Periodic Table

CLASSIFICATION OF ELEMENTS

Q: - What is periodic table?

Ans: - The table in which all the elements are classified according to their properties is called periodic table. In this table elements are arranged in the increasing order of their atomic number.In this arrangement elements with similar properties fall under same vertical column The aim of this arrangement of elements into the tabular form is to make their study easier.

HISTORY OF CLASSIFICATION OF ELEMENTS

1)DOBEREINER LAW OF TRIADS

Debereiner in 1829 made an attempt to arrange the elements in groups. He placed three elements in each group. He named those groups as triads. The three elements in a triad had similar chemical properties. He arranged the elements in a triad in the increasing order of their atomic masses. He generalized that the atomic mass of the middle element was nearly equal to the arithmetic mean of atomic masses of first and third element in a triad.

NEWLAND LAW OF OCTAVES

In 1864, John Newland and English chemist and a musician made an attempt to arrange the elements in order of their increasing atomic masses on the basis of eighth note of an octave in music.

According to this law if elements are arranged in the increasing order of their atomic masses, properties of every eighth element resembles with the first element.

Julius Lothar Meyer

In 1869 a German chemist, Julius Lothar Meyer plotted a graph of atomic volume (atomic mass/density) versus atomic mass for various elements. He noticed that elements with similar properties occupied similar positions on the curve. All the alkali metals occupied the peak positions. Thus Lothar observed a periodicity in the properties of elements with atomic masses. This was the first time that a definite pattern in the periodicity in the properties of elements was confirmed.

4 Mendeleev’s periodic table

In (1869)He proposed his periodic table on the basis of his periodic law which states that properties of elements are a periodic function of their atomic masses. It means when the elements are arranged in order of their increasing atomic masses, the elements with similar properties recur after certain regular intervals.

Russian Chemist Dmitri Mendeleev contribution in the field of classification of elements is a light post for the generations to come in chemistry.

He arranged the then known all 63 elements in the order of increasing atomic masses and similarities in their physical and chemical properties.

 

Among chemical properties,he focused on the compounds formed by elements with oxygen and hydrogen . The formulae of hydrides and oxides formed by an element were considered as one of the basic properties of  an element for arrangement in Mendeleev’s periodic table.

Further, the Mendeleev’s periodic law states –

“The properties of elements are the periodic function of their atomic masses.”

Arrangements ofMendeleev’s Periodic Table

It consists of –

(1) Six horizontal rows called periods

(2) Eight vertical columns called groups

In a group , the elements have similar properties and exhibit a clear trend down the group.

First seven groups were divided into two subgroups called ‘A’ and ‘B’

The elements which had kept in the left hand side of a group( group A) are called normal or representative elements.

The elements which were present in the right hand side (group B) are called >transition elements.

Group VIII had nine transition elements present in three sets, each containing three elements lied in the 4th, 5th and 6th period.

Q: - What is periodicity?

Ans: - The repetition in the properties of elements after certain regular intervals in the modern periodic table is called periodicity. It means when the elements are arranged in order of their increasing atomic numbers, the properties of the elements are repeated after certain regular intervals in the modern periodic table.

Q: - What is the cause of periodicity?

Ans: - The repetition of the same valence shell electronic configurations of the elements present in a group after certain regular intervals is the cause of periodicity.

Q :- What are magic numbers?

Ans :- The electronic configurations of alkali metals are repeated after regular intervals of 2,8,8,18,18 and 32. These numbers are sometimes called magic numbers.

Modern Periodic Table

 

Q: - Explain long form of modern periodic table in detail. 

 

 

Ans: -In 1921 Bohr constructed the long form of modern periodic table. For an easy and systematic study, the elements of the modern periodic table are classified into: 

 

1.PERIODS

 

2.GROUPS

 

3. BLOCKS

1) PERIODS

 The modern periodic table has been divided into seven horizontal rows called periods. The first, second and third periods are called short periods. They contain 2, 8 and 8 elements respectively .Fourth, fifth,sixth and seventh periods are called long periods .They contain 18, 18 ,32 and 32 elements respectively.

2) GROUPS

The modern periodic table has been divided into eighteen vertical columns called groups.

i) 1st group: - Elements of this group are called alkali metals. This group includes H, Li,Na,K,Rb,Cs and Fr.

ii) 2nd group :- Elements of this group are called alkaline earth metals and includes Be, Mg, Ca, Sr, Ba and Ra.

iii) 3rd to 12th groups :- The elements from group 3rd to group 12th are known as transition elements.

iv) 13th group :- This group is also called boron family because boron is the first element of this group. It includes B, Al, Ga, In and Tl.

v) 14th group :- This group is called carbon family. It includes C, Si, Ge, Sn and Pb.

vi) 15th group :- This group is also known as nitrogen family or pincogens. This group includes N, P, As, Sb and Bi.

vii) 16th group :- This group is known as oxygen family or chalcogens. It includes O, S, Se, Te and Po.

viii) 17th group :- The elements of this group are also called halogens . These are F, Cl, Br, I, and At.

ix) 18th group :- The elements of this group are known as inert gases or noble gases. These are He, Ne, Ar, Kr,Xe Rn.

3) BLOCKS

Elements in the modern periodic table have been divided into four blocks. This division is based upon the name of the orbital that receives the last electron. These blocks are called:

i) s-BLOCK :- The elements whose last electron enters into the s orbital of the outermost shell are called s-block elements. This block contains elements of group-I and group-II.

ii) p-BLOCK:- The elements whose last electron enters into the p-sub shell of the valence energy level are called p-block elements. This block contains elements from group number 13th to group number 18th .

Representative elements :- The s-block and p-block elements are collectively known as representative elements.

iii) d-BLOCK :- The elements whose last electron enters into the d-sub shell of the penultimate energy level are called d-block elements. This block contains elements from group 3rd to group 12th .

iv) f-BLOCK :- The elements whose last electron enters into the f-sub shell of the ante-penultimate energy level are called f-block elements. or 

The elements whose last electron enters into any one of the seven f-orbitals of the ante-penultimate shell are called f-block elements. This block consists of two series of elements .

 

1. Lanthanide

 

2. Actinide

These two series are placed at the bottom of the periodic table. Each series comprises fourteen elements. The first series is called Lanthanide and second series is known as Actinide. The f-block elements are also known as inner transition elements.

Q:- Why do elements in the same group have similar properties?

Ans:- Because they have similar outer shell electronic configurations. 

Q:- Why atomic number is batter base for classification of elements than atomic mass?

Ans :-Atomic mass is a nuclear property and it has very little effect on the physical and chemical properties of the elements . Atomic number is directly linked to the number of electrons present in the extra nuclear part of an atom. Physical and chemical properties depend upon the number of electrons. Therefore, atomic number is considered to be the better base for classification of elements.

CHARACTERISTICS OF s-BLOCK ELEMENTS

I. They are soft metals.

II. They have low melting and boiling points.

III. They have low ionization energies.

IV. They are highly reactive in nature.

V. They are electropositive in nature.

VI. They show oxidation states of +1 and +2.

VII. They form ionic compounds (except beryllium).

VIII. They impart characteristic colors to the flame (except Be & Mg).

IX. They are strong reducing agents

X. They are good conductor of heat and electricity.

CHARACTERISTICS OF p-BLOCK ELEMENTS

I. They include both metals and non-metals.

II. They exist in all the three physical states.

III. Their ionization energies are higher than s-block elements.

IV. They generally form covalent compounds.

V. Most of them show variable oxidation states.

VI. Their oxidizing character increases from left to right in a period.

VII. Reducing character increases from top to bottom in a group.

VIII. They possess higher values of Electronegativity.

CHARACTERISTICS OF d-BLOCK ELEMENTS

I. They are hard metals (except Hg).

II. They have high melting and boiling points.

III. They are good conductor of heat and electricity.

IV. They show variable oxidation states.

V. They form colored compounds

VI. They from complex compounds

VII. Most of them are used as catalysts

VIII. They form ionic and covalent compounds.

IX. Their ionization energies lie between s- and p-block elements.

X. They also form alloys with other metals.

CHARACTERISTICS OF f-BLOCK ELEMENTS

I. They are heavy metals

II. They have high melting and boiling points

III. They show variable oxidation states.

IV. They form complex compounds.

V. They form colored compounds

VI. Almost half of them are radioactive in nature.

VII. Most of them are paramagnetic in nature.

VIII. Most of them are synthetic.

PERIODIC PROPERTIES

Those properties which get repeated after a regular interval in a regular manner are called periodic properties.These properties vary gradually along a period from left to right or in a group from top to bottom. The important periodic properties are as follow:

1) Atomic and ionic radius

2) Ionization energy or enthalpy

3) Electron gain energy or enthalpy

4) Electronegativity

5) Melting point

6) Valency

7) Density

(I) ATOMIC RADIUS

The distance from the center of the nucleus to the outermost shell containing electrons. The atomic radius is a term used to describe the size of the atom.

 

 

There is no standard definition for this terrn. Atomic radius may refer to the ionic radius, covalent radius, metallic radius, or van der Waals radius.

COVALENT RADIUS

It is one-half of inter-nuclear distance between two covalently bonded atoms of the same element in a molecule. Or It is one-half of inter-nuclear distance between two covalently bonded atoms in a homo atomic molecule.

E.g. the inter-nuclear distance between two hydrogen atoms in hydrogen molecule is 74 pm. Hence the covalent radius of hydrogen is 74/2 = 37 pm. 

 

Similarly the inter nuclear distance between two chlorine atoms in chlorine molecules (Cl2) is 198 pm. Hence the covalent radius of chlorine atom is 198/2 = 99 pm. 

Van der waal’s radius

It is one- half of the inter-nuclear distance between two nearest atoms of two neighboring molecules in the solid state. E.g. the inter-nuclear distance between two adjacent or nearest hydrogen atoms of two neighboring molecules in the solid state is 240pm. Therefore its van der waal’s radius is 240/2 = 120 pm. This radius is called van der waal’radius because the two neighboring molecules are held together by weak van der Waal’s forces of attraction in the solid state. This radius is always greater than the covalent radius because a covalent bond is formed by the overlapping of atomic orbitals. The overlapping region shortens the inter-nuclear distance while there is no such overlapping in case of van der Waal’s forces. These forces are very weak hence their magnitude is very small in gaseous as well as liquid states of the substance. Hence this radius is determined in the solid state when the magnitude of these forces is maximum.

 

Q:- The atomic sizes of noble gases are measured in terms of van der waal’s radii. Why ?

Ans :- since noble gases do not form covalent bonds and therefore only weak physical forces are operative in the crystals of noble gases.

Q :- Why do noble gases have bigger atomic size than halogens?

Ans :- The atomic size of noble gases is measured in terms of van der waal’s radius and that of halogens in terms of covalent radius . The van der Waal’s radius is always greater than the covalent radius.

Q :- Covalent radius is smaller than van der waal’s radius . Why?

Ans :- A covalent bond is formed by the overlapping of atomic orbitals. The overlapping region shortens the inter nuclear distance while in case of van der waal’s forces there is no such overlapping and two nearest atoms of two neighboring molecules are relatively at larger distances. One half of their inter-nuclear distance will be more than the two similar atoms forming a covalent bond. Hence covalent radius is smaller than the van der Waal’s radius. 

METALLIC RADIUS

It is one-half of inter-nuclear distance between two nearest atoms in the metallic lattice.

Increasing order of different types of radii

Covalent radius < Metallic radius < Van der waal’s radius

Q: - Explain the variation of atomic radius in the periodic table?

Ans: - Variation of atomic radius in a period:- In general the atomic radius decreases as we move from left to right in a period. 

 

 

 

 

 

EXPLANATION OF REASON: - As we move left to right in a period the nuclear charge gradually increases by one unit at each step and at the same time one electron is added in the same energy level. Due to increased nuclear charge electrons of valence shell are pulled more strongly closer to the nucleus and thereby atomic radii keep on decreasing from left to right. E.g. In the third period there is gradual decrease in the atomic radii from sodium to chlorine.

VARIATION OF ATOMIC RADIUS IN A GROUP:-As we move from top to bottom in a group the atomic size increases.

 

 Reason or explanation: - This is due to increase in the number of shells and screening effect of inner shell electrons. These two factors dominate over the effect of increased nuclear charge. So, atomic radii increase as we move down the group. E.g. atomic radii of alkali metals increase as we move down the group from lithium to cesium.

Q: - What are ions?

Ans: - When a neutral atom losses or gains electrons, ions are formed.

They are of two types:

1) Cations: - The ions formed by the loss of electrons from a neutral atom are known as cations.

2) Anions: - The ions formed by the gain of electrons by a neutral atom are called anions. 

 

 

(II) IONIC RADIUS

It is the effective distance from the center of nucleus of an ion upto which it exerts an influence on its electron cloud. 

CHARACTERISTICS OF IONIC RADII

1) The size of the cation is always smaller than its neutral atom

2) Radius of anion is more than that of its parent atom.



Q :- The size of the cation is always smaller than its neutral atom ? or

Q :- The radius of the cation is always smaller than its parent atom?

Ans: - Cation is formed when a neutral atom loses electrons. In this cations the magnitude of nuclear charge remains the same as that in its parent atom whereas the number of electrons decreases .As a result of which the same nuclear charge exerts more attractive force on the remaining electrons and therefore electrons are pulled strongly towards the nucleus. This causes shrinkage of electron cloud and size of the cation decreases. E.g. atomic radius of sodium atom = 154 pm and sodium ion = 95 pm. 

Q: - Radius of anion is more than that of its parent atom. Why?

Ans: - Anion is formed when a neutral atom gains electrons. In this anion the magnitude of nuclear charge remains the same as that in its parent atom but the number of electrons increases. As a result of which the same nuclear charge exerts lesser attractive force on relatively larger number of electrons. This causes expansion of electron cloud and size of the anion increases as compared to its parent neutral atom. E.g. Atomic radius of chlorine =99 pm and chloride ion = 131 pm. 

 

Q: - Explain the variation of ionic radius within a period and  group in the periodic table?

1) Variation of ionic radii in a period

 2) Variation of ionic radii in a group 

 Variation of cationic radii in a group:

As we move down the group from top to bottom the cationic radius increases due to an increase in the number of shells .e.g. let us consider the ionic radii of cations of the 1st group . 

Variation of anionic radii within a group:

 

 The ionic radius of anions increases down the group due to an increase in the number of shells. Let us consider the example of anions of group 17.

(III) IONIZATION ENERGY:

The energy required to remove an electron from a neutral isolated gaseous atom or a cation is called ionization energy. Its units are Kj/mol. The process of removing an electron from an atom is always endothermic in nature. It means energy is supplied from surrounding to the system. It is represented by prefixing positive sign with enthalpy change .

M(g) +  IE  → M⁺ (g)  +  e^-

SUCCESSIVE IONIZATION ENERGIES: - The ionization energies required to remove 1st ,2nd and subsequent electrons from an isolated gaseous atom or a cation are collectively called successive ionization energies.

This concept is applicable where more than one electron need to be removed and then this is done one after the other that is in succession and not simultaneously.

FIRST IONIZATION ENERGY (IE1):- The energy required to remove first electron from an isolated neutral gaseous atom to form a univalent cation is called first ionization energy. 

 

M(g) +  IE₁  → M⁺ (g)  +  e^-

SECOND IONIZATION ENERGY (IE2) :- The energy required to remove 2nd electron from the univalent ion is called second ionization energy. 

 

M⁺ (g) +  IE₂  → M²⁺ (g)  +  e^-

THIRD IONIZATION ENERGY (IE3) :- The energy required to remove 3rd electron from the divalent ion is called third ionization energy.

M²⁺ (g) +  IE₃  → M³⁺ (g)  +  e^-

The increasing order of successive ionization energies is :

IE1 < IE2 < IE3

Reason: - When first electron is removed from neutral atom a monovalent cation is formed which has one electron less but same magnitude of nuclear charge as that of parent atom. As a result the attraction of nuclear charge increases over the remaining electrons. Hence the second ionization energy required to remove another electron will be higher than the first. Similarly the removal of second electron results in the formation of divalent positive ion and the attraction between the nucleus and the remaining electrons increases further hence the remaining electrons will be held more tightly. As a result the energy required to remove the third electron from the divalent positive ion will be even more than that required for the second electron. Let us consider the example of successive ionization energies of Aluminum atom.

M(g) +  IE₁  → M⁺ (g)  +  e^-  (577 kj/mol)


M⁺ (g) +  IE₂  → M²⁺ (g)  +  e^- (1795 kj/mol)


M²⁺ (g) +  IE₃  → M³⁺ (g)  +  e^- (2758 kj/mol)

FACTOR ON WHICH IONIZATION ENERGY DEPENDS

1) Atomic size: - IE decreases with the increase in atomic size.

2) Nuclear charge: - IE increases with the increase in magnitude of nuclear charge.

3) Penetration effect of electrons: - Greater the penetration effect more will be the IE . This means greater the probability of finding an electron closer to the nucleus, more will be its attraction towards nucleus and hence more energy will be required to remove that electron. Increasing order of penetration effect of different electrons .

f < d < p < s

4) Screening effect:-Greater the magnitude of screening effect lesser will be IE.

5) Symmetry of electronic configurations: - Greater the symmetry of electronic configuration more will be its stability and thus higher will be the value of IE.

PERIODIC TREND OF IONIZATION ENERGY

1) Variation along a period:-In general, as we move from left to right in a period, the ionization enthalpy increases.

Reason:-

2) Variation down the group: - As we move down the group, the ionization energy decreases.

Reason:-

IV) ELECTRON AFFINITY OR ELECTRON GAIN ENTHAPLY:-

The energy released when a neutral isolated gaseous atom accepts an electron to form a uninegaitve ion. Its units are kj/mol. Depending on the nature of element the process of adding an electron to the atom can be either endothermic or exothermic.

SUCCESSIVE ELECTRON GAIN ENTHALPIES

 The electron gain enthalpies involved for the addition of 1st ,2nd and subsequent electrons into an isolated gaseous atom or anion are collectively called successive electron gain enthalpies .This concept is applicable where more than one electron need to be added and then this is done one after the other that is in succession and not simultaneously.

First electron gain enthalpy for most of the elements is taken as negative and successive electron gain enthalpies are taken as positive. If the energy is released (exothermic) the electron gain enthalpy will be negative and if energy is supplied (endothermic) then it will be positive.

REASON: - When a neutral gaseous atom gains an electron, a uninegaitve ion is formed. The energy released is called first electron gain enthalpy. However the addition of second electron to the uninegaitve ion is strongly opposed due to inter electronic repulsions. Hence energy is supplied from outside for the addition of second electron to overcome these repulsive forces. That is why the second electron gain enthalpy has positive value.

Q: - The first electron gain enthalpy for most of the elements is negative. Why?

Ans: - . Energy considerations show that system with less energy is more stable than the system possessing more energy. Every system wants to acquire a state of less energy and more stability. Hence in order to acquire stable electronic configurations, these elements have strong attraction for the additional electron and their electron gain enthalpy becomes more negative. Actually acquisition of this state of stable electronic configurations by accepting extra electron is accompanied with the release of energy. Hence the first electron gain enthalpy for most of the elements is taken as negative that is the process is exothermic.

FACTORS ON WHICH ELECTRON GAIN ENTHALPY DEPENDS

1) Atomic size: - Electron gain enthalpy becomes less negative with the increase in atomic size. Because with the increase in atomic size, the attraction between the nucleus and the incoming electron decreases. E.g.

2) Nuclear charge: - The electron gain enthalpy becomes more negative with the increase in nuclear charge. Because the force of attraction between the nucleus and the incoming electron increases .

3) Symmetry of electronic configurations: - The electron gain enthalpies have positive values for the elements with symmetrical electronic configurations. Because these elements do not have any urge to take up extra electrons .As result energy has to be supplied to add an electron and therefore their electron gain enthalpies are positive.

PERIODIC TREND FOR ELECTRON GAIN ENTHALPY

1) VARIATION OF ELECTRON GAIN ENTHALPY ALONG A PERIOD:-In general, the electron gain enthalpies become more negative along a period.

REASON: - As we move across a period from left to right the atomic size decreases and nuclear charge increases. Both these factors result in greater attraction for the incoming electron and hence electron gain enthalpy becomes more and more negative.

2) VARIATION ELECTRON GAIN ENTHALPY IN A GROUP: - In general it becomes less negative as we move down the group.

REASON:-As we move down the group, both the atomic size and nuclear charge increase. But the effect of increase in atomic size is much more pronounced than the nuclear charge due to shielding effect. As a result, the attraction between the nucleus and incoming electron decreases and hence the electron gain enthalpy becomes less negative. 

(V) VALENCY

It is the ability of an element to react or combine with other element so as to form a chemical bond.That is why valency is called combining capacity of an element.The chemical properties of elements depend upon the number of these valence electrons .It is equal to the number of electrons gained, lost or shared in order to acquire the stable noble gas electronic configuration. It is determined by the number of valence electrons . In case of representative elements valency is either equal to the number of valence electrons or eight minus the number of valence electrons. Transition and inner transition elements show variable valency. The most common valency shown by transition elements is 2 and that of inner transition elements is 3. Charges are not associated with valency, i.e. valency is not + or -. It is always a whole number. For e.g. sodium has configuration (2, 8, 1) and tends to gain the configuration of neon, Ne (2, 8) by losing an electron. Hence, its valency is 1. 

 

Calculating the Valency of an Element (or Molecule)

Step 1

Consult the periodic table of the elements to determine the valency of an element. The periodic table is organized by groups in rows and columns, and the elements of groups I-VIII have the same valency as others in their group. All the elements in group VIII have eight electrons in their outer shells, and thus have a valency of zero (highly stable). Elements in group I just have one valent electron in their outer shells and thus have a valency of one, which means they are very reactive. Group IV/valency 4 elements like carbon are relatively stable. Group VI and VII elements like oxygen are also reactive as they seek electron pairs to complete their outer shell octet.

Step 2

Calculate the valency of an element using the total number of electrons. The valency of an atom is equal to the number of electrons in the outer shell if that number is four or less. Otherwise, the valency is equal to eight minus the number of electrons in the outer shell. The number of electrons in each shell of an atom is regular so if you know the number of electrons in the atom, then you can calculate the valency. All atoms (except hydrogen) have two electrons in the first electron shell, and up to eight electrons in each succeeding electron shell. For example, carbon has six electrons, two in the first shell, and four in the outer shell, giving it a valency of four. Oxygen has eight electrons, two in the first shell and six in the outer shell, giving it a valency of two (8 - 6 = 2).

Step 3

Calculate the valency of multi-element molecules using the same procedure. For example, to determine the valency of the ionic molecule phosphorus tetraoxide (PO4, four atoms of oxygen and one atom phosphorus) you multiply the total valency of the four oxygen atoms (valency 2) and subtract that from the valency of the phosphorus atom (valency 5). That reveals the valency of PO4 is 3.

Because of the ambiguity of the term valencey nowadays other notations are used in practice . oxidation state" is a more clear indication of the electronic state of atoms in a molecule. The "oxidation state" of an atom in a molecule gives the number of valence electrons it has gained or lost. In contrast to the valency number, the oxidation state can be positive (for an electropositive atom) or negative (for an electronegative atom). Elements in a high oxidation state can have a valence higher than four. For example, in perchlorates, chlorine has seven valence bonds and ruthenium, in the +8 oxidation state in ruthenium tetroxide, has eight valence bonds

TYPES OF VALENCY

1) FIXED VALENCY:-The valency of an element that does not change is called fixed valency. Elements with fixed valency possess only one valency. E.g. Carbon, Oxygen etc.

2) VARIABLE VALENCY:- The valency of an element is said to be variable if it is not fixed. Elements with variable valency possess more than one valency..E.g., Cu and Sn. The name of the element with lower valency in a compound end with a suffix ous while with higher valency ends with a suffix ic. 

 

PERIODIC TREND IN VALENCY

VARIATION OF VALENCY ALONG A PERIOD: - On moving along a period from left to right the valency of elements first increases from 1 to 4 and then again falls to 0. The noble gases have zero valency as they are chemically inert. E.g. lets us consider the elements of second period:

 

Reason: - . Valency is determined by the number of valence electrons. Hence this trend in the valency is due to the increase in the number of valence electrons as we move along a period from left to right In case of representative elements valency is either equal to the number of valence electrons or eight minus the number of valence electrons .

VARIATION OF VALENCY DOWN THE GROUP

All the elements in a group exhibit same valency. E.g. all members of the first group possess common valency that is one and all members of the second group possess same valency that is 2.

REASON :- This is due to the fact that all the elements in a group possess same number of valence electrons. That is why they exhibit a common valency.

(VI) ELECTRONEGATIVITY

Q: - What is Electronegativity?

Ans: - The tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond. Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7.

FACTORS THAT INFLUENCE ELECTRONEGATIVITY

I) ATOMIC SIZE: - Smaller the size of an atom, larger will be its Electronegativity value.

II) NUCLEAR CHARGE:-Electronegativity increases with the increase in nuclear charge of an atom.

PERIODIC TREND OF ELECTRONEGATIVITY

1) VARIATION OF ELECTRONEGATIVITY ALONG A PERIOD:-The Electronegativity values increase along a period from left to right. E.g. Halogens have highest value of electronegativity in their respective periods.

REASON: - This is because as we move along a period from left to right, the nuclear charge increases and atomic size decreases. As a result of which nucleus exerts a strong pull on the bonded electrons.

2) VARIATION OF ELECTRONEGATIVITY IN A GROUP:-The Electronegativity values decrease down the group. E.g. Among halogens fluorine has highest Electronegativity value.

REASON: - This is because when we move down the group, there is an increase in atomic size and screening effect. This over weighs the increased nuclear charge and thus the attraction between nucleus and bonded electrons decreases.

APPLICATIONS OF ELECTRONEGATIVITY

1) METALLIC AND NON-METALLIC CHARACTER: - With the increase in electronegativity the non-metallic character increases and metallic character decreases and vice versa.

2) TYPES OF COVALENT BONDS:- Depending upon the electronegativity of bonded atoms the covalent bond may be polar and non-polar. E.g HCl,Cl2,H2 etc.

(a) POLAR COVALENT BOND :-If a covalent bond is formed between two dissimilar atoms having different electro--negativity values then the shared pair is displaced towards more electronegative atom. Such a bond is called polar covalent bond. E.g. H—Cl,H—F ,H—O—H ,NH3 etc.

(b) NON-POLAR COVALENT BOND :-If a covalent bond is formed between the two identical atoms then the shared pair of electrons will lie exactly in the middle of both the atoms. Such a bond is called non-polar covalent bond. E.g. H2,Cl2,Br2,O2 etc.

3) COMPARISON OF POLARITY OF MOLECULES: -(it means out of the given polar molecules which one has more polar covalent bond.) The polarity of a molecule is measured in terms of its dipole moment.( Polar molecules always have some dipole moment. It helps to compare the relative polarities of the molecules.) Greater the value of dipole moment more will be the polarity of the molecule. (It is the product of magnitude of charge and bond length.) Higher the electronegativity difference between the two bonded atoms more is the dipole moment of the molecule and hence more will be the degree of polarity of a covalent bond. It keeps on decreasing as the electronegativity difference between the two bonded atoms decreases.

4) PERCENTAGE IONIC CHARACTER:- A polar covalent bond always possesses some percentage of ionic character. If the electronegativity difference between the two bonded atoms is 1.9, the bond is said to have 50% ionic and 50% covalent character. If this difference goes on increasing then percentage of ionic character is increased further. If this difference is less than 1.9, the bond is predominantly considered covalent.

(VII) DIAGONAL RELATIONSHIP

The resemblance in the properties of the elements present diagonally is called diagonal relationship. There are three pairs of elements in the periodic table which show diagonal relationship. These are Li—Mg, Be—Al and B—Si. The first three elements of the second period that is Li, Be and B are called bridge elements because they act as a bridge between the two groups by keeping diagonal similarity in properties. 

CAUSE OF DIAGONAL RELATIONSHIP: - This relationship occurs due to the opposing trends in periodic properties along a period and down the group. For example, atomic radii decrease along a period and increase down the group. On moving diagonally, the two different trends mutually cancel and therefore, these elements have similar properties.

CAUSE OF ANOMALOUS BEHAVIOUR OF FIRST ELEMENT OF EACH GROUP

1) Small size

2) Absence of vacant d-orbitals in the valence shell

3) High ionization energy

4) High electronegativity.

5) Ability to form pπ-pπ multiple bonds(e.g. C ,N and O)

CHEMICAL PROPERTIES OF THE ELEMENTS

The properties which can only be observed or measured by performing a chemical reaction. The chemical properties of elements depend upon their electronic configuration i.e. the number of valence electrons.

Important chemical properties are:

1) Electropositive nature :

2) Electronegative nature

3) Reducing character

4) Oxidizing character

5) Metallic character

6) Non metallic character

7) Reaction with oxygen

PERIODIC TREND IN THE CHEMICAL PROPERTIES OF THE ELEMENTS

ALONG A PERIOD

1) ELECTROPOSITIVE NATURE:-Elements with electropositive nature have the tendency to lose electrons and acquire positive charge. This character decreases as we move along a period from left to right. E.g. alkali metals placed at the extreme left of the periodic table are the most electropositive in nature.

2) ELECTRONEGATIVE NATURE: - Elements with electronegative nature have the tendency to gain electrons and acquire negative charge. This character increases as we move along a period from left to right. E.g. halogens placed on the extreme right of the periodic table are the most electronegative in nature.

3) REDUCING CHARACTER:-Elements which are electron donor are called reducing agents. This character decreases from left to right along a period. E.g. alkali metals placed at the extreme left of the periodic table are good reducing agents.

4) OXIDIZING CHARACTER: - Elements which are electron acceptor are called oxidizing agents. E.g. halogens placed on the extreme right of the periodic table are good oxidizing agents.

5) METALLIC CHARACTER: - Elements which are electropositive in nature are metals. This character decreases along a period. E.g. alkali metals placed at the extreme left of the periodic table are the most metallic elements.

6) NON-METALLIC CHARACTER: - Elements which are electronegative in nature are non-metals. Non-metallic character of the elements increases along a period from left to right. Hence halogens are most non-metallic elements.

7) REACTION WITH OXYGEN:-The binary compounds of elements with oxygen are called oxides. Along a period the basic character of oxides decreases and acidic character increases.

Following types of oxides are formed by the elements.

I) The elements present on the extreme left form basic oxides. E.g. Na2O, MgO.

II) Whereas those present on the extreme right form acidic oxides. E.g. P2O5, SO2,, Cl2O7.

III) Elements present in the center of the periodic table form amphoteric oxides. E.g. Al2O3 , SiO2 .



DOWN THE GROUP:-

1) ELECTROPOSITIVE NATURE:- Increases down the group.

2) ELECTRONEGATIVE NATURE: - Decreases down the group

3) REDUCING CHARACTER:- Increases down the group

4) OXIDIZING CHARACTER: - Decreases down the group

5) METALLIC CHARACTER: - Increases down the group

6) NON-METALLIC CHARACTER: - Decreases down the group

7) OXIDES:- Basic character increases down the group. 

CHEMICAL REACTIVITY

The chemical reactivity of the elements is directly related to the number of valence electrons and physical properties of the elements. As a result of which alkali metals placed at the extreme left are the most reactive metals and halogens placed on the extreme right of the periodic table are the most reactive non-metals. Chemical reactivity of the elements can be shown by their reactions with oxygen.

Write down the chemical reactions

Q:- Why do we call f-block elements as rare earth elements?

Ans:- Because they are rarely found in earth crust.



Q :- Why do we place lanthanides and actinides at the bottom of the periodic table?

Ans :- Because they resemble with each other but do not resemble with any other group elements.



Q :- Why are cations smaller than neutral atoms?

Ans :- Because their effective nuclear charge is greater than the neutral atoms



Q :- Why do anions have bigger size than neutral atoms?

Ans :- Because their effective nuclear charge is less than the neutral atoms.

Q :- What are iso-electronic or iso-electronic ions ?

Ans :- The chemical species or ions containing same number of electrons but different nuclear charge are known as iso-electronic or iso-electronic ions. E.g. O2-, Na+, F- and Al3+ are iso-electronic because each of them contains 10 electrons.

Q :- What is screening effect?

Ans: - The process of shielding or screening of valence shell electrons form the nucleus by the inner shell electrons is called screening effect. In other words the inner shell electrons act as a shield between the nucleus and valence shell electrons .As a result of which the outermost electrons do not feel the full charge of the nucleus. In simple words the force of attraction between the nucleus and valence shell electrons decreases due to the presence of inner shell electrons.

Q: - What is effective nuclear charge?

Ans: - The actual nuclear charge felt or experienced by the valence shell electrons due to the screening effect is called effective nuclear charge. It is denoted by the symbol Zeff or Z*. It is obtained by subtracting screening constant from the total nuclear charge as follow:

Zeff = Total nuclear charge (Z) – screening constant(S)



Here Z* is effective nuclear charge which is always less than the total nuclear charge. S is screening constant and it is the measure of magnitude of shielding effect of inner shell electrons. It is calculated using Slater’s rule. Greater the number of screening electrons greater will be the value of screening constant hence larger will be the screening effect.

(1) Numerical: - Calculate effective nuclear charge for the carbon atom.

Ans:- Formula to calculate effective nuclear charge is:



Zeff = Total nuclear charge (Z) – screening constant(S)

Now total nuclear charge for the carbon atom is Z = 6. As atomic = number of electrons or Number of protons in a neutral atom.

Now screening constant S can be calculated using Slater’s rule which is as follow:

1) Write electronic configuration and make groups. S and p orbitals of a shell are written together whereas d and f orbitals are written separately.

Electronic configuration of carbon = (1s2) (2s2 2p2)

2) For nth group each electron other than the one under consideration contributes = 0.35.

3) Each electron of n-1 group contributes = 0.85

S = (0.35*3)+(0.85*2) = 2.75

Z* = 6 -- 2.75

Z* = 3.25 Ans.



(2) Numerical :- Calculate effective nuclear charge for one of the outer electrons of oxygen atom.

Ans :- Electronic configuration of oxygen = (1s2) (2s2 2p4)

S = (0.35*5)+(0.85*2) = 1.75+1.70=3.45

Z = 8

Zeff = (Z) – (S)

Z* = 8 – 3.45

Z* = 4.55 Ans.-

 

NOMENCLATURE OF ELEMENTS WITH ATOMIC NUMBERS > 100

The naming of the new elements had been traditionally the privilege of the discoverer (or discoverers) 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 characterization, 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 discovering element 104. The Americans named it Rutherfordium 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 0 and numbers 1-9. These are shown in Table 3.4. The roots are put together in order of digits which make up the atomic number and “ium” is added at the end.

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 representatives from each country. The permanent name might reflect the country (or state of the 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 116 have been discovered. Elements with atomic numbers 113, 115, 117 and 118 are not yet known.

Q:-What would be the IUPAC name and symbol for the element with atomic number 120?

Ans:- the roots for 1, 2 and 0 are un, bi and nil, respectively. Hence, the symbol and the name respectively are Ubn and unbinilium.

BOND LENGTH: - The inter nuclear distance between the two bonded atoms is called bond length.


Heat of atomization: - the energy required to isolate an atom from solid is called heat of atomization.

ATOMIC VOLUME: - It is the volume occupied by one mole of the atoms of an element in the solid state. Its units are cc/mol.

DIPOLE

A polar molecule is one which contains polar covalent bond. It consists of two poles. Its one pole bears positive charge and other pole carries negative charge. Such a molecule with two poles is called dipolar molecule or dipole.

 In order to measure the polarity of a dipole or bipolar molecule or a polar molecule, a physical quantity called dipole moment is used. Polar molecules always have some value of dipole moment. It helps to compare the relative polarities of the molecules. Greater the value of dipole moment more will be the degree of polarity of the molecule that is in simple words more polar will be the covalent bond of the concerned molecule.

DEFINITION OF DIPOLE MOMENT: - It is the product of magnitude of charge and bond length(inter nuclear distance between the two bonded atoms). It is a vector quantity.It is denoted by the Greek letter Mu ( 𝞵).

𝞵 = q ⨯ d

It is represented by an arrow with its tail at the positive pole and head pointing towards the negative pole.

 UNITS :-Its units are Debye and denoted by the symbol D. The Debye   is a CGS unit (a non SI unit) of  dipole moment  which is named in honor of the physicist Peter J. W. Debye.

 

If the magnitude of charge is 1.0×10^-10 esu

and  bond length 1.0×10^-8 cm

μ = q × d

μ = 1.0×10^-10 esu  × 1.0×10^-8 cm

μ = 1.0 * 10^-18 esu cm

then the dipole moment is said to be one Debye hence

1 Debye = 1.0 * ^-18 esu cm 

 Relationship between CGS unit (esu cm)  of dipole moment and SI unit (C m) 

1 Debye = 1.0 * ^-18 esu cm

 

1 Debye = 1.0 *10 ^-18 esu cm  = 3.335 * 10 ^-30 C m

 

DIPOLE MOMENT FOR 100% IONIC MOLECULE

If a polar molecule is considered to be 100 % ionic then it will carry a magnitude of charge q :

q = 4.803 × 10^-10 esu .

Or q = 1.602 * 10 ^-19 C

This charge is equal to the charge on an electron.

Hence in order to calculate dipole moment assuming polar molecule 100 % ionic we will multiply this charge with bond length of the molecule:

μ = q × d  

 

RELATION BETWEEN ESU AND COULOMBS

 

One esu cm  = 3.335 * 10 ^-30 C m

 

 DIPOLE MOMENT NUMERICAL

1)In polar HCl molecule both H-atom and Cl-atom carry  8.08 × 10^-11 esu of magnitude of charges . The bond length between them is 1.275 × 10^-8 cm. Calculate dipole moment of HCl molecule.

Ans :- 1.03 * 10^-18  esu cm

 

 or 

 

1.03 * 10^-18   esu cm  × 1D/1.0 * 10^-18  esu cm  

 

=  1.03 D 

 

2) In polar LiH molecule the magnitude of charges on Li-atom and H-atom is 3.69 × 10^-10 esu . The bond length between them is 1.596 × 10^-8 cm. Calculate dipole moment of LiH molecule.

Ans :- 5.89 D

3) The magnitude of charges in KCl is 3.85 × 10^-10 esu and bond length is 2.6 × 10^-8 cm. Calculate dipole moment of KCl.

Ans :- 10 D

 

4) A covalent molecules A– B has a bond length of 150 pm. Calculate the dipole moment of this molecule if it were 100 % ionic.

Ans :- 7.2 D  

5) The observed dipole moment of a molecule A—B is 4.5 D and its calculated dipole moment is 7.2 D if it were 100 % ionic. Calculate its percentage ionic character? 

 

Ans= 62.5 %

 

6) The observed dipole moment of KCl is 10 D. The inter-atomic distance between K+ and Cl- in this molecule is 2.6 * 10^-8 cm . What will be dipole moment of KCl molecule if each atom carries opposite charges of one unit. Calculate the percentage ionic character of KCl.

Ans :- 80 %  

7) The observed dipole moment of LiH is 5.89 D. The inter-nuclear distance between Li and H in the molecule is 1.596 * 10^-8 cm . Calculate percentage ionic character. 

If LiH were 100% ionic then each end would carry charge equal to one unit = 1.602 ×10^-19 C , then its dipole moment would be :

μ = q × d  

=  1.602 ×10^-19 C  * 1.596 * 10^-10 m

 

= 2.577 * 10^-29 C m

 

Now its observed Dipole moment which  is given to us   = 1.964 * 10^-29 C m

 

Hence in order to find out percentage ionic character :

 

% ionic character = Observed dipole moment/Ionic dipole moment *100

 

= 1.964 * 10^-29 C m/2.577 * 10^-29 C m *100

 

Ans :- 76.8 %

 

 

8) The observed dipole moment of HCl is 1.03D. The bond length of HCl molecule is 1.275 * 10^-8 cm . Calculate the percentage ionic character of HCl.

If HCl were 100% ionic then each end would carry charge equal to one unit = 4.8 ×10^-10 esu , then its dipole moment would be :

μ = q × d  

=  4.8 ×10^-10 esu  * 1.275 * 10^-8 cm

 

= 6.12 * 10^-18 esu cm or 6.12 D

 

Now its observed Dipole moment which  is given to us  = 1.03 D 

 

Hence in order to find out percentage ionic character :

 

% ionic character = Observed dipole moment/Ionic dipole moment *100

 

= 1.03D/6.12 D *100

 

= 16.83 % Ans

9) The observed dipole moment of a molecule X—Y is 1.45 D and its bond length is 1.654 * 10^-8 cm .Calculate percentage ionic character in this bond.

Ans :- 18.3 %

 

DIPOLE MOMENT OF POLY ATOMIC  MOLECULES

For a molecule having only one bond , the dipole moment of the bond will be the dipole moment of the molecule. When a molecule consists of three or more atoms then the net dipole moment of the molecule is given by vectorial sum of individual dipole moments. Hence in order to find out the resultant or net dipole moment of the molecule it is necessary to know its geometry. For example  lets find out the dipole moment of CO2.

There are two C=O bonds in CO2 .

The dipole moment of each C=O bond is 2.3D.

CO2 molecule has linear shape or geometry.

The net dipole moment of CO2 will be given by the vectorial sum of dipole moments of two C=O bonds