Trends in Electronegativity

 (i).  Down The Group

Down the group atomic size increases and hence electronegativity decreases.

(ii).  Along The Period

Along the period, atomic size decreases and hence electronegativity increases.

IUPAC Nomenclature of Elements with Z > 100

IUPAC in 1997 suggested the rules for naming the elements with Z > 100. The main points of this nomenclature are;

1.  The names are derived directly from atomic numbers using numeral roots for 0 to 9 and adding suffix-ium. The roots for number 0 to 9 are;

Digit	0	1	2	3	4	5	6	7	8	9
Root	nil	un	bi	tri	quad	pent	hex	sept	oct	enn
Abbreviation	n	u	b	t	q	p	h	s	o	e

2.  The names are shortened if there is repetition of a word e.g. bi ium and tri ium are written as buim & trium etc.

3.  The symbol of element is obtained from first letters of the roots of number which make up atomic number of element. E.g. the name of element with atomic number 102 would be unnilbium. Thus its symbol is Unb (u for un, n for nil and b for bi).

Z

NAME	SYMBOL	Z	NAME	SYMBOL
101	Unnilunium	Unu	111	Unununium	Uuu
102	Unnilbium	Unb	112	Ununbium	Uub
103	Unniltrium	Unt	113	Ununtrium	Uut
104	Unnilquadium	Unq	114	Ununquadium	Uuq
105	Unnilpentuim	Unp	115	Ununpentium	Uup
106	Unnilhexium	Unh	116	Ununhexium	Uuh
107	Unnilseptium	Uns	117	Ununseptium	Uus
108	Unniloctium	Uno	118	Ununoctium	Uuo
109	Unnilennium	Une	119	Ununennium	Uue
110	Ununnilium	Uun	120	Unbinilium	Ubn

Factor’s Affecting Electronegativity

The various factors which influence electronegativity are
1. Atomic Radius
Electronegativity decreases with increase in size of atom.

2. Nuclear Charge
Electronegativity increases with increasing nuclear charge.

3. Screening Effect
Increase in number of inner electrons tends to decrease the electronegativity due to screening effect.

ElectroNegativity

“The tendency of an atom to attract electrons to itself when combined in a compound is called electronegativity of the atom of an element.”

Trends in Electron Affinity

VARIATION IN A PERIOD

On moving along the period, the atomic size decreases and hence electron affinity increases due to greater attraction for incoming electron. However, some irregularities are observed in a general trend. These are mainly due to the stable configurations of certain elements. Some important features of electron affinity are;

¤  Halogens have highest Electron affinity 

The electron affinity of halogens are the highest in their respective periods. This is due to small size and greater effective nuclear charge. They also need only electron to attain noble gas configurations. Thus, they have maximum tendency to accept an additional electron.

¤  Noble gases have zero Electron affinity

Noble gases have stable electronic configuration and hence they have no tendency to accept an electron. So, they have zero electron affinity.

¤  Electron affinity of Be, Mg are almost zero

Electron affinities of Be, Mg etc. are almost zero. This is because of the fact that their electronic configurations are stable because all the electrons are paired.

¤  Electron affinity of N & P are extremely low

Electron affinity of group 15 i.e. N, P etc. have electron affinity quite low, due to stable half -filled electronic configuration. 

VARIATION DOWN THE GROUP

On moving down the group the atomic size decreases and hence the incoming electron feels less attraction. Thus, electron affinity decreases down the group.

However in halogens, electron affinity of fluorine is smaller than chlorine. According to size Flourine should have higher value of electron affinity than chlorine. This is due to its small size. As a result of small size the repulsions among electrons in valence shell are relatively larger as compared to chlorine. This means that incoming electron in fluorine atom finds less attraction than that in chlorine atom. Consequently, electron affinity of chlorine is higher than flourine.

Factors Affecting Electron Affinity

The various factors which influence the electron affinity can be explained under following heads:-

1.  NUCLEAR CHARGE

Greater the nuclear charge, greater will be the attraction for the incoming electron and as a result larger will be the value of electron affinity.

2.  ATOMIC SIZE

Larger the size of an atom, larger will be the distance between the nucleus and the incoming electron. Thus, smaller will be force of attraction felt by incoming electron and hence smaller will be the value of Electron affinity.

3.  ELECTRONIC CONFIGURATION

Stable the configuration of an atom, lesser will be its tendency to accept an electron and hence lower the value of its electron affinity.

Electron Affinity

Electron affinity may be defined as, “The amount of energy released when an isolated gaseous atom accepts an electron to form monovalent gaseous anion.”

The process can be expressed as;

X (g)     +       e^–  −−−−−−>   X^– (g)   +   Energy (E. A.)

The value of Electron Affinity (E.A) measures the tendency of an atom to accept an electron.

SUCCESSIVE ELECTRON AFFINITIES

When the first electron is added to the gaseous atom, it forms a Uninegative ion and the energy released is called First Electron Affinity (E.A₁).Now if an electron is added to the uninegative ion, it experiences a repulsive force from an anion. As a result, the energy has to be supplied to overcome the repulsive force. Thus, in order to add the second electron energy is required rather than released. Therefore, the value of second Electron affinity is negative. Similarly, addition of 3^rd, 4^th electrons etc. also requires energy. Thus, the values of successive electron affinities (EA₂, EA₃ ……) are negative.

Trend in Ionisation Energy

Variation along Period
As we move along the period, the atomic size decreases due to increase in nuclear charge. Hence the ionisation energy also increases along the period. Thus alkali metals have smallest Ionisation energy.

Variation down the group
On descending a group, the atomic size increases and the nuclear attraction for the outer electrons decreases. Hence, it becomes easier to remove the outer electrons in a group. Thus, Ionisation energy decreases on descending a group. Thus in group 1, Cs has lowest Ionisation energy in the group as well as in periodic table.

Factors Influencing Ionisation Energy

The various factors which influence the Ionisation energy are;

¤    Size of atom

The Ionisation energy depends upon the size of atom, because with increase in size, the distance between nucleus and valence electrons increases and hence force of attraction between nucleus and valence electron decreases. As a result of which the valence electrons are loosely held and smaller energy is required to remove an electron. Thus Ionisation energy decreases with increase in size and increase with decrease in size.

¤    Nuclear charge

As the nuclear charge increases, the force of attraction between nucleus and valence electrons increases and hence it is difficult to remove an electron from valence shell. Thus with increase in nuclear charge, the Ionisation energy increases.

¤    Screening effect or Shielding effect

The inner electrons present in shells between nucleus and valence shell reduce the attraction between nucleus and the outermost electrons. This shielding effect depends upon the number of inner electrons. Larger the number of electrons in the inner shells, the greater is the screening effect. More the Shielding effect, easier will be to remove an electron and hence lesser will be ionisation energy.

¤    Penetration effect

s – electrons are closer to the nucleus than p-electrons, which is closer than d – electron and these inturn are closer than f-electrons of the same principle energy level. Hence s – electrons experience more attraction from the nucleus than p, d and f – electrons. Thus Ionisation energy to remove an electron from a given energy level decrease in order s > p > d > f.

¤    Electronic Configuration

Electronic Configuration plays vital role in determining the value of Ionisation energy. Atoms having stable configuration (i.e. fully filled or half filled) has least tendency to lose electron and hence have high value of I. E.

Successive Ionisation Energies

The energy required to remove an electron from a neutral gaseous atom is called First Ionisation Energy and is represented by I.E₁,

M (g)   +    I.E₁  → M⁺ (g)    +    e ^–

Once the first electron has been removed from the gaseous atom, it is possible to remove second electron from unipositive ion and then third electron from dipositive ion and so on. The energies required to remove first electron, second electron and third electron from neutral gaseous atom. Unipositive ion and dipositive ion are known as First Ionisation Energy (I.E₁), Second Ionisation Energy (I.E₂) and Third Ionisation energy (I.E₃)  respectively.

M ⁺   +   I.E₂   →  M²⁺    +    e ^–

M²⁺ +    I.E₃   →  M³⁺       +   e ^–

The second, third Ionisation energies and so on are collectively known as Successive Ionisation Energies

It may be noted that;

I.E₃   >    I.E₂   >    I.E₁

The order can be explained as follows;

After the removal of first electron, atom changes into Unipositive ion. In this ion the number of electrons decreases by one, but nuclear charge remains same as in neutral atom. As a result the force of attraction per electron increases and hence more energy is required to remove an electron. Hence value of I.E₂ is higher than I.E_1. Similarly removal of third electron requires more energy because the force of attraction per electron has further increased; hence I.E₃ is more than I.E_2.

Ionisation Energy

Ionisation energy may be defined as, “the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom.”

M (g)   +   Energy (I.E)   ----->    M⁺ (g)    +    e ^–

Ionisation energy is also called Ionisation Potential. It is most expressed in terms of electron volts (e.v) (1 e.v = 1.603 ´ 10 ^{– 19} J).

IsoElectronic Ions

The ions having same number of electrons but different magnitude of nuclear charge are called Iso-Electron Species or Ions.

Variation of size among Iso electronic ions

With the series of Iso-electronic ions, as the nuclear charge increases, the attractive forces between electrons and nucleus also increases. This results in the decrease of Ionic radius. In other words, size of Iso electronic ions decreases with increase in the magnitude of nuclear charge.

For example, N^3–, O^2–, F^–, Na⁺, Mg ²⁺, Al³⁺ are Iso electronic ions and have 10 electrons each. The order of their size is

Al ³⁺ <  Mg ²⁺  <  Na⁺ <  F^– <   O^2–  <  N^{3 –}

Al³⁺ have highest nuclear charge of 13 units and has smallest size, whereas N^3– has lowest nuclear charges and has largest size.      

Ionic Radius

Ionic radius may be defined as the, “effective distance from the nucleus of the ion to the point up to which it has an influence in the Ionic bond.”

The equilibrium distance between the nuclei of the two adjacent ions can be determined by X-ray analysis. Knowing the ionic radius of one of the ion, the ionic radius of other can be calculated.

The study of Ionic radii leads to two very important generalizations;

(i).  The size of cation is smaller as compared to that of present atom.

(ii). The size of anion is larger as compared to that of present atom.

(I).  RADIUS OF CATION IS SMALLER THAN PRESENT ATOM

A cation is formed, by loss of one or more electrons from gaseous atom. Thus, the nuclear charge (i.e. number of protons) remains same as that in parent atom, but the number of electrons becomes less in valence shell. As a result, the force of attraction on these electrons by nuclear charge increases due to increase in effective nuclear charge per electron. This causes shrinking of shell and hence size decreases.

(II). RADIUS OF ANION IS LARGER THAN PARENT ATOM

An Anion is formed by gain of one or more electrons by the gaseous atom. Thus the nuclear charge remains same (i.e. number of protons), but the number of electrons increases in valence shell. As a result, the force of attraction on these electrons decreases due to decrease in effective nuclear charge per electron. This causes increase in size.

Trend in Atomic Radii

Variation along Period

In a period, as we move from left to right the atomic radius decreases.

Reason

This is because with the increase in atomic number, (i.e., on moving from left to right along a period, the atomic number increases by one unit at each interval) the nuclear charge increases while the number of shells remain same. The electrons are attracted more strongly towards nucleus. The result is decrease in atomic radius. Thus in a given period, the alkali metals has the largest atomic radii and halogens, the smallest atomic radii.

Noble gases which are at the end of each period, are expected to have smallest radii, but they have comparatively larger atomic radii. This is due to the reason that in case of inert gases, the outer shell is complete and it has hence maximum electronic repulsion. Moreover in case of inert gases, the atomic size is expressed in terms of Vanderwaal’s radius because, noble gases do not form covalent bonds  while other form covalent bonds and hence their atomic radii is expressed in covalent radius. Thus as, Vanderwaal’s radius is greater than covalent radius, the atomic size of noble gases is larger than other elements of same period.

Variation down the group

As we move down the group, the atomic radius of the corresponding elements increases.

Reason

This is due to the reason that as we move down the group, the number of shells increases, though the nuclear charge also increases. The effect of increase in number of shells is more than the increase in nuclear charge, therefore, the effective nuclear charge decreases. Consequently, the distance of the outermost electron from the nucleus increases on going down the group.

On Short,

Atomic radii increase down the group.

Atomic radii decrease along the period.

Atomic Radii

Assuming atoms to be spheres, the atomic radii may be defined as, “The distance from the centre of nucleus of the atom to the outermost shell of electrons.” However, it is not possible to find precisely the radius of the atoms because of the following reasons;

(i).        Atom is too small to be isolated.

(ii).       Due to wave nature of electron, it is not possible to measure the exact distance from nucleus.

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

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

Thus, it is not possible to measures the absolute value of the atomic radius of an element. However, it   may expressed in three different forms depending upon the nature of bonding in the atoms. These are:

¤    Covalent radii

¤    Vanderwaal’s radii

¤    Metallic radii

¤    COVALENT RADIUS

 “The covalent radius of an atom is equal to half the internuclear distance between two identical atoms that are joined by a covalent bond”.

The covalent bond should essentially be a single covalent bond. If it is double or triple bond, the value of covalent radius changes.

The distance between the centres of the two bonded atoms (inter nuclear distance) can be determined by X-ray diffraction methods, Thus,

Covalent radius = Inter Nuclear distance in bonded atoms/2

Covalent radius is always shorter than the actual radius because overlapping of orbitals involved in Covalent bond formation decreases the internuclear distance.

¤   VANDERWAAL’S RADIUS

Vanderwaal’s radius may be defined as, “one half of the inter–nuclear distance between two similar adjacent atoms belonging to two neighbouring molecules of the same substance in the solid state.

The Vanderwaal’s forces of attraction are quite weak forces. Their magnitude is small in the gaseous as well as in liquid state of substance. Therefore, Vanderwaal’s radius is determined in solid state when their magnitude is maximum. The Vanderwaal’s radius is always more than the corresponding value of the covalent radius e.g. Vanderwaal’s radius of chlorine is 180 pm while its covalent radius is 99 pm.

¤   METALLIC RADIUS

Metallic radius may be defined as “the distance between the centres of the nuclei of two adjacent atoms in the metallic crystal.”

Metallic bond is weaker than Covalent bond, therefore, inter-nuclear distance between two atoms in Metallic bond is more than as in the covalent bond. But metallic bond is stronger than Vanderwaal’s radius Thus, Vanderwaal’s radius > Metallic radius > Covalent radius. Out of these, covalent radius is easier to determine. Therefore the atomic radius of an element is normally expressed as its covalent radius except in noble gases which do not form covalent bonds in their atoms. In noble gases, Vanderwaal’s radius represents the atomic radius.

Valency

The valency of an element may be defined as the combining capacity of element. It is generally expressed in terms of number of hydrogen atoms or the number of chlorine atoms or double the number of oxygen atoms that combine with an atom of the element or replace it. The electrons present in the outermost shell are called Valence Electrons and these electrons determine the valency of the atom.

Variation of Valency

(a).  Variation along period: – Valency of normal elements i.e. s and p-block elements first increases from 1 to 4 and then decreases from 3 to zero. But d-block elements show variable valencies. f-block elements exhibit only electropositive valency, the most common being 3.

(b).  Variation down the group :– On moving down a group, the number of valence electrons remains same and therefore, all the elements in a group exhibit same valency. For example, all the elements of group 1 have valency equal to one and those of group 2 have valency equal to two.

CTM:–  Valency can be determined as follows:-

If no. of electrons in valence shell is 1 to 4 then valency of atom is equal to no. of electrons in valence shell. If no. of electrons in valence shell is 5 to 8, then valency is equal to valence electron -8

Periodic Properties

The properties of elements are periodic function of their atomic numbers. Hence when elements are arranged in the order of increasing atomic number, the characteristic properties of elements reappear at regular intervals. Such characteristic properties of elements which reoccur at definite intervals in periodic table are called Periodic Properties. There are many periodic properties like valency, atomic size, ionic size, ionisation energy, electron affinity, electronegativity etc. These periodic properties can be discussed in the next post.

Division of Periodic Table

Long form of periodic table can be divided into four main blocks. These blocks are called s, p, d, & f. The divisions into four blocks depend upon the type of the orbital into which the last electron of atom has entered.

s – BLOCK 

The elements in which the last electron enters the s – subshell of their outer – most energy level are called s–Block elements. Since s–subshell can accommodate only two electrons so, s – block consists of only two groups i.e. group 1 and group 2. The general electronic configuration of s–block is ns^1–2, where ‘n’ represents the outermost shell. The elements of group 1 are called Alkali Metals & group 2 is called Alkaline Earth Metals.

The general characteristics of s–Block elements are;

(i).          They are soft Metals.

(ii).         They are highly electropositive.

(iii).        They have low ionisation energies.

(iv).        They are very reactive & form ionic compounds.

(v).         They show oxidation state of +1 or +2.

(vi).        They are good reducing agents. 

p – BLOCK ELEMENTS

The elements in which the last electron enters the p–subshell of their outermost energy level are called p–Block elements. The exception is Helium (1_S²). Since p–subshell can accommodate six electrons, so p–block consists of six groups i.e. 13, 14, 15, 16, 17, 18. Most of these elements are non-metals, some are metalloids and a few others are heavy elements which exhibit metallic character. The general configuration of p–block is ns²np^1-6.

The general properties of p–block elements are;

(i).          They show variable oxidation states.

(ii).         They form ionic as well as covalent compounds.

(iii).        Most of them are non-metals.

(iv).        Most of them are highly electronegative.

(v).         Most of them form acidic oxides. 

d – BLOCK ELEMENTS

The elements in which last electron enters the d-sub-shell of the penultimate shell are called d-block elements since d-subshell can accommodate 10 electrons, so d-block consists of ten groups from 3 to 12. The general configuration is (n–1) d ^1–10 ns ^1–2, when ‘n’ is outer most shell. However, the exception is ₄₆Pd (Palladium) whose configuration is 4d¹⁰ 5s⁰, d-block consists of three complete rows and fourth is incomplete. The three rows are called First, Second and Third Transition Series.

The general characteristics of d-block elements are;

(i).          They are hard, high melting points.

(ii).         They show variable oxidation states.

(iii).        They form coloured complexes.

(iv).        They form ionic as well as covalent compounds.

(v).         Most of them exhibit paramagnetism.

(vi).        Most of them possess catalytic properties.

f – BLOCK ELEMENTS

The elements in which the last electron enters the f – subshell of anti-penultimate (third to the outer – most shell) shell are called f-block elements. Their general configuration is (n–2) f ^1–14,     (n–1)d^0–1, ns², where n represents the outermost shell. They consist of two series elements each containing 14 elements placed at the bottom of the periodic table. The first series follow Lanthanium, La and is called Lanthanide Series or simply Lanthanides. The second series follow actinium, ₈₉AC and are called Actinide Series or Actinides. 

The general characteristics of f – block elements are;

(i).          They show variable oxidation states.

(ii).         They are high melting metals.

(iii).        They have high densities.

(iv).        They form coloured compounds.

(v).         Most of the elements of Actinide series are radioactive.

Elements of s-block and p-block are called Representative Elements, elements of d-block are called Transition Elements and f-block elements are called Inner-Transition Elements.

Groups

A vertical column of the periodic table is called a Group. A group consists of a series of elements having similar configuration of the outer energy shell. There are eighteen vertical columns in the long form of periodic table. According to IUPAC, these groups are now numbered from 1 to 18. It may be noted that the elements belonging to the same group are said to constitute a FAMILY. For example, elements of group17 constitute halogen family. Similarly group 16 constitute Oxygen family or Chalcogen Family.

Structural Features of Long Form of Periodic Table

PERIODS

Each horizontal row in the Modern Periodic table is called a Period. In all there are seven periods denoted by numbers 1 to 7. Each period starts with filling of ns orbitals and ends when ‘np’ orbitals are complete. The number of period represents the principle quantum number of valence shell of element present in it.

·                     1^st Period

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

·                     2^nd Period

The second period starts with electrons beginning to enter the second energy shell (n = 2). Since there are only four orbitals in this shell (2s, 2p_x, 2p_y, 2p_z) to be filled which can accommodate eight electrons. Thus, second period has eight elements.

·                     3^rd Period

In this period, filling of third shell (n = 3) takes place. Third shell consists of nine orbitals. But the five 3d orbitals have higher energy than 4s orbitals. As such only four orbitals (one 3s & three 3p) of third shell are filled before filling of fourth shell. Hence there are only eight elements in third shell).

·                     4^th Period

The fourth period corresponds to n = 4. It starts with the filling of 4s – orbitals & 4p orbitals. However, after 4s orbitals but before 4p orbitals, there are five 3d orbitals to be filled. Thus in all nine orbitals (one 4s, five 3d & three 4p) are to be filled and as such there are Eighteen elements in fourth period.

·                     5^th Period

The fifth period corresponds to n = 5. It is similar to fourth period. There are nine orbitals (one 5s, five 4d and three 5p) to be filled. Therefore there are Eighteen elements as well in fifth period.

·                     6^th Period

This period starts with filling of 6s orbitals and ends with 6p. But in between that 4f-orbitals and 5d–orbitals are to be filled. So, there are sixteen orbitals in all (one 6s, seven 4f, five 5d & three 6p) to be filled. Thus, there are Thirty-two elements in sixth period.

·                     7^th Period

This period starts with filling of 7s and ends with 7p. But in between 5f, 6d are to be filled. So, this too consists of sixteen orbitals to be filled. So, it should contain thirty two elements. But it is still incomplete.

It may be noted that periods 2 and 3 contain 8 elements each and are called Short Periods. There are 18 elements each in 4^th & 5^th period and they are called Long Periods. Sixth period containing 32 elements is called Longest Period while seventh is incomplete.

The sixth & seventh periods contains a group of 14 elements called Lanthanides & Actinides respectively. If all Lanthanides and actinides are placed horizontally in the same group, there will be undue expansion of the periodic table. Hence they are placed in the separate horizontal row at the bottom of Main periodic table.