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CBSE Class 11 Chemistry Notes For Chapter 3 Mendeleev’s Periodic Table

Based on his periodic law, Mendeleev arranged the then-known elements in the form of a table (consisting of several rows and columns) which is known as Mendeleev’s periodic table. Mendeleev’s original periodic table (1871) contained only 63 elements known at that time. There were no places for inert gases because these were not discovered at the time of publication of the table.

Mendeleev, however, left several blank places in the table and predicted that there must be some unknown elements which would be discovered in due course of time. He even predicted their properties based on the properties of the adjacent elements. Mendeleev’s predictions were proved to be astonishingly correct when these elements were discovered later. Mendeleev’s table, now in use, is a modified version ofthe table originally designed by him. Important features of the modified form of Mendeleev’s periodic table are discussed below.

Periods and Groups:

In Mendeleev’s periodic table, the elements were arranged in the increasing order of their atomic weights (but in the modified form these were arranged in increasing order of their atomic numbers) into several horizontal rows. These horizontal rows were placed one below the other in such a way that chemically similar elements fell in the same vertical column. The horizontal rows are called periods and the vertical columns are called groups or families.

There is a gradual change in the properties of the elements with an increase in atomic mass across a period. However, elements belonging to the same group exhibit close chemical similarities. In the modern version of Mendeleev’s table, there are seven periods (1 to 7) and nine groups (I to VIII and 0). Gr-0 consists of the inert gases (Mendeleev’s original table did not contain this group)

CBSE Class 11 Chemistry Notes Chapter 3 Mendeleev’s Periodic Table

Main features of Mencleleov’s Periodic Table

The first period contains only 2 elements (II and He).

• This is called the shortest period. The second period contains only 8 dements (Li-Bc-B-C-NO-F-Nc), beginning with alkali metal Li and ending with Inert gas No. This Is called the first short period.
• The period also contains elements (Na-Mg-Al-Si-P-SCl-Ar), beginning with the alkali metal Na and ending with the Inert gas Ar.
• This is called the second period. The elements of these two short periods occur in nature in very large amounts and they typify the properties of all the other members of the group to which they belong. So they are called typical elements.
• The fourth period contains 18 elements. It begins with the alkali metal K and ends with the inert gas Kr.
• This period is called the first long period. The fourth period contains 10 additional elements than the second and third periods.
• These 10 elements (Sc to Zn) are called the transition elements. This period consists of two series (the even and the odd series).
• The fifth period also contains 18 elements.
• It begins with the alkali metal Rb and ends with the inert gas Xe. This period is called the second long period.
• 10 elements from Y to Cd are called transition elements. The fifth period also consists of two series (the even and the odd series).

The sixth period contains 32 elements, and so it is called the longest period. It begins with the alkali metal Cs and ends with the inert gas Rn. This period contains 10 transition elements (La and Hf to Hg) and 14 lanthanide elements (Ce to Lu ). These 14 elements are also called rare earth elements because these elements were believed to be present in nature in negligible amounts.

The sixth period also consists of two series (the even and the odd series). The seventh period may contain a maximum of 32 elements (beginning with Fr), but all the elements have not yet been discovered. Till now 28 elements have been discovered. So it is an incomplete period.

All elements of this period are radioactive. The elements from Francium (Fr) to Uranium (U) are naturally occurring, while the elements beyond uranium are man-made. In this period, the 14 elements beyond actinium, Ac [i.e., the elements from thorium (Th) to lawrencium (Lr) ] are called actinides, while the elements beyond uranium (U) are called transuranic elements.

Even and odd series: Elements belonging to each of the 4th, 5th and 6th periods are divided into two series:

The even and the odd series. The three even series begin with the alkali metals K, Rb and Cs, while the three odd series begin with the coinage metals Cu, Ag and Au respectively.

Subgroups:

Except for the Gr-VHI and Gr-0, each of the other groups (Gr-I to VII) is divided into two subgroups designated as ‘A’ and ‘B! In long periods (4th, 5th and 6th), the elements of the even series are placed in subgroup A and those of the odd series are placed in subgroup B.

In short periods (2nd and 3rd), elements of Gr-I and Gr-II are placed in subgroup-A, while those of the other groups are placed in subgroup-B. Within the same group, the properties of the elements of subgroups and B are altogether different, except for their valencies. However, elements of the same subgroup exhibit more or less similar properties.

For example: 

Alkali metals of Gr-IA are closely alike. However, Gr-IA metals differ remarkably from the coinage metals of Gr-IB (Cu, Ag and Au), although they have a common valency of ‘1’.

Additional pieces of information about groups and sub-groups:

• Elements of subgroup A are more electropositive than those of subgroup B.
• For example, Gr-VIIA elements (Mn, Tc, Re) are electropositive, while Gr-VIIB elements (F, Cl, Br and I) are electronegative characters
• .Gr-VIII has no subgroups. It contains a total of 9 elements, belonging to periods 4, 5 and 6.
• These nine elements—Fe, Co, Ni (period-4); Ru, Rh, Pd (period-5) and Os, Ir, Pt (period-6) are arranged in this manner due to similarity in their properties and their, atomic weights are also close to v each other. Each grip, of three elements, is called Mendeleev’s triad elements. Mendeleev coined the term ‘transitional element’ for these elements.

Gr-0 has no subgroups. It contains inert gases:

• He, Ne, Ar, Kr, Xe and Rn. These elements are chemically inert and do not exhibit any tendency to combine with other elements.
• So they are zero-valent elements and placed in Gr-‘0′. This group acts as a bridge between highly electronegative halogens (VIIB) and highly electropositive alkali metals (IA).
• Due to their similarity in chemical properties, La and 14 elements from Ce -Lu are placed together in Gr-3A of the 6th period.
• The 14 elements from Ce to Lu are called lanthanoids. For similar reasons, Ac and 14 elements from Th-Lr are placed together in Gr-3A ofthe 7th period. The 14 elements from Th to Lr are ( called actinoids.

Importance & usefulness of Mendeleev’s periodic table Systematic Study of the elements:

• Mendeleev, for the first time, arranged a vast number of elements in such a way that the elements with similar chemical properties are placed in the same group.
• This made the study of elements quite systematic because if the properties of one element (and its compounds) in a particular group are known, then the properties of the rest of the elements (and their compounds) can be predicted.

Prediction of new elements:

• Mendeleev left some gaps in the periodic table to accommodate new elements to be discovered in future. he even predicted the properties of those unknown elements based on their positions in the table.
• When these elements were discovered, their properties were found to be similar as predicted By Mendeleev. For example, Mendeleev left two vacant places below b and al in gr-3 and one vacant place below.
• He named those elements eka-boron, ca-aluminium and ca-silicon respectively as he predicted that the properties of these elements would be similar to that of boron, aluminium & silicon.
• In 1075, de Baisbaudron discovered eka-aluminium and named it gallium. In 1079, n. l.
• Nilson discovered eka-boron and named it scandium. In 1006, Winkler discovered eka-silicon and named it as germanium.

Mendeleev’s Periodic Table Class 11 Notes NCERT

It was observed that these newly discovered elements had properties similar to those already predicted by Mendeleev before their discovery.

Correction of doubtful atomic weights:

• With the help of Mendeleev’s periodic table, doubtful atomic weights of some elements are rectified.
• For example: Be was assigned an atomic weight of 13.5 based on its equivalent weight (4.5) and valency (wrongly taken as ‘3’ because Be had certain similarities with trivalent metal Al).
• With an atomic weight of 13.5, Be should be placed between carbon (At. weight 12) and nitrogen (At. weight 14).
• But no vacant place was available In between C and N. Mendeleev asserted that Be must be bivalent because of its similarity with Mg, Ca etc. Thus he corrected its atomic weight as 4.5 X 2 = 9.0.

Defects Of Mendeleev’s Periodic Table

Discrepancy or anomaly in periodicity:

• Mendeleev arranged the elements in increasing order of their atomic weights.
• But he violated this principle in certain cases to give appropriate positions to some elements based on their properties i.e., he laid more emphasis on the properties of those elements rather than their atomic weights.
• In the following four pairs of elements, elements with higher atomic weight have been placed before elements with lower atomic weight.

Position of hydrogen:

• The controversial position of hydrogen in the periodic table also hints at discrepancies within the table.
• Like the alkali metals ofGr-IA, it exhibits univalency, high reactivity, electropositive character, strong affinity for non-metals and reducing character.
• On the other hand, like the halogens of Gr- VIIB, it has atomicity, high ionisation energy, non-metallic character, existence in the gaseous state at normal temperature and pressure ability to combine with milk-forming hydrides (Example: Nall).
• Since hydrogen exhibits similarities as well as dissimilarities with both the alkali metals and the halogens, the placement of hydrogen in any one of these two groups will naturally create difficulties.
• So it is desirable to fix a separate position for hydrogen in Mendeleev’s periodic table.

Placement of similar elements in different groups and dissimilar elements in the same group:

In some cases, elements with almost similar properties have been placed in different groups.

Example:

• Cu and Hg resemble in properties but Cu is Gr-1B while Hg has been placed in Gr-IIB.
• Likewise, elements like Ba (Gr-2A) and Pb (Gr – 4B) have been placed in different groups.
• Again, some elements with dissimilar properties have been placed in the same group.
• Highly reactive alkali metals such as Li, Na, K etc., have been placed together with almost inactive coinage metals such as Cu, Ag and Au in Gr-1.
• Likewise, Mn, Te and Re having no similarity with F, Cl, Br etc. have been placed together.

Lack of separate positions for Gr-8 elements:

• No proper place has been allotted to nine elements belonging to Gr-8 although they have many similarities in properties.
• These are arranged in three triads, one in each of the 4th (Fe, Co, Ni ), 5th(Ru, Rh, Pb) and (Os, Ir, Pt)periods.

Lack of suitable positions for Lanthanoids and Actinoids:

• The 14 elements following La from Ce to Lu (lanthanoids) and the 14 elements following Ac from Th to Lr (actinoids) have not been allotted separate positions in the main skeleton of the periodic table.
• They have been placed in two separate rows at the bottom of the table. Besides, the number of elements in the lanthanoid and actinoid series cannot be determined from Mandeleev’s periodic table.

Position of isotopes:

• Isotopes of an element have different atomic weights. So they should be placed at different positions in the periodic table.
• However, all the isotopes of any specific element are placed in a single position (i.e., same period and same group)in Mendeleev’s periodic table.

Mendeleev’s Periodic Table Chapter 3 NCERT Notes

Moseley’s experiment:

The atomic number determines the fundamental property of an element. In 1913, Moseley measured the frequencies of X-rays emitted by different metals when bombarded with high-speed electrons.

• He observed that the frequencies ofthe prominent X-rays emitted by different metals were different but for each metal, there was a fixed value.
• He observed further that the square root of the frequency (v) of the X-rays emitted by a metal was proportional to the atomic number but not to the atomic mass of the metal, Le., Jv = a(Z- b) where ‘a’ is the proportionality constant and is a constant for all the lines in a given series of X-rays.
• Thus a plot of Tv vs Z gave a straight line but a plot of Jv vs atomic mass does not bear such a linear relationship.

This led Moseley to conclude that atomic number was a better fundamental property of an element than atomic mass.

Modification Of Mendeleev’s Periodic Law

• Mendeleev regarded atomic weight as the fundamental property of an element and so he considered atomic weight as the basis of periodic classification of elements.
• But Moseley, from his experimental results, showed clearly that atomic number is a better fundamental property of an element than its atomic weight.
• This led Moseley to suggest that atomic number (Z) should be the basis of the classification of elements. This gave birth to a new periodic law known as the modem periodic law.

Modern Periodic Law:

• The physical and chemical properties ofthe elements are aperiodic functions of their atomic numbers.
• This implies that, if elements are arranged in order of increasing atomic numbers, the elements with similar chemical properties are repeated after certain regular intervals.

Rectification of the discrepancy in periodicity with the help of modern periodic law:

• The original periodic law, based on atomic weight, was violated in the case of four pairs of elements [(Ar, K), (Co, Ni), (Te, I), (Th, Pa)].
• In each pair, an element with a higher atomic weight is placed before the element having a lower atomic weight.
• In the modern periodic table (based on the atomic number), this discrepancy disappears because the atomic numbers of K, Ni, I and Pa are greater than those of Ar, Co, Te and Th respectively.
• Placement of all the isotopes of any specific element in the same position of the periodic table is quite justified as the isotopes of elements have the same atomic number although they have different atomic weights.

Theoretical justification of modern periodic law:

• Only nuclear electrons (or more specifically valence shell electrons) take part in chemical reactions, while the atomic nucleus remains unaffected.
• So it is understandable that the properties of the elements will depend upon their atomic numbers (equal to the number of electrons) rather than their atomic weight or mass numbers (equal to the total number of protons and neutrons).

Periodicity of elements

The periodic repetition of elements having similar properties after certain regular intervals when the elements are arranged in the increasing order of their atomic numbers is called periodicity.

Cause of periodicity:

• According to modern periodicals, there is a repetition of properties of the elements after certain regular intervals when they are arranged in order of their increasing atomic numbers.
• Again from a close study of electronic configurations of various elements, it is observed that with successive increases in atomic number, there occurs a repetition of similar outermost shell electronic configuration (valence shell electronic configuration) after certain regular intervals.
• By correlating these two observations, it can be concluded that periodicity in properties is due to the recurrence of similar valence shell electronic configuration after certain regular intervals when the elements are arranged in order of increasing atomic numbers.

NCERT Class 11 Chemistry Mendeleev’s Periodic Table Solutions

This can be illustrated by the following examples.

1. Elements of Gr-1A have outermost electronic configuration ns1 (where n = outermost principal quantum number).
2. These elements exhibit similar chemical properties due to their similarity in the valence shell electronic configuration.
3. Elements of Gr-7B have outermost electronic configuration ns²np⁵.
4. All the halogens exhibit similar chemical properties due to their similarity in valence shell electronic configuration.
5. Inert gases belonging to this group possess similar chemical properties because they have similar valence shell electronic configurations (ns²np⁶).
6. It should be noted that properties of elements get repeated only after intervals of 2, 8,18 or 32 in the atomic numbers of the elements because similar electronic configurations recur only after such intervals.
7. The numbers 2, 8, 18 and 32 are called magic numbers. These numbers are very useful in locating elements with similar properties.

Moderntableor Long Form Of Periodic Table Bohr’s Table:

This is an improved form of the periodic table based on modern periodic law. It is also called Bohr’s table since it follows Bohr’s scheme for the classification of the element based on the outermost electronic configuration governed by the Aufbau principle. It consists of periods and 18 groups.

CBSE Class 11 Chemistry Notes For Chapter 3 Long Form Of Periodic Table

Description of periods:

• Like Mendeleev’s modified table, it also consists of numbers from as1 to 7 from top to bottom.
• The period number is equal to the value of i.e., the principal quantum number corresponding to the outermost shell of the atoms of the elements belonging to that period.
• Each period begins with the filling of electrons in a new energy level. Several elements in each period are twice the total number of atomic orbitals available in the energy level that are being filled.

First period:

This period begins with the filling of the first energy level (n = 1). Since the first shell has only one orbital [i.e., Is), which can accommodate a maximum of two electrons, there can be only two elements in the first period. These are hydrogen (Is¹) and helium (Is²).

Second period:

• It starts with the filling of the second energy level (n = 2). Since the second shell contains four orbitals (one 2s and three 2p), it can accommodate a maximum of (2 × 4) = 8 electrons. So, there are eight elements in the second period.
• It begins with lithium (Li) in which 1 electron enters the 2s -orbital (₃Li: 2s¹) and ends up with neon (Ne) in which the second shell gets filled (₁₀Ne: 2s²2p⁶).

Third period:

• The third period begins with the filling of the third energy level (n = 3). This energy level contains nine orbitals (one 3s, three 3p and five 3d).
• According to the Aufbau principle, 3d -orbitals will be filled up only after filling the 4s -orbital.
• Consequently, the third period involves filling only four orbitals (one 3s and 3p ) which can accommodate a maximum of (2 × 4) = 8 electrons. So, there are 8 elements in the third period from ₁₁Na (3s¹) to ₁₈Ar (3s²3p⁶).

Long Form of Periodic Table Class 11 Notes NCERT

Fourth period:

• This period corresponds to the filling of the fourth energy level (n = 4). Out of 4s,4p,4d and 4f-orbitals belonging to this shell, filling of 4d -and 4f-orbitals does not occur in this period since their energies are higher than that of even 5s -orbital.
• It must however be remembered that after filling the 4s -orbital, the filling of five 3d -orbitals begins since the energy of the -orbital is greater than that of the 4s orbital but less than that of the 4p -orbital.
• So the fourth period involves filling of or 9 Orbitals (one 4s, five 3d and three 4p), which can accommodate (2 × 9) = 18 electrons.
• Therefore, the fourth period contains 18 elements from potassium (₁₉K: 4s¹) to krypton (₃₆Kr: 4s²3d¹⁰ 4p⁶).
• This period contains 10 elements more than the third period corresponding to filling off 3d -orbitals. These 10 elements [₂₁Sc(3d¹4s²) to ₃₀Zn(3d¹⁰4s²)] are called the first series of transition elements.

Fifth period:

• The fifth period corresponds to the filling of electrons in the fifth energy level (n = 5). Like the fourth period, it also accommodates 18 elements since only nine orbitals (one 5s, five 4d and three 5p) are available for filling with electrons.
• It starts with rubidium in which one electron enters 5s -orbital (₃₇Rb: 5s¹) and ends up with xenon in which the filling of 5p -orbital is complete (₅₄Xe: 5s²4d¹⁰5p⁶). 10 elements from ₃₉Y(5s²4d¹) to ₄₈Cd(5s²4d¹⁰) corresponding to filling of five 4d orbitals are called second series oftransition elements.

Sixth period:

• The sixth period corresponds to the filling of electrons in the sixth energy level (n = 6).
• This period involves the filling of sixteen orbitals (one 6s, seven4f, five 5d and three 6p) which can accommodate a maximum of (2 × 16) = 32 electrons. So there are 32 elements in the sixth period.
• It begins with caesium (Cs) in which one electron enters 6s -orbital (₅₅Cs: 6s1) and ends up with radon in which filling of 6p -orbital is complete (₈₆Rn: 4f¹⁴5d¹⁰6s²6p⁶).
• Filling up of 4f-orbitals begins with cerium(₅₈Ce) and ends with lutetium (₇₁Lu). These 14 elements constitute the first inner-transition series, also called lanthanoids or rare earth elements.
• These are separated from the main frame periodic table and are placed in a horizontal row at the bottom of the table Again, 10 elements lanthanum (₅₇La), hafnium (₇₂Hf) to mercury (₈₀Hg), corresponding to successive filling of (10) 5d -orbitals, constitute the third transition series.

Seventh period:

• This period corresponds to the filling of electrons in the seventh energy level (n = 7). Like the sixth period, it is expected to accommodate 32 elements corresponding to the filling of 16 orbitals (one 7s, seven5f, five 6d and three 7p ).
• However, at present this period is incomplete consisting of 28 elements. The last element of this period will have an atomic number of of118 and will position the inert gas family.
• In this period, after filling of 7s -orbital [₈₇Fr: 7s¹ and ₈₈Ra: 7s²], the next two electrons enter the 6rf-orbital (this is against the Aufbau principle) corresponding to the elements 8gAc and goTh.
• Thereafter, the filling up of 57- orbital begins with ₉₁Pa and ends from Th to Lr are commonly called actinoids, which constitute the second Inner-transition series, Although Th does not contain any electron.

‘if orbital, it is considered to be a member of the actinoid series. Like lanthanoids, 14 members of the actinoid series are placed separately in a horizontal row at the hotter of the periodic table.

Number of elements in different periods and types of orbitals being filled up:

Description of groups:

Each of the 18 groups in the long form of the periodic table consists of many elements whose atoms have similar electronic configurations ofthe outermost shell (valence shell). The members of each group exhibit similar properties. Successive members in a group are separated by magic numbers of either 8 18 or 32. According to the recommendation of IUPAC (1988), the groups are numbered from 1 to 18.

Designations of these groups in different systems are presented in the following table

Designations of different groups [including outer electronic configuration:

Specific names of the elements of certain groups:

Long form of the periodic table

The superiority of the long form of the periodic table over Mendeleev’s periodic table

• In the long form of the periodic table, it is easy to remember and reproduce all the elements more easily in a sequence of atomic numbers.
• It relates the positions of the elements in the table to their electronic configurations more dearly.
• Gradual change In properties along the periods or similarity in properties along the groups can be interpreted by considering electronic configurations of the elements.

For example:

Elements of the same group exhibit marked similarities due to similar outer electronic configurations.

• Splitting the periodic table into s-,p-, cl- and f-blocks has made the study of the elements easier.
• The maximum capacity of each period to accommodate a specific number of elements is related to the capacity of different electronic shells to accommodate the maximum number of electrons.
• Due to the elimination of sub-groups, dissimilar elements do not fall In the same group. Each vertical column (group) accommodates only those elements which have similar outer electronic configurations, thereby, showing similar properties.
• Group-8 elements (involving triads) of Mendeleev’s table, have been provided separate positions in groups 8, 9 and 10.
• Elements belonging to 1, 2, and 13-17 groups are classified as representative elements, while those belonging to 3-12 groups are classified as transition elements.

Elements are further classified as active metals (belonging to groups 1 and 2), heavy metals (belonging to groups 3-12) and non-metals (belonging to groups 13-18).

• Transition elements of the 4th, 5th, 6th and 7th periods are assigned appropriate positions in this periodic table.
• The completion of each period with an inert gas element is more logical. In a period as the atomic number increases, the quantum shells are gradually filled up until an inert gas configuration is achieved at group 18.
• It thus eliminates the even and odd series belonging to the periods 4, 5 and 6 of Defects of the long form of the periodic table.

If the Position of hydrogen:

The position of hydrogen is not settled. It can be placed along with alkali metals in group 1 or with halogens in group 17, as it resembles the alkali metals as well as the halogens.

Position of helium:

Based on electronic configuration, He (Is²) should be placed in group 2. But, it is placed in group 18 along with the p -p-block elements. No other p -p-block element has the electronic configuration ofthe type ns².

Position of lanthanoids and actinoids:

Lanthanoids and actinoids have not been accommodated in the main frame of the periodic table.

Position of isotopes:

Isotopeshavenotgot separate places. Properties of isotopes of heavier elements are more or less the same, but isotopes of lighter elements differ drastically in their physical, kinetic and thermodynamic properties. So it is not desirable to place the isotopes in the same position. Despite these limitations, the long form of the periodic table, based on electronic configurations, is much more scientific and thus finds extensive use.

Classification Of Elements Into Different Blocks

Amount In tin long form of the periodic table it has been divided Into four blocks viz, s -block, p -block, d -block and f- block. It Is done based on the nature of atomic orbitals into which the Inst electron (the differentiating electron) gets accommodated. Elements of s and p -blocks except Inert gases, are called representative elements, and d -block elements, on the other hand, are called transition elements.

S -block elements

Elements In which the last electron enters the -subshell of their outermost energy level (n) are called s – s-block elements.

Since s- s-subshell can accommodate a maximum of 2 electrons, only two groups are included in this block.

Elements of group-1 (alkali metals) and group-2 (alkaline earth metals) which have outermost electronic configurations rui1 and ns2 respectively constitute the s -block. This block is situated at the extreme left portion of the periodic table.

Outermost electronic configuration of s-block elements:

ns¹‾² Inert gas element, helium (He, Is²) Is also considered as an s -s-block element.

Characteristics of s-block elements:

In the case of these elements, all shells except the outermost one, are filled with electrons.

• Except for H, all other elements of this block are metals. Because of their low ionisation potential, these metals are very reactive and do not occur freely in nature.
• All the metals of this block are good reducing agents because of the value of ionisation potential.
• They are good conductors of heat and electricity.

They are soft metals. They have low melting points, boiling points and low densities as compared to the adjacent transition elements.

• Cations of group-IA and group-DA elements are diamagnetic and colourless since their orbitals do not contain odd electrons.
• Except for Be and Mg, -block elements impart specific colour to the flame (flame test).
• Salts of these elements except dichromate, permanganate arid chromate, are colourless.
• Compounds of these elements are mainly ionic (only Li and Be can form covalent compounds in many cases).

They form stable oxides with oxygen (Na₂O, CaO), produce chlorides with chlorine (NaCl, CaC₂) and also form salt-like hydrides (NaH, KH, CaH₂) with hydrogen.

• Hydroxides of these elements [except Ca(OH)₂, Mg(OH)₂ and Be(OH)₂] are soluble in water at ordinary temperature.
• The non-luminous flame of the Bunsen burner is rich in electrons. During the flame test, metal ions are converted into short-lived neutral atoms by accepting electrons from the flame.
• Valence electrons of these neutral atoms absorb energy from the flame and get promoted to higher energy levels.

When the electrons return to lower energy levels, the absorbed energy is emitted in the form of radiation of different wavelengths in the visible range and as a consequence, different colours, depending.

• The wavelengths of emitted light radiations are Imparted to the flame.
• For instance, the generation of golden-yellow flame during the flame test with sodium salt is due to the transition of one electron of Na -atom from 3s -orbital to 3p -orbital and its return to 3sorbltal after a very short interval.
• The ionisation potentials of Be and Mg are sufficiently high because of their smaller size.
• So, their electrons cannot be excited to higher energy levels by absorbing energy from the flame. As a result, they fail to respond to the flame test.

Class 11 Chemistry Chapter 3 Long Form of Periodic Table Notes

P -block elements

Elements in which the last electron enters p -subshell of their outermost energy level (n) are called p-block elements.

Since p -p-subshell can accommodate a maximum of six electrons, 6 groups are included in this block.

Elements of group-13, 14, 15, 16, 17 and 18 (excluding helium) having the outermost electronic configurations:

ns²np1, ns²np², ns²np³, ns⁶np⁴, ns²np⁵ and ns²np⁶ respectively, constitute the p -block. This block is situated at the extreme right portion of the periodic table

The elements of group 18 are balled noble gases or inert gases. They have the shell electronic configuration ns²np⁶ in the outermost shell. Group-17 elements are called halogens (salt producers)’ and group-16 elements are called chalcogens (ore-forming).

These two groups of elements have high electron-gain enthalpies (high negative values of ΔH) and hence readily accept one or two electrons respectively to attain the stable noble gas configuration thereby forming negative and negative anions respectively.

The elements of s- and p -blocks taken together are called representative normal or main group elements.

Outermost electronic config. of-flock elements: ns²np⁶

Based on electronic configuration, helium (Is2) should not be considered as a p -p-block element, but from the standpoint of its chemical inertness (owing to the presence of a filled valence shell) it is justified to place group-18 along with other noble gas elements.

Characteristics of block elements:

Ionisation enthalpies of p -p-block elements are higher as compared to those of —-block elements. Most of the p -p-block elements are non-metals, some are metals and a few others are metalloids and inert gases.

• The metallic character increases from top to bottom within a group non-metallic character increases from left to right along a period.
• Hence, metals exist at the bottom ofthe left side ofthe p -block whereas non-metals lie at the top of the right ofthe p -block. Metalloids (B, Si, Ge, As, Sb ) stand midway between them.
• The oxidising character of -p-block elements increases from left to right in a period and the reducing character increases from top to bottom in a group.
• Most of them form covalent compounds, although ionic character increases continually down the group.
• Elements of this block are non-conductors of heat and electricity, except metals and graphite.

Elements of this block are mostly electronegative.

• Some of them exhibit variable oxidation states or valence states. Oxidation states may be both positive and negative.
• Non-metallic elements of this block form acidic oxides. They can form both coloured and colourless compounds.
• 4th, 5th and 6th-period elements can form complex compounds by coordinate covalency due to the presence of vacant d-orbitals.
• Some of the p -block elements Fe.g.-Q Si, P, S, B, Ge, Sn, As etc.) show the phenomenon of allotropy.

Carnation property is shown by some- block elements (Example: C, Si, Ge, N, S etc.)

d-block elements (Transition elements) Elements in which the last electron enters d -the subshell of their penultimate shell (i.e., the second from the outermost) are called d -block elements, d -subshell can accommodate a maximum of 10 electrons.

Therefore, ten groups are included in d -the block. Elements of group-3 [(n-1)d¹ns²], 4, 5, 6, 7, 8, 9, 10, 11 and 12 [(n- 1)d¹⁰ns²] constitute the d -block.

Atoms of the elements belonging to these groups usually contain or 2 (sometimes zero) electrons in the s -s-orbital of their outermost shell (i.e., n -th shell), while the differentiating electrons are being progressively filled in, one at a time, in the d -subshell of their penultimate shell [i.e., (n- 1) -th shell].

Electronic configuration of outer shell: (n-1) d^1-10ns^1-2

This block is situated in between s -and p -blocks. D -block elements form a bridge between the chemically active metals of groups 2 on one side and the less reactive elements of groups 13 and 14 on the other side.

Hence, d -block elements are called transition elements These elements have been divided into four series called the first, second, third and fourth transition series.

First transition series or 3d-series:

The first transition series consists of 10 elements, belonging to the 4th period, from scandium ₂₁Sc) to zinc (₃₀Zn) in which 3d -orbitals are being progressively filled in. Zn is not a transition element.

Second transition series or 4d-series:

The second transition series also consists of 10 elements, belonging to the 5th period, from yttrium (₃₉Y) to cadmium (₄₈Cd) in which 4d -orbitals are being progressively filled in. Cd is not a transition element.

Third transition series or 5d-series:

Third transition series also consists of 10 elements, belonging to the 6th period. These are lanthanum (₅₇La) and elements from hafnium (₇₂Hf) to mercury (₈₀Hg). In all these elements, 5d orbitals are being successively filled in. Hg is not a transition element.

Fourth transition series (6d-series):

The fourth transition series is formed from a part of the seventh period and it contains 10 elements. These, are actinium (₈₉Ac) and elements from rutherfordium (104Rf) to ununbium (₁₁₂Uub), in which 6d orbitals are being progressively filled in.

All d -block elements are not transition elements. Only those d-block elements in which atoms in their ground state or any stable oxidation state contain incompletely filled subshells are considered transition elements.

Characteristics of d-block elements:

• All d -block elements are metal. Their ionisation potential lies mid-way between those of s and p-block elements.
• Elements of the 5d series (especially Pt. Au and Hg) are inert under ordinary conditions. Thus, they are known as noble metals.
• Elements of this block exhibit variable oxidation states and valencies because ofthe presence of partially filled d orbitals in their atoms, ns -electrons and different numbers of(n-1)d electrons participate in bonding at the time of reaction with atoms of other elements.
• They are solids (except Hg), hard and have high melting and boiling points. They can form both ionic and covalent compounds.
• They exhibit paramagnetic character due to the presence of one or more unpaired electrons in their atoms or ions (exception-Sc³⁺, Ti⁴⁺, Zn²⁺, and Cu⁺ which do not contain odd electrons and are diamagnetic). Fe and Co can be converted into magnets and hence, they are ferromagnetic.
• They frequently form coloured ions in solids or solutions. With the change in their oxidation numbers, there also occurs a change in the colour ofthe formed ions.
• d -block elements exhibit a very distinctive property of forming coloured coordination complexes.
• This tendency may be ascribed to the small size of the atom or ion, a high nuclear charge of the ion and the presence of an incomplete d -d-orbital, capable of accepting electrons from the ligands.
• They are less electropositive than s -s-block elements but more electropositive than p -p-block elements.
• Several transition metals such as Cr, Mn, Fe, Co, Ni, Cu etc., and their compounds are used as catalysts. Many transition metals form alloys.

F-block elements (Inner-transition elements)

Elements in which the differentiating electron (i.e., the last electron) enters the f-subshell of their antepenultimate shell (i.e., the 3rd from the outermost) are called f-block elements.

All the F-block elements belong to group 3 (3B) of the periodic table. In these elements, s -orbital last shell (n) is filled, d -subshell of the penultimate shell [i.e., (n- 1) th shell] contains 0 or 1 electron, while f-subshell of the antepenultimate shell [i.e., (n-2)th shell] gets progressively filled in.

General electronic config.: (n-2)f^1-14(n-1)d^0-1ns²

Lanthanide series or 4f-Series:

• The first series follows lanthanum (La) in the 6th period and consists of 14 elements from cerium (₅₈Ce) to lutetium (₇₁Lu).
• These 14 elements are collectively called lanthanoids because they closely resemble lanthanum in their properties.
• These are also called rare-earth elements since most of these elements occur in very small amounts in the earth’s crust.

Actinoid series or 5f-series:

The second series follows actinium (sgAc) in the 7th period and consists of 14 elements from thorium (goTh) to lawrencium (103Lr).

These 14 elements are collectively called actinoids because they closely resemble actinium in their properties.

All the actinoids are radioactive elements. 4fand 5f-series of elements are also called inner-transition elements because they form transition series within the transition elements of d -block.

Characteristics of f-block elements:

• They are all heavy metals.
• They exhibit variable valency. +3 oxidation state is most common. Few elements are found to occur in +2 and +4 oxidation states.
• Some members exhibit paramagnetism due to the presence of odd electrons.
• They form complex compounds, most of which are coloured.
• They have high densities.
• They generally have high melting and boiling points.
• Within each series, the properties of the elements are quite similar. It is very difficult to separate them from a mixture.
• Actinoids are radioactive. The first three members (Th, Pa, U ) occur in nature, while the others are man-made.
• The elements after uranium are called transuranic elements.

NCERT Solutions Class 11 Chemistry Chapter 3 Long Form of Periodic Table

Stair-step diagonal

• The right side of the long form of the periodic table is composed of p -block elements belonging to groups 13 (3A), 14(4A), 15(5A), 16(6A), 17(7A) and 18 (WlA or 0).
• This segment includes four types of elements viz., metals, nonmetals, metalloids and inert gases.
• There is no sharp line of demarcation to classify the metals and non-metals, but the zig¬ zag diagonal line (looking like stair-steps) running across the periodic table from boron (B) to astatine (At) is considered as a separation between the metals and non-metals.

This line is called the stair-step diagonal. The elements B, Si, Ge, As, Sb and Te bordering this line and- running diagonally across the periodic table are 8 known as metalloids (which exhibit properties that are characteristics of both metals and non-metals).

The elements (except Al ) lying between the stair-step diagonal line and the d -block elements are referred to as post-transition elements.

Determination Of The Position Of An Element In Long Form Of Periodic Table

Since there is a close relationship between the long form of the periodic table and the electronic configuration of elements, the serial numbers of periods and groups and the type of block to which an element belongs can be predicted by following the guidelines given below:

1. Period:

Serial number of the period = principal quantum number (n) of the valence shell.

Example:

Mg (ls²2s²2p⁶3s²) belongs to the third period because the principal quantum number of its valence shell is 3.

2. Block:

The publicly into which the differentiating electron [i.e., the last electron) enters, represents the block to which the given element belongs (except He ).

Example:

Sc (ls²2s²2p⁶3s²3p⁶4s²3d¹) belongs to d -block because the last electron [i.e., the 21st electron) enters the 3d -subshell.

3. Group:

The group to which an element belongs can be predicted based on the number of electrons present in the outermost [i.e., fth) and the penultimate [i.e., [n —1) th] shell.

For s -s-block elements: Group-number = Number of valence electrons i.e., no. of electrons in the ns -orbital.

For p -p-block elements: Group-number = 10 + no. of valence electrons = 10 + no. of ns -electrons + no. of np electrons.

For d -d-block elements: Group-number = no. of ns- electrons + no. of (n- l)d -electrons.

For f- f-block elements: Group number= 3 (fixed).

Examples:

Determination of the position of the elements with the following electronic configurations in the long form of the periodic table—

• ls²2s²2p⁶3s¹
• ls²2s²2p⁴
• ls²2s²2p⁶3s²3p⁶3d²4s²
• ls²2s²2p⁶3s²3p⁶4s²
• ls²2s²2p⁶3s²3p⁶3d¹⁰4s¹

1.

• The given element belongs to s -the block because the differentiating electron (i.e., the 11th electron) enters the 3s orbital.
• For this -block element, group-number = no. of electrons in the 3s -orbital =1.
• Serial no. of the period = principal quantum number ofthe valence shell =

2.

• The differentiating electron [i.e., the 8th election) enters the p-subshell.
• So, tile given element belongs to p block, [b] Serial no. of the period = principal quantum number of the valence shell
• = For this p -block element, group number= 10+ no. ofvalence electrons = 10 + number of ns electrons + no. of np -electrons =10 + 2 + 4 =16.

3.

• The differentiating electron [i.e., the 22nd electron) enters the 3d -subshell.
• So, the given element belongs to d -the block,
• Serial no. of the period = principal quantum number of the valence shell = 4.
• For die d block element, group-number = no. of ns -electrons + no. of (n- 1)d -electrons = 2 + 2 = 4.

4.

• The differentiating electron {l.e., the 20th electron) enters the 4s -subshell. So, the given element belongs to s -block,
• Serial no. of the period = principal quantum number of the valence shell =4.
• For this -block element, group-number = no. of electrons in outermost shell = 2.

5.

• The differentiating electron [i.e., the 29th electron) enters the 3d -subshell.
• So, the given element belongs to d -block, [b] Serial no. of the period = principal quantum number of the valence shell = 4.
• For this d -d-block element, group no. = no. of ns electrons + no. of[n- 1)d -electrons = 1 + 10 = 11.

CBSE Class 11 Chemistry Notes For Chapter 3 Valency

The valency of an element is defined as the combining capacity of that element. The valency of an element is usually expressed in terms of the number of H-atoms that combine with an atom of the element.

• The chemical properties of an element depend upon the number of electrons present in the outermost shell ofthe atom.
• Electrons present in the outermost shell are called valence electrons and these electrons determine the valency ofthe atom.
• In the case of the representative elements the valency of an atom is generally equal to either the number of valence electrons or equal to eight minus the number of valence electrons,
• However, transition and inner-transition elements exhibit variable valency involving electrons of the outermost shell as well as d- or f-electrons present in penultimate or antepenultimate shells.

Variation of valency in a period: 

In the case of the representative elements, the number of valence electrons increases from 1 to 7 from left to right in a period.

• Oxygen-based valency increases from 1 to 7 and it becomes a zero noble gas series (because of its inertness).
• The maximum valency of ‘8’ is shown only by Os and Ru in OsO₄ and RuO₄ respectively. These two elements (transition elements) belong to a group- 8 (VmB)in the periodic table.
• However, hydrogen-based and chlorine-based valency of representative elements along a period first increases from group-1 to 4 (valency= group no.) and then decreases from group-4 to 0 (valency= 8- group no).

Valency of the elements of the third period concerning oxygen and hydrogen:

Valency of elements of the second period to chlorine:

Variation of valency in a group:

On moving down a group, the number of valence electrons remains the same. Therefore, all the element groups exhibit the same valency.

Valency Class 11 Notes NCERT

Example:

All the elements of group-IA (Li, Na, K, Rb etc.) have alency T’ and that of group-2A (Mg, Ca, Sr etc.) exhibit avalencyof’21Noble gases present in group-8A are zerovalent since these elements are chemically inert.

Ionisation Enthalpy or inonisation potential.

If energy is supplied to an atom, electrons may be promoted to higher energy states. If sufficient energy is supplied, one or more electrons may be removed completely from the atom leading to the formation of a cation. This energy is referred to as ionisation energy or orionisation enthalpy (ΔH_i).

Ionisation Enthalpy or ionization potential Definition:

Ionisation enthalpy or more accurately first ionisation enthalpy of an element is defined as the amount of energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom existing in its ground state to form a cation in the gaseous state.

Explanation:

If ΔH₁(or I₁) is the minimum amount of energy required to convert any gaseous atom in its ground state into gaseous ion M+, then the ionisation enthalpy or more accurately first ionisation enthalpy of M is ΔH₁(or I₁)

Importance:

• The ionisation enthalpy of an element gives an idea about the tendency of its atoms to form gaseous cations.
• Energy is always required to remove electrons from an atom and hence, ionisation enthalpies are always positive.

Units:

It is expressed in kj per mole of atoms (kj. mol^-1).

Formerly, it was expressed in electron-volt per atom (eV- atom-1) or kcal per mole of atoms (kcal. mol^-1)

1ev per atom =23.06 kcalmol^-1=96.5 kl-mol¯¹.

Ionisation enthalpy :

‘Ionisation enthalpy is also called ‘ionisation potential’ because it is the minimum potential difference required in a discharge tube to remove the most loosely bound electron from an isolated gaseous atom to form a gaseous cation.

Successive ionisation enthalpies:

Like the removal of the first electron from an isolated gaseous atom, it is possible to second, third etc., electrons successively from cations one after another.

The minimum amount of energy required to remove the second, third etc., electrons from unipositive, dipositive etc., ions to form M²⁺, M³⁺ etc., ions of the element are called second ΔH₂(or I₃), tlirid ΔH₃ (or I3 ) etc., ionisation enthalpies respectively.

⇒ \( \mathrm{M}^{+}(g)+\Delta H_2\left(\text { or } I_2\right) \rightarrow \mathrm{M}^{2+}(g)+e \)

⇒  \( \mathrm{M}^{2+}(g)+\Delta H_3\left(\text { or } I_3\right) \rightarrow \mathrm{M}^{3+}(g)+e\)

The second ionisation enthalpy is higher than the first ionisation enthalpy as it is more difficult to remove an electron from a cation than from a neutral atom. ] Similarly, the third ionisation enthalpy is higher than the second and so on i.e.,

ΔH₁(or I₁) < ΔH₂(or l₂) < ΔH₃(or l₃)

If not mentioned, the term ‘ionisation enthalpy is always used to mean the first ionisation enthalpy of an element.

Formerly, first, second, third etc, ionisation enthalpies were denoted by the symbols I₁, I₂, I₃ etc. Such symbols will be used in many places in this book.

Factors governing ionisation enthalpy:

Atomicsize:

• Ionisation enthalpy decreases as the atomic size increases and vice-versa.
• The attractive force between the electron (to be removed) and the nucleus is inversely proportional to the distance between them.
• Thus, as the size of the atom increases, the hold of the nucleus over valence electrons decreases and consequently ionisation enthalpy decreases. For example, l₁(Li)>l₁(Na)>I₁(K).

The magnitude of nuclear charge:

• Ionisation enthalpy increases with an increase in nuclear charge and vice-versa.
• This is due to the fact the force of attraction between the valence electron (to be removed) and the nucleus increases with an increase in the nuclear charge provided that the outermost electronic shell remains the same.

Screening effect of inner-shell electrons:

• As the screening effect or shielding effect of the inner electrons increases, the ionisation enthalpy decreases
• In multi-electron atoms, the inner electronic shells act like a screen between the nucleus and the outermost electronic shell.
• As a result, the nuclear attractive force acting on the electrons in the outermost shell is somewhat reduced i.e., the effective nuclear charge gets reduced to some extent.
• Thus, the inner-electronic shells shield the electron (to be removed) from the nuclear attractive force, resulting in a reduction of ionisation enthalpy.

If other factors do not change, the ionisation enthalpy decreases with an increase in the number of inner electrons.

• In multi-electron atoms, the ability of the electrons present in the inner shells to shield or screen the outer electrons from the attractive force of the nucleus is called the shielding effect or screening effect.
• Naturally, the magnitude of the screening effect depends on the number of electrons present in the inner shells.
• In a particular energy level, the screening effect of the electrons presenting different subshells follows the sequencers >p> d> f.
• Due to the screening effect, the valence shell electrons do not feel the full charge ofthe nucleus.
• The actual nuclear charge experienced by the valence shell electrons is called the effective nuclear charge.
• This is given by the relation, Effective nuclear charge (Z)= total nuclear charge (Z) – screening constant (cr) where the screening constant (cr) takes into account the screening effect ofthe electrons present in the inner shells.

Penetration effect of electronic subshells:

Ionisation enthalpy increases as the penetration effect ofthe electron (to be removed) increases. It is known that in the case of multielectron atoms, the electrons in the s -s-orbital have the maximum probability of being found near the nucleus. In a given quantum shell this probability goes on decreasing in the sequence s->p-> d-> f.

This means that in a given shell, the penetration power of different subshells decreases in the order: of s->p-> d-> f-

• Now, if the penetration power of an electronic is greater, it is closer to the nucleus and held more firmly by it.
• So it is more difficult to remove such an electron from the atom and consequently, ionisation enthalpy will be high.
• Thus for the same shell, the energy needed to knock out an s -s-electron is greater than that required for a p-electron, which in turn will be more than that required to remove a d-electron and so on. In other words, ionisation enthalpy follows the sequence, s>p> d> f.

Effect of half-filled and filled subshells:

• It is known that half-filled and filled subshells have extra stability associated with them.
• Therefore, the removal of electrons from such subshells (having extra stability) requires more energy than expected.
• Consequently, atoms having half-filled or filled subshells in their valence-shell have higher values of ionisation enthalpies.

Example:

Be(ls²2s²) has higher ionisation enthalpy than B(ls22s22p1) because ionisation of Be requires the removal of one electron from its filled 2s -orbital in the valence shell.

For similar reasons Mg(ls²2s²2p⁶3s²) has higher ionisation enthalpy than Al(ls²2s²2p⁶3s²3p¹)

N(ls²2s²2p³) has higher ionisation enthalpy than O(ls²2s²2p⁴) because ionisation of nitrogen requires the removal of one electron from its half-filled 2p -the valence shell. Similarly, the ionisation enthalpy of P(3s²3p³) is greater than that of S(3s²3p⁴).

CBSE Class 11 Chemistry Valency Notes PDF

Effect of the electronic configuration of the outermost shell:

• Atoms, having the outermost electronic configuration ns²np⁶, are exceptionally stable because of their filled octet.
• Removal of an electron from an atom having such a stable electronic configuration requires a large amount of energy.
• Consequently, the noble gases He, Ne, Ar, Kr, Xe etc. (with outermost electronic configuration ns2np6) have very high ionisation enthalpy.

Variation of ionisation enthalpy in the periodic table:

• The periodic trends of the first ionisation energy of the elements are quite striking as seen from the graph.
• The graph consists of several maxima and minima. In each period maxima occur at the noble gases which have filled the octet with the electronic configurations (ns²np⁶).

• Due to very high ionisation enthalpies, these elements are almost inert and show extremely low chemical reactivity.
• In each period minima occur at the alkali metals which have only one electron in the outermost s -orbital. Due to very low ionisation enthalpies, these elements are highly reactive.
• Variation of ionisation enthalpy across a period: For representative elements (s and p -p-block elements), ionisation enthalpy usually increases with increasing atomic number across a period.

This is because as we move from left to right across a period—

The nuclear charge increases regularly, several shells remain the same and the addition of different electrons occurs in the same shell, and atomic size decreases.

• As a result of a gradual increase in nuclear charge and a simultaneous decrease in atomic size, the valence electrons are more and more tightly held by the nucleus.
• Therefore, more and more energy is needed to remove one valence electron and hence, ionisation enthalpy increases with an increase in atomic number across a period.

In any period, alkali metal has the lowest ionisation enthalpy and inert gas has the highest ionisation enthalpy.

On careful examination of ionisation enthalpy values, some irregularities in the general trend are noticed. Can each period be explained based on different factors governing ionisation enthalpy?

Examples:

1. I₁ of Be>I₁ of B:

• Forionisation of boron (1s²2s²2p¹), one electron is to be removed from the singly filled 2p orbital
• And this requires lesser energy, while for the ionisation of beryllium (1s²2s²) one electron is to be removed from the more penetrating For For For filled For 2s For orbital.
• Furthermore, the Removal of an electron from Batom gives B+ a stable electronic configuration with a filled 2s -subshell (ls²2s²) and so it requires a smaller amount of energy.
• On the other hand removal of an electron from the filled 2s -orbital of Be -atom to give Be+ (1s²2s¹) requires a greater amount of energy.
• Consequently, the first ionisation enthalpy of B is less than that of Be. of

2. I₁ of  N  > I₁of O:

• The electronic configuration of nitrogen (ls²2s²2p³) in which the outermost 2p -subshell is exactly half-filled is more stable than the electronic configuration of oxygen (ls²2s²2p⁴)in which the 2psubshell is neither half-filled nor filled.
• Removal of 1 electron from the O -atom gives 0+ with a stable electronic configuration having a half-filled 2p -subshell (ls²2s²2p²³), but this is not so in the case of the N -atom because N+ has the electronic configuration ls²2s²2p².
• In other words, the removal of an electron from the O -atom gives a cation with a more stable electronic configuration than that obtained by the removal of one electron from the N -atom. Thus, the first ionisation enthalpy of oxygen is less than that of nitrogen.

⇒ \(\mathrm{N}\left(1 s^2 2 s^2 2 p^3\right) \stackrel{-e}{\longrightarrow} \mathrm{N}^{+}\left(1 s^2 2 s^2 2 p^2\right)\)

⇒ \(\mathrm{O}\left(1 s^2 2 s^2 2 p^4\right) \stackrel{-e}{\longrightarrow} \mathrm{O}^{+}\left(1 s^2 2 s^2 2 p^3\right)\)

3. The very high I₁ value Ne:

The exceptionally high value of the first ionisation enthalpy of neon (noble gas) amongst the elements of the 2nd period is due to its stable electronic configuration(ns²np⁶) ofthe outermost shell.

Variation of ionisation enthalpy down a group:

For representative elements, ionisation enthalpy decreases regularly with an increase in atomic number on moving down a group from one element to the other.

NCERT Class 11 Chemistry Valency Solutions

Explanation: The regular decrease in ionisation enthalpy (1.E.) may be attributed to the following factors:

On moving down a group, the atomic size increases successively due to the addition of one new electronic shell at each succeeding element.

• Thus, the distance of valence shell electrons from the nucleus increases.
• Consequently, the nuclear attractive force on the valence electrons decreases and this, in turn, decreases the ionisation potential.
• There is an increase in the shielding effect on the outermost electrons due to an increase in the number of inner electronic shells. This increased shielding effect tends to decrease the ionisation potential on moving down a group.
• On moving down a group, the nuclear charge increases regularly and this increases the force of attraction of the nucleus on the valence electrons; this tends to increase the ionisation potential.
• The combined effect of the increase in size and the shielding effect outweighs the effect of the increased nuclear charge.
• Consequently, the ionisation enthalpies of the elements decrease regularly on going down a group.

This is evident from the values of the first ionisation enthalpies of the elements of group-1 (alkali metals) as given in the adjacent table.

Periodic variation of first ionisation enthalpies (eV) of the elements is evident from the following table.

Some important facts about ionisation enthalpy:

• The ionisation enthalpy of representative elements (s and block elements) increases from left to right across the period.  Exceptions are observed for some pairs of elements,
  • Example:
    • I₁ (Be)>I₁ (B)
    • I₁ (Mg)>I₁ (Al)
    • I₁ (N)>I₁ (O)
•  In any period, alkali metal has the least ionisation enthalpy. Cesium (Cs) has the lowest value of I. All the elements.
•  In each period, inert gas elements show the highest value of first ionisation enthalpy. Helium (He) has the maximum value ofI.E. of all the elements.
•  Among the representative elements, metals have low I.E., while non-metals have high values of I.E.
• Generally, first ionisation enthalpies of transition elements (d -block elements) increase slowly from left to right in a period. This is partly due to the poor screening effect of d orbitals and partly due to electron-electron repulsive forces.
•  f-block elements also show a small change in their ionisation enthalpies on increasing atomic number.
• . From Pd to Ag, from Cd to In and also from Hg to Tl, there is a sudden decrease in ionisation enthalpy even though the atomic number increases.

Pair of elements

Electron-gain enthalpy or electron affinity:

Energy is released when an electron is added to an isolated gaseous atom to convert it into a negative ion.

This energy is called electron-gain enthalpy. Electron-gain enthalpy of an atom is thus a measure of its tendency to form an anion. ΔH_egor EA denotes it.

Electron-gain enthalpy or electron affinity Definition:

Electron-gain enthalpy is defined as the enthalpy change involved when an electron is added to an isolated gaseous atom in its lowest energy state (ground state) to form a gaseous ion carrying a unit negative charge.

Explanation:

If q is the amount of energy released when an electron is added to the isolated gaseous atom ‘X’ in its ground state to convert to the gaseous ion X-, then the electron-gain enthalpy (electron affinity) of X is given by, ΔH_eg= -q

⇒ \(\mathrm{F}(g)+e\longrightarrow \mathrm{F}^{-}(g)+328 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \)

Or, \(\mathrm{F}(g)+e\longrightarrow \mathrm{F}^{-}(g), \Delta H=-328 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \)

⇒ \(\mathrm{Be}(g)+e\longrightarrow \mathrm{Be}^{-}(g)-66 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}\)

Or, \(\mathrm{Be}(g)+e\longrightarrow \mathrm{Be}^{-}(g), \Delta H=66 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}\)

When an F -atom combines with an electron to form an F- ion, energy is released. So enthalpy change has a negative value. Thus electron-gain enthalpy of fluorine is given by AHeg = be supplied to convert a Be -atom to a Be ion.

So enthalpy change has a positive value. Thus electron-gain enthalpy of beryllium is given, by ΔH_eg= +66 kj. mol^-1

Class 11 Chemistry Valency Definition and Examples

Points to remember:

• Electron-gain enthalpy of an atom is a measure of its tendency to form an anion.
• Electron-gain enthalpy has usually a negative value, but it may also have a positive value, especially for noble gases.
• The numerical value of the ionisation enthalpy of an I negative ion (X-) is equal to the electron-gain enthalpy of the neutral atom (X).

However, energy is usually evolved during the process of electron acceptance but energy is usually absorbed during the expulsion of electrons from an atom. So electron gain enthalpy of X and ionisation enthalpy of X- have opposite signs.

• Electron-gain enthalpy with a -ve sign indicates that energy is released when the neutral atom accepts an electron (only numerical values are taken for comparison when periodicity or other properties are considered).
• The high value of electron-gain enthalpy indicates that an added electron is strongly bound, while a low value indicates that a new electron is weakly bound to the atom.

Units:

Electron-gain enthalpies are expressed in kilojoule per mole (kj. mol-^-1) or in electronvolt (eV) per atom.

Successive electron-gain enthalpies:

Like the second and higher ionisation enthalpies, second and higher electron-gain enthalpies are also possible. However, the addition of a second electron to a negative ion (X^–) is opposed by the electrostatic force of repulsion.

So energy is to be supplied for the addition ofthe second electron. Thus, the second electron-gain enthalpy of an element is positive, and so is the third, and so on.

For example, when an electron is added to an oxygen atom to form an O-ion, energy is released. However, when another electron is added to the O^– ion to form the O^2- ion, energy is absorbed.

First electron-gain enthalpy:

⇒ \(\begin{array}{r}
\mathrm{O}(g)+e \longrightarrow \mathrm{O}^{-}(g), \quad \Delta H_{e g}=-141 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \\
\text { Energy released }
\end{array}\)

Second electron-gain enthalpy:

⇒ \(\begin{array}{r}
\mathrm{O}^{-}(\mathrm{g})+e \longrightarrow \mathrm{O}^{2-}(\mathrm{g}), \quad \Delta H_{e g}=\underset{ }{780 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}} \\
\text { Energy absorbed }
\end{array}\)

Similarly, the first and second electron-gain enthalpies of sulphur are -200 kj mol^-1 and +590 kj mol^-1 respectively.

Factors governing electron-gain enthalpy:

In general, the factors favouring the ionisation process disfavour the electron-gain process.

Effective nuclear charge:

As effective nuclear charge (Z+) increases, the force of attraction between the nucleus and the incoming electron increases and hence, the numerical value of electron gain enthalpy increases.

Thus, the numerical magnitude ofelectrongain enthalpyof carbon (Z = 6, IE = -122 kj.mol^-1 ) is greater than that ofboron(Z = 5, IE = -27 kj-mol^-1 ).

Atomic size:

As the size of the atom increases, the distance between the nucleus and the outermost shell (which receives the incoming electron) increases.

• If the effective nuclear charge (Z+) per electron at the periphery is more or less the same for different species
• Example: In a group of representative elements), the force of attraction towards the nucleus of the electrons at the periphery is less for the larger species.
• Consequently, the numerical magnitude of electron-gain enthalpy decreases as the atomic size increases. Thus for representative elements, the numerical value of electron-gain enthalpy decreases as the atomic number increases on moving down a group.

Nature of the orbital into which new electron gets accommodation:

• Orbitals which can penetrate more towards the nucleus are more suitable to accommodate the incoming electron.
• Thus the ease of accommodation of Incoming eLectron follows the order ns> np> nd > nf, as the penetration effect of different orbitals follows this sequence.
• So the numerical magnitude of electron-gain enthalpy decreases in the sequence ns > np> nd> nf.

Nature of the outer electronic configuration:

If the atoms of an element bear extra stability due to either the half-filled or full-filled subshell in their outermost level, then such atoms are very much reluctant to accept the incoming electron.

• On the other hand, if the newly added electron creates a half-filled or full-filled subshell, then the process is favoured.
• Thus some ofthe elements of Gr 2A(ns²), Gr-2B[(n- 1)d¹⁰ns²], Group-VA (ns²np³) and all the noble gas elements (ns2np6) have positive electron-gain enthalpies (ΔH_eg).
• On the other hand, elements of Gr-7A have very high electron-gain enthalpies with negative signs, because they can attain inert gas configuration accepting one electron.
• Variation of electron-gain enthalpy across a period: On moving from left to right in a period, effective nuclear charge, Z nuclear charge (Z)- shielding effect of the inner shells) increases and size decreases with the increase in atomic number.
• Both these factors tend to increase the nuclear attraction experienced by the incoming electron and hence, the numerical value of electron-gain enthalpy.

In general, increases in a period from left to right. It reaches a maximum value at Gr-7A (halogens).

Due to some characteristic electronic configuration, the general trend is violated in some cases.

For example:  Be and N in the 2nd period; Mg and P in the third period).

Variation of electron-gain enthalpy down a group:

For the representative elements, on moving down a group, the effective nuclear charge Z per electron at the periphery (outermost shell) remains more or less constant because the effect of increased nuclear charge is counterbalanced by the shielding effect of the inner electronic shells.

• However, the atomic size gradually increases due to the addition of new quantum levels.
• Thus the nuclear attractive force experienced by any added electron (incoming electron) decreases as the atomic number increases, and consequently, the numerical value of electron-gain enthalpy decreases down a group.

Some typical trends in electron-gain enthalpy & their explanation:

Though the electrostatic attractive pull towards the nucleus favours the addition of an electron to the smaller-sized F -atom, the added electron, however, creates an unfavourable effect.

• The added electron Experiences significant electron-electron repulsion from the other electrons present in the small-sized 2p -subshell.
• On the other hand, in a chlorine atom, the added electron goes to the large-sized 3p -subshell. Hence, it experiences less electron-electron repulsion.
• Another factor that favours the uptake of electrons by the Cl -atom, is that there is the possibility of the delocalisation of the increased electron density in the vacant 3d -orbital of Cl-atom.
• This mechanism is not operative in F-atom because of the absence of d -orbital in the second shell. Consequently, the numerical value of electron-gain enthalpy of Cl is greater than that of F.

Electron-gain enthalpy is greater than that of O:

The reason for this anomaly is similar to that of Cl versus F.

• The added electron experiences considerable electron-electron repulsion from the other electrons present in the small-sized 2p -subshell of O.
• This repulsion outweighs the increased attractive force of the nucleus acting on the added electron. In the S-atom, the added electron goes to the large-sized 3p -subshell.
• Hence, it experiences less electron-electron repulsion. Another factor that favours the uptake of electrons by S -S-atoms is that there is a possibility ofthe delocalisation of the increased electron density in the vacant-3d -orbital of S-atom.
• This mechanism is not operative in the O -atom because of the absence of any d orbital in the 2nd shell.
• Consequently, the numerical value of electron-gain enthalpy S is greater than that of O.

Gr-2A metals (Example: Be, Mg etc.) have lower electron-gain enthalpies than Gr-IA metals (Example: Li, Na, K etc:

Gr- A metals have outer electronic configuration ns². Hence, the addition of an extra electron brings the configuration ns²np¹.

This process is disfavoured in two ways:

1. The addition of a new electron destroys the full-filled subshell structure and accommodation of the new electron occurs in the p -orbital which is less penetrating. For alkali metals (ns¹), however,
2. Accommodation of the new electron occurs in the ns -subshell giving rise to a filled ns² configuration.

Thus, the electron-gaining process is more favourable for Gr-1 A elements compared to Gr-2A elements. Be and Mg of Gr-2A have positive electron-gain enthalpy.

Valency Chapter 3 NCERT Class 11 Important Topics

Halogens have the highest electron-gain enthalpies:

• This is because of the valence-shell electronic configuration of the halogensis ns²np⁵ and so they require only one more electron to acquire the stable inert gas-like electronic configuration (ns²np⁶).
• As a result, halogen atoms have a strong tendency to accept an additional electron. Consequently, the numerical values of their electron-gain enthalpies are very high.

Phosphorous (3s²3p³) has relatively low electron-gain enthalpy:

• This is because the P -atom has a relatively stable outer electronic configuration with exactly half-filled p -orbital.
• Hence, it is reluctant to accept an extra electron. Consequently, it has low electron-gain enthalpy.

The electron-gain enthalpy of noble gas is high and positive:

• The atoms of noble gases have a very stable outermost electronic configuration with filled subshells (ns²np⁶).
• Any additional electron would have to be placed in an orbital ofthe next higher energy level.
• The shielding effect of the inner electrons and the large distance from the nucleus makes the addition of an electron highly unfavourable. So, noble. High and acquisitive values electron-gain enthalpy.

Electronegativity

This topic will be discussed elaborately in the chapter ‘Chemical Bonding’ Here we will briefly discuss only the definition of electronegativity and its periodicity.

Electronegativity Definition:

Electronegativity is defined as the tendency of an atom to attract the shared pair of electrons towards its nucleus when the atom isovalently bonded in a molecule.

• Consequently, the more electronegative atom withdraws the bonding electron cloud more towards its nucleus giving rise to an accumulation of negative charge on it.
• The electronegativity of an element is not its inherent property. It depends on its surrounding environment in the molecule in which the electronegativity of the element is being considered.
• Thus the electronegativity of S is different in different compounds such as H₂S, SO₂, SF₆ etc.
• Further, it is to be noted that unlike ionisation enthalpy and electron-gain enthalpy, electronegativity is not a measurable quantity.

Factors controlling electronegativity: Electronegativity of the elements depend on—

• The atomic number of an element, i.e., the total quantity of positive charge in the nucleus of an atom,
• Size of atom Or Atomicradius, number of electronicshes in an atom, oxidation state ofthe atom, state of hybridisation of the atom in the molecule under consideration.
• Note the electronegativity of elements.

Variation of electronegativity across a period:

• As Cs 0.7 At 2.2 Ford -block element, on moving down from 3d- to from left to right along a period, nuclear charge Increases while the atomic radius or size decreases.
• Hence the attraction between the outer (or valence) electrons and the nucleus increases with increasing atomic number.
• Consequently, the electronegativity of the atom increases from left to right across a period.
• Thus alkali metals of group-1 on the extreme left have the lowest electronegativity whereas, the halogens in group-17 on the right have the highest values of electronegativity in their respective periods.

This is evident from the electronegativity values ofthe elements of the second and third periods.

Variation of electronegativity down a group:

As we move down a group, atomic size (radius) as well as the magnitude of nuclear charge increases but the effect of increased nuclear charge on the outer electrons is mostly counterbalanced by the screening effect of a larger number of inner electronic shells.

• Hence the nuclear pull on the outer (valence) electrons decreases due to the increase of atomic size on moving down a group.
• Consequently, the electronegativity of an atom decreases from top to bottom in a group.
• This is evident from the electronegativity values of the alkali metals of group-1 (IA) and halogen elements of group-17(8A)

Electronegativity values down a group:

Group – 1 Elements:

Group – 17 Elements:

Ford -block element, on moving down from 3d- to 4d -series, electronegativity falls slightly but on reaching 5d series, electronegativity increases due to lanthanide contraction.

Relationship between electronegativity and non-metallic (or metallic) character of elements:

Non-metallic elements have a strong tendency to gain electrons. So, electronegativity is directly related to the metallic character elements.

• It can be further extended to say that electronegativity is inversely related to the metallic character of elements.
• Thus the increase in electronegativity along a period is accompanied by an increase in non-metallic character (or decrease in metallic character) of elements.
• Likewise, the decrease in electronegativity down a group is accompanied by a decrease in the non-metallic character (or increase in metallic character) of elements.

All these periodic trends are summarised in the given figure:

Direction arrows indicate an increasing trend of the respective properties

Periodicity in density, melting point and boiling point

Different elements exhibit periodicity in various physical properties such as density, melting point, boiling point etc.

Periodic variation of density:

On moving along a period from left to right, the density of representative elements first increases reaches the maximum value at group-3A or 4A and then decreases with an increase in atomic number.

This trend is observed particularly in the case of representative elements. In a group, density generally increases from top to bottom with a rise in atomic number.

Periodic variation of melting and boiling points:

On moving along a period from left to right, the melting and boiling points of representative elements first increase, reach maximum values at group 4A and thereafter go on decreasing.

Minimum melting and boiling points are shown by the noble gas in the respective period.

Periodicity In Properties Of Oxides And Hydrides

Nature of oxides of the elements:

On moving from left to right across a period, the basic properties and electrovalent character of oxides of elements decrease while their acidic property and covalent character gradually increase.

• On the other hand, in a group, the basic property of oxides increases from top to bottom.
• The nature of the oxides of transition metals depends on the oxidation state of the metals.
• With the increase in the oxidation state of transition metals, the acidic properties of their oxides increase.

Basic and acidic nature of oxides of different elements:

Example:

CrO is a basic oxide, Cr₂O₃ is amphoteric and CrO₃  is an acidic oxide. In the case of oxides of the elements of the second period, it is observed that lithium oxide (Li₂O) is strongly basic. It reacts with water to produce a strong base namely lithium hydroxide (LiOH)

⇒ \(\mathrm{Li}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons 2 \mathrm{Li}^{+}+2 \mathrm{OH}^{-} \rightleftharpoons 2 \mathrm{LiOH}\)

BeO is an amphoteric oxide. It reacts with both acids and bases to form salt and water.

⇒ \(\text { Basic property: } \mathrm{BeO}+2 \mathrm{HCl} \rightarrow \mathrm{BeCl}_2+\mathrm{H}_2 \mathrm{O}\)

⇒  \(\text { Acidic property: } \mathrm{BeO}+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{BeO}_2+\mathrm{H}_2 \mathrm{O}\)

B₂O₃ is an acidic oxide though it possesses a slight basic property. It reacts with water to form orthoboric acid and with alkali to yield borate salt.

CO₂ is an acidic oxide and it reacts with alkali to produce carbonate salt. N₂O is an acidic oxide. It reacts with alkali to produce salt and water.

⇒ \(\mathrm{B}_2 \mathrm{O}_3+3 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{H}_3 \mathrm{BO}_3 ; \mathrm{CO}_2+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{CO}_3+\mathrm{H}_2 \mathrm{O}\)

⇒  \(\mathrm{N}_2 \mathrm{O}_5+2 \mathrm{NaOH} \rightarrow 2 \mathrm{NaNO}_3+\mathrm{H}_2 \mathrm{O}\)

Class 11 Chemistry Chapter 3 Valency Notes

Nature of hydrides of elements:

As we move from left to right across a particular period, the tendency of the elements to form hydrides and the thermal stability, covalent character, acidic property, and volatility of the hydrides increases while the reducing property progressively decreases.

• The hydrides of the strongly electropositive metals towards the left of a period are ionic having high melting points.
• On ionisation, they produce hydride ions (H-). Again, hydrides of non-metals towards the right of the period are covalent and have low melting and boiling points.
• On moving down a group, the tendency of the elements to form hydrides decreases.
• The stability of the hydrides also decreases in the same sequence. Variation of other properties along any group depends on the group to which the hydride-forming element belongs.

Hydrides of elements and variation in their properties:

Hydrides of alkali metals in group 1A and alkaline earth metals in group 2 are salt-like polar or ionic.

• These compounds comprise positive metallic ions and negative hydride ions (H¯). On electrolysis of these ions and negative hydride ions (H-).
• On electrolysis these are discharged at the cathode and anions (H- ions) at the anode;

Example:

Electrolysis of molten sodium hydride leads to the formation of metallic sodium at the cathode and H2 gas at the anode.

Hydrides of the elements of groups 4A to 7A are covalent and nonpolar;

Example:

CH₄, SiH₄, PH₃ etc., are gaseous and insoluble in water. NH₃ and H₂S are gaseous but soluble in water. An aqueous solution of NH₃ is feebly basic and the aqueous solution of H₂S is weakly acidic.

On the other hand, HCl, HBr and HI dissociate almost completely despite being covalent compounds more soluble in water and dilute aqueous solutions.

⇒ \(\mathrm{HX}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{X}^{-} \quad[\text { where } \mathrm{X}=\mathrm{Cl}, \mathrm{Br}, \mathrm{I}]\)

Aqueous solution of HX is strongly acidic. In electrolysis of their aqueous solutions hydrogen ions (H+) are liberated at the cathode and halide ions (X-) at the anode.

Example:

When an aqueous solution of hydrochloric acid is electrolysed, H₂ gas is evolved at the cathode and Cl₂ gas at the anode.

The trend of variation in properties of different elements in the periodic table from left to right across a period and from top to bottom in a group is shown in the given table-

NCERT Solutions Class 11 Chemistry Chapter 3 Valency

Diagonal Relationship Definition:

Some elements of certain groups in the second period show similarity in properties with the diagonally opposite elements of the third period, and such similarity in properties is referred to as a diagonal relationship.

Reason for diagonal relationship:

The Reason for the diagonal relationship is due to opposing trends in periodic properties along a period from left to right and down the group.

• For example, the atomic and ionic radius of elements decrease a periodic but increase down a group Ionisation enthalpy, electron gain enthalpy and electron negativity increase along a period but decrease down a group.
• On moving diagonally, two opposite trends mutually cancel, so the elements of the period- 2 and 3 listed above are related to each other diagonally and they show similar chemical properties. Thus Li resembles Mg; Be resembles Al; and B resembles Si.
• The diagonal relationship is also explained based on polarising power ofcation. On moving along the period from left to right, the charge on the cation increases, while ionic size decreases and hence polarising power increases.
• Again on moving down a group, the charge on the cation remains the same, while ionic size increases.
• Hence polarising power decreases. So on moving diagonally, polarising power remains more or less the same and the elements exhibit similar properties.

Absence of diagonal relationship in case of long periods:

• Because of the intervening d – and f-series, the diagonal relationship does not hold well for long-period elements (4th, 5th… period elements).
• Because the group trend of many properties in the transition series is opposite compared to that in the representative elements. However, the trend along the periods remains the same for both the representative and d -d-block elements.

CBSE Class 11 Chemistry Notes For Chapter 3 Position Of Hydrogen And Inert Gases In The Periodic Table

The position of hydrogen in the periodic table is controversial. Given its chemical analogy with both the elements of group and that of group- 7A, it can either be placed in group 2A or group 7A. Resemblances of hydrogen with the element.

Arguments in favour of placing hydrogen in group 1A

1. Valency:

The electronic configuration of hydrogen is Is¹ and the general electronic configuration of the elements of group-1A is ns1, i.e., like the elements of group-1A, hydrogen has only one valence electron and its valency is 1.

2. Electropositive character:

Like all group-1A elements, hydrogen tends to form cations by losing one electron.

⇒ \(\mathrm{Na}-e \longrightarrow \mathrm{Na}^{+} ; \mathrm{K}-e \longrightarrow \mathrm{K}^{+} ; \mathrm{H}-e \longrightarrow \mathrm{H}^{+}\)

Like elements of group-1A, hydrogen reacts with electronegative elements such as chlorine, oxygen and sulphur to produce similar types of compounds

Examples: HCl , H₂O , H₂S ; NaCI , Na₂O , Na₂

CBSE Class 11 Chemistry Notes Chapter 3 Position of Hydrogen

3. Electrolysis of chloride compounds:

Electrolysis of molten NaCl results in the deposition of metallic sodium at the cathode. Likewise, when an aqueous solution of HCl is electrolysed, H₂ gas is liberated at the cathode.

NaCl ⇌ Na⁺  + Cl^–

Cathode : Na⁺ + e → Na

Anode : Cl^–-e →Cl ; CI + Cl→Cl₂↑,

HCl  ⇌ H⁺ + Cl^–

Cathode : H+ + e →H ; H + H→H₂↑

Anode : Cl^–-e →Cl ; Cl + Cl→Cl₂↑

4. Reducing property:

Like the elements of Gr-IA, hydrogen loses electrons easily and exhibits a reducing property.

5. Formation of alloy:

Hydrogen dissolves in metals like Pd, Pt etc., by adsorption. This occlusion of hydrogen is comparable to the formation of alloys by elements of group IA.

6. Mutual displacement:

Hydrogen atom(s) of hydrochloric, sulphuric or nitric acids can be displaced by the same number of atoms of group-IA elements. Again, atoms of the group-IA elements can be replaced by hydrogen atoms from the salts produced.

7. Formation of stable oxide:

Oxides of group-IA elements are highly stable

Example: Na₂O, K₂O etc.). Similarly, oxide of hydrogen (H20) is also highly stable.

8. Formation of peroxide:

Like the elements of group IA, hydrogen also forms peroxide (H₂O₂). The analogous peroxides of group-IA elements are Na₂O₂, K₂O₂ etc.

9. Electron affinity:

• Hydrogen and the group-IA elements have comparable values of electron affinity.
• In light of the above similarities between the elements of group IA and hydrogen, hydrogen can be placed along with the elements of group 1A.
• However, the placement of hydrogen in group 1A leaves six vacant places in between H and He in the first period.

Position of Hydrogen and Inert Gases in Periodic Table Class 11 Notes

Arguments in favour of placing hydrogen in group 7A

1. Electronic configuration:

The electronic configuration of hydrogen is 1s¹ and the electronic configuration of the outermost orbit of the elements of group- 7A is ns²np^5- i.e., the outermost orbit of both hydrogen and elements of group-6LA has 1 electron less than the electronic configurations of the nearest inert gas. So, their valency 1.

2. Non-metallic character and atomicity:

Like the dements of group-7A, hydrogen is also non-metal and forms a diatomic molecule.

3. Formation of anion:

Like the elements of group 7A, the hydrogen atom also tends to attain the electronic configuration of Its nearest Inert gas (Me) by accepting an I electron and forming an anion (H^– );

Example:

⇒ \(\mathrm{H}\left(1 s^1\right) \stackrel{+e}{\longrightarrow} \mathrm{H}^{-}\left(1 s^2\right) ; \mathrm{X}\left(n s^2 n p^5\right) \stackrel{+e}{\longrightarrow} \mathrm{X}-\left(n s^2 n p^6\right)\)

Both 7A elements and hydrogen form electrovalent halide and hydride respectively. During the electrolysis of metallic hydrides, like halogens, hydrogen is also liberated at the anode.

NaCl ⇌ Na⁺ + Cl^–

Cathode : Na⁺ + e  → Na

Anode:  Cl^–-e → Cl; Cl+ Cl→ Cl₂ ↑

NaH (molten) ⇌ Na⁺ + H^–

Cathode : Na⁺ + e ιNa

Anode: H^–– e → H; H + H→H₂↑

4. Formation of covalent compounds:

Just like elements of group VIIA, hydrogen reacts with different non-metals to produce covalent compounds with analogous formulas.

Compoundsinvolving H: CH₄, NH₃, H₂O, HF, SiH₄.

Compounds involving Cl: CCl₄, NCl₃, Cl₂O, CIF, SiCl₄

5. Substitution by halogens:

H-atoms of the hydrocarbons can be substituted by Gr-7A elements, partially or completely.

Ionisation potential:

Just like the elements of group 7A, the ionisation, potential of hydrogen is very high but the ionisation potential of alkali metals is quite low. The following table of ionisation potentials shows the comparative picture of ionisation potential quite explicitly.

Maintenance of continuity in the periodic table:

• If H is placed in group 7A, no vacant space remains between H and H. So, continuity in the periodic table is not disturbed.
• From the above discussion, it is apparent that hydrogen is a unique element characterised by peculiar and distinctive unique element characterised by peculiar and distinctive position in the periodic table.
• It is reasonable to set aside a separate position for hydrogen in the periodic table. In the modern periodic table, hydrogen has been J given a completely separate place, at the top of the table.

NCERT Class 11 Chemistry Chapter 3 Position of Hydrogen and Inert Gases

Position of Inert Gases In The Periodic Table

• Inert gas elements have very stable electronic configurations of their outermost or valence shell (ns² for He and ns²np⁶ for others).
• For this reason, these elements show little or no tendency to lose or gain electrons to form ions to give electrovalent bonds or do not share electrons with other elements to form covalent bonds.
• So the combining capacity or valency of these elements is zero.
• Thus they are placed in group ‘zero’ ofthe periodic table.
• This group forms a bridge between the most electropositive alkali metal elements of group-7A.

CBSE Class 11 Chemistry Notes For Chapter 3 Periodic Trends In Properties Of The Elements

The properties of elements can be divided into two categories:

1. Properties of individual atoms: The properties such as atomic and ionic radii, ionisation energy, electron affinity, electro¬negativity and valency are the properties of the individual atoms & are directly related to their electronic configurations.
2. Properties of groups of atoms: The properties such as melting
   point, boiling point, heat of fusion, heat of vaporisation, atomic volume, density etc. are the bulk properties i.e., the properties ofa collection of atoms and are related to their electronic configurations indirectly.

All these properties which are directly or indirectly related to the electronic configurations of the elements are called atomic properties. Since electronic configurations of the elements are periodic functions of their atomic numbers, these atomic properties are also periodic functions of atomic numbers of the elements. Thus, atomic properties are also called periodic properties.

Remember that, when we descendin a group, the chemical properties ofthe elements remain more or less the same due to same valence shell configuration, but there is a gradual change in physical properties due to gradual change in the size ofthe atoms owing to addition ofnewelectronic shells.

Atomic size or atomic radius

If an atom is assumed to be a sphere, the atomic size is given by the radius ofthe sphere, called the atomic radius.

Atomic size or atomic radius Definition:

The distance from the centre of the nucleus to the outermost shell containing the electrons is called the atomic radius

Difficulties in precise measurement of atomic radius:

• It is not possible to isolate a single atom for the measurement of its radius.
• The electron cloud surrounding the nucleus does not have a sharp boundary as the probability of finding an electron can neverbe zero even at alarge distance from the nucleus.
• The magnitude ofatomic radius changes from one bonded state to another.

CBSE Class 11 Chemistry Notes Chapter 3 Periodic Trends

Types of atomic radius:

As already mentioned, the size of an atom varies from one environment to another. Therefore, several kinds of atomic radii have been defined.

These are

1. Covalent radius
2. Metallic radius
3. Van der
4. Waals radius.

1. Covalent radius:

It is defined as one-half of the distance between the centres of two atoms of the same element bonded by a single covalent bond

Thus for homonuclear diatomicmolecules, covalent radius (r) = ½ internuclear distance.

Example: In H₂ molecule, the internuclear distance is = 0.74 A

= 74pm. So, the covalent radius of hydrogen atom = ½  × 0.74

= 0.37 Å= 37pm.

1Å = 10^-10m, 1 pm = 10^-12m,

2. Metallic radius:

It is defined as one-half of the internuclear distance between two adjacent atoms in a metallic lattice.

Example:

The distance between two adjacent copper atoms in solid copper is 2.56 Å(determined by X-ray diffraction).

Hence, the metallic radius of copper is 1.28 Å. Similarly, the metallic radii of sodium and potassium have been determined as 1.86 Å and 2.31 Å respectively.

Note: Covalentradii ofNa and K are 1.54 Å and 2.03 Å

A metallic radius is always greater than a covalent radius.

3. Van der Waals radius:

It is defined as one-half of the distance between the nuclei of two non-bonded neighbouring atoms belonging to two adjacent molecules of an element in the solid state.

Periodic Trends in Properties of Elements Class 11 Notes

Example:

The distance between nuclei of two adjacent Cl -atoms of two adjacent chlorine molecules in solid states is 3.6  Å. So, the van der Waals radius of Cl-atom is

⇒ \(\frac{3.6}{2}\)= 1.8 A.

Since the van der Waals force of attraction is weak even in the solid state, the internuclear distances between the atoms of two adjacent molecules held by van der Waals forces are much larger than those between covalently bonded atoms (which involve mutual overlap of atomic orbitals).

Van der Waals radii are always greater than covalent radii. For example, the van der Waals radius of chlorine is 1.80 Å while its covalent radius is 0.99 Å

Sequenceofthree types of atomic radii: van der Waals radius >Metallic radius > Covalentradius

Variation of atomic radii or sizes in the periodic table

1. Variation of atomic radii or sizes across a period:

While moving from left to right across a period in the periodic table, atomic radii or atomic sizes progressively decrease.

Atomic radii across a period  Explanation:

The principal quantum shell remains unchanged in the same period. So, the differentiating electrons enter the same shellbutdue to an increase in the number of protons, the positive charge of the nucleus also increases. So, attractive force of the nucleus for electrons in the outermost shell also increases. Hence, the atomic sizes, rather than the atomic radii of elements in the same period, gradually decrease with an increase in atomic number. In any period, atomic size of the element of group-1A is maximum and that of the halogen of group VILAisminimum.

Variation of atomic [covalent] radii of the elements of the third period (n = 3]:

2. Variation of atomic radii or size down a group:

On moving down along any group of the periodic table, the atomic sizes rather than the atomic radii of the elements increase remarkably.

Atomic radii or size down a group Explanation:

On moving down a group, a new electronic shell is added to each succeeding element, though the number of electrons in the outermost shell remains the same. This tends to increase the atomic size. At the same time, there is an increase in nuclear charge with an increase in atomic number. This tends to decrease the size.

However, the effect of the increased nuclear charge is partly reduced by the shielding effect of the inner electronic shells.In practice,it is found that the effect of the addition of a new electronic shell is so large that it outweighs the contractive effect of the increased nuclear charge. Hence, there occurs a gradual increase in atomic radii on moving down a group in the periodic table.

3. Variationofsizein agroupforheavierelements:

On moving down a group, the relative rate of increase of atomic radio is slow for heavier elements. So, the differences in size for heavier species such as Cs (6s¹) and Fr(7s¹) of group-1 or Ba(6s²) and Ra(7s²) of group-2 are very small. This is due to the presence of electrons in the inner d and f-orbitals having a poor screening effect. D and f-electrons do not screen the outer electrons effectively from the pull of the nucleus. There is only a small increase in size in spite of the addition of a new electronic shell.

Example: Na(1.54 Å), K(2.03Å), Rb(2.16 Å), Cs(2.35Å) etc

In the case of transition elements, such a decrease in the atomic size is relatively less. Here the differentiating electron instead of entering the outermost orbit, goes to the penultimate (n- f )d -subshell which is closer to the nucleus. However, due to the poor shielding effect of the additional d -electrons, there is a small increase in effective nuclear charge and hence, the decrease in the size with an increase in the atomic number is relatively small.

In the case of lanthanide elements, the differentiating electrons enter the (n-2)f-orbital having a very poor shielding effect. So, in such cases, the increase in the effective nuclear charge is greater than that in the case of transition elements. Thus, the decrease in atomic sizes of the lanthanides is much more regular and distinct as compared to the transition elements. So, the difference in the extent of the decreased atomic sizes of these elements is due to their relative inability to screen the outermost electrons from attraction of the nucleus.

Periodic Trends in Properties of Elements Class 11 NCERT Solutions

Covalent radii (pm] of representative element:

 Variation of atomic radii with atomic numbers across the second period:

Variation of atomic radii with atomic numbers for group-1 metals:

Screening effect or shielding effect:

In multi-electron atoms, the electrons present in the inner shells shield the electrons in the valence shell from the pull of the nucleus. It means that the electrons of the inner shells act as a screen between the nucleus and the electrons in the valence shell. This is known as the screening effect shielding effect.

Screening effect or shielding Definition:

The ability of the inner electronic shells to shield the outer electrons from the attraction of the nucleus is called the screening effect or shielding effect.

Magnitude of the screening effect of electrons belonging to different subshells follows the sequence: s>p> d>f.

Periodic Trends Chapter 3 Class 11 NCERT Notes

Important points regarding atomic (covalent) radius:

• The alkali metals occupying positions at the extreme left side ofthe periodic table have the largest size in period
• The halogens occupying positions at the right side of the periodic table have the smallest size in each period.
• The noble gases present at the extreme right of the periodic table have larger atomic radii than those of the preceding halogens because van der Waals radii (but not covalent radii) are taken into consideration for noble gases.
• In transition series {d -block elements), there is only a small decrease in size with successive increases in atomic number because the differentiating electrons enter into (n-1)d subshell, which partially shields the increased nuclear charge acting on the valence electrons. This is known as d contraction.
• For inner-transition series ( /-block elements), the decrease in atomic radii wide increase in atomic number is relatively greater and more regular as compared to the transition series elements. This is so because the differentiating electrons enter into (n- 2)/ -subshell having a very poor shielding effect.

In a group of representative elements, there is a continuous increase in atomic radii with an increase in atomic number.

On going down a group of transition elements, there is an increase in size from first member to second member as expected, but for higher members, the increase in size is quite small. This is due to lanthanide contraction. For example, Cu(1.28 Å), Ag(1.44 Å), Au (1.44 Å)etc