CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 3

Periodic Classification Of Elements

##### Periodicity In Chemical Properties

CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 3

Periodic Classification Of Elements

Classification means identifying similar species and grouping them together.

Need for Classification – There are 118 elements and each of them has innumerable compounds. It is very difficult to study and understand all the elements and their compounds individually. Hence there is a need for classification or grouping of elements based on some common properties, so that the study of elements and their compounds could be made easy.

Lavoisier divided elements into two main types known as metals and non-metals.

For some groups of three elements,

• When the three elements in a triad were written in the order of increasing atomic masses; the atomic mass of the middle element was roughly the average of the atomic masses of the other two elements.

For example atomic masses of Li, Na and K are respectively 7, 23 and 39, thus the mean of atomic masses of 1st and 3rd element is equal to the mass of the second one.

 Group A Elements Atomic Masses Group B Elements Atomic Masses Group C Elements Atomic Masses Li 7 Ca 40 Cl 35.5 Na 23 Sr 88 Br 80 K 39 Ba 137 I 127

He could identify only a few such triads and so the law could not gain importance. In the triad of Fe, Co, Ni, all the three elements have a nearly equal atomic mass and thus do not follow the above law.

Newland’s Law of Octaves

According to this law

• Every eighth element has properties which are a repetition of those of the first, if arranged in order of increasing atomic mass, like the eight notes of musical scale.
 sa (do) re (re) ga (mi) ma (fa) pa (so) da (la) ni (ti) H Li Be B C N O F Na Mg Al Si P S Cl K Ca Cr Ti Mn Fe Co, Ni Cu Zn Y In As Se Br Rb Sr Ce, La Zr — —

Drawback of Newland’s Law of Octaves:

• It was found that the Law of Octaves was applicable only upto calcium. After calcium every eighth element did not possess properties similar to that of the first.

• It was assumed by Newland that only 56 elements existed in nature and no more elements would be discovered in the future. But, later on, several new elements were discovered, whose properties did not fit into the Law of Octaves.

• In order to fit elements into his Table, Newland adjusted two elements in the same slot, but also put some unlike elements under the same note. Cobalt and nickel are in the same slot and these are placed in the same column as fluorine, chlorine and bromine which have very different properties than these elements. Iron, which resembles cobalt and nickel in properties, has been placed far away from these elements.

Thus, Newland’s Law of Octaves worked well with lighter elements only.

Mendeleev’s Periodic Table

Mendeleev arranged 63 elements, known at that time, in the periodic table. According to Mendeleev

• The properties of the elements are a periodic function of their atomic masses.

Basis of Classification

The formulae of the hydrides and oxides formed by an element were treated as the basic properties of an element for its classification.

He selected hydrogen and oxygen because they are very reactive and formed compounds with most elements. If the formula for hydride was RHx, then the element R will be placed in xth group.

Structure of Mendeleev’s Periodic Table

• In the periodic table, the elements are arranged in vertical rows, called groups and horizontal rows, called periods. There are eight groups indicated by Roman Numerals I, II, III, IV, V, VI, VII, VIII.

• The elements belonging to first seven groups have been divided into sub-groups designated as A and B on the basis of similarities. The elements that are present on the left hand side in each group constitute sub-group A while those on the right hand side form sub-group B. Group VIII consists of nine elements arranged in three triads.

• There are six periods (numbered 1, 2, 3, 4, 5 and 6). In order to accomodate more elements, the periods 4, 5, 6 are divided into two halves. The first half of the elements are placed in the upper left corners and the second half into lower right corners in each box.

• At some places the order of atomic mass was changed in order to justify the chemical and physical properties.

Achievements of Mendeleev’s Periodic Table

1. The arrangement of elements in groups and periods made the study of elements quite systematic. That is, if properties of one element in a particular group are known, those of the others in the group can be easily predicted.

2. Prediction of new elements and their properties - Many gaps were left in this table for undiscovered elements and properties of these elements could be predicted in advance from their expected position. Mendeleev named them by prefixing a Sanskrit numeral, Eka (one) to the name of preceding element in the same group. For instance, scandium, gallium and germanium, discovered later, have properties similar to Eka–boron, Eka–aluminium and Eka–silicon, respectively.

3. Correction of doubtful atomic masses: Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties.

4. When inert gases were discovered they were placed in a new group without disturbing the order.

Limitations of Mendeleev’s Classification

1. He could not assign a correct position to hydrogen in his periodic table, as the properties of hydrogen resemble both with alkali metals as well as with halogens.

2. The isotopes of the same element will be given different position if atomic mass is taken as basis. But the properties of isotopes of the same element are same.

3. The atomic masses do not increase in a regular manner in going from one element to the next. At some places the order of atomic masses in the table was changed in order to justify the chemical and physical nature. For example, cobalt (atomic mass 58.9) appeared before nickel (atomic mass 58.7).

4. Some dissimilar elements were grouped together while some similar elements were placed in different groups. For example Li, Na, K are grouped with Cu, Ag, Au and Cu and Hg are placed in different groups.

5. Group viii elements have been put in three triads without any justification.

6. Lanthanoids and Actinoids have not been provided any position in the table and put separately without any justification.

Modern Periodic Law

This law was given by Henry Moseley in 1913.

Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. A plot of $\sqrt{\mathbit{\nu }}$ (where ν is frequency of X-rays emitted) against atomic number (Z) gave a straight line and not the plot of $\sqrt{\mathbit{\nu }}$ vs atomic mass. He thereby concluded that the atomic number is a more fundamental property of an element than its atomic mass.

The law states,

• Properties of the elements are the periodic function of their atomic numbers.

Modern Periodic Table/Long Form of Periodic Table or Bohr’s Table

Moseley proposed the modern periodic table also known as long form of the periodic table.

1. There are 18 groups and 7 periods in the table. The groups are numbered 1, 2, 3 …. 18 from left to right and the periods from 1 to 7 from top to bottom

2. The elements belonging to a particular group make a family and usually named after the first member.

3. In a group all the elements contain the same number of valence electrons.

4. In a period all the elements contain the same number of shells, but as we move from left to right the number of valence shell electrons increases by one unit.

5. The maximum number of electrons that can be accommodated in a shell is 2n2 where ‘n’ is the number of the given shell from the nucleus.

Cause of periodicity

Periodicity may be defined as the repetition of the properties of the elements separated by certain definite gap of atomic numbers.

• The properties of elements are due to the periodic repetition of similar electronic configuration of elements as the atomic number increases.

Chemical bonding depends on the number of electrons in the valence shell of an element. Due to same number of electrons in the outermost shell of elements in the same group, the chemical properties are similar.

Since nucleus (mass) resides deep inside the atom, it does not take part in chemical reactions. On the other hand the electrons are outside, in the orbits, hence they interact with other atoms.

Nomenclature of Element with Z > 100

 Digit Name Abbreviation 0 nil n 1 un u 2 bi b 3 tri t 4 quad q 5 pent p 6 hex h 7 sept s 8 oct o 9 enn e
 Atomic Number Name according to IUPAC Symbol IUPAC Official Name IUPAC Symbol 101 Unnilunium Unu Mendelevium Md 102 Unnilbium Unb Nobelium No 103 Unniltrium Unt Lawrencium Lr 104 Unnilquadium Unq Rutherfordium Rf 105 Unnilpentium Unp Dubnium Db 106 Unnilhexium Unh Seaborgium Sg 107 Unnilseptium Uns Bohrium Bh 108 Unniloctium Uno Hassium Hs 109 Unnilennium Une Meitnerium Mt 110 Ununnillium Uun Darmstadtium Ds 111 Unununnium Uuu Rontgenium Rg 112 Ununbium Uub Copernicium Cn 113 Ununtrium Uut * – 114 Ununquadium Uuq Flerovium Fl 115 Ununpentium Uup * – 116 Ununhexium Uuh Livermorium Lv 117 Ununseptium Uus * – 118 Ununoctium Uuo * –

Trends in Modern Periodic Table

The trends observed in some important properties of the elements in moving down the group (from top to bottom of the table) and across a period (from left to right in a period) are discussed below.

Electronic Configurations of Elements and The Periodic Table

• An electron in an atom is characterised by a set of four quantum numbers (n, l, m, s), and the principal quantum number (n) defines the main energy level known as shell.

• The filling of electrons into different subshells/orbitals (s, p, d, f ) in an atom is done as per certain rules and order.

• The distribution of electrons into orbitals of an atom is called its electronic configuration.

• An element’s location in the Periodic Table reflects the quantum numbers of the last orbital filled.

Electronic Configurations across Periods

The period indicates the value of n for the outermost or valence shell. (Take examples)

Electronic Configurations in Groups

Elements can be classified into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. (Take examples)

s-block elements: Group-1 (Alkali metals) and Group-2 elements (Alkaline earth metals) which respectively have ns1 and ns2 outermost electronic configurations

p-block elements: They belongs to group 13 to 18. The outer most electronic configuration is ns2 np1-6.

Exception - He (1s2) is an s-block element but is positioned with the group 18 elements (ns2 np6) because it has completely filled valence shell and as a result, exhibits properties characteristic of other noble gases.

Group 15 – Pnictogen

Group 16 – Chalcogens

Group 17 – Halogens

Group 18 – Noble gases

d-block elements (Transition elements) are the elements of group 3 to 12 having outer electronic configuration (n-1)d1-10 ns1-2. Four transition series are 3d, 4d, 5d and 6d. The 6d series is incomplete.

##### f-block elements (Inner- transition Series)

Lanthanoids charecterised by the filling of 4f orbitals, are the elements following lanthanum from 58Ce to 71Lu. The are called 4f-inner transition series. Actinoids characterised by filling of 5f orbitals, are the elements following actinium from 70Th to 103Lr. Characteristic outer electronic configuration is (n-2)f1-14(n-1) d0-1ns2. Thesy are called 5f-inner transition series.

Noble Gases: The gaseous elements of group 18 are called noble gases. The general outermost electronic configuration of noble gases (except He) is ns2 np6. He exceptionally has 1s2 configuration. Thus the outermost shell of noble gases is completely filled.

• The elements after uranium are called Transuranium Elements.
• FILL THE TABLE

Metallic and non-metallic character

• In a period from left to right metallic character decreases while non-metallic character increases.

• In a group metallic character increases from top to bottom while non-metallic character decreases.

A zig-zag line separates metals from non-metals. The borderline elements – boron, silicon, germanium, arsenic, antimony, tellurium and polonium – are intermediate in properties and are called metalloids or semi-metals.

As the effective nuclear charge acting on the valence shell electrons increases across a period, the tendency to lose electrons will decrease. Down the group, the effective nuclear charge experienced by valence electrons is decreasing because the outermost electrons are farther away from the nucleus. Therefore, these can be lost easily. Hence metallic character decreases across a period and increases down a group.

These trends also help to predict the nature of oxides formed by the elements because it is known that the oxides of metals are basic and that of non-metals are acidic in general.

The distance between the centre of nucleus of an isolated atom to its outermost shell containing electrons, is called its atomic radius.

Half the inter-nuclear distance between two similar atoms of any element which are covalently bonded to each other by a single covalent bond is called covalent radius.

Half the inter-nuclear separation between two similar adjacent atoms belonging to the two neighbouring molecules of the same substance in the solid state is called the van der waals’radius of that atom.

Half the distance between the nuclei of the two adjacent metal atoms in a close packed lattice of the metal is called its metallic radius.

Trend

• The atomic radius generally decreases on moving from left to right along a period.

This is due to an increase in nuclear charge which tends to pull the electrons closer to the nucleus and reduces the size of the atom.

• Size of atoms of inert gases is the largest in the period due to inter electron repulsion. Also since no bonding is involved, van der Waal’s radius is taken into account.

• Atomic size generally increases in a group from top to bottom due to increase in number of shells.

Exceptions

 N O F 70pm 74pm 72pm
1. In case of N to O, inter electron repulsions outweigh the effect of increased nuclear charge, hence r(N) < r(O).

2. In case of O to F, nuclear charge increases by 1. 2 of the p-orbitals now have 2 electrons each, which repel each other. However the increased nuclear charge outweighs the repulsion, hence r(F) < r(O).

Cation is formed by removal of an electron from an atom.

Anion is formed by gain of an electron by an atom.

• The ionic radius is the distance between centres of cations and anions in ionic crystals.

• A cation is smaller than the parent atom, because number of protons is more than the number of electrons.

• An anion is anion is larger than the parent atom, because of inter electron repulsion.

Na+ (64 pm) → F- (136 pm)

Iso-Electronic Species

Atoms or ions which have same number of electrons are called iso-electronic species.

For example

 Species O2- F- Na+ Mg2+ No. of Protons 8 9 11 12 No. of Electrons 10 10 10 10 Size of Mg2 < Na+ < F- < O2-
• In case of iso-electronic species, the cation with greater positive charge has smaller radius but anion with greater negative charge has the larger radius.
##### Ionisation Enthalpy

The energy required to remove an electron from an isolated gaseous atom (X) in its ground state, is called ionization enthalpy.

M(g) + ΔiH1 → M+(g) + e-

The successive ionization enthalpies of loosing electrons are called 1st iH1), 2ndiH2) and 3rdiH3) ionization enthalpies.

M+(g) + ΔiH2 → M2+(g) + e-

The term ionization enthalpy is taken for the first ionization enthalpy.

SI unit is kJ mol-1.

Ionisation Enthalpy is always +ve.

Shielding Effect

Effective nuclear charge (Zeff) experienced by valance electrons will be less than the actual charge due to shielding or screening.

Shielding is more if the inner shells are completely filled. Hence Alkali (1st group) experience greater shielding and can give away valance electrons easily.

Shielding effect of s-orbital is more than that of p-orbital.

Trend

• Generally the ionization enthalpies follow the order (there are few exceptions),

(∆iH1)  <  (∆iH2 ) <  (∆iH3)

• The ionization enthalpy increases on moving from left to right in a period.

Exceptions

Be-B

Boron (Z=5) has slightly less electron enthalpy than that of Beryllium because shielding by s-orbital electrons is more than by p-orbital electrons due the shape and size.

N-O

In case nitrogen the 3 p electrons reside in different orbitals whereas in case of oxygen 2 of them share one p-orbital, resulting in inter electron repulsion. Hence ionization enthalpy of oxygen is smaller than that of nitrogen.

• The ionization enthalpy decreases on moving top to bottom in a group.
##### Electron Gain Enthalpy

The electron gain enthalpy ((∆egH) is the molar enthalpy change when an isolated gaseous atom or ion in its ground state adds an electron to form the corresponding anion.

X(g)  +  e- → X-(g) + ∆egH

The ∆egH may be positive or negative.

The successive values for the addition of second, third etc. electron, are called second (∆egH2), third (∆egH3) etc. electron gain enthalpies. For example,

X(g) + e- → X-(g),  ∆H = ∆egH1 is called first electron gain enthalpy

X(g) + e- → X2-(g),   ∆H = ∆egH2 is called second electron gain enthalpy

X2-(g) + e- → X3- (g)   ∆H = ∆egH3 is called third electron gain enthalpy

• The electron gain enthalpy increases (becomes more negative) from left to right in a period.

• The electron gain enthalpy decreases (becomes less negative) from top to bottom in a group.

Exceptions

egH(O) is less negative than ∆egH(S) and ∆egH(F) is less negative than ∆egH(Cl).

Explantion – The electrons added to O and F occupy 2p orbital which is smaller orbit and inter electron repulsion is higher, whereas in case of S and Cl the orbitals are 3p, the size of orbitals is much larger and inter electron repulsion is lesser.

Factors Affecting Electron Gain Enthalpy

• Attraction of electrons towards the nucleus

• Repulsion of electrons from each other

Electronegativity

The relative tendency of an atom to attract the shared electron pair of electrons towards itself is called electronegativity.

Electronegativity is not a measureable quantity. However there are numerical scales developed, e.g. Paling, Mullikan etc.)

Electronegativity of an element is not fixed. It varies depending on the element to chich it is bound.

It provides the nature of force between bonding atoms.

Electronegativity is directly related to that non-metallic properties of elements and inversely related to the metallic properties (or electropositivity) of elements.

Trend

• The electro-negativity increases from left to right in a period.

• The electro-negativity decreases from top to bottom in a group.
##### Valence Electrons

The electrons present in outermost shell are called valence electron.

• The number of valence electrons in a group is same, whereas the number of valence electrons within a period increases from left to right.

##### Periodicity In Chemical Properties

Valency or Oxidation State

The number of electrons which an atom loses or gains or shares with other atom to attain the noble gas configuration is termed as its valency.

• The valency first increases then decreases from left to right in a period.

• The valency remains constant from top to bottom in a group

• There are many elements which exhibit variable valence. This is particularly characteristic of transition elements and actinoids.

Anomalous Properties of Second Period Elements

The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other members of their respective group.

For example, Lithium and Beryllium form covalent compounds, whereas the other members of these groups form ionic compounds.

The behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.

Reasons of Anomalous Behaviour/Diagonal Relationship

• Their small size,

• High electro negativity,

• Non-availability of d-orbitals in their valence shell. Hence they show maximum valency of 4 whereas other members of the group can show valnecy greater than 4.
• The first member of each group of p-block elements displays greater ability to form pp-pp multiple bonds to itself (e.g. C=C, C≡C O=O, N≡N) and to other second period elements (e.g. C=O, C≡N, N=O) compared to subsequent elements in the group.

Periodic Trends and Chemical Reactivity

• All chemical and physical properties are a manifestation of the electronic configuration of elements

• The elements at the extreme left and right (barring the noble gases) are highly reactive due to high metallic and non-metallic character respectively.

• Metals form basic oxides, non-metals acidic and the elements in between form amphoteric oxides.