Class 10 Science Periodic Classification of Elements Exam Notes

This Class 10 Science exam notes page covers the chapter Periodic Classification of Elements. It explains the need for classification, early attempts by Lavoisier, Dobereiner, Newlands and Mendeleev, the modified periodic table with groups and subgroups, the defects of Mendeleev's table, the modern periodic law and the long form of the periodic table, and the periodic trends in atomic size, ionisation energy, electron affinity and metallic and non-metallic character.

Topics Covered

1. Introduction and important terms
2. Lavoisier's classification
3. Dobereiner's law of triads
4. Newlands' law of octaves
5. Mendeleev's periodic law and periodic table
6. Modified version of Mendeleev's periodic table
7. Discrepancies in Mendeleev's periodic table
8. Modern periodic table and long form
9. Periodicity in properties
10. Metallic and non-metallic character
11. Points to remember

Introduction

One hundred and fifteen elements are known till date, and many more may be known in future. All elements differ in their properties due to the presence of different kinds of atoms in them. Elements combine to form millions of compounds.

Important Terms and Concepts

1. Need of classification: It is difficult to study each and every element individually and to know its properties and uses. Therefore, elements have been classified into groups on the basis of their similarities in properties.
2. Classification: Classification means grouping of elements on the basis of similarities in properties. For example, all of you belong to Class X because you resemble each other in some properties: all of you have passed Class IX and all of you are in the age group of 13 to 15 years.
3. Basis of classification: Classification is done on the basis of similarities in properties so that a systematic study can be made of the elements.

EARLY ATTEMPTS OF CLASSIFICATION

Lavoisier's Classification

Lavoisier classified elements into metals and nonmetals. This classification was based on certain distinctive physical properties such as hardness, malleability and lustre. On the basis of these properties, sodium and lead were grouped together as metals.

Limitations
(1) Hardness, malleability and lustre were found to be the only common properties of sodium and lead. Otherwise, the two elements were entirely different.
(2) In this classification there was no place for elements with properties resembling those of both metals and nonmetals.

Therefore, Lavoisier's classification was found to be inadequate.

Dobereiner's Classification (Law of Triads)

In 1817, the German chemist Johann Dobereiner classified elements having similar chemical properties into groups of three. These groups were called triads. He proposed a law known as Dobereiner's law of triads. According to this law, when elements are arranged in the order of increasing atomic mass in a triad, the atomic mass of the middle element is approximately equal to the arithmetic mean of the atomic masses of the other two elements.

The classification of elements into triads was very successful in predicting the atomic mass and properties of the middle element. Further, this classification showed that there exists some relationship between the properties of elements and their atomic masses. This paved the way for future attempts at classification of elements.

Limitation: All the elements could not be grouped into triads.

Newlands' Classification (Law of Octaves)

In 1864, John Newlands, an English chemist, showed that when elements are arranged in the order of their increasing atomic masses, the eighth element, starting from a given element, is a kind of repetition of the first one, like the eighth note in an octave of music.

sa re ga ma pa dha ni sa, where the first and the eighth note are the same.

A part of Newlands' classification is given below, where the figures under the symbols show the atomic masses.

Octaves of music and Newlands' arrangement of elements (column labels)

The eighth note repeats the first (sa, or do, again). Each row of the following table starts again under sa/do.

Starting from lithium (Li), the eighth element is sodium (Na). The eighth element starting from sodium is potassium. The properties of lithium, sodium and potassium are similar. The properties of beryllium, magnesium and calcium are similar too.

Limitations
(i) This law worked well for lighter elements (up to calcium), but it could not be applied to heavier elements, because starting from calcium every eighth element was found to have properties different from those of the first element.
(ii) Newlands emphatically said that only 56 elements exist in nature and no more elements are likely to be discovered in future. This concept was later found to be untrue with the discovery of many new elements.
(iii) In arranging elements in the form of a table, Newlands clubbed two elements together at the same place and in the same column. He also placed some dissimilar elements in the same column. For example, cobalt (Co) and nickel (Ni) were clubbed together in the column of fluorine (F), chlorine (Cl) and bromine (Br) (under sa/do). Cobalt and nickel have properties entirely different from those of fluorine, chlorine and bromine. Cobalt and nickel have properties similar to those of iron, but iron (Fe) was placed in a column (under ni/ti) different from the column of cobalt and nickel.

However, this law supported the idea that the properties of elements depend upon their atomic masses. It also showed that the properties of elements are repeated after a certain interval, i.e., the properties of elements are periodic in nature.

MENDELEEV'S PERIODIC LAW AND PERIODIC TABLE

While working systematically on the physical and chemical properties of elements, Dmitri Ivanovich Mendeleev noticed that the properties of elements varied regularly with their atomic masses. He arranged the 63 elements then known in a table on the basis of similarities in properties. It was found that most of the elements occupied places in the table in order of their increasing atomic masses. In 1869, Mendeleev formulated a law, now known as the periodic law.

Mendeleev's Periodic Law: The properties of elements are periodic functions of their atomic masses. This means that if the elements are arranged in order of increasing atomic masses, then those with similar properties are repeated at regular intervals.

On the basis of the periodic law, Mendeleev presented his classification in the form of a table, now known as Mendeleev's periodic table. A simplified version of this periodic table is given below. In this table, copper, silver and gold find places in groups I as well as VIII.

This table consists of vertical columns called groups and horizontal rows called periods. There are only eight groups in the table. Mendeleev left some vacant places (shown by question marks) for the yet undiscovered elements. Noble gases were not discovered at that time, so he did not provide any place for them in his periodic table.

Mendeleev's idea was remarkable in that he used a fundamental atomic property (atomic mass) as the basis of classification. While classifying elements, he laid special emphasis on two factors:
1. Similar elements were grouped together.
2. Elements were arranged in order of increasing atomic masses.

Modified Version of Mendeleev's Periodic Table

The elements which were undiscovered, and for which Mendeleev had left vacant places, were discovered later. Some of these are scandium (Sc), gallium (Ga) and germanium (Ge). These elements were accommodated in their proper places in the table. The elements helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn) became known only towards the end of the nineteenth century. These elements, called noble gases, were placed in the table as a separate group, called group 0. The periodic table had to be modified then. The modified version of the table is shown below.

Lanthanide series* (along with lanthanum): Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu

Actinide series** (along with actinium): Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr

Features of the Modified Version of Mendeleev's Periodic Table

1. Groups into subgroups: Each group of this periodic table is further divided into two subgroups, A and B. The properties of elements within a subgroup resemble each other more markedly, but they differ from those of the elements of the other subgroup. For example, lithium (Li), sodium (Na), potassium (K), etc., of subgroup IA have a close resemblance of properties, but they have hardly any resemblance to the coinage metals (Cu, Ag and Au) of subgroup IB.
Mendeleev allowed the subgroups to be represented within the same group.

2. Prediction of errors: This periodic table could predict errors in the atomic masses of some elements on the basis of their position in the periodic table. For example, when the periodic table was published, the experimental value of the atomic mass of beryllium (Be) was supposed to be 13.65 and its valency 3. So, the position of Be should have been somewhere else, but Mendeleev placed it at its appropriate position on the basis of its properties. He further suggested that the atomic mass of Be needed correction. Mendeleev predicted its atomic mass to be 9.1 and its valency 2. Later investigations proved him right.
Similarly, the atomic mass of uranium was corrected from 120 to 240. Corrections were also made in the atomic masses of gold, platinum, etc.

3. Prediction of properties of higher and undiscovered elements: Mendeleev classified the elements in order of their increasing atomic masses. However, this order had to be ignored at some places to make sure that elements with similar properties fell in the same group. In doing so, he left some vacant places in the table. These vacant places were kept reserved for elements not yet discovered. Mendeleev was confident that these elements would be discovered later and would occupy these vacant places. He also predicted the properties of these undiscovered elements on the basis of his study of the properties of the neighbouring elements. When the missing elements of Mendeleev's periodic table were discovered subsequently, their properties were found to be very similar to those predicted by Mendeleev.

The elements scandium, gallium and germanium were not known in 1871, but their existence was predicted by Mendeleev. He named these elements eka-boron, eka-aluminium and eka-silicon. When these elements were discovered, they were named scandium, gallium and germanium respectively, and their properties were found to be in good agreement with those predicted by Mendeleev. The properties of eka-aluminium (predicted by Mendeleev) and those of gallium (discovered later) are given in the table below.

Considering its atomic mass, titanium (Ti) should have been placed below aluminium in the periodic table, but Mendeleev placed it below silicon (Si) because the properties of titanium were similar to those of silicon. Thus, a gap was left below aluminium in the periodic table. This gap was filled up by gallium, which was discovered later. The properties of gallium (Ga) were found to be similar to those of boron and aluminium.

4. Basic features intact: All the basic features of Mendeleev's periodic table are intact even today. When a new class of elements, the noble gases, was discovered, they found a place in a separate group called the zero group. The existing order of the periodic table was not disturbed.

Discrepancies in Mendeleev's Periodic Table

Mendeleev's periodic table has the following defects.

1. Position of hydrogen: The position of hydrogen in the periodic table is anomalous. Hydrogen resembles the alkali metals (Li, Na, K, etc.), so it may be placed in the group of the halogens (VII A) as well.

2. Position of lanthanides and actinides: The elements from atomic number 57 to 71 are collectively known as lanthanides. They do not have a proper place in the periodic table. They have all been placed at the same position, in group III and period 6. Similarly, the actinides (atomic numbers 89 to 103) also have no proper place in the periodic table. These elements have also been placed at the same position, in group III and period 7.

3. Some similar elements are separated, while some dissimilar elements are placed in the same group: Some similar elements are separated in the periodic table. For example, copper (Cu) and mercury (Hg), silver (Ag) and thallium (Tl), and barium (Ba) and lead (Pb). On the other hand, some dissimilar elements have been placed together in the same group. For example, copper (Cu), silver (Ag) and gold (Au) have been placed in group I along with the alkali metals. Similarly, manganese (Mn) is placed in the group of the halogens.

4. Presence of some anomalous pairs of elements: In Mendeleev's periodic table, elements are arranged in order of increasing atomic mass. In some places, this order has been ignored.
(a) The atomic mass of argon is 40 and that of potassium is 39, but argon is placed before potassium in the periodic table.
(b) The positions of cobalt and nickel are not in proper order. Cobalt (atomic mass 58.9) is placed before nickel (atomic mass 58.6).
(c) Tellurium (atomic mass 127.6) is placed before iodine (atomic mass 126.9).
(d) Thorium (atomic mass 232.12) is placed before protactinium (atomic mass 231).

5. Position of isotopes: The isotopes of an element have no place in the periodic table.

The failure of Mendeleev's periodic law to explain the wrong order of atomic masses of some elements, and the position of isotopes, led scientists working in this field to conclude that atomic mass cannot be the basis for the classification of elements. There must be a more fundamental property of elements which can be the basis of classification.

MODERN PERIODIC TABLE

Henry Moseley, an English physicist, found that the atomic number (Z) is the fundamental property of an element, and not the atomic mass, for the classification of elements.

Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers, i.e., the number of protons or electrons present in the neutral atom of an element.

Long Form of Periodic Table: The elements are arranged in increasing order of their atomic numbers. The prediction of the properties of elements and their compounds can be made with precision. All drawbacks of Mendeleev's periodic table vanish when the elements are arranged on the basis of increasing atomic numbers.

Elements in a Group

(1) They show similar chemical properties due to similar outer electronic configuration, i.e., the same number of valence electrons.
(2) They show a gradation in properties due to the gradually varying attraction of the nucleus on the outer valence electrons as we go down the group.

Main Features of the Long Form of the Periodic Table

(1) It shows the arrangement of elements based on the modern periodic law.
(2) There are 18 vertical columns known as groups.
(3) There are 7 horizontal rows known as periods.
(4) Elements having similar outer electronic configurations, i.e., the same number of valence electrons, are placed in the same group. For example:

Group 1 (K L M N O P Q shells)
H (1): 1
Li (3): 2, 1
Na (11): 2, 8, 1
K (19): 2, 8, 8, 1
Rb (37): 2, 8, 18, 8, 1
Cs (55): 2, 8, 18, 18, 8, 1
Fr (87): 2, 8, 18, 32, 18, 8, 1

Group 2 (K L M N O P Q shells)
Be (4): 2, 2
Mg (12): 2, 8, 2
Ca (20): 2, 8, 8, 2
Sr (38): 2, 8, 18, 8, 2
Ba (56): 2, 8, 18, 18, 8, 2
Ra (88): 2, 8, 18, 32, 18, 8, 2

Group 13 (K L M N O P shells)
B (5): 2, 3
Al (13): 2, 8, 3
Ga (31): 2, 8, 18, 3
In (49): 2, 8, 18, 18, 3
Tl (81): 2, 8, 18, 32, 18, 3

Group 14 (K L M N O P shells)
C (6): 2, 4
Si (14): 2, 8, 4
Ge (32): 2, 8, 18, 4
Sn (50): 2, 8, 18, 18, 4
Pb (82): 2, 8, 18, 32, 18, 4

Group 15 (K L M N O P shells)
N (7): 2, 5
P (15): 2, 8, 5
As (33): 2, 8, 18, 5
Sb (51): 2, 8, 18, 18, 5
Bi (83): 2, 8, 18, 32, 18, 5

Group 16 (K L M N O P shells)
O (8): 2, 6
S (16): 2, 8, 6
Se (34): 2, 8, 18, 6
Te (52): 2, 8, 18, 18, 6
Po (84): 2, 8, 18, 32, 18, 6

Group 17 (K L M N O P shells)
F (9): 2, 7
Cl (17): 2, 8, 7
Br (35): 2, 8, 18, 7
I (53): 2, 8, 18, 18, 7
At (85): 2, 8, 18, 32, 18, 7

Group 18 (K L M N O P shells)
He (2): 2
Ne (10): 2, 8
Ar (18): 2, 8, 8
Kr (36): 2, 8, 18, 8
Xe (54): 2, 8, 18, 18, 8
Rn (86): 2, 8, 18, 32, 18, 8

(5) In periods, elements in which the number of electrons in the outermost shell increases gradually in steps of one are placed. For example:

Period 1 (K shell): H (1) 1 and He (2) 2
Second period (K, L shells): Li (3) 2, 1; Be (4) 2, 2; B (5) 2, 3; C (6) 2, 4; N (7) 2, 5; O (8) 2, 6; F (9) 2, 7; Ne (10) 2, 8

(6) Each group in the table signifies an identical outer shell electronic configuration, i.e., the same valence electrons. For example, group 1 has 1 valence electron, group 2 has 2 valence electrons, group 13 has 3, and group 14 has 4 valence electrons.

(7) Each period starts with the filling of a new shell. For example:
1st period: the K shell (1st shell) starts filling with hydrogen and ends at helium.
2nd period: the L shell (2nd shell) starts filling from Li (3) up to Ne (10).
3rd period: the M shell (3rd shell) starts filling from Na (11) up to Ar (18).
4th period: the N shell (4th shell) starts filling from K (19) up to Kr (36), and so on.

(8) The periodic table is divided into four blocks:
(a) s-block elements: Group 1 and 2 elements are called s-block elements.
(b) p-block elements: Group 13 to 18 elements are called p-block elements.
(c) d-block elements: Group 3 to group 12 are called d-block elements or transition elements (in between s-block and p-block elements).
(d) f-block elements: The elements placed at the bottom of the periodic table are known as f-block elements. The fourteen elements after La (57, lanthanum) are called lanthanoids, and the fourteen elements after Ac (89, actinium) are called actinoids.

Naturally Occurring and Synthetic Elements

Elements up to atomic number 92 occur in nature, except technetium, Tc (Z = 43) and promethium, Pm (Z = 61), which are formed from radioactive elements. Here Z represents the atomic number.

Elements beyond atomic number 92 are man-made elements. They are also called synthetic elements.

Groups

(1) Elements in a group have the same number of valence electrons.
(2) The chemical properties of elements in a group are similar because of the same number of valence electrons. For example, all the group 1 elements have 1 valence electron. They form positively charged ions by losing one electron when the required amount of energy is supplied, i.e., Li⁺, Na⁺, K⁺.

Group 1 elements are called alkali metals. Group 2 elements are called alkaline earth metals. Group 2 elements have 2 valence electrons in the outermost shell. They can lose both valence electrons to form dipositive cations, i.e., Be²⁺, Mg²⁺, Ca²⁺, etc. Positively charged ions are called cations.

Group 13 elements belong to the boron family, group 14 to the carbon family, group 15 to the nitrogen family and group 16 to the oxygen family.

Group 16 elements contain 6 valence electrons in their outermost shells, i.e., two electrons less than the maximum number of electrons that can be present in the outermost shell. They can gain 2 electrons more easily than lose 6 electrons. They change into dinegative ions such as O^2- and S^2-. Group 17 elements, called halogens, contain 7 valence electrons. They can gain one electron to acquire a stable electronic configuration, i.e., 8 electrons in the outermost shell, and form uninegative (single negative) ions such as F^-, Cl^-, Br^- and I^-. Negatively charged ions are called anions.

Group 18 elements, called noble gases, have their outermost shell completely filled. The elements of this group have no tendency to lose or gain electrons. Thus, these elements have zero valency and are almost unreactive. Hence they are called noble gases. However, nowadays, compounds of Kr, Xe and Rn have been prepared.

In any particular group, the number of shells increases but the number of valence electrons remains the same.

Periods

(1) The horizontal rows in the periodic table are called periods.
(2) There are 7 periods in the long form of the periodic table.
(3) The first period contains 2 elements, hydrogen and helium. They have only one shell.
(4) The second period contains 8 elements: lithium Li (3), beryllium Be (4), boron B (5), carbon C (6), nitrogen N (7), oxygen O (8), fluorine F (9) and neon Ne (10). The second period has 2 shells (K and L), and the L shell is progressively filled.
(5) The elements of the 3rd period are Na (11), Mg (12), Al (13), Si (14), P (15), S (16), Cl (17) and Ar (18).

In the 3rd period, the 3rd shell (M shell) is progressively filled, and there are three shells.
(6) The 4th period has 18 elements.
(7) The 5th period has 18 elements.
(8) The 6th period has 32 elements.
(9) The 7th period has 32 elements.
In periods, the number of valence electrons increases from left to right in s- and p-blocks.

Long form of the Periodic Table (atomic number, symbol, atomic mass, name). Key: white cells = metals, grey cells = non-metals, black cells = metalloids. Groups 1-2: s-block elements; groups 3-12: d-block elements; groups 13-18: p-block elements; group 18 = noble gases; lanthanide and actinide series at the bottom = f-block elements.

PERIODICITY IN PROPERTIES

The properties of elements depend upon their electronic configuration, which changes along a period and down a group in the periodic table. There is periodicity in properties, i.e., repetition of properties after a regular interval, due to similarity in electronic configuration.

Atomic Size (Atomic Radius)

Atomic size means the radius of an atom. It is defined as the distance between the centre of the nucleus and the outermost shell (valence shell) of an isolated atom.

Covalent radius: It is defined as half of the distance between the centres of nuclei of two atoms (bond length) bonded by a single covalent bond. For example, the bond length in the H-H (hydrogen molecule) is 74 pm.
Covalent radius = 1/2 x 74 pm = 37 pm (picometre)
[1 pm = 10^-12 m]
It can be measured in the case of diatomic molecules of non-metals.

Metallic radius: It is defined as half of the internuclear distance between two metal ions in a metallic crystal. It is measured in the case of metals.

Variation of atomic size in a group: Size generally increases from top to bottom in a group.
Reason: It is due to the addition of a new shell, i.e., the number of shells goes on increasing.

Group 1 elements and their atomic radii:

Variation of atomic size along a period: Atomic size goes on decreasing along a period from left to right.
Reason: It is due to the increase in nuclear charge (number of protons in the nucleus), which pulls the electrons towards it. The force of attraction between the nucleus and the valence electrons increases, therefore the atomic size decreases.

Elements of the 2nd and 3rd periods and their atomic radii:

IONISATION ENERGY AND ELECTRON AFFINITY

The chemical nature and reactivity of an element depend upon the ability of its atoms to donate or accept electrons. This can be measured quantitatively with the help of ionisation energy and electron affinity of its atoms.

Ionisation energy: It is defined as the energy required to remove an electron completely from an isolated gaseous atom of an element. The energy required to remove the first electron is called the first ionisation energy.
A(g) + I.E₁ → A⁺(g) + e^-

Second ionisation energy: The energy required to remove an electron from a unipositive ion is called the second ionisation energy.
A⁺(g) + I.E₂ → A²⁺(g) + e^-
The second ionisation energy is greater than the first ionisation energy due to the increase in positive charge, hence the increase in the force of attraction between the nucleus and the valence electron.
1st I.E. < 2nd I.E. < 3rd I.E.

Variation of ionisation energy down a group: Ionisation energy goes on decreasing down a group.
Reason: As the atomic size increases down a group, the distance between the valence electrons and the nucleus increases. The force of attraction between the nucleus and the valence electrons decreases, therefore the energy required to remove the electron, i.e., the ionisation energy, goes on decreasing.

Example:

Variation of ionisation energy along a period: It generally goes on increasing along a period from left to right, with a decrease in atomic size.
Reason: Due to the decrease in atomic size, the force of attraction between the valence electrons and the nucleus increases, and therefore more energy is required to remove an electron.

Example, ionisation energy of elements of the 2nd period:

There is a decrease in ionisation energy from Be to B and from N to O. The reason for this will be studied in higher classes.
Group 18 elements (noble gases) have the highest ionisation energy in their respective periods due to their stable electronic configuration, i.e., 8 electrons in their valence shells, except He, which has 2 electrons.

Electron Affinity

It is the energy change when an electron is gained by a neutral gaseous atom, converting it into a negatively charged ion. It is a measure of the attraction or affinity of the gaseous atom for an extra electron to be added.
Cl(g) + e^- → Cl^-(g) + E.A.

Factors: It depends upon atomic size as well as electronic configuration.

Variation down a group: Electron affinity generally goes on decreasing down a group.
Reason: Due to the increase in atomic size, the force of attraction between the nucleus and the electron to be added becomes less.

Variation along a period: Electron affinity increases from left to right in a period.
Reason: It is due to the decrease in atomic size, which leads to an increase in the force of attraction between the nucleus and the electron to be added.

Example, electron affinity of group 17 elements:

However, deviations from this rule are observed in the variation of electron affinity.

METALLIC AND NON-METALLIC CHARACTER

Groups 1 to 12 are metals. Groups 13 to 18 comprise non-metals, metalloids and metals.

Metalloids: Elements which resemble both metals and non-metals are called metalloids. They are also called semi-metals, e.g., boron, silicon, germanium, arsenic, antimony, tellurium and polonium.

Properties of metals:
(i) They are malleable.
(ii) They are ductile.
(iii) They are good conductors of heat and electricity.
(iv) They generally have 1 to 3 valence electrons.
(v) They have the same or fewer electrons in their outermost shell than the number of shells.
(vi) They are mostly solids.

Properties of non-metals:
(i) They exist in solid, liquid or gaseous state.
(ii) They are generally brittle.
(iii) They are non-conductors.
(iv) They have 4 to 8 valence electrons.

Variation of metallic character: Metallic character increases down a group due to the decrease in ionisation energy. It decreases along a period due to the increase in ionisation energy from left to right.

Variation of non-metallic character: Non-metallic character decreases down a group because of the decrease in electron affinity, which is due to the increase in atomic size. Along a period, non-metallic character increases from left to right due to the increase in electron affinity, which is due to the decrease in atomic size.

POINTS TO REMEMBER

Dobereiner's law of triads states that in a given set of three elements (triad), the atomic mass of the middle element is approximately equal to the average of the atomic masses of the other two elements.

Newlands' law of octaves states that if the elements are arranged in order of increasing atomic mass, the eighth element starting from a given element shows a repetition of the properties of the first element.

According to Mendeleev's periodic law, the physical and chemical properties of elements are a periodic function of their atomic masses.

Mendeleev's periodic table (i) helped correct the wrongly assigned values of atomic masses of some elements, and (ii) predicted the properties of some undiscovered elements.

In Mendeleev's periodic table, (i) isotopes had no place, and (ii) some similar elements are separated, while some dissimilar elements are placed in the same group.

In Mendeleev's periodic table there are eight groups (I to VIII). Each group is divided into subgroups A and B, except group VIII. In the modern periodic table there are 18 groups.

The vertical columns of the periodic table are called groups. The horizontal rows of the periodic table are called periods.

Modern periodic law: The physical and chemical properties of elements are a periodic function of their atomic numbers.

The modern periodic table (long form) is based on atomic number.

There are seven periods in the periodic table. The first six periods consist of 2, 8, 8, 18, 18 and 32 elements respectively. The seventh period is still incomplete.

Elements in a group have similar properties.

The regular occurrence of similar chemical properties of elements with increasing atomic number is known as chemical periodicity.

A series of transition elements starting from actinium are called actinides or actinoids.

A family of elements headed by helium in the periodic table constitutes the noble gases. They are unreactive.

An element that is intermediate between metals and non-metals is known as a metalloid.

The elements of group 1 are called alkali metals.

The elements of group 2 are called alkaline earth metals.

Question. Which of the following pairs of elements does not belong to same group ?
(A) Cl, Br
(B) N, P
(C) Mg, Ca
(D) Al, Si
Answer : D

Question. Which of the following statement is False?
(A) In modern periodic table each block contains a number of columns equal to the number of electrons than can occupy that sub-shell.
(B) The greatest increase in ionization enthalpy is experienced on removal of electron from core noble gas configuration.
(C) In 13th group, values of electronegativity increases regularly from top to bottom
(D) Anything that influences the valence electrons will affect the chemistry of the element and the valence shell is not affected by nuclear mass.
Answer : C

Question. The atomic number of an element x is 12. The element left to it and right to it in same period of periodic table are -
(A) Be, Ca
(B) Al, P
(C) K, Li
(D) Na, Al
Answer : D

Question. An element has atomic number 13. Which of the following statements concerning this element is wrong ?
(A) the atom has 13 protons
(B) the element is placed in the third period of the periodic table
(C) the element is placed in group IV in the periodic table.
(D) the formula of its oxide is X₂O₃
Answer : C

Question. Which one of the following belongs to representative group of elements in the periodic table ?
(A) Lanthanum
(B) Argon
(C) Chromium
(D) Aluminium
Answer : D

Question. Elements of which group form anions most readily ?
(A) Oxygen family
(B) Nitrogen family
(C) Halogens
(D) Alkali metals
Answer : C

Question. Mendeleev classified elements in -
(A) increasing order of atomic number.
(B) increasing order of atomic masses.
(C) decreasing order of atomic masses.
(D) decreasing order of atomic number
Answer : B

Question. Hydrogen has the tendency to gain one electron to acquire helium configuration. In this respect it resembles -
(A) alkali metals
(B) carbon
(C) alkaline earth metals
(D) halogens
Answer : D

Question. Elements A, B, C, D and E having the following electronic configuration -
(i) 1s², 2s², 2p¹ (ii) 1s², 2s², 2p⁶, 3s², 3p¹
(iii) 1s², 2s², 2p⁶, 3s², 3p³ (iv) 1s², 2s², 2p⁶, 3s², 3p⁵
(v) 1s², 2s², 2p⁶, 3s², 3p⁶
Electronic configurations belonging to same group are -
(A) (i) and (iii)
(B) (i) and (iv)
(C) (i) and (ii)
(D) (i) and (v)
Answer : C

Question. The most electropositive element in periodic table is:
(A) F
(B) Na
(C) K
(D) Cs
Answer : D

Question. Modern periodic law was proposed by-
(A) Mendeleev
(B) Henry Moseley
(C) Werner
(D) Bohr and Bury
Answer : B

Question. The long form of periodic table consists of -
(A) seven periods & eight groups.
(B) seven periods & eighteen groups.
(C) eight periods & eighteen groups.
(D) eighteen periods & eight groups.
Answer : B

Question. Which of the following is smallest in size -
(A) O^2–
(B) C^4–
(C) F^–
(D) N^3–
Answer : C

Question. The element with atomic number 56 is likely to have the same outer shell configuration as the element with atomic number :
(A) 12
(B) 18
(C) 14
(D) 30
Answer : A

Question. Which of the following statement is false ?
(A) The metallic character on moving from left to right in a period, decreases.
(B) Metallic character increases on going down a group.
(C) atomic radius decreases on moving from left to right in a period.
(D) The chemical reactivity of metals, on moving from left to right in a period increases.
Answer : D

Question. The ionisation energy of isotopes of an element will be -
(A) same
(B) different
(C) dependent on atomic masses
(D) dependent on the number of neutrons present in the nucleus
Answer : A

Question. In modern periodic table, the number of verticle columns are :
(A) 18
(B) 07
(C) 08
(D) 09
Answer : A

Question. The early attempts to classify elements were based on -
(A) atomic number
(B) atomic mass
(C) electronic configuration
(D) None of these
Answer : B

Question. Cl, Br, I, if this is a Dobereiner’s triad and the atomic masses of Cl and I are 35.5 and 127 respectively, then the atomic mass of Br is -
(A) 162.5
(B) 91.5
(C) 81.25
(D) 45.625
Answer : C

Question. According to Newlands’ law of octaves 3rd element will resemble in its properties to -
(A) ninth element
(B) eighth element
(C) tenth element
(D) eleventh element
Answer : C

Question. Which of the following is the most reactive halogen ?
(A) F
(B) Cl
(D) Br
(D) I
Answer : A

Question. Which of the following pairs of elements belong to the same period of the periodic table ?
(A) C, Mg
(B) N, Ar
(C) Ca, Cl
(D) K, Cu
Answer : D

Question. The number of elements in the 4th period of periodic table are -
(A) 8
(B) 18
(C) 10
(D) 32
Answer : B

Question. The difference between ions and atoms is of -
(A) relative size
(B) electronic configuration
(C) presence of charge
(D) All of these
Answer : D

Question. Newlands could classify elements correctly only upto -
(A) copper
(B) chlorine
(C) calcium
(D) chromium
Answer : C

Question. All the members in a group of long form of periodic table have the same -
(A) valency
(B) number of valence electrons
(C) chemical properties
(D) All of these
Answer : D

Question. An element M has an atomic number 9 and atomic mass 19. Its ion will be represented by-
(A) M
(B) M⁺²
(C) M^–
(D) M^–2
Answer : C

Question. In the third period of the periodic table, the element having smallest size is -
(A) Na
(B) Ar
(C) Cl
(D) Si
Answer : C

Question. Which of the following elements has maximum metallic character ?
(A) Li
(B) N
(C) Na
(D) P
Answer : C