# MP Board Class 11th Chemistry Important Questions Chapter 3 Classification of Elements and Periodicity in Properties

## MP Board Class 11th Chemistry Important Questions Chapter 3 Classification of Elements and Periodicity in Properties

### Classification of Elements and Periodicity in Properties Important Questions

Classification of Elements and Periodicity in Properties Objective Type Questions

Question 1.

Question 1.
Variable valency is represented by:
(a) Metallic elements
(b) Normal elements
(c) Transitional elements
(d) Non – metallic elements.
(c) Transitional elements

Question 2.
Of the following, whose size is the largest:
(a) Al
(b) Al+
(c) Al2+
(d) Al3+
(a) Al

Question 3.
Electronic structure of the last shell of highly electronegative element is:
(a) ns2p3
(b) ns2 np4
(c) ns2p5
(d) ns2 np6
(c) ns2p5

Question 4.
Electron affinity enthalpy depends on:
(a) On atomic size
(b) On nuclear charge
(c) On atomic number
(d) On both atomic size and nuclear charge
(d) On both atomic size and nuclear charge

Question 5.
Electron gain enthalpy of noble gases is:
(a) Less
(b) Nearly zero
(c) High
(d) Very high.
(b) Nearly zero

Question 6.
Alkali metals behave:
(a) As a good oxidising agent
(b) As a good reducing agent
(c) As a good hydrolyser
(d) None of these.
(b) As a good reducing agent

Question 2.
Fill in the blanks:

1. Electron affinity of noble gases is …………………..
2. Electron affinity of Be and Mg of second group is …………………….
3. In a period, on moving from left to right, ionisation energy …………………………
4. Periodic law made by taking atomic mass as the basis is called …………………………….
5. …………………. is called Eka Boran
6. Na+ is …………………….. than Na.
7. Li shows diagonal relationship with ………………………..

1. Zero
2. Zero
3. Increases
4. Mendeleev’s periodic law
5. Aluminium
6. Smaller
7. Mg.

Question 3.

1. Which scientist gives Modem periodic law?
2. Which element have highest electronegativity?
3. An element having atomic number 20, will be kept in which group in periodic table?
4. What is electron affinity of Noble gases?
5. What is the change in ionization energy on moving left to right in periodic table?
6. How many periods and groups in modem periodic table?
7. What is the other name of Eka – silicon?
8. How many elements are their in first period of modem periodic table?
9. What is the reason for the similarity in properties of Li and Mg?
10. Which oxidation state of A1 is most stable?

1. Mosle (1913)
2. Fluorine
3. In second group
4. 0 (Zero)
5. Increases
6. Period = 7 and Group = 18,
7. Germanium
8. 2 (Two)
9. Diagonal relationship
10. +3.

Question 4.
Match the following:

1. (c)
2. (a)
3. (b)
4. (e)
5. (d)

Classification of Elements and Periodicity in Properties Short Answer Type Questions – I

Question 1.
What is Mendeleev’s periodic law?
Mendeleev’s periodic law:
“The physical and chemical properties of the elements are a function of their atomic masses”. This means that the physical and chemical properties of the elements get repeated after a fixed interval.

Question 2.
What is Modern periodic law?
Modern periodic law: “Physical and chemical properties of elements are periodic functions of their atomic number”. This means when the elements are arranged in their increasing atomic numbers then after a fixed interval the elements having similar properties are repeated.

Question 3.
This law was given by Doebereiner. According to this law atomic mass of the middle element was an arithmetic mean of the atomic masses of the other two.
Li7 Na23 K39
That is $$\frac{39 + 7}{2}$$ = 23.

Question 4.
Prove on the basis of quantum numbers that the 6th period of periodic table should have 32 elements?
The sixth period begins with the filling of principal quantum number n = 6. In this period, the electrons enters in 6s, Af 5 d and 6p. In this sub – shells total 16 (1 + 7 + 5 + 3) orbitals are present. According to Pauli’s exclusion principle every’ orbital have maximum two electrons. So 16 orbitals should have only 32 electrons. So, in 6th period 32 elements are present.

Question 5.
How many periods and groups were present in Mendeleev’s periodic table? What are the number of elements in each period?
In Mendeleev’s periodic table 7 periods and 9 groups are present. In first period 2 elements, in second and third period 8 – 8 elements, fourth and fifth period, 18 – 18 elements, in sixth period 32 elements and seventh period is incomplete. In this period elements of atomic number 90 to 103, including actinides were present.

Question 6.
What change takes place on moving from left to right in period?
On moving from left to right in period, ionization energy and electron affinity increases, but metallic property, basicity of oxides and atomic radii decreases.

Question 7.
What change occur on moving from top to bottom in a group?
On moving top to bottom in a group atomic radii and ionic radii increases but ionization potential, electron affinity and melting point decreases.

Question 8.
What is diagonal relationship? Explain giving example?
Diagonal relationship:
It has been observed that some of the elements of second period show similarities with the elements of third period present diagonally to each other, though belonging to different groups. For example, lithium resembles with magnesium of (group 2) and beryllium resembles with aluminium (of group 3) and so on. This similarity in properties of elements present diagonal is called diagonally relationship.

Question 9.
What is the difference between electron gain enthalpy and electro negativity?
Electron gain enthalpy is the amount of energy released when an electron is added to the neutral gaseous atom. But electronegativity is the power of an element to attract covalent electron towards it.

Question 10.
Tell in the form of periods and group that Z = 114 will be placed?
By writing the electronic configuration the maximum value of n will show the period.
114Z = 86[Rn]7s2, 5f14, 6d10, 7p2.
(∵ n = 7) period = 7
Group = 14(10 + 2 + 2).

Question 11.
Why the ionization potential of noble gases are high?
Due to fully filled orbitals and complete octet, the electronic configuration of noble gases are stable. So to remove an electron from outer orbit high energy is required.

Question 12.
Why the electron affinity of halogen elements are high?
Halogens have the large (-ve) value of (∆egH) electron affinity because of their largest effective nuclear charge (Zeff) and smallest size in their respective period. Also they have strong tendency to gain one electron to attain the stable noble gas configuration i.e., from ns2, np5 to ns2, np6 configuration. Thus, the electron gain process is highly exothermic for halogens. However, as we move from Cl to I, the electron enthalpies becomes less and less negative due to corresponding increase in their atomic size.

Question 13.
Whose electron affinity is more among F and Cl and why?
Electron affinity of F is less than Cl because of small size of F, there is more repulsion among the electrons of 2p orbitals. So additive electron does not stabilize the atom.

Question 14.
Why the electron affinity of noble gases is zero?
Noble gases have zero electron gain enthalpy. All noble gases have fully – filled valence shells (ns2, np6). Due to their highly stable (ns2, np6). Configurations, noble gases have absolutely no tendency to take an additional electron. Hence, noble gases have zero electron gain enthalpy.

Question 15.

1. Atomic radius gives us idea about size of the atom. Atomic radius may be taken as the distance between the centre of the nucleus and the outermost shell of the atom. It can be measured either by X – ray or by spectroscopic methods.

1. Ionic radius tells us about size of the ion. It is defined as the distance between centre of the nucleus and the point upto which the ion has influence in the ionic bond.

Question 16.
Why the value of second IE2 is more than IE1 for any element?
IE1 value is the energy needed to remove first electron from the gaseous atom. With the loss of the electron. Therefore, Z eff increases for the remaining electron. Therefore, greater energy is needed to remove the second electron from the gaseous atom. Therefore, IE2 value of an element is always more than its IE1 value.

Question 17.
Or, Size of cation is smaller than the related atom. Why?
Cations are positively charged ions and are formed when a neutral atom loses ‘ one or more valency electrons. Thus a cation possesses the same nuclear charge but less number of electrons as compared to the parent atom. Therefore effective nuclear charge on remaining electrons increases. This causes the decrease in size. Hence, the radius of a cation is always smaller than the radius of the atom from which it is formed.

Question 18.
Anions are negatively charged ions and are formed when one or more electrons are added to the valence shell of a neutral atom. Thus, an anion contains the same nuclear charge but more electrons as compared to its parent atom. This decreases the effective nuclear charge i.e., nucleus exert less influence on the valence electrons. The valence electrons move away from the nucleus and the size increases due to expansion of electron cloud. Hence, the radius of an anion is always larger than that of parent atom.

Question 19.
Why the value of electron affinity of N and Be is near zero?
The 2s orbital of Be is complete so it is stable, so it does not allow the entry of any electron. Similarly the p orbitals of N is half filled and so stable and does not accept any electron. That is why the value of electron affinity for Be and N is near zero.

Question 20.
The size of Mg+2 ion less than O-2 ion whereas the electron configuration of both is same. Explain?
The electronic configuration of Mg+2 ion and O2- ion is same. Both have eight electrons in valence shell but the nuclear charge of Mg+22 ion is +12 whereas that of O2- ion is +8. That is why the nuclear charge for outer electron in Mg+2 is more than O2-, so the size of Mg+2 ion is less than O2- ion.

Question 21.
The value of electron affinity increases on moving left to right in a period. Why?
On moving across a period, the size of the atom decreases and nuclear charge 1 increases. Due to this the attraction for incoming electron increases. That is why the value of electron affinity increases in a period.

Question 22.
Explain the trend of metallic character of elements in periodic table?
Metallic and non – metallic character:
Metallic character of an element is measured in terms of tendency of that element to lose electron i.e., electropositivity. Similarly, non – metallic character of an element is measured in terms of tendency of that element to gain electron i.e., electronegativity.

We know that ionisation enthalpies as well as electronegativities of the element increases along a period from left to right. This implies that the metallic character is maximum on extreme left (alkali metals) while the non – metallic character is the maximum on extreme right (halogens).

Question 23.
Among the O and N whose electron gain enthalpy is more?
Among O and N, the electron gain enthalpy of O is more than N, because in N the valence orbital 2s22p3 is half filled and so stable. So the tendency to accept the incoming electron is more in O than in N. Other than this the size of O atom is smaller than N atom so due to high nuclear charge the incoming electron added fastly.

Classification of Elements and Periodicity in Properties Short Answer Type Questions – II

Question 1.
The first IE1 of Al is less than Mg, Why?
The electronic configuration of A1 and Mg is:
Al13 – 1s2, 2s2, 2p6, 3s2, 3p1
Mg12 – 1s2, 2s2, 2p6, 3s2
The atom will be stable if it have full – filled or half – filled orbitals. In Al the p – orbital is vacant whereas in Mg the s orbital is fully filled. That is the orbital of Al is less stable than Mg. So, the energy required to remove an electron from Al atom is less in compare on to Mg atom, that is why the IE1 of Al is less than Mg.

Question 2.
What do you mean by Shielding effect?
Electrons have negative charge and they repel each other and this force of repulsion decreases the attractive force from nucleus to outer shell, due to this the electrons present in valency shell bounded loosely with the nucleus. In this way, the electrons present between nucleus and outer electrons causes shield and this effect is called Shielding effect.

Question 3.
Discuss the trend in ionization potential from moving left to right in a periodic table?
On moving left to right in a periodic table the nuclear charge increases and atomic radius decreases, due to this the outer electrons are attracted more towards the nucleus and to remove an electron from outer shell more energy is required. So the ionization energy increases on moving left to right.

Question 4.
What do you think that the second electron gain enthalpy of O will be positive, more negative or less negative than first electron gain enthalpy? Give reason for your answer:
O(g) + eg → Og; ∆egH = – 141 kJ mol-1
O(g) + eg → O2-g; ∆eg H = + 780 kJ mol-1
When, on gaining one electron O atom forms O ion than energy is released. So first electron gain enthalpy is negative. But when an electron added in O ion than O2- ion forms and it feels a strong repulsive force and so energy is required for addition of electron. Due to this the second electron gain enthalpy of O is positive.

Question 5.
The electron gain enthalpy of F is less negative than Cl, Why?
Electron gain enthalpy of F is less – ve than that of Cl because when an electron is added to F, the added electron goes to the smaller n = 2 quantum level and suffers repulsion from other electrons present in this level. In case of Cl, the added electron goes to the larger n = 3 quantum level and suffer much less repulsion from other electrons.

Question 6.
Why the long form of periodic table is better than Mendeleev’s periodic table?
The long form of periodic table is better than Mendeleev’s becasuse it has several advantages over Mendeleev’s table. The important advantages are:
1. The arrangement of elements is easy to remember and reproduce.

2. The elements have been classified on the basis of atomic number which is more fundamental than atomic weight of the elements.

3. This periodic table is closely connected to the electronic configuration of ele-ments. Therefore, the position of an element in the periodic table can easily be justified. The electronic configuration of an element can be predicted if its position in the periodic table is known.

4. The resemblances and differences of the properties of elements in periods and groups are explained on the basis of electronic arrangement in various shell i.e., electronic configuration. Thus it reflects trends in physical and chemical properties of the elements.

Question 7.
Write the characteristics of Modern periodic table?
Characteristics of modern periodic table:

1. This periodic table is based on the electronic configuration of elements.
2. It reflects the sequence of filling the electrons in order of sub – energy levels s, p, d an Af
3. It gives clear division of elements, transitional elements and inner transition elements.
4. Attempt has been made to separate metals from non – metals. The elements on the left – hand side of the periodic table are metals and more non – metallic elements have been placed on right side.

Question 8.
Explain, why all transition elements are d-block elements but all d – block elements are not transition elements?
Those elements whose outer electron goes in d – orbital are known as J – block elements or transition elements. The general outer electronic configuration of these elements is
(n – 1)d1-10ns0-2. The electronic configuration of Zn, Cd and Hg is (n – 1 )d10nsd2, but they do not show the properties of transition elements. In ground state and in general oxidation state the d – orbitals of these elements are completely filled, therefore they will not be considered as transition elements. So on the basis of properties all transition elements are d – block elements, but on the basis of electronic configuration all d – block elements are not transition elements.

Question 9.
Among the pairs whose ionization energy is less. Why?

1. Cl or F
2. Cl or S
3. K or Ar
4. Kr or Xe.

1. Among Cl and F ionization energy of Cl is less than F because size of F is small.
2. Among Cl and S, the ionization energy of S is less as the size of S is bigger than Cl.
3. Among K and Ar the ionization energy of K is less as in the valence shell of K only one electron is present which can be easily donated and K+ is formed but Ar has complete octet and so stable and so more energy is required for removal of electron.
4. Among Kr and Xe, the ionization energy of Xe is less because the size of Xe is bigger than Kr.

Question 10.
What are s – block elements? Write their main properties?
Elements in which the last electron enters into s – orbital are known as s – block elements. It includes Alkali metals of group – I and Alkaline earth metals of group – II.
General properties of s – block elements:

1. In outermost orbit the electronic configuration is ns1 or ns2.
2. These elements have definite positive oxidation number which is +1 and +2 in group 1 and group 2.
3. Except hydrogen, all are metals, ionization potentials are low. Thus, these are strong electropositive, strong reducing agent and strongly metallic in nature.
4. Elements of 1st group are alkali metals and elements of 2nd group are called alkaline earth metals.
5. Oxides of these elements are basic in nature. Oxides of 1st group dissolve in water to give alkalis.
6. These elements form electrovalent compounds.
7. These elements provide colour on heating with flame.
8. These elements are very reactive. Reacts with water and acids displacing their hydrogen as H2 gas.

Question 11.
What are p – block elements? Write their main properties?
The elements in which the last electron outers in p – block are called p – block. General characteristic of p – block elements:

1. These elements have 2 electrons in 5 sub – shell and 1 to 6 electrons in p sub-shell in outermost orbit. Only in zero group elements, outer shell is completed.
2. Definite positive or negative oxidation numbers are represented by the elements. Some of the elements show variable valencies.
3. These elements form simple ions as well as complex ions of CO32-, NO3.
4. These are generally non – metals and metalloids. Some of the elements are heavy elements. e,g., Pb, Bi, etc.
5. Oxides are acidic in nature. Some oxides are amphoteric, e.g., PbO, SnO etc.
6. These elements form covalent compounds with each other but with the elements of s – block form electrovalent compounds.

Question 12.
Which elements are called transition elements? Throw light on their properties?
Transition elements:
The elements which are in between s – block and p – block are called as transition elements or elements in which d – orbitals are partially filled are called transition elements. In this, last electron goes to d – orbital of orbit inner to outer so called d – block elements. Boiling point, melting point and densities of these elements are high.
Example: Cr, Mn, Fe, Cu, etc.
Properties:

1. In this two outer orbits are incomplete
2. Electronic configuration is (n – 1) d1 to 10 ns1 to 2
3. It has metallic properties
4. These elements show variable oxidation state
5. These form coloured ions
6. These form complex salt
7. These are good catalysts
8. These form complex compounds
9. These are generally diamagnetic and
10. These also form with non-metal compounds.

Question 13.
What are inner transition elements? Write their general properties?
In these elements three outer orbitals are incomplete. In all these elements, the s – orbital of the last shell (n) is completely filled, the d – orbitals the the penultimate (n – 1) shell invariably contains zero or one electron but the f – orbital of the antipenultimate (n – 2) shell (being lower in energy than d – orbitals of the penultimate shell) gets progressively filled. Hence the general configuration of f – block elements is (n – 2) f1-14 (n – 1) d0 – 1ns2
Lanthanides: The last electron enters in 4f subshell.
Actinides: The last electron enters in 5f subshell.

Question 14.
Explain ionization energy and electron affinity?
Ionization Energy:
The minimum amount of energy which is needed to remove the most loosely bound electron from a neutral isolated gaseous atom in its ground state to form a gaseous cation.
Mg + Energy → M+g + eg
Electron Affinity:
It is equal to the change in enthalpy when an isolated gaseous atom accepts an electron to form a monovalent anion. It is denoted by ∆egH.
In this way, if energy is released by addition of an electron than electron affinity is positive. But if energy is required to add an electron to a negative ion than electron affinity will be negative.

Question 15.
The size of an atom is usually expressed in terms of its radius called atomic radius. It is very important property because several physical and chemical Cation properties are related to it. If the atom is assumed to be spherical than the term atomic radius means the distance from the centre of the nucleus to the outermost shell of electrons. According to quantum mechanical model, the atomic radius is defined as, the distance from the centre of the nucleus to the point up to which the density of electron cloud (i. e., probability of finding the electron) is maximum.

The ionic radius of an ion may be defined as the distance, from its nucleus to the point up to which the nucleus has influence on the electron cloud of the ion.

Question 16.
In a periodic table how the size of the atom changes? Explain?
The distance from the centre of the nucleus to the outermost electron is called atomic radius.
In a period:
As we move left to right in a period atomic number increases by one unit in each successive element. As the addition of electron takes place in the same principal shell, they do not screen each other from the nucleus. It means the increased nuclear charge is not neutralize by the extra valency electron, cause a decrease in the size of atom.

In a group:
In moving down the group, nuclear charge is increasing with the increase in atomic number and we expect that the size of atom should decrease. But at the same time while going from one atom to another, there is increase in the number of electron shells. The effect of increased nuclear charge is reduced by shielding effect (screening effect) of the electrons present in the inner shells. Therefore, the effect of increase in the electron shell is more pronounced than the effect of increase in nuclear charge. Consequently, the atomic size or atomic radius increases down the group.

Question 17.
What is the change in reduction and oxidation properties of elements in periodic table?
Reducing agents:
Those elements which form positive ion by removal of electron, are called reducing agents. The reducing power of any element depends upon the tendency of it to donate electron. The ionization potential of alkali elements are less due to their large size. Therefore they can easily donate electrons, so alkali elements are strong reducing agents. On moving from left to right in a periodic table in a period the reducing power decreases and on moving top to bottom in a group reducing power increases.

Oxidizing agents:
Those elements which accept electrons will work as oxidizing agents. The oxidizing power of any element depends upon its electron accepting property. Halogen easily accept electron so they are strong oxidizing agent. On moving left to right in a period the oxidizing power increases and on moving top to bottom the oxidizing power decreases.

Question 18.
Differenctiate the Modern periodic table and Mendeleev’s periodic table?
Differences between Modem periodic table and Mendeleev’s periodic table:
Modem Periodic Table:

1. In Modem periodic table the elements are arranged in increasing order of atomic number.
2. The elements are kept in 18 groups.
3. The metals and non – metals are separately kept.

Mendeleev’s Periodic Table:

1. In Mendeleev’s periodic table the elements are arranged in increasing order of atomic mass.
2. The elements are kept in 9 groups and 7 sub – groups.
3. The metals and non-metals are not separated.

Classification of Elements and Periodicity in Properties Long Answer Type Questions – I

Question 1.
How many types of elements are there on the basis of electronic configuration? Give example of each?
Or Describe the different types of element on the basis of electronic arrangement?
The elements have been divided into four main classes as follows:

1. Inert gases or Noble gases or Aerogens:
The elements which belongs to the group 18 of the periodic table are known as noble gases or inert gases. They have 8 electrons in their outermost shell (ns2 np6). Therefore their combining capacity or valence is zero. Hence they are inert in nature. All the members are gaseous in nature.

2. Representative elements:
All the elements of s and p – block with exception of noble gases are called representative elements or normal elements. Representative elements includes

These elements have the outer electronic configuration from ns1 to ns2np5. Elements of this type use only thier outer shell electrons in the bonding with the other atom.

3. Transition elements (d – block elements):
These are the elements in which the last shell is not completely filled. These elements have configuration (n – 1)d1-10ns1-2 i.e., these elements contain 1 – 9 electrons in the penultimate shell and 1 – 2 electrons in the valence shell. The elements having partially filled d – orbitals are named transition elements because these represent a transition (change) from the most electropositive elements to the most electronegative elements. Strictly speaking, elements of group 12 (Zn, Cd, Hg) cannot be included in transition elements, although these elements chemically resemble the transition elements in many respects. All these elements closely resemble each other due to the presence of the same number of electrons in the outermost shell.

4. Inner transition elements (f – block elements):
Lanthanides (58Ce – 71Lu) and actinides (90Th – 103Lw) are collectively known as inner transition elements. In these elements, three outermost shells are not completely filled. The last electron in those enters in the (n – 2) subshell. These elements generally have the same number of electrons in the last two shells, (n -1) and n – orbital. The properties of these elements are so close to each other that it is not easy to separate lanthanides from one another.

Question 2.
Write the applications of Mendeleev’s periodic table?
The important contributions of Mendeleev’s periodic table to chemistry are as follows:
1. Systematic study of chemistry:
The Mendeleev’s periodic table categorised the elements for the first time in a systematic way. This helped immensely in the study of the chemistry of elements and their compounds and made the study much easier. If the properties of a group are known, one can easily guess the properties of all the elements (and their compounds) placed in that group.

2. Prediction of new elements:
At the time of Mendeleev’s 56 elements were known. While arranging these elements, he left some gaps. These gaps represented the undiscovered elements. Further Mendeleev predicted the properties of these undiscovered elements f on the basis of their positions. For example, both gallium (Ga) and germanium (Ge) were not known when Mendeleev gave the periodic table. On the basis of their expected positions, he named these elements as Eka – aluminium and Eka – silicon (from the Sanskrit word eka meaning next) because he believed that these would be similar to aluminium and silicon respectively. When these elements were actually found, their properties were the same as predicted by him.

3. Correction of atomic masses:
Mendeleev also corrected the atomic masses of certain elements with the help of their expected positions and properties. For example, be – ryllium was assigned an atomic weight of 13 – 5 on the basis of its equivalent weight (4 – 5) and valency (wrongly calculated as 3). As such, it should have been placed between carbon (atomic mass 12) and nitrogen (atomic mass 14).

But no vacant place was available in between C and N and further more properties of beryllium did not justify such a position. Therefore valency 2 was assigned to beryllium which gave it an atomic weight of 4.5 × 2 = 9, as atomic weight = equivalent weight × valency and it was placed in proper position between lithium (atomic weight 7) and boron (atomic weight 11). In a similar way atomic masses of many other elements were corrected.

4. Useful in research:
Mendeleev’s periodic table is very useful to further research and study of properties of different elements.

Question 3.
What do you mean by electronegativity of an element? How it is different from electron affinity? How electronegativity changes in periodic table?
1. Electronegativity:
It is defined as the relactive tendency of an element to attract the shared pair of electrons in a covalent bond towards itself.

2. Electron affinity:
Electron affinity is equal to the change in enthalpy when an isolated gaseous atom accepts an electron to form a monovalent anion.
The value of first electron affinity is negative as the process is exothermic but in second electron affinity energy is absorbed because the negative ion repels the upcoming electron.

3. Difference:
Electron affinity is the property to bind extra electron whereas electro negativity is the ability of an atom of a molecule to attract electron.
Trend of Electronegativity:

4. In a period:
As we move left to right in a period electronegativity increases due to decrease in size and a corresponding increase in effective nuclear charge.

5. In a group:
While moving down a group the increased nuclear charge is neutralized by screening effect. In general, the electronegativity decreases down the group due to in-creasing size.

Question 4.
The atomic number of an element is 17. By giving the electronic configuration give its place in periodic table?
The electronic configuration of atom having atomic number 17 = Is2, 2s2 2p6, 3s2, 3p5 = 2, 8, 7.
1. Its outer configuration is 3s23p5, so it is a p – block elements.

2. Group:
If one orbital is vacant, than their present in outer shell show the group i.e., the element is 7th group.

3. Sub – group:
If the last electron enter in s or p-sub-shell than the element is of a sub – group.

4. Period:
No. of shells show the no. of periods. This element is member of third period. group – 7, sub – group – A, period – 3.

Question 5.
What do you mean by isoelectronic species? Write the name of any one species which is isoelectronic with following atoms and ions:

1. F
2. Ar
3. Mg2+
4. Rb+

In isoelectronic species the number of electrons are same but the nuclear charge is different. In such species by increasing nuclear charge, the size of the atom increases.

1. Number of e in F = 9 + 1 = 10
2. Number of e in Ar = 18
3. Number of e in Mg+2 = 12 – 2 = 10
4. Number of e in Rb+ = 37 – 1 = 36

N3-, O2-, Ne, Na+ and Al+3 species, F and Mg2+ are isoelectronic species. P3-, S2-, Cl, K+ and Ca2+ species, Ar is isoelectronic species.
Similarly Br, Kr and Sr2+ species and Rb+ is isoelectronic species.

Classification of Elements and Periodicity in Properties Long Answer Type Questions – II

Question 1.
What is Modern Periodic law? Give the description of periodic table based on this law?
Modern Periodic Law:
“The physical and chemical properties of elements are the periodic functions of their atomic numbers”.

Description of Periodic table:
Characteristics of Modern Periodic table:

1. Metals and non – metals are kept separate.

2. Strongly electropositive elements (s – block elements) are kept on the left side of transitional element (d – block elements) in groups 1 and 2, transitional elements are kept in the middle in groups 3 to 12 and non-metallic elements (p – block elements) are kept on the right side of transitional element (d – block elements) in groups 13 to 17.
Thus, s – block
1 – 2 groups
d – block
3 – 12 groups
p – block
13 – 17 groups

3. On the basis of filling of electrons, elements are classified into s – block, p – block, d – block and f – block elements.

4. Transitional elements are kept apart from normal elements. These elements lie in between s – and p – block. In them the last electron enters into (n – 1) d orbital therefore, these elements show similarly in properties rather than gradation.

5. Rare earth elements (Lanthanides and Actinides) are kept away from the main table in a suitable place.

6. Sub – groups of a group are eliminated.

7. Fe, Co and Ni are kept with transitional elements which is the suitable place.

8. Hydrogen is kept in group 1.

9. This classification relates the atomic configuration of elements with their position in the periodic table due to which their study is very easy.

Question 2.
What are the defects of Mendeleev’s periodic table ? How they can be remove by Modern Periodic Law?
Defects of Mendeleev’s Periodic table:
Inspite of remarkable contribution made to chemistry, Mendeleev’s periodic table had certain defects. These defects are as follows:

1. Position of Hydrogen:
Hydrogen is a unique element and resembles with the elements of group IA (alkali metals) as well as with those of VII A group (Halogens) in its properties. Therefore, it should have been placed in both IA and VII A groups in Mendeleev’s periodic table. Therefore, the position of hydrogen in the periodic table is anomalous or controversial.

2. Position of Isotopes:
Isotopes are the atoms of same element having same atomic number but different atomic masses. Therefore, according to Mendeleev’s classification, these should be placed at different places depending upon their atomic masses. However, isotopes have not been given separate places in the periodic table.

3. Position of Lanthanides and Actinides:
Position of lanthanides (14 elements following lanthanum, atomic number 58 – 71) and actinides (14 elements following actinium, 1 atomic number 90 – 103) is also anomalous in Mendeleev’s periodic table. In this periodic table, all these elements are supposed to be placed together in III group which is not in accordance to the periodic law.

4. Similar elements placed in different group:
The elements like silver and thallium, barium and lead, copper and mercury show similar properties, yet they are placed in . different groups in Mendeleev’s periodic table.

5. Eighth group:
This group is full of anomalies. Except osmium, no other elements of this group shows the group valency i.e., no other element is octavalent. Further they all are arranged in three triads without any justification.

6. Cause of periodicity:
Mendeleev could not explain the cause of periodicity among the elements.

Modern periodic table removes a lot of defects of Mendeleev’s periodic table like:

1. Since, this table is based on electronic configuration therefore hydrogen is placed in group 1.

2. To put element of high atomic mass before the element of low atomic mass:
In Mendeleev’s periodic table, elements of higher atomic mass precedes elements of lower atomic mass. But their position is justified in Modem periodic table because element with higher atomic mass has lower atomic number, e.g., cobalt (atomic mass 58.93, atomic number 28).

3. Position of Isotopes:
Isotopes of an element possess same atomic number, therefore, there is no necessity of giving them separate positions in the periodic table.

4. To place elements with different properties at the same place: In Mendeleev’s periodic table, elements of different properties are placed at same place. For example, elements of sub – group IA and IB. But in Modem periodic table, this defect is removed by separating the sub – groups.

5. Position of Noble gas:
Noble gases have been placed between VIIIB electromotive elements and IA electropositive elements.

6. Diagonal relationship:
It can be explained on the basis of electronic configuration and atomic radius.

Question 3.
Compare the Mendeleev’s periodic table and Modern periodic table?
Comparison between Mendeleev’s periodic table and Modem periodic table:
Mendeleev’s periodic table:

1. Elements are arranged in order of their increasing atomic weight.
2. There are 9 horizontal blocks known as groups.
3. Zero group is added later on.
4. Elements of different properties kept in same group.
5. Every isotope don’t have different place.

Modem periodic table:

1. Elements are arranged in order of their increasing atomic number.
2. There are 18 horizontal blocks called groups.
3. Nobel gases are at the end of each period.
4. Elements of different properties kept in different groups.
5. There is no need to give different place to isotopes.

Question 4.
What is ionisation energy? Explain the factors affecting ionisation energy?
Ionization energy:
The amount of energy which is required to separate the outermost electron from atom is called ionic energy or ionic potential i.e., to convert atom into positive ion the necessary energy required is called as ionic energy. Its unit is kJ mole-1.
So, M(g) + Energy → M+(g) + e
M(g) – Electron → M+(g) + Ionisation energy.
Factors affecting ionization energy:
1. Size of atom or ion:
Greater the size of atom or ion, weaker are the forces of attra-ction and lower is the value of ionization energy.

2. Nuclear charge:
Greater the nuclear charge, more is the attraction for electrons and hence, greater is the value of ionization energy.

3. Shielding effect:
In multielectron atom with increase in atomic number, shielding effect increases due to which valence shell electron feels lesser attraction and hence, value of ionization energy is lower.

4. Penetration effect:
Simple the shape of orbitals, more is the penetration of it for the nucleus. That is it experience greater attraction. Thus, value of ionization energy is higher. It follows the order s > p > d > f

5. Electronic configuration:
Completely filled and half-filled orbitals are more stable than any other arrangement. Thus, value of ionization energy is higher for it.

6. Trend in periodic table:

(a) In period:
Ionisation energy generally increases from left to right in a period. This is due to gradual increase in nuclear charge and decrease in atomic size of the elements.

(b) In a group:
There is a gradual decrease in ionisation energy moving top to bottom. This is due to increase in the number of the main energy shell i.e., size of the atom increases.

Question 5.
What do you mean by Electron affinity? Explain the factors affecting it?
Electron affinity:
The electron affinity of an atom is the energy released, when an electron is added to a neutral atom. It is expressed in kcal/mol. Electronegativity is a relative number on an arbitrary scale while electron affinity is expressed in the units of energy (kcal/mol).
Ag + e → Ag + E1 (Exothermic)
Ag + e → A2- – E2 (Endothermic)
A2-g + e → A3-g – E3 (Endothermics)
This way in the addition of an electron to neutral isolated atom energy is released and electron affinity is positive, but in an anion energy is required to add an electron because the anion opposes the intrance of an electron. Thus, electron affinity for charged ions is negative.

Factors affecting electron gain enthalpy: Some important factors affecting electron gain enthalpy are:

1. Atomic size
2. Nuclear charge
3. Electronic configuration.

1. Atomic size:
As the atomic size increases, the distance between the nucleus and the incoming electron increases. This results in lesser attraction. Consequently, electron gain enthalpy becomes less negative.

2. Nuclear charge:
Greater the nuclear charge, greater will be attraction for the incoming electron and as a result, AH becomes more negative.

3. Electronic configuration:
Atoms having stable electronic configuration have lesser tendency to accept the electron. Due to this the value of electron gain enthalpy becomes less negative.

4. Half – filled and ful – filled orbitals:
Half – filled and ful – filled orbitals are stable and in this condition they cannot easily accept the electron. So the value of electron affinity for such elements are zero.

5. Periodicity:

1. In periods: On moving left to right in a period the electron affinity increases as the nuclear charge increases.
2. In group: On moving down the group the electron affinity decreases due to increase in size of atom.