n-2)f orbital The mutual shielding effect of (n-2)f electrons is very little because the shape of f-subshell is very much diffused Thus the effective nuclear charge increases in comparison to the mutual shielding effect of (n-2) f electrons The outer electrons are attracted more by the nucleus Consequently the atomic

and ionic radii decreases from La (57) to Lu (71)

This type of contraction also occurs in actinides The jump in contraction between the consecutive elements in the

actinides is greater than lanthanides This is due to the lesser shielding of 5f-electrons which are therefore pulled more

strongly by the nucleus

In a group

(i) The atomic radius of elements increases moving from first transition series (3d) to second transition series (4d) This

is due to the increase in number of shells with the increase in atomic number

(ii) The atomic radii of second (4d) and third (5d) transition series in a group is almost same except Y(39) and

La (57)

In third transition series, there are fourteen lanthanides in between La (57) of III B and Hf (72) of IV B groups,

so the atomic radius of Hf(72) decreases much due to lanthanide contraction in lanthanides The difference in the nuclear charge in the elements of a group in first and second transition series is + 18 units while this difference in second and third transition series is + 32 units except Y (39) → La(57) Due to the increase of + 32 units in the nuclear charge there is a sizable decrease in the atomic radius which balances the increase in size due to the increase

in number of shells

So in a group moving from second to third transition series, the atomic radii of the elements almost remain the same except IIIB The difference is about 0.02Å

Worked-out Examples

(a) N < Be < B (b) F– < O2– < N3– (c) Na < Li < K (d) Fe3+ < Fe2+ < Fe4+

Ans As we go from F to O to N, the positive charge in the nucleus decreases This causes decrease in

nucleus-electron attraction causing the nucleus-electron cloud to spread more in space

Also, larger negative charge involves larger electron-electron repulsion causing spreading of electron cloud in space

3 According to the Periodic law of elements, the variation in properties of elements is related to their

(AIEEE 2003)

(c) Nuclear neutron-proton number ratio (d) Atomic masses

Ans (b) According to 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

4 The reduction in atomic size with increase in atomic number is a characteristic of elements of (AIEEE 2003)

(c) radioactive series (d) high atomic masses

Ans (a)

5 The set representing the correct order of ionic radius is:

(a) Li+ > Be2+ > Na+ > Mg2+ (b) Na+ > Li+ > Mg2+ > Be2+

(c) Li+ > Na+ > Mg2+ > Be2+ (d) Na+ > Mg2+ > Li+ > Be2+

Ans (d) Follow the periodic trends

6 The radii of Ar is greater than the radii of chlorine Why?

Solution: In chlorine, the radii means the atomic or covalent radii which is actually half the intermolecular distance between 2 atoms whereas in Argon the radii means the Van der Waal’s radii as Argon is not a diatomic molecule Van der Waal’s radii is actually half the distance between adjacent molecule So Van der Waal’s radii being larger than atomic radii, Argon, has got a larger radii than chlorine

7 Berilium and Al are placed in different periods and groups but they show the similar properties Explain

Solution: On moving across a period, the charge on the ions increases and the size decreases, causing the rising power to increase On moving down a group, the size increases and polarising power decreases On mov-ing diagonally i.e., from Be to Al these two effects partly cancel each other and so there is no marked change in properties

pola-Ionization Potential

n The energy required to remove the most loosely bound electron from the outermost orbit of one mole of isolated

gas-eous atoms or ion of an element, is called ionization energy (IE) This ionization is an endoergic or energy-absorbing

process The corresponding potential difference to be applied to provide ionization energy to an electron is known as

Ionization Potential (IP) Hence, the ionization Energy of hydrogen is 13.6 eV and the ionization Potential is 13.6 V.

n An electron cannot be removed directly from an atom in solid state For this purpose, the solid state is converted to

gaseous state and the energy required for this is called sublimation energy

n A(g) →( )IP1 A+1(g) →( )IP2 A+2(g) →( )IP3 A (g) +3

n The energy required to remove one electron from a neutral gaseous atom to convert it to monopositive cation, is called first ionization potential (I IP) The energy required to convert a monopositive cation to a dipositive cation is called second ionization potential (II IP)

n I IP < II IP < III IP because as the electrons go out of the atom, the effective nuclear charge increases and the ionic size goes on decreasing Thus the forces of attraction on valence shell electrons increases and hence the order

Factors affecting ionization potential

(i) Number of shells: With the increase in number of shells the atomic radius increases i.e., the distance of outer most

shell electron from the nucleus increases and hence the ionization potential decreases

(ii) Effective nuclear charge: Atomic size decreases with increase in effective nuclear charge because, higher the

effective nuclear charge stronger will be the attraction of the nucleus towards the electron of the outermost orbit and higher will be the ionization potential

(iii) Shielding effect: The electrons of internal orbits repel the electron of the outermost orbit due to which the attraction

of the nucleus towards the electron of the outermost orbit decreases and thus atomic size increases and the value of ionization potential decreases

(iv) Stability of half-filled and fully filled orbitals: The atoms whose orbitals are half-filled (p3, d5, f7) or fully-filled (s2, p6, d10, f14) have greater stability than the others Therefore, they require greater energy for removing an electron However, stability of fully filled orbitals is greater than that of the half-filled orbitals

(v) Penetration power: In any atom, the s orbital is nearer to the nucleus in comparison to p, d and f orbitals Therefore,

greater energy is required to remove an electron from s orbital than from p, d and f orbitals Thus the decreasing order

of ionization potential of s, p, d and f orbitals is as follows:

s > p > d > f

Periodic trends in ionization potential

(a) In a Period: The value of ionization potential normally increase on going from left to right in a period, because

ef-fective nuclear charge increases and atomic size decreases

Exceptions

n In second period, ionization, potential of Be is greater than that of B, and in the third period ionization potential of Mg

is greater than that of Al due to high stability of fully filled orbitals

n In second period, ionization, potential of N is greater than O and in the third period ionization potential of P is greater than that of S, due to stability of half-filled orbitals

n The increasing order of the values of ionization potential of the second period elements is

Li < B < Be < C < O < N < F < Ne

The increasing order of the values of ionization potential of the third period elements is

Na < Al < Mg < Si < S < P < Cl < Ar

Ionization Potential of Transition elements

n In transition elements, the value of ionization potential has very little increase on going from left to right in a period because the outermost orbit remains the same but electrons get filled up in the (n–1)d orbitals resulting in very little increase in the values of ionization potential

n In transition element series, the first ionization potential normally increases with increase in atomic number on going from left to right, but this periodicity is not uniform The value of ionization potential of transition elements depends on the following two important factors

(a) The value of ionization potential increases with increase in effective nuclear charge

(b) The value of ionization potential decreases with increase in shielding effect when the number of electrons increases

in (n–1)d orbitals

n In the first transition element series, the first ionization potential normally increases on going from left to right from Sc

to Cr because shielding effect is much weaker in comparison to effective nuclear charge The value of first ionization potential of Fe, Co and Ni remains constant, because shielding effect and effective nuclear charge balance one another

The value of ionization potential shows slight increase from Cu to Zn because they have fully filled s and d orbitals The value of first ionization potential of Mn is maximum because it has maximum stability due to fully filled s and half filled orbitals.

Ionization Potential of Inner Transition elements

The size of inner transition elements is greater than that of d block elements Therefore the value of ionization potential of

f block elements is smaller than that of d block elements and due to almost constant atomic size of f block elements in a period the value of their ionization potential remains more constant than that of d block elements

In a group

n The value of ionization potential normally decreases on going from top to bottom in a group because both atomic size and shielding effect increase

Exception:

n The value of ionization potential remains almost constant from Al to Ga in the 13th A group (B > Al , Ga > In)

n In 4th group i.e., Ti, Zr and Hf the I.P of Hf is higher than that of Zr due to Lanthanide contraction Thus the I.P of 4thgroup varies as Ti > Zr < Hf

Some Noteworthy Points

n In the periodic table the element having highest value of ionization potential is He

n The values of ionization potential of noble gases are extremely high, because the orbitals of outermost orbit are filled (ns2 , np6) and provide great stability

fully-n In a period, the element having least value of ionization potential is an alkali metal (group Ist ) and that having highest value is inert gas (Group 18th)

applications of ionization potential

n The elements having high values of ionization potential have low reactivity, e.g., inert gases

n The value of ionization potential decreases more on going from top to bottom in a group in comparison to a period Therefore, reactivity increases and the atom forms a cation by loss of electron

n The elements having low value of ionization potential readily lose electron and thus behave as strong reducing agents

n The elements having low value of ionization potential readily lose electron and thus exhibit greater metallic property

n The elements having low value of ionization potential readily lose electron and thus oxide and hydroxides of these ments have basic property

ele-Worked-out Examples

8 The atomic numbers of vanadium, (V), chromium (Cr), manganese (Mn) and iron (Fe) are respectively 23, 24, 25

and 26 Which one of these may be expected to have the highest second ionization enthalpy? (AIEEE 2003)

Hence, the correct order of atomic size is.

Solution: All the above four species have one electron each, H has least number of proton Therefore, H will have

least value of ionization potential

14 Which of the following should have less than 11.0 eV difference between the values of first and second ionization

potentials?

Ans (1)

Solution: Mg exhibits higher oxidation state (Mg+2) than Na (Na+1) and K (K+1)

15 The first I.P of nitrogen is greater than oxygen while the reverse is true for their second I.P values Explain

Solution: The first I.P corresponds to the removal of first electron Since nitrogen is already half-filled So more energy is required to remove the electron But once the electron is removed from oxygen it gains half-filled stability and therefore the 2nd I.P becomes high

16 The ionization energy of the coinage metals fall in the order Cu > Ag < Au Explain

Solution: In all the 3 cases an s-electron in the unpaired state is to be removed In the case of Cu, a 4s electron is

to be removed which is closer to the nucleus than the 5s electron of Ag So I.P decreases from Cu to Ag However, from Ag to Au the 14 f electrons are added which provide very poor shielding effect The nuclear charge is thus enhanced and therefore the outer electron of Au is more tightly held and so the IP is high

eLeCTRON aFFINITy

Atoms not only lose electrons to form positive ions but also gain electrons to form negative ions Electron affinity, as the

name indicates, is a measure of an atom’s tendency to gain an electron The higher an atom’s electron affinity the more likely it is to gain an electron Quantitatively, electron affinity is the nega tive of the enthalpy change associated with the addition of an electron to a mole of gaseous atoms to form gaseous ions with a 1– charge Thermal energy is released when

an electron is added to most atoms to form a 1– ion

A(g) + e– → A–(g) + thermal energy where A represents an atom of an element The change shown in equation is exothermic The energy change associated with an exothermic change always has a-sign The higher an element’s electron affinity, the more thermal energy is given off when an electron is added to an atom of the element

Fig 1.4 Electron affinities of the elements, kJ/mol

The most recent figure shows that electron affinity is a rather irregular periodic function of atomic number In general, electron affinity increases going from left to right across a row of the periodic table, reaching a maximum with the halogens

in Group VIIA The halogens have a much greater tendency to gain an electron than the alkali metals in Group IA Both actual and effective nuclear charges increase across a period Atomic radius decreases and an electron added to the outer shell is closer to a larger positive charge in atoms on the right side of the periodic table Thus, more energy is released when

an electron is added

As you can see from figure, only nitrogen and the elements of Groups 2, 12, and 18 have negative electron affinities Energy is required to force electron into atoms of nitrogen, the alkaline earth metals (Group 2), the elements of Group 12, and the noble gases (Group 18) Addition of an electron to nitrogen and to atoms of the elements of the groups listed is an endo-thermic process Because the noble gases have both high ionization energies and low electron affinities, the noble gases are unreactive

In atoms of elements of Groups 2, 12, and 18, all subshells that contain: electrons are filled; the next electron must go into a higher energy subshell The irregularities of the Group 15 elements, such as N and P, are also explained by electron configuration An electron added to an atom of a Group VA element must go into a p orbital that already contains an elec-tron The electron already present in the orbital repels the second electron As a result, less energy than expected is released when an electron is added to a Group VA element

In general, electron affinity decreases going down groups in the periodic table, for example, Cl > Br > I > At (see figure below) Going down a group, the radii of the shells increases because the principal quantum number is higher The added electron is further from the positive charge in the nucleus Therefore, the quantity of energy released when an electron is added is smaller

Fluorine appears out of line compared to the other members of Group VIIA The electron affinities of oxygen and nitrogen also seem low These three atoms are small and have high effective nuclear charges It is surprising that so little energy is given off when an electron is added to nitrogen, oxygen, and fluorine atoms However, the second shell is small and repulsions are large be tween the electrons already present in the second shell and the electron that is entering In third (and higher) period elements, the electrons already present in the outer shell are spread out through a larger volume Repulsions between them and the electron that is entering are smaller.

%U

5E U &G ;H

&O )

/L

3 1D

$U 0J 1H 1

%H +H

$WRPLFQXPEHU

Fig 1.5 Variation of electron affinities of the elements with atomic number Unfortunately, the

electron affinities of the lanthanides and of a few other elements have not yet been measured

Many properties of the first members of groups of representative elements are out of line compared with the properties

of the larger members of the group In several cases, the properties of elements in the second period are more like the erties of the third period element in the next-higher-numbered group than they are like the properties of the other members

prop-of the same group Thus, lithium is similar to magnesium, beryllium to aluminum, and boron to silicon

Worked-out Examples

17 The formation of the oxide ion O2– (g) requires first an exothermic and then an endothermic step as shown below:

O(g) + e– = O–(g), ∆H° = – 142 kJ mol–1

O–(g) + e– = O2–(g), ∆H° = 844 kJ mol–1

(a) O– ion will tend to resist the addition of another electron

(b) Oxygen has high electron affinity

(c) Oxygen is more electronegative

(d) O– ion has comparitively larger size than oxygen atom

Ans (a) O– ion exerts a force of repulsion on the incoming electron The energy is required to overcome it

18 In which of the following arrangements the order is not according to the property indicated against it?

(AIEEE 2005)

(a) Al3+ < Mg2+ < Na2+ < F– increasing ionic size

(b) B < C < N < O– increasing first ionization enthalpy

(c) I < Br < F < Cl– increasing electron gain enthalpy (with negative sign)

(d) Li < Na < K < Rb– increasing metallic radius

Ans (b) In a period the value of ionization potential increases from left to right with breaks where the atoms have

some what stable configuration In this case, N has half-filled stable orbital Hence has highest ionization energy Thus the correct order is B < C < O < N

19 For the addition of one mole of electrons to one mole of gaseous chlorine atoms

to an O– ion to form an O2– ion requires 708 kJ/mol of energy:

O–(g) + e– + 708 kJ → O2–(g)

The energy needed to force an electron into an O– ion is greater than the energy given off when an electron is added to

an oxygen atom to form the O– ion:

However, oxide ions are common in solids In solids, the negatively charged oxide ions are close to positively charged ions The attraction between oppo sitely charged ions is large and the combination of anions and cations is very exothermic

Periodicity in electron affinity:

(i) In general electron affinity value increases on moving from left to right in a period because effective nuclear charge increases

Exceptions

(a) The electron affinity value of alkaline earth metals of IIA group is zero

(b) Electron affinity value of alkali metals of IA group is also approximately zero because these elements have the

tendency of losing the electron instead of gaining the electron

(c) Electron affinity values of nitrogen and phosphorous (15) are lesser than the electron affinity values of carbon

and silicon respectively It is due to the comparatively stable half-filled configuration (np3) of nitrogen and phorus and the tendency to acquire the stable np3 configuration by the gain of one electron in carbon and silicon (np)2

(d) The theoretical value of the electron affinity of zero group inert gas elements is zero due to stable s2p6

configura-tion

(ii) In a group moving from top to bottom the electron affinity value of elements decreases because the atomic size increases

Exceptions

(a) Electron affinity values of second period elements are smaller than the electron affinity values of third period

elements of group 15 to 17 This unexpected behaviour can be explained by the very much high value of charge

densities, of second period elements due to much smaller size The electron being added experiences tively more repulsion and the electron affinity value decreases.

(b) The electron affinity of fluorine (Second period) is less than the electron affinity of chlorine (third period)

2p-orbitals in fluorine are much more compact than 3p-orbitals of chlorine So the electron being added in 2p-orbitals experiences comparatively more repulsion and the electron affinity value decreases

electrovalent or covalent—the concept of electronegativity

Electronegativity is a measure of the power of an atom to attract electrons; it involves the two terms, ionization energy and electron affinity Mulliken has defined electronegativity as the arithmetical mean of the first ionization energy and the electron affinity of an atom Pauling’s definition is some what different and it is his electronegativity values that are given in Table Fluorine, the most electronegative element, is given an arbitrary value of 4.0 and the electronegativities of the atoms

of other elements are related to it

Table 1.5 The Pauling Electronegativity Values of Some Elements

H 2.1

Li 1.0 1.5Be 2.0B 2.5C 3.0N 3.5O 4.0F

Na 0.9 Mg 1.2 1.5Al 1.8Si 2.1P 2.5S 3.0Cl

K 0.8 1.0Ca — Ge 1.7 As 2.0 2.4Se 2.8Br

Rb 0.8 1.0Sr — 1.7Sn 1.8Sb 2.1Te 2.4I

Cs 0.7 0.9Ba

As the table shows, electronegativity decreases down a particular group This is because the atom becomes sively larger and the inner elec tron shells decrease the attraction between the positive nucleus and the peripheral electrons Electronegativity increases across a particular period from alkali metal to halogen, since the progressive increase in nuclear charge exerts a contracting effect on the electron shells (in the same period, additional electrons reside in the same shell)

progres-A bond formed between two atoms of similar electronegativity will be essentially covalent Increase in negativity of one atom will result in that atom having greater control over the bonding pair of electrons, i.e., the co-valent bond will be polarized Further increase in electronegativity will result in increased polarity of the bond until eventually the electron pair can be considered to reside almost entirely on one atom, i.e., an electrovalent bond will be established, thus:

20 Electron affinity of SF5 is among the highest known but that of SF6 is quite modest

Solution: Sulphur in SF6 is saturated in terms of maximum covalency and maximum co-ordination number and

so it has very little tendency to attract electron In SF5 which is actually a free radical so it has got a very strong tendency to attract an electron to fulfill its covalency and co-ordination number Therefore the electron affinity of

SF5 is high

It now becomes clear why Group 1 and 2 elements react with Group 16th and 17th elements to give essentially trovalent compounds, caesium fluoride, Cs+ F-, being the most ‘ionic’ compound On the other hand, non-metals react by forming essentially covalent bonds.

elec-The reason why aluminium fluoride is an electrovalent compound, Al3+(F-)3, whereas aluminium chloride, Al2Cl6, is covalent is probably because the fluorine atom is more electronegative than the chlorine atom

An explanation in terms of lattice energy would run as follows:

Since the fluoride ion is smaller than the chloride ion, the lattice energy of the Al3+(F–)3 crystal structure would be greater than that of a similar Al3+(Cl–)3 structure by virtue of the closer approach of the oppositely charged ions (other things being equal) In this instance, the two explana tions amount to practically the same thing, since the fluorine atom is more electronegative than the chlorine atom because it is smaller and con sequently gives rise to a smaller anion

electronegativity scale

Some arbitrary scales for the quantitative measurement of electronegativities are as under:

(i) Pauling’s scale: Pauling related the resonance energy(∆AB) of a molecule AB with the electronegativities of the atoms A and B If xA and xB are the electronegativities of atoms A and B respectivey then,

In an ionic molecule AB, EA–B(experimental) is more than EA–B(Theoretical)

Pauling assumed the electronegativity value of fluorine 4 and calculated the electronegativity values of other elements from this value

(ii) Mulliken’s electronegativity: In 1934, Mulliken suggested an alternative approach to electronegativity based on the

ionization energy and electron affinity of an atom Consider two atoms A and B If an electron is transferred from

A to B, forming ions A+ and B–, then the energy change is the ionization energy of atom A (IA) minus the electron affinity of atom B (EB) that is, IA – EB Alternatively, if the electron was transferred the other way to give B+ and

A– ions, then the energy change would be IB – IA If A+ and B– are-actually formed,

then this process requires less energy, and

Mulliken used I and E values measured in electron volts, and the values were about 2.8 times larger than the Pauling values We now measure I and E in kj mol–1 The energy 1 eV/molecule = 96.48 kJ mol–1, so the commonly accepted Paul-ing values are more nearly obtained by perform ing this calculation (I + E)/(2 × 2.8 × 96.48) or (I – E)/540

This method has a simple theoretical basis, and also has the advantage that different values can be obtained for different oxidation states of the same element It suffers from the limitation that only a few electron affinities are known It is more usual to use the approach based on bond energies

aLLReD aND ROChOw SCaLe

In 1958, Allred and Rochow considered electronegativity in a different way, and worked out values for 69 elements They defined electronegativity as the attractive force between a nucleus and an electron at a distance equal to the covalent radius

This force F is electrostatic, and is given by:

2

effective 2

pal quantum number (the shell that the electron occupies), and the type of electron, s.p, d or f Screening factors have been

worked out by Slater, so this provides a convenient method of calculating electronegativity values These F values may be converted to electronega tivity values on the Pauling scale of values using an empirical relationship:

c = 0.744 + effective

2

0.359 Zr

The electronegativity values so obtained agree quite closely with those obtained by Pauling and Mulliken

As the oxidation number of an atom increases, the attraction for the electrons increases, so the electronegativity should also increase Allred and Rochow’s method gives slightly different values:

Mo(II) 2.18 Fe(II) 1.83 Tl(I) 1.62 Sn(II) 1.80

Mo(III) 2.19 Fe(III) 1.96 Tl(III) 2.04 Fe(IV) 1.96

Z = The actual number of charge present in the nucleus i.e., number of protons and σ = Shielding constant

Factors affecting the Magnitude of electronegativity

(i) Atomic radius: As the atomic radius of the element increases the electronegativity value decrease.

1Electronegativity

Atomic radius α

(iii) Effective nuclear charge: The electronegativity value increases as the effective nuclear charge on the atomic nucleus

increases

Electronegativity a Effective nuclear charge (Zeff)

(iii) Oxidation state of the atom: The electronegativity value increases as the oxidation state (i.e., the number of positive

charge) of the atom increases

(iv) Hybridization state of an atom in a molecule: If the s- character in the hybridization state of the atom increases

electronegativity increases because s-electrons are comparatively nearer to the nucleus For example, the egativity values of C-atom in various hybridization states are as under:

Periodicity in electronegativity

(i) In a period, moving from left to right, the electronegativity increases due to the increase in effective nuclear charge (ii) In a period, the electronegativity value of gp 1 alkali metal is minimum and that of gp 17 halogen is maximum (iii) In a group moving from top to bottom, the electronegativity decreases because atomic radius increases

(iv) The electronegtivity value of F is maximum and that of Cs is minimum in the periodic table

(v) The electronegativity of Cs(55) should be more than Fr(87) but it is less This is due to the increase of +32 units in

nuclear charge of Fr which makes the effective nuclear charge comparatively high

(vi) On moving from second to third transition series in a group [except Y(39) → La (57)], electronegativity increases due to the increase of +18 units in nuclear charge

(vii) The electronegativity of inert gas elements of zero group is zero due to stable s2p6 configuration Inert gases are monoatomic molecules and the electronegativity is of bonded atoms

applications of electronegativity

(i) Partial ionic character in covalent bond: The ionic character of a covalent bond increases as the electronegativity

difference of bonded atoms increases According to Haney and Smith, if the electronegativity difference of bonded atoms is ∆x then percentage ionic character of the bond = 16∆x+3.5∆x2

If the value of ∆x is 2.1 then ionic character percentage is about 50 For example the order of ionic character

in H–X bond is as follows:

H–F>H–Cl>H–Br>H–I

Because the electronegtivity difference of bonded atoms (∆x) decreases

(ii) Bond strength: If the electronegativity difference of covalently bonded atoms (∆x) increases, the bond energy of

the covalent bond also increases For example, the order of the H–X bond strength is,

H – F > H – Cl > H – Br > H – I

As the bond strength is decreasing the acid strength is increasing So order of increasing acid strength is,

HF < HCl < HBr < HI

(iii) Acidic and basic nature of oxides of normal elements in a period: The acidic nature of the oxides of normal

ele-ments increases as we move from left to right in a period In a period from left to right the electronegativity of the elements increases So the difference of the electronegativities of oxygen and the elements (xO –xE) decreases If the (xO – xE) values is about 2.3 or more then oxide will be basic If (x0 – xE) values is less than 2.3 the oxide will be acidic The oxides of the IIIA elements are amphoteric

The order of acidic or basic nature of the oxides of third period elements may be given as under:

Na2O MgO SiO2 P2O5 SO3 Cl2O7

The value of xO – xE is decreasing

Basic nature is decreasing

Acidic nature is increasing

(iv) Metallic and non metallic properties of elements:

(a) The metallic character decreases as the electronegativity of the element increases

(b) On moving from left to right in a periods, the electronegativity of the elements increases So the metallic

char-acter decreases

(c) On moving down a group, the electronegativity of the elements decreases So the metallic character increases

(v) Basic nature of the hydroxides of elements: A hydroxide MOH of an element M may ionize in two ways in water.

If the ionization is according to eqn (2) then it is basic This is only possible when ionic character of O –H bond

is less than M–O bond i.e (xO – xH) < (xO – xM)

Chemical bond

A chemical bond is defined as the force holding together two atoms or groups of atoms forming an aggregate of ions, or molecular species such that there occurs a lowering of energy Such definition permits us to discuss chemical bonds under

the following three classes:

1 The ionic bond (also called electrovalent or electrostatic bond),

2 The covalent bond and

3 The metallic bond.

It will be proper to warn at the outset that these three bonds represent only three extreme types of forces between atoms There will be many chemical compounds whose properties would indicate bonds of intermediate types The discussion on the above bond types will be followed by a short resume on the various intermolecular forces, often expressed as van der Waals forces

The ionic bond

It is wellknown that sodium and chlorine react vigorously to yield crystals of sodium chloride Crystallographic studies have shown that there is no discrete NaCI molecule in the crystal lattice; instead each sodium ion (Na+) is surrounded by six chloride ions (Cl–) and vice versa This packing of two ions is repeated throughout in a non-ending pattern The two oppositely charged ions are held together by means of electrostatic forces of attraction Each neutral sodium atom Na° loses one electron to form a monopositive sodium ion (Na+) Similarly, a neutral chlorine atom (Cl°) gains one electron to form a mononegative chloride ion (Cl–) The two ions Na+ and Cl– then build up a close packed type ionic crystal structure Conceptually, we find a transfer of electron from sodium atom to the chlorine atom:

and electronegativity also influence the formation of ionic compounds It is possible to draw the following conclusions:

(a) The lower the charge on the cation the better is the chance that an ionic compound will be formed The lower the charge on the cation the lower is the ionization potential/energy

(b) The formation of ionic compounds is favoured by large size of the metal atom or ion since an increase in size is sociated with a decrease in ionization potential/energy

(c) Small size and low charge on the anion will favour ionic compounds because these factors are associated with creasing electron affinity and electronegativity.

in-The above discussion gives us a rough guide as to when to expect an ionic compound Unexpected results are times observed For instance, the second electron affinity of oxygen is negative, yet it forms stable, ionic compounds with alkalies or alkaline earths Clearly, there must be other factors which play significant roles in the formation of ionic compounds

some-Lattice Energy

Mere formation of oppositely charged ions does not lead to an ionic compound In an idealised ionic compound there is considerable electrostatic force operating between the cations and anions considered to be hard spheres When these cat-ions and anions are brought together in large and equal (or equivalent) numbers they will arrange themselves in a regular pattern so as to attain a close packed type structure For example the structure of sodium chloride reveals that each sodium ion is surrounded by as many as six chloride ions and vice versa In caesium chloride each caesium ion has eight chloride ions as near neighbours and vice versa While attaining such close packed structure energy is liberated and thus the system

gets stabilized The energy liberated when one mole of the ionic crystal is formed from an assembly of free gaseous ions from infinite separation is called the lattice energy.

M+

(g) + X–

(g) → MX(s) + lattice energyConversely, work has to be done on the ionic crystal to disperse the crystal into an assembly of widely separated ions Formation of the lattice is an exothermic process while lattice dispersion is an endothermic process

Table 1.6 Lattice Energies of Some Alkali Halides

Alkali halide A–X distance in A (pm) Calculated lattice energy kcal/

mole (kJ/mole) Lattice energy from Born-Haber cycle kcal/mole (kJ/mole)

Energetics of Formation of Ionic Substances

The energy included in the formation of an ionic compound from its constituent elements may be considered as shown by the Born-Haber Cycle for the formation of one mole of sodium chloride from sodium and chlorine

Sublimation – (s) S (g) I (g)

+ +

A

-Dissociation Addition of e – 2(g) 1/2D* (g) E (g)

1 Cl Cl Cl

2 → +  + →

Crystal formation –

(g) (g) U (S)

+

Where, S = heat of sublimation of sodium metal

I = ionization energy of sodium

D = heat of dissociation of molecular chlorine

Ea = electron affinity of chlorine, and

U = lattice energy of sodium chloride

The amount of heat liberated in the overall reaction is the heat of formation of sodium chloride From the above,

∆Hf = S + I + 12D – Ea – U

The most important of these energy terms are I, Ea and U, since these are considerably greater than the remaining terms

S and D

More the negative value of the heat of formation, greater would be the stability of the ionic compound produced Thus

on the basis of the above equation, formation of an ionic compound is favoured by,

(a) low ionization energy (I) of the metal

(b) high electron affinity (E A) of the other element

(c) higher lattice energy (U) of the resulting compound

Solvation Energy

When a cation or an anion in the gaseous state reacts with a solvent in the liquid state, energy is liberated This liberated energy is termed salvation energy When the solvent concerned is water the liberated energy is called hydration energy:

Z±x (g) + ∝ H2O (l) → Z±x (aq) + hydration energy

Water is a polar solvent (each hydrogen is δ+ while each oxygen is δ–) A gaseous cation is attracted by the negative

oxygen end of the water molecule while the anion is attracted by the hydrogen end The greater the charge/surface area i.e., the charge density of the ion the greater will be the amount of energy liberated This explains why Li+ has a higher solvation energy than either Na+ or K+ Again Mg2+ has a higher hydration energy than Ba2+ or Li+ For the same reason hydration energy decreases from F– to I– ion.

Interestingly, there is a considerable difference in the hydration energies of F– and K+ ions (Table), although they have the same ionic radius (;1.36A; 136 pm) This is due to the fact that F– interacts with the small hydrogen atom (with

δ+ charge) of H2O while K+ has to interact with the much larger oxygen atom (with δ– ) More hydrogen atoms can get around a F– ion compared to oxygen atoms around the K+ ion Thus, there is scope of strong hydrogen bonding in the case

of F– but not in the case of K+

Table 1.7 Hydration Energy of Some Common ions

Ion Hydration energy kcal/mole (kJ/mole) Ion Hydration energy kcal/mole (kJ/mole)

The high charge density of H+ leads to high hydration energy of the ion, Covalently bonded hydrogen chloride dissolves

in water to form hydrated ions In water there is an interaction between the dipoles of HCI and H2O The polar hydrogen end of HC1 moves close to the negative oxygen end of water while the chlorine end orients itself towards the hydrogen end of water Combined hydration energies of H+ and Cl(–1091 –364 = –1455 kJ/mole) is so high that it leads to splitting

of the covalent bond in HCl in favour of formation of the hydrated ions

Properties associated with Ionic bond

1 Ionic crystal lattice is built up of oppositely charged ions.

2 Since electrostatic forces of attraction are omnidirectional i.e., extend in all directions, each ion tends to drag as many

of opposite kind around itself A non-ending array of alternate positive and negative ions exists Hence no isolated discrete molecule exists in the crystal lattice; instead “giant molecules” are formed Ionic bonds are non-directional

3 Ionic compounds are hard, and have high melting points and boiling points This is because strong electrostatic forces

have to be overcome in order to provide the necessary freedom of movement that characterises a liquid state

So long as the ionic charge is not too large, hardness of ionic crystals increases with increasing ionic charge and

decreasing ionic radius:

M2+radius: A (pm) 1.10(110) 0.99(99) 0.65(65)

(In Mohs scale hardness of diamond is taken as 10)

4 Ionic compounds are soluble in polar solvents and insoluble in nonpolar solvents When an ionic compound is

dissolved the crystal lattice is ruptured and the ions are held farther apart in solution than in the crystal During the interaction of ionic compound with solvent, the lattice energy is lost but a solvation energy is gained due to the ions getting solvated The negative end of a dipolar solvent molecule, attracts the cation and the positive end faces the

anions A polar solvent like water can penetrate the crystal lattice of an ionic compound and can effect dissolution

The solvent character is decided by the dielectric constant of the solvent, and the higher the dielectric constant the better it is as a solvent for ionic compounds Conversely a non-polar solvent does not possess the necessary dipoles

to initiate dissolution of an ionic compound It has a low dielectric constant

5 Ionic compounds are conductors in fused state or in solution In solid crystalline state such compounds are only poor

conductors In fused state an ionic compound is composed of ions which are not confined to fixed positions as in the crystal lattice The freely moving ions can therefore permit electric conductance

6 The reactions of ionic compounds are reactions of their ions and these are instantaneous in solution Common

precit-piation reactions used in analysis belong to the class of ionic reactions

Some Noteworthy Points

n Ionic bond was proposed by Kossel

n Elements of group 1 and group 2 on combining with halogens, oxygen and sulphur generally, form ionic bonds

n Beryllium generally forms covalent/polar covalent bonds

n Bonding in compounds of transition metals (in lower oxidation state) is ionic with partial covalent character

n No bond is 100% ionic in character It has some percentage of covalent character which is explained on the basis of Fajan rules (discussed later)

n The lattice energy of bi-bivalent ions > uni-bi, or bi-uni > uni- univalent ions.

n PbSO4, AgCl, AgBr and Agl are ionic compounds but only sparingly soluble (almost insoluble) in water because of high value of lattice enthalpy

n The polar solvents have high values of dielectric constant Water is one of the best solvents as it has high value of electric constant Due to high value of dielectric constant, the electrostatic force of attraction between the ions decreases and these ions get separated and ultimately get solvated by the solvent molecules The non-polar solvents having very low value of dielectric constants are not capable of dissolving ionic compounds

di-n The dielectric constant of D2O is greater than H2O

n Dissolution of most of the ionic solids like NaCl, KC1, KBr, NH4NO3 etc., in water is endothermic i.e., ∆Hsoln is tive This clearly indicates that in all such cases;

Dissolution in such cases is possible due to large positive value of ∆Ssoln. which, makes the process spontaneous by ing and ∆Gso1n negative (∆G = ∆H — T∆S)

Worked-out Examples

21 BaSO4 being an electrovalent compound does not pass into solution state in water Explain

Ans The hydration energy of BaSO4 being lesser than the lattice energy and thus ions are not separated

22 Why do electrovalent compounds exist as solid?

Ans Strong non-directional coulombic forces between ions result in higher lattice energy

23 Sodium conducts electricity in solid state while sodium chloride does not?

Ans Sodium being metal involves metallic bonding between atoms in solid state having mobile electrons which

conducts current In NaCl, ions are held in lattice and not free to move

Valence Bond (VB) Theory of Covalent Bond: We have so far not attempted to answer the logical question as to

why at all sharing of electrons gives rise to the formation of a stable bond The exact answer lies in quantum mechanics

We will endeavour to offer an explanation in popular language

Fig 1.6 Potential energy curve for hydrogen molecule

Let us consider the covalent bonding in the diatomic hydrogen molecule Each hydrogen atom consists of a one-proton

nucleus and a single 1s electron We assume that nucleus A and electron A constitute the hydrogen atom Ha and that nucleus

B and electron B make up the second hydrogen atom HB As the two atoms approach each other in addition to the lombic attraction between each nucleus and its own electron, new Coulombic attractive forces begin to operate between nucleus A and electron B and between nucleus B and electron A Again Coulombic repulsive forces also exist between the two electrons, and between the two nuclei

Cou-The Covalent bond

It is easy to have an understanding of an ionic bond which is formed as a result of electrostatic attraction between positely charged ions Conceptually, it is, however, much difficult to gain an understanding of covalent bonding It is unreasonable to suppose that the bond holding identical non-metal atoms to each other, as in H2 or F2 molecule, could be formed in the way sodium chloride is formed Electron transfer seems more reasonable when two atoms with widely dif-ferent electron affinity, such as hydrogen and fluorine, combine to form the compound HF The properties of HF, however, differ widely from those characteristic of ionic compounds Hydrofluoric acid is a non-conductor of electricity, is low boil-ing etc Lewis suggested that such non-ionic molecular compounds (H2, F2, HF etc.,) are formed as a result of sharing of electrons between atoms whereby each atom can attain a noble gas configuration (that is either two electrons as in helium,

op-or an octet in the outermost shell like the other noble gases) A covalent bond may be defined as a fop-orce holding together atoms through sharing of electrons

Hypervalent and Subvalent Compounds: There are many compounds where the number of electrons around an atom exceed an octet Such compounds are called hypervalent compounds Examples are SF6, POC13, SOC12, etc On the other hand, there are compounds where the number of electrons around an atom falls short of an octet e.g., NO, NO e.g., etc These latter are known as subvalent compounds

Worked-out Example

24 Which properties of the elements depend on the electronic configuration of atoms and which do not?

Ans Chemical and physical properties depends on electronic configuration; nuclear properties dot not.

At large distances of separation, the system consists of two noninteracting hydrogen atoms As the atoms approach each other the electrical potential energy is converted into kinetic energy As the distance decreases, the conver sion is more rapid A critical internuclear dis tance is ultimately reached, at which the poten tial energy of the two-atom system reaches a minimum At this distance (0.74 A) (74 pm) the attractive forces balance the repulsive forces and a stable

H2 molecule results When the internuclear distance falls below 0.74 A (74 pm) the repulsive forces predominate and the molecule becomes unstable There is a proton-proton attractive force only within the very short nuclear range of 10–13 cm Below 0.74 A (74 pm) the forces are still electrostatic in nature and because of the proximity of the two protons the proton-proton repulsion very much out-weighs the attractive forces of one proton over the electron of the other atom The result is dissociation of H2 molecule into atoms

At the optimum internuclear distance of 0.74 A (74 pm) each nucleus exerts an attractive forces on its own electron as also on the electron of the other atom Consequently, the two electrons forming the covalent bond are shared by the two atoms Quantum mechanically we say that there is a very high probability of finding the pair of electrons in between the two atoms The higher the probability the stronger is the bond That is, the greater is the energy necessary to break the

bond In terms of orbitals of the two electrons, there occurs a significant overlap of the two 1s orbitals of the two hydrogen

atoms, such that one electron originally fonfined to a single orbital can now spread itself over the other orbital Thus the two electrons of HA and HB now have a common orbital and allowing according to Pauli principle no more than two elec-trons per orbital, it follows that the electron spins of HA and HB have to be opposite if they are to share a common orbital

In the case of two with their electron spins parallel there is no minimum in the potential energy curve i.e., no H2 molecule

is formed Therefore in order that a covalent bond may be formed between two atoms a necessary condition is that each of the atoms must possess an unpaired spin.

at oms, and between one hydrogen atom and a fluorine atom since each fulfils the criterion of having one unpaired electron

in its outer orbital But no such molecule formation can take place between two helium atoms since helium has a tion of 1s2 that is, a filled is shell with no unpaired spin

configura-Covalency

It is defined as the number of electrons contributed by an atom of the element for sharing with other atoms as to achieve noble gas configuration It can also be defined as the number of covalent bonds formed by the atom of the element with other atoms The usual covalency of an element except hydrogen (which has covalency 1) is equal to (8-number of the group to which an element belongs) This is true for the elements belonging to 14, 15, 16 and 17 groups

Element Group (8–Group number) Covalency

7ZR

7KUHH

S

These four elements do not possess d-orbitals in their valency shell However, the elements having vacant d-orbitals

in their valency shell like P, S, Cl, Br, I, show variable covalency by increasing the number of unpaired electrons under excited conditions, i.e., unpairing the paired orbitals and shifting the electrons to vacant d-orbitals [Such a shifting is not

possible in the case of H, N, O and F because d-orbitals are not present in their valency shell].

Phosphorus shows 3 and 5 covalencies

Sulphur atom shows 2, 4 and 6 covalencies

Sulphur atom in ground state

7ZRXQSDLUHGHOHFWURQV

FRYDOHQF\ 

Sulphur atom in excited state,

(a) When p-orbital is unpaired (first excited state.)

(second excited state.)

Chlorine shows 1, 3, 5 and 7 covalencies

Chlorine atom in excited state,

(a) First excited state when p-orbital is unpaired

two p-orbitals are unpaired

Thus, variable covalency is shown by those elements whose atoms have vacant d-orbitals in their valency shell.

Covalent bond may be single, double or a triple bond Double and triple covalent bonds are called multiple covalent

bonds Single covalent bond is formed by sharing of only one electron pair This bond is represented by single dash (—)

Dou-ble and triple covalent bonds are formed when atoms bonded together share two or three electron pairs, respectively These bonds are represented by double dash (=) and triple dash (=) respectively Some examples of covalent bonding are given below:

(i) Formation of hydrogen molecule: In the forma tion of hydrogen molecule, each hydrogen atom contributes one

electron and then the pair is shared between two atoms Both the atoms acquire stable configuration of helium Thus, the molecule consists of one single covalent bond