Chemistry Part I Chapter 3 Classification Of Elements And Periodicity In Properties

(i) p-block: Six groups
(ii) d-block: Ten groups.

These are called alkali elements.

As the d-subshell can have a maximum of ten electrons.

(i) Transition elements:
     (n - 1)d^1-10  ns^1-2.
(ii) Inner transition elements:
      (n - 2)f^1-14   (n - 1)d^0-1  ns².

Tips: -

(i) The classification simplifies the learning of chemistry.

(ii) Periodic table explains clearly why elements in a group display similar properties. The periodic table serves as a guide for scientific investigations.

In the third period the filling up of only 3s and 3p-orbitals occur. Therefore, in this period there are only two s- and six p-biock elements. Since the third period (n = 3) starts with Z = 11 and ends at Z = 18, hence the element is chlorine (CI) with atomic number (Z) = 17 which is present in the third period and seventeenth group of the periodic table.

Covalent radius of chlorine atom
=  136 pm - 37 pm = 99 pm.

It increases from left to right in a period.

The element with highest IE₁ value: Kr (Noble gas).
The element with lowest IE₁ value: K (Potassium)

(i) I^-  (ii) Br has the greater ionisation enthaply.

Boron (B) has lowest Ionization energy.

Higher the value of ionisation enthalpy of an atom, greater will be its electronegative and non-metallic character. 

Lower the Ionisation enthalpy, greater will be the reducing power and basic nature of the element.

Helium (He).

More enthalpy is requried to remove an electron held more firmly by the unipositive ion.

Generally increases.

Noble gases elements occupy the maxima of ionisation potential curve.

O>C>B>N.

+116 kJ mol^-1.

The size of an isoelectronic species increases with a decrease in the nuclear charge (Z). For example, the order of the increasing nuclear charge of F^–, Ne, and Na⁺ is as follows:

       F^– < Ne < Na⁺

Z     9     10     11

Therefore, the order of the increasing size of F^–, Ne and Na+ is as follows:

Na⁺ < Ne < F^–

(i) Element having lowest negative electron gain enthalpy is C.
(ii) Element having highest I.E₁ value is C.

Valency of chlorine = (-1)
Valency of water = 0
Let the oxidation state of Al be x.
           x -1         0
       [Al Cl   (H₂O)₆]²⁺
        x - 1 + 6(0) = +2
or                         x = +3
Thus the oxidation state of Al is +3. Since six legends (which can denote a pair of electrons) i.e. H2O molecules are attached to the alumninum atom, thereofre, its covalency is 6.

In chemistry 114 element are known. The most of elements take a part in variety of chemical transformations to number of substances and it is also very difficult to study the properties of all these element and their compound individually. Thus in order to solve this problem it is necessary to classify the element.

Different scientists tried to group elements with similar elements.
(i) Dobereiner Triads (1822). Dobereiner arranged similar elements in a group of three so that the atomic mass of the central element was nearly equal to arithmetic mean of the atomic mass of the other two elements. The group of three elements was named as triad. For example,
Triad                    Lithium     Sodium          Potassium           Mean atomic mass
Atomic masses       7                23                    39                     
Triad                  Chlorine        Bromine              Iodine             Mean atomic mass
Atomic masses       35.5            80                      127                    
But this concept of triad could be applied to a limited number of elements. It was rejected soon.
(ii) Newland's law of octaves (1864). Newland observed that if the known elements were arranged in order of their increasing atomic masses, similar properties recurred in every eighth element like nodes in a musical scale. This generalisation was known as Law of octaves. This system worked very well for lighter elements.

 

But this law failed in case of heavier elements and was rejected.
(iii) Lother Meyer’s work (1870). Lother Meyer plotted a graph between atomic volumes of elements in the solid state and their atomic masses. He found that the elements with similar physical properties occupied similar positions on the curve. By taking into account this fact, he drew up a periodic table in which elements were arranged in order of increasing atomic masses.

Mendeleev arranged all the known elements (63 at that time) in horizontal rows and vertical columns in order of increasing atomic masses in the form of table known as Periodic table. Mendeleev’s table has few vertical columns called groups and a few horizontal rows called periods. Mendeleev left certain gaps in his table and predicted that new elements would be discovered to fill up the gaps.
The Mendeleev's modified periodic table consists of:
(i) Nine vertical columns called groups. These are numbered from I to VIII and zero (The members of the zero group were not discovered at the time of Mendeleev). Each group from I to VII, is further sub-divided into two sub-groups-A and B. Group VIII consists of three sets, each containing three elements. Group zero consists of noble gases.

(ii) Seven horizontal rows are called periods. These are numbered from I to VII. First period contains two elements, second and third periods contain eight elements each. These periods are called short periods. Fourth and fifth periods contain eighteen elements each. These periods are called long periods. Sixth period contains thirty two elements and is called longest period. Seventh period is incomplete and contains nineteen elements (At. No. 87-105).

The periodic table has following merits and applications:
(i) Prediction of new elements: Mendeleev left various vacant places in his table which provided a clue for the discovery of new elements.
(ii) Determination of correct atomic masses. This table has also been useful in correcting the doubtful atomic masses of some elements.
(iii) In research work. Periodic table provided useful information which helped in stimulating research work.
(iv) Prediction of properties. The properties of elements and their compounds can be predicted from their positions in the table.

Drawbacks of Mendeleev's table:
(i) Position of hydrogen: Hydrogen is placed in group I-A of the periodic table. However, it resembles the elements of both group I A (alkali metals) and group VII A (halogens). Therefore, the position of hydrogen in the periodic table is not clear.
(ii) Anomalous pairs of elements: In some cases, elements of higher atomic masses are placed before those having lower atomic masses. Examples:
  Ar        and         K;
39.95                     39.1
Co          and        Ni;         Te       and    I
58.93                  58.71    127.6           126.9
(iii) Metals and non-metals: No attempt has been made to place metals and non-metals separately in it.
(iv) Lanthanoides and Actinodes: Fifteen elements are placed in one position.
Group III B
6th period:  La + 14 lathanoide elements. (₅₈Ce to ₇₁Lu).
7th period:   Ac + 14 actinoide elements. (₉₀Th to ₁₀₃Lr)
Thus, lanthanoids and actinoides have not been provided separate and proper places in the Mendeleev's periodic table.
(v) Position of isotopes:  Isotopes of elements are placed in the same position in the table though according to their atomic masses they should have been placed in different positions.
(vi) Similar elements separated in the table: Certain elements such as copper and mercury or gold and platinum which possess similar chemical properties are placed in different groups.
(vii) Dissimilar elements placed together in the same group: Alkali metals such as Li, Na, K etc. (group IA) are grouped together with coinage metals such as Cu, Ag and Au (group IB) though their properties are quite different.
(viii) Number of elements in the periods. The presence of only 2 elements in the first period, 8 in the second and third etc. can not be explained.

The repetition of the properties of elements after certain intervals when the elements are arranged in order of increasing atomic numbers is known as periodicity.

Cause of periodicity: We know that the physical and chemical properties of the elements depend upon the number of electrons present in their outermost shell. Elements having similar outer electronic configurations have similar properties. When the elements are arranged in the periodic table on the basis of their increasing atomic numbers, similar valence shell electronic configurations are repeated after certain regular intervals of atomic numbers i.e. 2, 8, 8, 18, 18, 32 (magic numbers). This repetition of similar valence shell electronic configurations of the elements in a group is the cause of periodicity of the elements in the periodic table. For example, all the members of alkali metal family, which have definite gaps of atomic numbers (8, 8, 18, 18, 32), have similar physical and chemical properties because they have similar valence shell electronic configuration i.e. ns¹ as shown below:

In a similar manner, all the halogens i.e. elements of group 17 have similar outer electronic configuration (ns² np⁵) and as such possess similar properties.

Long form of the periodic table (also called Bohr's table) is constructed on the basis of repeating configuration of the atoms when they are arranged in order of increasing atomic numbers.

Salient structural features:  This table has the following features:
(i) All the elements have been arranged in the increasing order of atomic numbers.
(ii) Elements having similar electronic configurations have been placed together at one place and those having different electronic configurations have been placed at different places in the periodic table.
(iii) Table consists of seven horizontal rows called periods and each period starts with different principal quantum number. The first, second and third periods are called short periods. They contain 2.8, 8 elements respectively. Fourth, fifth and sixth periods are called long periods. They contain 18, 18 and 32 elements respectively. Seventh period is incomplete and contains 19 elements (2+14+3).
(iv) The series of 14 elements just after lanthanum (₅₇La) i.e. from cerium (₅₈Ce) to lutetium (₇₁Lu) is collectively called lanthanoides and are placed at the bottom of the periodic table to avoid undue side-wise expansion of the periodic table.
(v) The series of 14 elements just after actinium (₈₉Ac) i.e. from thorium (₉₀Th) to lawrencium (₁₀₃Lr) is collectively called actinoides and are placed at the bottom of the periodic table to avoid undue side-wise expansion of the periodic table.
(vi) Lanthanoids and actinoids are also called Inner Transition elements.
(vii) The elements beyond uranium (atomic number 93 to 105) which have been prepared artificially by nuclear reactions are also called Transuranic elements or man made elements.
(viii) Table consits of eighteen vertical columns called groups or families.
(ix) Elements in each column have the greatest electronic similarities on the basis of their electronic configuration, the elements have been divided into s-block, p-block, d-block and f-block.
(x) According to 1UPAC system, groups are numbered as 1-18:
(i) Alkali metals (Group 1)
(ii) Alkaline earth metals (Group 2)
(iii) In between groups 2 to 13 there are ten columns which constitute group 3 to 12. The elements present in these groups are known as Transition elements.
(iv) Pincogens (group 15) (v) Chalcogens (Group 16) (vii) Halogens (group 17) (viii) Aerogens, Noble gases, inert or rare gases (group 18).
(ix) The elements constituting groups 1, 2 and group 13 to 18 are called normal or representative elements.

Long form periodic table is regarded as better than the Mendeleev’s table due to the following reasons:

1. Position of hydrogen. Hydrogen shows similarities with both metals(alkali metals) and non-metals (halogens) but it is placed with alkali metals.
2. Position of helium. It is kept along with p-block elements. 
3. Position of lanthanoids and actinoids. They are kept outside the main body of the periodic table. 

In 1997, IUPAC recommended the approved official names for elements with atomic numbers 104 to 109. The names for elements with atomic number 110 and beyond have yet to be officially announced. This nomenclature is based on the Latin words for the atomic number of the element.
(i) The names are derived directly from the atomic number using the numerical roots for 0 and numbers 1to 9.

(ii) The roots are strung together in the order of digits which make up the atomic number and “ium” is added at the end.
For example let us write the IUPAC name and symbol for the elements having atomic number 112. The roots for 1, 1, 2 are un, un and bi respectively. Hence the name and the symbol respectively are ununbium and Hub.

Element with atomic number 120:
The roots for 1 is un
The roots for 2 is bi
The roots for 0 is nil
Hence the IUPAC name for the element with atomic number 120 is unbinilium and its symbol is Ubn.

The long form of periodic table has been divided into four blocks. These are called:
(i) s-block (ii) p-block
(iii) d-block , (iv) f-block.
Division of elements into s, p, d and f-blocks: This division is based upon the name of the orbital which receives the last electron.
1. s-block elements. The elements in which the last electron enters the s-orbital of the outermost enthalpy level are called .s-block elements. The general electronic configuration of s-block elements is ns^1-2 where n stands for the outermost shell. In s-block, the alkali metals of group 1 and alkaline earth metals of group 2 are included.
2. p-block elements: The elements in which the last electron enters the p-orbital of the outermost enthalpy level are called p-block elements. The general electronic configuration of p-block elements in ns²np^1-6 where n stands for the outermost shell. The elements of Groups 13, 14, 15, 16, 17, 18 having 3, 4, 5, 6, 7 and 8 electrons respectively in the outermost enthalpy levels constitute p-block elements.

3. d-block elements:The elements in which the last electron enters the d-orbital of the penultimate enthalpy level (last but one shell) are called d-block elements. General electronic configuration of d-block elements is (n - 1) d^{1-10 n}s^1-2 where n represents the valence shell, d-block elements have three complete series of ten elements in each whereas the fourth series is incomplete.
Note, Exception is ₄₆Pd whose configuration is 4d¹⁰ 55⁰:
4. f-block elements:The elements in which the last electron enters the f-orbital of the anti-penultimate (third to the outermost shell) enthalpy level are called f-block elements. The general electronic configuration of f-block elements is (n -2) f^1-14 (n-1) d^0-1 ns²where n represents the outermost shell. Such type of elements have two series. First is 4f series which is also called lanthanoides series whereas second is 5f series which is also described as actinoide series. These two series are placed separately at the bottom of the periodic table. The elements included in these two series are called inner transition elements.

These are:
(i) By classifying the elements into s, p, d and f blocks, we can study their physical and chemical properties in a very organised and systematic manner.
(ii) By knowing the block of the element, it is very easy to study the general behaviour and important properties of the elements. For example if the element belongs to s-block, it means it must have 1 or 2 electrons in outermost shell. Hence it must have low value of IE, strongly metallic and electropositive character and form cations easily.
(iii) This classification of elements also explains why properties of d-block elements vary slightly within the same period. This is because inner (n -1) d orbital is being progressively filled up.
(iv) In f-block elements there is progressively filling of (n - 2) f orbitals and this classification also explains why f-block elements are nearly identical in their chemical properties.

General outer electronic configuration of s-, p-,d-and f- block elements:
(i) s- Block elements:
ns^1-2 where n = 2 to 7
(ii) p-Block elements: ns² np^1-6 where n = 2 to 6
(iii) d-Block elements:
(n - 1)d^1-10 ns^0-2 where n = 4 to 7
(iv) f-block elements:
(n - 2)f^0-14 (n -1)d^0-1 ns² where n = 6 to 7.

i) Lawrencium (Lr) with Z = 103  is an actinide which involves the filling of 5f orbital andBerkelium (Bk) with Z = 97  is an actinide which involves the filling of 5f orbital (both of them are f block elements).

ii) Seaborgium (Sg) with Z = 106 is a d- block element present in 7th period &  6th column of periodic table.

We know that the gaps of atomic numbers in a group are 8, 8, 18, 18, 32. Hence the element which proceeds the element with Z = 114 in the same group must have an atomic number equal to 114 - 32 = 82. This represents lead (Pb) which is present in the 6th period and belongs to group 14 of the p-block. This means the element with Z = 114 must also belong to group 14 (second p-block element) of 7th period. Thus the location of the element with Z = 114 in the periodic table is:
period = 7th, Block = p-Block, Group = 14.

The transition elements are those elements which represent a change of state from the highly metallic elements of s-block to the non-metals of p-block i.e. they lie between the elements of s-and p-blocks. There are four transition series each marking the filling of 3d, 4d, 5d and 6d orbitals respectively as:
(i) First transition series (3d). Scandium (₂₁Sc) to Zinc (₃₀Zn).
This series lies in the fourth period of the periodic table and consists of 10 elements.
General electronic configuration is 3d^1-104s^1-2.
(ii) Second transition series (4d). Ytterium (₃₉Y) to Cadmium (₄₈Cd).
This series lies in the fifth period of the periodic table and consists of 10 elements.
General electronic configuration is 4d^1-10 5s^0-2
(iii) Third transition series (5d). Lanthanum (₅₇La), Hafinium ₇₂Hf) to Mercury (₈₀Hg).
This series lies in the sixth period of the periodic table and consists of 10 elements.
General electronic configuration is 5d^1-10 6s^1-2.
(iv) Fourth transition series (6d). Actinium (₈₉Ac), Kurchatovium (₁₀₄Ku) to atoic number 112 have been synthesised. This series lies in the seventh period of the periodic table and consits of 10 elements.
Hence d-block elements consist of 40 elements.

The important charateristics of transition elements:
(i) These elements are divided into four transition series each marking the filling of 3d, 4d, 5d and 6d orbitals respectively.
(ii) They are hard and brittle metals having high melting and boiling points.
(iii) They are good conductors of heat and electricity.
(iv) They show variable oxidation states.
(v) They are generally paramagnetic due to the presence of unpaired electrons.
(vi) They form coloured ions.
(vii) These elements have a strong tendency to form complexes.
(viii) Most of the transition elements (Mn, Co, Ni, Cr) and their compounds are used as catalysts.

The number of elements in each block are:
(i) s-block: There are 13 elements in s-block. Except the first period (having only one element), each other period of the block has 2 elements.
(ii) p-block: There are 31 elements. With the exception of first period (having only one element), each other period has got 6 elements.
(iii) d-block: This block has 40 elements. Each period i.e. fourth, fifth, sixth and seventh of the block has got 10 elements.
(iv) f-block: This block has 28 elements in the form of 2 horizontal rows of 14 elements each.
Lanthanoides = 14 elements.
Actinoides = 14 elements

In the long form of the periodic table, each period starts with the filling up of a new principal shell. Clearly, the 4th period will begin with filling of 4th shell i.e. N shell. Thus, fourth period starts with the filling up of 4s orbtial (n = 4). After filling 4s orbital, the filling of 3rd and then 4p takes place. It is so because enthalpy of 3d sub level in between 4s and 4p sub levels. As,
4s = 2 electrons
3d = 10 electrons             
4p = 6 electrons
        _______________
        18 electrons
Therefore, 18 elements (not eight) are present in the 4th period. The 4d and 4f sublevels are higher in enthalpy than 5s and hence are filled up in the next periods.

 

In lanthanoids (₅₈Ce to ₇₁Lu) 4f-subshell is being filled and in actinides (₉₀Th to₁₀₃Lr), 5f-subshell is being filled. All the 14 lanthanoids and all the 14 actinoids have similar properties. They have been placed at the bottom of the periodic table in two separate horizontal rows because of the following reasons:
(i) to save space,
(ii) to keep the elements having similar properties in a single column,
(iii) to avoid undue sidewise expansion of the periodic table.

The period group or block of an element can be predicted from the electronic configuration of the elements.
(i) Period of the element corresponds to the principal qauantum number of the valence shell.
(ii) Block of the element corresponds to the sub-shell which receives the last electron.
(iii) Group is predicted from the number of electrons in the outermost (n) or penultimate (n -1) shell as follows:
For s-block elements, group number is equal to the number of valence electrons (ns electrons).
For p-block elements, group number is equal to 10+ number of valence electrons (ns and np).
For d-block elements, group number is equal to the sum of the numbers of (n - 1)d and ns electrons.
For f-block elements group number is 3.
 

(i) 3s² 3p⁵ : p-block
(ii) 3d¹⁰ 4s² : d- block
(iii) 3p⁶ 4s² : s-block
(iv) 6s² 4f⁰ : s-block
(v) 4s¹ 3d⁵ : d-block

(i) The electronic configuration of the element is 3s²3p⁴. Thus the element belongs to p-block, third period ( ∵ n = 3) and group 16.
(ii) The electronic configuration of the element is 3d²As². Thus the element belongs to d-block, fourth period (∵ n = 4) and group 4.
(iii) The electronic configuration of the element is 4f⁷ 5d¹6s². Thus the element belongs to f-block, sixth period (n = 6) and group 3.

(i) The third alkali metal :The atomic number is 19 and the electronic configuration is, 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
(ii) The second transition element: The atomic number is 22 and the electronic configuration is, 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d².
(iii) The fourth noble gas: The atomic number is 36 and electronic configuration is, 1s²2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶.
(iv) The first halogen: The atomic number is 9 and electronic configuration is, 1s² 2s²2p⁵.

(i) A(Atomic number 7):
Electronic configuration: 1s² 2s² 2p³
Period = 2nd
Group = 15
Block  = p-block
(ii) B(Atomic number 19):
Electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
Period = 4th
Group = one (1)
Block = s-block

(iii) C (Atomic number 29):
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
Period = 4th
Group = Eleven (11)
Block = J-block

(iv) D (Atomic number 36)
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
Period = 4th
Group = Eighteen (18)
Block = p-block

The element belonging to 3rd period and with four valence electrons:
(i) belongs to Group 14
(ii) is silicon (Si) with atomic number 14.
(iii) The electronic configuration of ₁₄Si is

(i) The electronic configuration of element with Z = 117 can be represented as

This element belongs to Group 17 (halogen family).
(ii) The electronic configuration of element with Z = 120 can be represented as

This element belongs to Group 2 (alkaline earth metals).

Metals: Metals are generally solids at room temperature (mercury is an exception), lustrous, malleable, ductile, high melting and boiling points and good conductors of heat and electricity. Metals comprise more than 75% of all known elements and appear on the left side of the perodic table. For example iron, copper, aluminium, lead etc. The elements become more metallic as we go down a group.
Non-metals: Non-metals are usually solids or gases at room temperature, non-lustrous, low melting and boiling points and poor conductors of heat and electricity. They are located at the top right hand side of the periodic table. For example oxygen, hydrogen, carbon, sulphur, etc.

The non-metallic character increases as we go from left to right across the periodic table.

Metalloids or semimetals: These are the elements which show common characteristic of metals and non-metals. For example, germanium, silicon, arsenic, antimony and tellurium.

Periodic properties are those properties which get repeated after a regular interval in a regular manner. The important periodic properties are:
1. Atomic and ionic radius
2. Ionization enthalpy
3. Electron gain enthalpy
4. Electronegativity
5. Melting points
6. Density
7. Valency.

Assuming atom to be spherical, the distance between the nucleus and the outermost orbit of the atom is called the atomic radius of the element. However, it can not be determined precisely due to the following reasons:
(i) Atom is too small to be isolated.
(ii) According to the probability picture of electrons, an atom does not have a well-defined boundary.
(iii) The probability distribution of an atom is also affected by the presence of other atoms in its neighbourhood.
(iv)The size of the atom also changes from one bonding state to another.
Thus, it is not possible to determine the absolute value of atomic radius practically. In other words, atomic radius has no meaning.

where 198 pm = Internuclear distance between two chlorine atoms (Cl—Cl) in a molecule of Cl₂.
In case of heteronuclear molecules (containing different atoms) such as HCl, the covalent radius may be defined as: “The difference of the distance between the nuclei of two bonded atoms and covalent radius of one of the atoms in a heteronuclear molecule.”
Covalent radius = [Internuclear distance of two bonded atoms] -[Covalent radius of one of the two atoms in heteronuclear molecules]
For example,
Covalent radius of H-atom (in HCl)
= Internuclear distance of HCl - covalent radius of Cl atom
= 136 pm - 99 pm = 37 pm.

Variation of atomic radii in a group. As we move down a group in the periodic table, the atomic radii go on increasing.
Reasons :
(i) Down a group, the nuclear charge goes on increasing. As a result, atomic radii must decrease.

(ii) As we move down a group, a new enthalpy shell is added at each succeeding element though the number of electrons in the valence shell the remain the same. As a result, atomic radii must increase.
The effect of the progressive addition of a new shell outweighs the effect of increased nuclear charge. Hence atomic radii of elements increase with the increase in atomic number as we move from top to bottom down a group.

Variation of atomic radii in a period. As we move from left to right in a period, atomic radii go on decreasing.

Reason: As we move from left to right in a particular period, the atomic number i.e. nuclear charge increases by one unit in each succeeding element but the corresponding addition of electron takes place in the same enthalpy shell. As a result, electrons are pulled little closer to the nucleus thereby making each individual shell smaller and smaller.

Vander Waal's radius is defined as one half of the distance between the nuclei of two adjacent atoms belonging to two neighbouring

molecules of an element in the solid state, vander Waal’s radii are determined by X-ray diffraction method.
Comparison of covalent radius and vander Waal’s radius.
vander Waal’s radii are always larger than covalent radii. This is because for the formation of covalent bond, atoms move to come closer to each other due to the overlapping of orbitals.
On the other hand, vander Waal’s forces operating between atoms are weak, therefore, adjacent atoms belonging to two neighbouring molecules are at relatively larger distances. Clearly, one half of internuclear distance between adjacent atoms belonging to two neighbouring molecules will be more than one half of the distance between the nuclei of two like atoms forming a single covalent bond. Consequently, vander Waals’ radii are always larger than covalent radii. For example, covalent radius of chlorine is 99 pm whereas vander Waal’s radius of chlorine is 180 pm.

Anion (or negative ion) is formed by the gain of one or more electrons by the gaseous atom. In the anion, nuclear charge is the same as that in parent atom but the number of electrons has increased. As a result, the same nuclear charge acts on relatively larger number of electrons. Thus, effective nuclear charge per electron is decreased which causes the electron cloud to expand. Therefore, anion is always larger than its parent neutral atom.
For example,
Atomic radius of Cl = 99 pm (nuclear charge = - 17 ;e = 17)
Atomic radius of Cl^- = 131 pm (nuclear charge = 17 ; e = 18)

Atomic radius means the size of the atom i.e. the distance from the centre of the nucleus of the atom to the outermost shell of electrons. In the case of non-metals, atomic radius is called covalent radius. It is defined as one-half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond. In the case of metals, atomic radius is called metallic radius. It is defined as one-half of the distance between the two adjacent atoms in the crystal lattice.
Ionic radius means the size of the ions. It may be defined as the effective distance from the nucleus of the ion to the point to which it has an influence in the ionic bond. The size of the cation is always smaller as compared to that of the parent atom. The size of the anion is always larger as compared to that of the parent atom.

Isoelectronic ions are those which have a same number of electrons and hence same electronic configuration but the different magnitude of nuclear charge. A set of species (cations or anions) having the same number of electrons is known as iso-electronic series, e.g. N^3-, O^2-, F^-, Na⁺, Mg⁺, Al³⁺ are isoelectronic species.
Within the series of isoelectronic ions, as the atomic number or nuclear charge increases, the attractive force between the nucleus and a same number of electrons also increases. This results in the decrease of ionic radius. 

For example,

 Ion                          =N^3-    O^2-    F^-     Na⁺    Mg⁺       Al³⁺

No. of electrons   =  10      10     10      10       10      10

At. No                      =  7        8      9       11       12      13

Since, the size of the iso-electronic ion decreases with the increase in atomic number, the order of decreasing size is

N^3- > O^2- > F^- > Na⁺ > Mg²⁺ > Al³⁺

Ions of different elements which have the same number of electrons but different magnitude of the nuclear charge are called isoelectronic ions.

(i) Since F^- has 10 electrons, therefore N^3- (7 + 3), O^2- (8 + 2), Ne (10 + 0),  Na⁺ (11 - 1), Mg²⁺ (12 - 2): Al³⁺ (13 - 3) are isoelectronic, as each one of these contains 10 electrons.

(ii) As Ar has 18 electrons, therefore P^3- (15 + 3), S^2- (16 + 2), Cl^- (17 + 1), K⁺ (19 - 1), Ca²⁺ (20 - 2) are isoelectronic with it as cach one of each contains 18 electrons.

(iii) As Mg²⁺ has 10 electrons, therefore N³⁺ F , Ne, Na⁺, Al³⁺ are isoelectric with it as each one of these contains 10 electrons.

(iv) As Rb⁺ has 36 electron, therefore Br^- (35 + 1), Kr (36 + 0) and Sr²⁺ (38 -2) are isoelectronic with it as each one of these contains 36 electrons.

(a) Each one of these ions contains 10 electrons and hence all are isoelectronic ions.
(b The ionic radii of isoelectronic species increase with a decrease in the magnitudes of nuclear charge. The arrangement of the given species in order of their increasing nuclear charge is as follows:

N^3-< O^-2<F^-<Na⁺<Mg²⁺<Al³⁺

(i) Na, Na⁺ pair. The size of Na atom is greater because the cation i.e. Na⁺ is formed by the loss of one electron from the atom.
(ii) Br, Br^- pair. The size of Br^- ion is greater as anion is formed by the gain of one electron by bromine atom.
(iii) Br, I pair. The size of 1 is greater because it comes after Br in the halogen family and atomic size increases down the group.

(i) Na⁺: Mg²⁺ (10 electrons)
(ii) Cl^-: Ar  (18 electrons)
(iii) Ca²⁺: K⁺ (18 electrons)
(iv) Rb⁺: Kr (36 electrons)

Atomic radii decrease across a period, cations are smaller than their parent atoms. Among isoelectronic ions, the one with the larger positive nuclear charge will have a smaller radius.
Hence the largest species is Mg and the smallest one is Al³⁺ .

(i) K or K⁺: K atom has large size than K⁺ ion because the size of cation (or radius) is always smaller as compared to the parent atom.
(ii) Br or Br^- : Br^- ion has large size because it is an anion and it has always large size as compared to the parent atom.
(iii) O^2- or F^-: Both the anions are isoelectronic species. Since the nuclear charge of O^2-ion is small, its size is, therefore, more than that of F^-ion.
(iv) Li⁺ or Na⁺: Both the cations are formed from the atoms which are present in group 1. As Na atom is larger than Li atom, therefore, Na⁺ ion is also bigger in size than Li⁺ion.
(v) P or As: Both these atoms belong to group 15 of the periodic table. Since As is placed after P, therefore its size is large because atomic size increases down the group.
(vi) Na⁺ or Mg²⁺: Both the cations are isoelectronic in nature with 10 electrons each. Since the nuclear charge of Na⁺ ion is more than that of Mg²⁺ ion, its size is, therefore, large.

The main factors are:
(i) Size of the atom: As the size of the atom increases, the distance between the nucleus and outermost electrons increases. So the attraction between the nucleus and the outermost electron decreases and the enthalpy required to remove the electron also decreases. Thus, with an increase in atomic size, the ionisation enthalpy decreases.

(ii) Nuclear charge: With the increase in nuclear charge, the force of attraction between the nucleus and the valence electrons increases. Consequently, more enthalpy is required to remove a valence electron. Hence ionisation enthalpy increases with the increase in nuclear charge.

(iii) Screening effect of inner shells: The electron shells present in between the outermost shell and the nucleus act as screens. These shells reduce the attraction of nuclear charge over the outermost electrons. This reduction in force of attraction by the shells present in between the nucleus and valence electrons is called screening or shielding effect. As the number of inner shells increases, the screening effect also increases. As the screening effect increases, the ionisation enthalpy decreases.

(iv) Penetration effect of electrons: It is well known that s-electrons are more penetrating towards the nucleus than p-electrons and the penetration power decreases in a given shell (same value of n) in the order s >p >d>f.
Now if the penetration of the electron is more, it will be closer to the nucleus and held firmly. Hence ionisation enthalpy will be high. Thus, for the same shell, it is easier to remove p-electron in comparison to s-electron. Therefore, for the same value of n, ionisation enthalpy decreases in the order s > p > d > f

(v) Stability of the electronic configuration: More enthalpy is required to remove an electron from an element having configuration in which all the orbitals of the same sublevel are exactly half filled (p³ or d⁵ : stable configuration) or completely filled (p⁶or d¹⁰: stable configuration). Thus, more stable the electronic arrangement, greater is the ionisation enthalpy.

As we move down a group. the ionisation enthalpy goes on decreasing.
Reason. As we move down the group.
(i) nuclear charge regularly increases.
(ii) The size of the atom regularly increases due to the addition of a new shell every time and
(iii) shielding effect also go on increasing due to the increase in the number of electrons in the inner shells. 
The effect of the increase in atomic size and shielding effect is much more than the effect of the increase in nuclear charge. As a result, an electron becomes less and less firmly held to the nucleus as we move down the group. Hence there is the gradual decrease in the value of ionisation enthalpy on moving from top to bottom in a group.

Within the main group elements, the ionization enthalpy decreases regularly as we move down the group due to the following two factors:
(i) Atomic size: As we move down the group, the atomic size of the atom regularly increases due to the addition of one new principal energy shell at each succeeding element.
(ii) Screening effect: With the addition of new shells, the number of inner electron shells, which shield the valence electron increases. In other words, shielding effect or screening effect increases.
As a result of both these factors, the force of attraction of the nucleus for the valence electrons decreases. Thus, the electron becomes less and less firmly held to the nucleus as we move down the group. Hence there is a regular decrease in the value of ionization enthalpy on moving down the group.

In general, ionisation enthalpy increases as we move from left to right across a period.
Reason:
We know that on moving across the period from left to right.

(i) the nuclear charge increases,
(ii) progressive addition of electrons occurs in the same enthalpy level and
(iii) atomic size decreases.

Thus, due to the gradual increase in nuclear charge and the simultaneous decrease in atomic size, the valence electrons are more and more tightly held by the nucleus. Therefore, more and more enthalpy is needed to remove the electron and hence ionisation enthalpy keeps on increasing.

The electronic configuration of Be (Z = 4) and B(Z = 5) are,

(i) The electronic configuration of Be in which all the occupied orbitals are fully filled is more stable than boron.
(ii) It is relatively easier to remove electron from 2p subshell in boron than removal of electrons from 2s subshell in Be as 2p orbital is less penetrating than 2s orbital. Therefore, 2p electron of B is less firmly held by the nucleus and so, it is relatively easily removed.

This is due to the fact that electronic configuration of nitrogen in which all 2p orbitals,

are half filled, is much more stable than that of oxygen,

in which all the 2p orbitals are neither half filled nor fully filled. Therefore. IE₁ of N is more than that of oxygen.