CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES

THE PERIODIC TABLE OF THE ELEMENTS

PERIODIC TABLE

It is a table of elements in which the elements with similar properties are placed together.

DOBEREINER’S LAW OF TRIADS

This was the classification of elements into groups of three elements each with similar properties such that the atomic weight of the middle element was the arithmetic mean of the other two e.g. Ca, Sr, Ba; Cl, Br, I etc.

The difference between atomic weights of any two successive elements is nearly constant.

NEWLAND’S LAW OF OCTAVES

This was an arrangement of elements in order of increasing atomic weights in which it was observed that every eighth element had properties similar to those of the first just like the eight node of an octave of music.

Prouts Hypothesis : Atomic weight of an element is simple multiple of atomic weight of hydrogen.

MENDELEEV’S PERIODIC TABLE

This is based upon Mendeleev’s periodic law which states that the physical and chemical properties of the elements are a periodic function of their atomic weights.

Mendeleev named Gallium and Germanium elements as eka-aluminium and eka-silicon respectively.

STRUCTURAL FEATURES OF THE MENDELEEV’S PERIODIC TABLE

1. Mendeleev’s original periodic table consists eight vertical columns are called groups I-VIII & seven horizontal rows are called periods 1-7. But modified Mendeleev’s periodic table contains nine groups, i.e., I-VIII and zero (of noble gases).
2. All groups except VII and zero have been further divided into two sub-groups called A and B. A groups of left hand side consist of normal elements while groups B of right hand side consist of transition elements.
3. Elements of group IA are called alkali metals while those of group IB (i.e, Cu, Ag, Au) are called coinage metals.

MODERN PERIODIC LAW

Moseley formed the basis of the modern periodic law. He discovered that the square root of the frequency of the more prominent X-rays emitted by a metal was proportional to the atomic number and not the atomic weight of the atom of the metal. Hence atomic number should be the basis of classification of the elements.

Modern periodic law states, “that the physical and chemical properties of the elements are a periodic function of their atomic number.”

It was observed that the elements with similar properties reoccurred at regular intervals of 2, 8, 8, 18 or 32. These numbers (2, 8, 8, 18 and 32) are called magic numbers, and cause of periodicity in properties.

STRUCTURAL FEATURES OF THE LONG FORM OF THE PERIODIC TABLE

1. GROUPS -
1. The 18 vertical columns, of the periodic table, are called groups
2. Elements of groups 1, 2, 13-17 are called normal or representative elements.
3. Elements of groups 3-12 are called transition elements
4. The elements belonging to a particular group is known as a family and is usually named after the first element. For example, Boron family (group 13). In addition to this, some groups have typical names. For example,

Elements of group 1 are called alkali metals

Elements of group 2 are called alkaline earth metals

Elements of group 3 are called pnicogens

Elements of group 16 are called chalcogens

Elements of group 17 are called halogens

Elements of group 18 are called noble gases

2. PERIODS - The 7 horizontal columns (or rows) are called periods. The seven periods of periodic table are

Shortest period - 1st period (1H 2He) contains 2 elements. It is the shortest period.

Short periods - 2nd period (3Li 10Ne) and 3rd period (11Na  18Ar) contains 8 elements each. These are short periods.

Long periods - 4th period (19K 36Kr) and 5th period (37Rb 54Xe) contain 18 elements each and are called long periods.

Longest period - 6th period (55Cs – 86Rn) contains 32 elements and is the longest period.

Incomplete period - 7th period (87Fr –) is, however, incomplete and contains at present only 26 elements.

3. BLOCKS - The periodic table is divided into four main blocks (s, p, d and f) depending upon the subshell to which the valence electron enters into.
1. s-Block - Elements of group 1 and 2 constitute s-block.
2. p-Block - Elements of group 13 to 18 constitute p-block.
3. d-Block - Elements of group 3 to 12 constitute d-block.
4. There are three complete series and one incomplete series of d-block elements. These are:

1st or 3d-transition series which contains ten elements with atomic number 21-30 (21Sc–30Zn).

2nd or 4d-transition series which contains ten elements with atomic number 39–48  (39Y–48Cd).

3rd or 5d-transition series which also contains ten elements with atomic n umber 57 and 72-80 (57La, 72Hf–80Hg).

4th or 6d-transition series which contains only ten elements.

5. The f-block elements comprise two horizontal rows placed at the bottom of the periodic table to avoid its un-necessary expansion and make the symmetrical nature of periodic table. The two series of f-block elements containing 14 elements each. Lanthanides - The 14 elements from 58Ce–71Lu in which 4f-subshell is being progressively filled up are called lanthanides or rare earth elements. Actinides - Similarly, the 14 elements from 90Th –103Lr in which 5f-subshell is being progressively filled up are called actinides.
6. Elements of s and p-blocks are called normal or representative elements, those of d-block are called transition elements while the f-block elements are called inner transition elements.
7. The 11 elements with Z = 93–103 (93Np – 103Lr) which occur in the periodic table after uranium and have been prepared by artificial means are called transuranics. These are all radioactive elements.

NOMENCLATURE OF ELEMENTS WITH ATOMIC NUMBER > 100

According to IUPAC the nomenclature can be derived using numerical roots for 0 and numbers 1-9 for atomic numbers of elements.

The roots are put together in order of digits which make the atomic number and ‘ium’ is added at the end. Use the following table

Example : Name the IUPAC name of the element of atomic  number 108 : Name will be Unniloctium and symbol - UnO

Example – Name the element with atomic number 115.

Name will be - Ununpentium and symbol UuP

DIAGONAL RELATIONSHIP

Certain elements of 2nd period i.e., Li, Be, B. show similarity with their diagonal elements in the 3rd period i.e., Mg, Al, Si, as shown below:

This is called diagonal relationship and is due to the reason that these pairs of elements have almost identical ionic radii and polarizing power. (i.e. charge/size ratio).

PERIODIC PROPERTIES

Properties which show a regular gradation when we move from left to right in a period or from top to bottom in a group are called periodic properties. These properties are atomic size, ionisation energy electron affinity etc.

ATOMIC SIZE

It refers to the distance between the centre of nucleus of atom to its outermost shell of electrons. The absolute value of atomic radius cannot be determined because

• It is not possible to locate the exact position of electrons in an atom as an orbital has no sharp boundaries.
• It is not possible to isolate an individual atom.
• In a group of atoms, the probability distribution of electrons is influenced by the presence of neighbouring atoms.

Since absolute value of atomic size cannot be determined, it is expressed in terms of the operational definitions such as covalent radius, vander waal’s radius, ionic radius and metallic radius.

Covalent radius : It is half of the distance between the nuclei of two like atoms bonded together by a single covalent bond, hence it is also known as single bond covalent radius (SBCR). Thus, covalent radius .

where d = internuclear distance between two covalently bonded like atoms.

Van-der Waal’s radius : It is one-half of the distance between the nuclei of two adjacent atoms belonging to two neighbouring molecules of an element in the solid state.

The covalent radius is always smaller than the corresponding van der waal’s radius.

Metallic radius : It is half of the distance between two successive nuclei of two adjacent metal atoms in the metallic closed packed crystal lattice. Metallic radius of an element is always greater than its covalent radius.

Ionic radius :  It is the effective distance from the nucleus of the ion upto the electrons in the outer shell to which it can  influence the ionic bond.

An atom can be changed into a cation (by loss of electron) which is always much smaller than the corresponding atom, or to an anion (by gaining of electrons) which is always larger than the corresponding atom.

FACTORS INFLUENCING COVALENT RADIUS

1. Multiplicity of bond : Covalent radii depends on the multiplicity of bonds. e.g.,

     Bond length     Radius of C-atom

 1.54Å 

 1.34Å 

2. Percentage of ionic character : Covalent radius of H in HCl, HBr and  HI are different.
3. Effective nuclear charge : Greater the effective nuclear charge, more tightly is the hold with nucleus and hence smaller the radius.

PERIODIC VARIATION OF ATOMIC RADII

1. On moving down the group the valence shells become far away from the nucleus and thus the atomic radius increases.
2. On moving along the period, the effective nuclear charge increases and thus the electron cloud is attracted more strongly towards the nucleus resulting in the contraction of atomic radius.

ISOELECTRONIC IONS OR SPECIES

These are ions of the different elements which have the same number of electrons but different magnitude of the nuclear charge. The size of isoelectronic ions decreases with the increase in the nuclear charge.

IONISATION ENERGY (I.E.)

The amount of energy (work) required to remove an electron from the last orbit of an isolated (free) atom in gaseous state is known as ionisation potential or energy or better first ionisation potential of the element, i.e.,      

1. The amount of energies required to remove the subsequent electrons (2nd, 3rd, ...) from the monovalent gaseous cation of the element one after the other are collectively called successive ionisation energies. These are  designated as I.E1, I.E2, I.E3, I.E4 and so on. It may be noted that.  I.E4 > I.E3 > I.E2> I.E1  (for a particular element)

IE is expressed in eV/atom or kcal mol–1 or kJ mol–1

Note that eV atom–1 = 23.06 kcal mol–1 = 96.3 kJ mol–1

2. In general, the first I.E. increases along the period from left to right. However there are some exceptions to the general trend -
1. I.E. decreases from elements of group 2 3.
2. I.E. decreases from elements of group 15 16.
3. In a group of the periodic table, the ionisation energy decreases from top to bottom.
4. The factors which affect the ionisation energy are
1. Atomic size or radius : I.E. decreases as the atomic size increases so the attractive force decreases.
2. Number of electrons in the inner shell (screening effect) :  On moving down a group, the number of inner shells increases which increase the screening effect and hence the ionisation potential tends to decrease.
3. Nuclear charge : On moving along the period, effective nuclear charge increases due to addition of electrons in same valence shell and hence ionisation energy increases.
4. Stable configuration : Half filled or completely filled subshells possess extra stable nature and thus it is more difficult to remove electron and hence more is I.E.
5. Penetration effect : More penetrating (i.e. more closes) are subshells of a shell to the nucleus, more tightly the electrons are held towards the nucleus and more is I.E.

I.E. : s > p > d > f for a given shell

Penetration power  Surface area of a subshell

5. In second period elements, the correct increasing order of ionisation energies is

IE1 : Li < B < Be < C < O < N < F < Ne

IE2 : Be < C < B < N < F < O < Ne < Li

6. In third period elements, the correct increasing order of ionisation is

IE1 : Na < Al < Mg < Si < S < P < Cl < Ar

IE2 : Mg < Si < Al < P < S < Cl < Ar < Na

ELECTRON AFFINITY (EA)

It is the amount of energy released when a gaseous atom accepts the electron to form gaseous anion .

1. EA values are expressed in eV/atom or kcal/mol or kJ/mol.
2. The energy change accompanying the addition of first, second, third etc. electrons to neutral isolated gaseous atoms are called successive electron affinities and are designated as EA1, EA2, EA3 etc.
3. The first EA is always taken as positive. However, the addition of second electron to the negatively charged ion is opposed by coulombic repulsion and hence required (absorbed) energy for the addition of second electron. Thus, second electron affinity (EA2) of an element is taken as negative. For example,

(Exothermic)

(Endothermic)

4. Electron affinity increases in moving along the period from left to right due to increase in charge. But the values are unexpectedly low in elements of group 2, 15 and 18 due to stable electronic configurations of exactly half-filled and completely filled orbitals.
5. Within a group of the periodic table the electron affinities decreases from top to bottom.
6. In general, electron affinity follows the following trend:

Halogens > Oxygen family > Nitrogen family > Metals of groups 1 and 13 and Non-metals of group 14 > Metals of group 2.

7. The electron affinities are indirectly measured with the help of Born-Haber Cycle,  i.e.,

Where, S, D, IE, EA and U are the heat of sublimation, bond dissociation energy, ionization energy, electron affinity and lattice energy respectively.

8. Electron affinity depends upon:-
1. Effective nuclear charge : Greater the effective nuclear charge, more is the attraction of nucleus towards the electron and  hence higher will be the value of E.A.
2. Atomic size : Greater the atomic radius of the atom, less will be the attraction of the nucleus to the electron to be added and hence lower will be the value of E.A.
3. Penetrating power : Due to penetrating power, E.A. for addition of electron show the order s > p > d > f
4. Electronic configuration : Half filled and fully filled subshell are extra stable and thus oppose the addition of electron which leads to lower, E.A. value e.g. EA, of C > EA, of N.

ELECTRONEGATIVITY (EN)

It is the tendency of an atom in a molecule to attract the bonded shared pair of electrons, towards itself

1. There are several electronegativity scales:-
1. Mulliken scale : On the Mulliken scale, electronegativity X is taken as average of IE and EA, i.e.,

 where IE and EA are expressed in electron volts

or where IE and EA are expressed in kJ mol–1

or  where IE and EA are expressed in kcal mol–1

2. Pauling scale : This is the most widely used scale and is based upon bond energy data. According to Pauling, the difference in electronegativity of two atoms A and B is given by the relationship as

where XA and XB are electronegativities of the atoms A and B respectively while.

where EA–B, EA–A and EB–B represent bond dissociation energies of the bonds A-B, A-A and B-B respectively. The Pauling and the Mulliken scales are related to each other by the relation,

2. In a period, EN increases from left to right due to decrease in size and increase in nuclear charge of atoms.
3. In a group, EN decreases from top to bottom due to increase in atomic size.
4. Electronegativity depends on:

• Atomic size
• Nuclear charge
• Shielding effect
• Oxidation state - EN increases as the positive oxidation state increases.
• Hybridization - Greater is the s-character in a hybrid orbital more is electronegativity.

5. If electronegativity difference is greater than 1.7 the bond is ionic otherwise covalent.
6. In general, greater is difference of EN between two atoms smaller is the bond length.
7. Smaller the electronegativity, larger is the atomic size.

VALENCY

Valency of an element is the number of electrons gained or lost or shared with other atoms in the formation of compounds.

Valency of group 1 and 2 elements is equal to the number of electrons in the outermost shell, while that for groups 13 to 14 is group number -minus 10 and that for group 15–18 is 8 -minus the number of electrons in the outermost shell.

ATOMIC VOLUME

It may be defined as the volume occupied by one mole atoms of the element at its melting point in solid state.

1. It is obtained by dividing the gram atomic mass of the element by its density.
2. It decreases along the period, reaches a minimum in the middle and then starts increasing, because of different packing arrangement of atoms in different elements in the solid state, i.e., P4, S8 etc.
3. In moving down the group atomic volume goes on increasing gradually

ACID-BASE BEHAVIOUR OF OXIDES AND HYDROXIDES

The oxides or hydroxide of an element may act either as base or an acid depending upon its ionization energy.

1. If the IE is low, it acts as a base and if the IE is high, it acts as an acid.
2. The IE of alkali metals is the lowest, therefore, their oxides and hydroxides are the strongest bases. The basic character of their hydroxides increases in the order:

CsOH > RbOH > KOH > NaOH > LiOH

3. The IE of halogens is quite high, therefore, their oxides are the strongest acids. The acidic character of their oxides and hydroxides decreases in the order:

HClO4 > HBrO4 > HIO4

4. Within a period, the ionization energies of the elements usually increase and hence their oxides and hydroxides show a gradual variation from strongly basic through amphoteric to strongly acidic character.
5. The non-metallic character, oxidising character and acidic nature of oxides of the elements increases from left to right in a period and decrease from top to bottom in a group. The stability of the metal increases and activity decreases from left to right in a period whereas the stability decreases and activity increases down the group.

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