# Periodic Table — Periodic Properties & Variation of Properties

## 2009

#### Question 1

Among Period-2 elements — Lithium; Carbon; Fluorine; Neon - State the one which has high electron affinity.

Fluorine has the highest electron affinity.

#### Question 2

Consider the section of the periodic table given below:

Group
numbers
1A
1
IIA
2
IIIA
13
IVA
14
VA
15
VIA
16
VIIA
17
O
18
Li D  OJNe
AMgESi HK
BC FG  L

Some elements are given in the above table in their own symbol and position in the periodic table, while others [shaded] are shown with a letter. With reference to the table:

(i) Which is the most electronegative?

(ii) How many valence electrons are present in G?

(iii) Write the formula of the compound between B and H.

(iv) In the compound between F and J, what type of bond will be formed?

(v) Draw the electron dot structure for the compound formed between C and K.

(i) J is the most electronegative.

(ii) G has 5 valence electrons.

(iii) B2H is the name of the compound between B and H.

(iv) Covalent bond will be formed between F and J.

(v) Electron dot structure for the compound formed between C and K is shown below:

#### Question 3

Define the following term - Ionization Potential.

Ionization Potential is the amount of energy required to remove an electron from the outer most shell of an isolated gaseous atom.

## 2010

#### Question 1

(i) The number of electrons in the valence [outermost] shell of a halogen is :

1. 1
2. 3
3. 5
4. 7

(ii) Electro negativity across the period - increases/decreases.

(iii) Non metallic character down the group - increases/decreases.

(i) The number of electrons in the valence shell of halogen is 7.

(ii) Electro negativity across a period - increases.

(iii) Non metallic character down the group - decreases.

#### Question 2

Atomic number of an element is 16. State :

(i) To which period it belongs.

(ii) The number of valence electrons in the element.

(iii) Is the element a metal or a non metal.

(i) It belongs to the 3rd period.

(ii) The number of valence electrons are 6.

(iii) It is a non metal.

#### Question 3

Define the following terms:

(i) Ionization potential

(ii) Electron affinity

(i) Ionization Potential is the amount of energy required to remove an electron from the outer most shell of an isolated gaseous atom.

(ii) Electron affinity is the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.

## 2011

#### Question 1

Give reasons - The oxidizing power of elements increases from left to right along a period.

The oxidizing power of elements increases from left to right along a period because electro-negativity and the non metallic character increases from left to right. As oxidizing power depends on tendency to gain electrons and non-metals are good oxidizing agents hence oxidizing power of elements increases across a period.

#### Question 2

(i) Across a period, the ionization potential .......... (increases, decreases, remains same).

(ii) Down the group, the electron affinity .......... (increases, decreases, remains same).

(i) Across a period, the ionization potential increases.

(ii) Down the group, the electron affinity decreases

#### Question 3

(i) In the periodic table, alkali metals are placed in the group ...........

1. 1
2. 11
3. 17
4. 18

(ii) Which of the following properties do not match with elements of the halogen family?

1. They have seven electrons in their valence shell.
2. They are highly reactive chemically.
3. They are metallic in nature.
4. They are diatomic in their molecular form.

(i) 1
Reason — In the periodic table, alkali metals are placed in the group 1 as they have one electron in the outer most shell.

(ii) They are metallic in nature.
Reason — Halogens are non-metallic in character.

#### Question 4

State the group and period of the element having three shells with three electrons in the valence shell.

The element having three shells with three electrons in the valence shell is in group 13 [III A] and period 3.

## 2012

#### Question 1

Select the element in period 3 whose electron affinity is zero.

1. Neon
2. Sulphur
3. Sodium
4. Argon

Argon
Reason — Electron affinity of argon is zero as Argon is in 3rd period and is an inert gas. Electron affinity of inert gases is zero.

#### Question 2

Give reasons:

(i) Ionization potential of element increases across a period.

(ii) Alkali metals are good reducing agents.

(i) The ionization potential of element increases across a period because the atomic size decreases due to an increase in nuclear charge and electrons in the outermost shell are more strongly held because of which greater energy is required to remove the electron.

(ii) Alkali metals have one electron in their valence shell. In order to be stable, they easily lose this electron and get oxidized. Hence, they are good reducing agents.

#### Question 3

There are three elements E, F and G with atomic numbers 19, 8 and 17, respectively.

Classify the elements as metals and non-metals.

The electronic configuration of the given elements is as follows :

E = 19 = 2, 8, 8, 1

F = 8 = 2, 6

G = 17 = 2, 8, 7

We observe that E has 1 electron in the outer most shell, hence it will try to lose it's electron and attain a stable state. Therefore, it is a metal.

On the other hand, F and G will try to gain 2 and 1 electron respectively in order to attain a stable state. Hence, they are non-metals.

#### Question 4

Name : A metal present in period 3, group 1 of the periodic table.

Sodium is a metal present in period 3 and group 1 of the periodic table.

## 2013

#### Question 1

Among Period-2 elements — Lithium; Carbon; Chlorine; Fluorine — State the one which has high electron affinity.

Fluorine has the highest Electron affinity.

#### Question 2

Group
numbers
1 - IA2 - IIA13 - IIIA14 - IVA15 - VA16 - VIA17 - VIIA18 - 0
2nd periodLi D  OJNe
3rd periodAMgESi HM
4th periodRTI Qu y

In the above table, H does not represent hydrogen.

Some elements are in their own symbol and position in the periodic table, while others are shown with a letter.

Identify :

(i) The most electronegative element.

(ii) The most reactive element of group I.

(iii) The element from period 3 with least atomic size.

(iv) The noble gas of the fourth period.

(v) How many valence electrons are present in Q.

(vi) Which element from group 2 would have the least ionization energy?

(vii) In the compound between A and H, what type of bond is formed and give it's molecular formula.

(i) J is the most electronegative element.

(ii) R is the most reactive element of group 1.

(iii) M is the element from period 3 with least atomic size.

(iv) y is the noble gas of the fourth period.

(v) 5 valence electrons present in Q.

(vi) T is the element from group 2 which has least ionization energy.

(vii) ionic bond is formed between A and H.

The molecular formula is :

2A + H ⟶ A2H

#### Question 3

Identify : The element which has the highest ionization potential.

Helium has highest ionization potential.

## 2014

#### Question 1

(i) Ionization potential increases over a period from left to right because the:

1. Atomic radius and nuclear charge increases
2. Atomic radius and nuclear charge decreases
3. Atomic radius increases and nuclear charge decreases
4. Atomic radius decreases and nuclear charge increases

(ii) An element A belonging to Period 3 and Group II will have :

1. 3 shells and 2 valence electrons
2. 2 shells and 3 valence electrons
3. 3 shells and 3 valence electrons
4. 2 shells and 2 valence electrons

(i) Atomic radius decreases and nuclear charge increases
Reason — The ionization potential of element tends to increase across a period from left to right because the atomic size decreases due to an increase in nuclear charge and electrons in the outermost shell are more strongly held because of which greater energy is required to remove the electron.

(ii) 3 shells and 2 valence electrons
Reason — Period of an element determines the number of shells and group determines number of valence electrons.

#### Question 2

Atomic number of element Z is 16

(i) State the period and group to which Z belongs.

(ii) Is Z a metal or a non-metal?

(iii) State the formula of the compound between Z and hydrogen.

(iv) What kind of a compound is this?

(i) The element belongs to 3rd period and 16 [VI A] group.
Reason — The element Z has atomic number 16 and so the electronic configuration will be 2, 8, 6. The number of shells present in an atom determines it's period. Hence, Z will belong to 3rd period as it has three shells. The number of valence electrons determine the group of the element. Hence, Z will belong to group 16 [VI A] as it has 6 electrons in the valence shell.

(ii) Z is a non metal.

(iii) It forms H2Z

(iv) It is a weak acid.

#### Question 3

In the activity series of metals — M is a metal above hydrogen in the activity series and it's oxide has the formula M2O. M2O when dissolved in water forms the corresponding hydroxide which is a good conductor of electricity. In the above context, answer the following :

(i) What kind of combination exists between M and O?

(ii) State the number of electrons in the outermost shell of M ?

(iii) Name the group to which M belongs.

Given, M is a metal

(i) Electrovalent bond exits between M and O because the bond is formed between a metal and non-metal due to oppositely charged ions.

(ii) Number of electrons in the outer most shell of M is 1. It is so because the valency of O is -2 and as 2 atoms of M combine with O to form M2O, hence we can say that M has 1 valence electron.

(iii) M belongs to group 1 [1A] because there is 1 electron in the outer most shell.

#### Question 4

Give a phrase for : Amount of energy released when an atom in the gaseous state accepts an electron to form an anion.

Electron affinity is the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.

#### Question 5

Match the option - A : Metal or B : Iron - with the statements (i) and (ii) :

(i) The metal which forms two types of ions

(ii) An element with electronic configuration 2, 8, 8, 3

(i) The metal which forms two types of ions — B : Iron

(ii) An element with electronic configuration 2, 8, 8, 3 — A : Metal

## 2015

#### Question 1

The element with the least electronegativity is :

1. Lithium
2. Carbon
3. Boron
4. Fluorine

Lithium
Reason — Electronegativity increases from left to right. As Lithium is on the left side in periodic table hence it is the least electronegativity among the given options.

#### Question 2

Arrange the elements as per the instructions :

(i) Cs, Na, Li, K, Rb (increasing order of metallic character).

(ii) Mg, Cl, Na, S, Si (decreasing order of atomic size).

(iii) Na, K, Cl, S, Si (increasing order of ionization energy)

(iv) Cl, F, Br, I (increasing order of electron affinity)

(i) Li < Na < K < Rb < Cs

(ii) Na > Mg > Si > S > Cl

(iii) K < Na < Si < S < Cl

(iv) I < Br < F < Cl

#### Question 3

Select a covalent oxide of a metalloid from the following — SO2 , SiO2, Al2O3, MgO, CO, Na2O.

SiO2 is a covalent oxide of a metalloid.

#### Question 4

The metals of Group 2 in the periodic table from top to bottom are — Be, Mg, Ca, Sr, and Ba.

(i) Which one of these elements will form ions most readily. Give reasons.

(ii) State the common feature in the electronic configuration of all these given elements.

(i) Ba - Elements at the bottom of a group are most metallic, have large atomic size and lowest ionisation potential. So, the outer electrons are loosely held and will form ions from metals most readily and thus are more reactive.

(ii) As the elements belong to group 2 thus they all have 2 electrons in the valence shell.

## 2016

#### Question 1

Select the answer from A, B, C and D : An element with the atomic number 19 will most likely combine chemically with the element whose atomic number is:

1. 17
2. 11
3. 18
4. 20

17
Reason — The element with atomic number 19 has electronic configuration of (2, 8, 8, 1) and that with atomic number 17 has electronic configuration of (2, 8, 7). This element with atomic number 17 needs 1 electron to complete its octet whereas element with atomic number 19 has an extra electron. Hence, an element with atomic number 19 is most likely to combine with the element with atomic number 17.

#### Question 2

Identify the term in each of the following :

(i) The tendency of an atom to attract electrons to itself when combined in a compound.

(ii) The electrons present in the outermost shell of an atom.

(i) Electronegativity.

(ii) Valence Electrons.

#### Question 3

Write the correct symbol > (greater than) or < (less than) in the statements :

(i) The ionization potential of potassium is .......... that of sodium.

(ii) The electronegativity of iodine is .......... that of chlorine.

(i) The ionization potential of potassium is < (less than) that of Sodium.

(ii) The electronegativity of iodine is < (less than) that of chlorine.

#### Question 4

Use the letters only written in the Periodic Table given below to answer the questions :

IIIIIIIVVVIVIIO
1       L
2Q EGJZM
3R
4T

(i) State the number of valence electrons in atom J.

(ii) Which element shown forms ions with a single negative charge?

(iii) Which metallic element is more reactive than R?

(iv) Which element has it's electrons arranged in four shells?

(i) Atom J has 5 valence electrons as it belongs to the fifth group in the periodic table.

(ii) M forms ions with a single negative charge as it has 7 electrons in the valence shell and obtaining one more electron completes it's octet.

(iii) Metal T is more reactive than R because elements at the bottom of the group are more reactive.

(iv) T has it's electrons arranged in 4 shells as it belongs to 4th period.

#### Question 5

Fill in the blanks by selecting the correct word :

(i) If an element has a low ionization energy then it is likely to be .......... (metallic/non metallic).

(ii) If an element has seven electrons in it's outermost shell then it is likely to have the .......... (largest/smallest) atomic size among all the elements in the same period.

(i) If an element has a low ionization energy then it is likely to be metallic.

(ii) If an element has seven electrons in it's outermost shell then it is likely to have the smallest atomic size among all the elements in the same period.

## 2017

#### Question 1

Select the correct answer — The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion is called .......... (electron affinity, ionization potential, electronegativity)

Ionization Potential

#### Question 2

Match the atomic number 2, 4, 8, 15 and 19 with each of the following:

(i) A solid non-metal belonging to the third period.

(ii) A metal of valency 1.

(iii) A gaseous element with valency 2.

(iv) An element belonging to Group 2.

(v) A rare gas.

(i) 15

(ii) 19

(iii) 8

(iv) 4

(v) 2

#### Question 3

Arrange as per the instruction :

(i) He, Ar, Ne (Increasing order of the number of electron shells)

(ii) Na, Li, K (Increasing Ionization Energy)

(iii) F, Cl, Br (Increasing electronegativity)

(iv) Na, K, Li (Increasing atomic size).

(i) He < Ne < Ar
Reason — Number of electron shells increases as we move down the group.

(ii) K < Na < Li
Reason — Ionization Potential decreases as we move down the group.

(iii) Br < Cl < F
Reason — Electronegativity decreases as we move down the group.

(iv) Li < Na < K
Reason — Atomic Size increases as we move down the group.

## 2018

#### Question 1

Give one word or a phrase for the following statement: The energy released when an electron is added to a neutral gaseous isolated atom to form a negatively charged ion.

Electron affinity

#### Question 2

Give reasons :

(i) Inert gases do not form ions.

(ii) Ionization potential increases across a period — left to right.

(i) Inert gases have completely filled octet which makes them extremely stable. They neither lose, nor gain electrons. Hence, they do not form ions.

(ii) The ionization potential of an element increases across a period because the atomic size decreases due to an increase in nuclear charge and electrons in the outermost shell are more strongly held because of which greater energy is required to remove the electron.

#### Question 3

In Period 3 of the Periodic Table, element B is placed to the left of element A. On the basis of this information, choose the correct word from the brackets to complete the following statements:

(i) The element B would have (lower/higher) metallic character than A.

(ii) The element A would probably have (lesser/higher) electron affinity than B.

(iii) The element A would have (greater/smaller) atomic size than B.

If the element are placed as B and then A in the 3rd period of the periodic table then

(i) The element B would have higher metallic character than A as metallic character decreases across a period.

(ii) The element A would probably have higher electron affinity than B as electron affinity increases across a period.

(iii) The element A would have smaller atomic size than B as atomic size decreases across a period.

## 2019

#### Question 1

Choose the correct answer from the options A, B, C and D given :

The most electronegative element from the following elements is :

1. Magnesium
2. Chlorine
3. Aluminium
4. Sulphur

Chlorine
Reason — Chlorine is the most electronegative element among the given options because electronegativity increases left to right in a period.

#### Question 2

Fill in the blank:
In Period 3, the most metallic element is …………….. (sodium / magnesium / aluminium)

In Period 3, the most metallic element is Sodium because metallic character decreases from left to right across a period.

#### Question 3

Give the appropriate term defined by the statement given :

The tendency of an atom to attract electrons towards itself when combined in a covalent compound.

Electronegativity

#### Question 4

Arrange the following:

(i) Li, K, Na, H (In the decreasing order of their ionization potential)

(ii) F, B, N, O (In the increasing order of electron affinity).

(i) H > Li > Na > K
Reason — Ionization Potential decreases as we move down the group.

(ii) B < N < O < F
Reason — Electron Affinity increases from left to right across a period.

#### Question 5

Study the extract of the Periodic Table given below and answer the questions. Give the alphabet corresponding to the element in question. Do not repeat an element. State which element :

A
C DE
B     GF

(i) Forms an electrovalent compound with G.

(ii) Is non-metallic and has a valency of 2?

(iii) Is an inert gas?

(iv) State the ion of which element will migrate towards the cathode during electrolysis.

(i) B forms an electrovalent compound with G.

(ii) E is non-metallic and has a valency of 2 as it belongs to group 16.

(iii) F is an inert gas as it belongs to the zero group.

(iv) A has positive ions which migrate towards cathode.

## 2020

#### Question 1

Choose the correct answer from the given options given :

The element with highest ionization potential is :

1. Hydrogen
2. Caesium
4. Helium

Helium
Reason — Helium has highest ionization potential because Ionization Potential increases from left to right across a period and decreases down a group.

#### Question 2

Give one word or phrase for : The tendency of an atom to attract electrons to itself when combined in a compound.

Electronegativity

#### Question 3

The question represents the elements P, Q. R and their atomic numbers.

Answer the following using only the alphabets given — P = 13; Q = 7; R = 10

(i) Which element combines with hydrogen to form a basic gas.

(ii) Which element has an electron affinity zero.

(iii) Name the element, which forms an ionic compound with chlorine.

(i) Q combines with hydrogen to form a basic gas.

(ii) R has electron affinity zero as it's octet is complete and it is a stable element.

(iii) P forms an ionic compound with chlorine as it has +3 valency.

#### Question 4

Name the element : An alkaline earth metal present in group 2 and period 3.

Magnesium is an alkaline earth metal which is present in group 2 and period 3.

#### Question 1

State the fundamental property on which the modern periodic table or long form of periodic table is based.

Modern periodic table or long form of periodic table is based on the Modern Periodic Law which states that —
Physical and chemical properties of elements are periodic function of their atomic number. Hence, atomic number is the fundamental property of an element on which the periodic table is based.

#### Question 2

State the important salient features of the modern periodic table. State how separation of elements and periodicity of elements forms an important feature of the modern periodic table.

Salient Features of Modern Periodic Table are:

1. Classification — Physical and chemical properties of elements are periodic function of their atomic number.
2. Position — Correlates position of elements with it's electronic configuration.
3. Methodical Arrangement — Arrangement of elements is in increasing order of atomic numbers in:
1. Seven horizontal rows called called periods and
2. Eighteen vertical columns called groups.
4. Periods — Completion of each period is logical since each period:
1. begins with an element having one electron in outermost shell
2. ends with zero group element having completely filled outer shell.
3. A transition from metallic to non-metallic is seen across a period.
5. Groups — Each vertical column accommodates elements with the same outer electronic configuration hence having similar properties.
1. 18 vertical columns consists of groups 1 to 17 and 18 [zero group]
2. Group 1, 2 and 13 to 17 [IA to VII A] are called Normal elements.
3. Group 3 to 12 [IB to VII B and VIII] are called Transition elements.
4. Group 18 [zero] at extreme right contains Noble or Inert gases.
6. Separation of elements — Modern periodic table provides separation of elements with similar properties in the following ways:
1. Reactive metals are placed in group 1 [IA] and 2 [IIA]
2. Transition elements (i.e., metals) are placed in the middle.
3. Non-metals are placed in the upper right corner of the periodic table.
7. Periodicity of Elements — Gradual change in properties is seen with increase in atomic number in the periodic table. Recurrence of properties is observed with elements belonging to the same subgroup in the periodic table after a difference of 2, 8, 18 or 32 in the atomic numbers due to recurrence of similar valence shell electronic configuration.

#### Question 3

What are 'periods'. State the correlation of a period number with the elements of that period.

Periods are the seven horizontal rows in the modern periodic table.
The number of shells in an atom determines it's period. Hence, all the elements of a particular period have same number of shells as its period number.

#### Question 4

Name the elements in correct order of their increasing atomic numbers present in the first, second and third short periods of the periodic table. State each elements electronic configuration.

Elements of first, second and third period in increasing order of their atomic number are mentioned below (atomic number of each element is mentioned in brackets):

First period — H(1), He (2)

Second period — Li (3), Be (4), B (5), C (6), N (7), O (8), F (9), Ne (10)

Third period — Na (11), Mg (12), Al (13), Si (14), P (15), S (16), Cl (17), Ar (18)

Electronic configuration of the above mentioned elements is as follows —

H — 1

He — 2

Li — 2, 1

Be — 2, 2

B — 2, 3

C — 2, 4

N — 2, 5

O — 2, 6

F — 2, 7

Ne — 2, 8

Na — 2, 8, 1

Mg — 2, 8, 2

Al — 2, 8, 3

Si — 2, 8, 4

P — 2. 8, 5

S — 2, 8, 6

Cl — 2, 8, 7

Ar — 2, 8, 8

#### Question 5

Give a reason why :

(a) completion of each period is logical.

(b) period-2 elements are called 'bridge elements'.

(a) Completion of each period is logical, since each period begins with an element having one electron in the outer most shell and ends with zero group element having completely filled outer shell.

(b) The elements of the second period show resemblance in properties with the elements of the next group of third period, due to very less electronegativity difference. This leads to a diagonal relationship, viz. Li and Mg, Be and Al, B and Si. These elements are hence called Bridge elements.

#### Question 6

State the property trends in general on moving from left to right in a period of the periodic table.

The property trends in general on moving from left to right in a period of the periodic table are :

1. Number of valence electrons — increases by one
3. Ionization potential — increases.
4. Electron affinity — increases (exception, electron affinity is zero for noble gases).
5. Non-metallic character — increases.
6. Metallic character — decreases
7. Electronegativity — increases (with the exception of noble gases, they have complete octet so do not attract electrons to itself).

#### Question 7

State :

(i) the bonding and state of chlorides of period-3 — group 1 [IA], 15[VA], 16 [VIA] and

(ii) the bonding and character of oxides of period-3 — group 1 [IA], 13[IIIA] and 16[VIA].

(i) The element of group 1 [IA] of period 3 is Na and it's chloride forms ionic bond and is in solid state.
The element of group 15 [VA] of period 3 is P and it's chloride forms covalent bond and is in liquid / solid state.
The element of group 16 [VIA] of period 3 is S and it's chloride forms covalent bond and is in liquid state.

(ii) The element of group 1 [IA] of period 3 is Na and it's oxide forms electrovalent bond and is strongly basic in character.
The element of group 13 [IIIA] of period 3 is Al and it's oxide forms electrovalent bond and is amphoteric in character.
The element of group 16 [VIA] of period 3 is S and it's oxide forms covalent bond and is acidic in character.

#### Question 8

What are 'groups' of the Modern Periodic Table ? What does the 'group number' signify.

Modern periodic table has 18 vertical columns. Each vertical column accommodates elements with the same outer electronic configuration (i.e., valence electrons), hence have similar properties.
Group number signifies the same outer electronic configuration and similar properties.

#### Question 9

State the type of elements present in :

(a) group 1[IA]

(b) group 2 [IIA]

(c) group 3 to 12 [IB to VIIB and VIII]

(d) group 13 to 16 [IIIA to VIA]

(e) group 17 [VIIA]

(f) group 18 [0].

(a) Group-1 [IA] — Alkali metals

(b) Group-2 [IIA] — Alkaline earth metals

(c) Group-3 to 12 IB to VII B, VIII — Transition elements - inner transitional elements.

(d) Group 13-16 [III B to VI A] — Post Transition elements.

(e) Group 17 (VII A) — Halogens.

(f) Group 18 (0) — Noble / Inert gases.

#### Question 10

What are transition elements and inner transition elements. State the position of the inner transition elements. State why noble gases are considered unreactive elements.

Transition Elements — Elements belonging to Groups 3 to 12 are called transition elements. The valence electron of these elements is in d orbital. They are metals and lie between strongly electropositive metals on the left and least electropositive elements on the right. They all have similar properties.

Inner Transition Elements — Elements belonging to Group 3 in 6th and 7th period are called Inner transitional elements. The valence electron of these elements is in f orbital. They consists of two horizontal rows of elements which are placed at the bottom of the table. These rows are called Lanthanides (rare earth) and Actinides (radio active). Each row has 14 elements.

Noble gases are unreactive because their octet is complete and they have a stable electronic configuration.

#### Question 11

State the characteristics which remain similar and those which show a transition on moving down a sub-group.

Characteristics which remain similar on moving down a sub-group are as follows:

1. Valency of electrons.
2. Chemical properties.

Characteristics which show a transition on moving down a sub-group are as follows:

1. Metallic character increases down the sub-group.
2. Number of electron shells increases by one as we move down each sub-group.

#### Question 12

Compare the properties of the elements of group 1[IA] i.e. alkali metals and group 17 [VIIA] i.e., halogens.

Below table shows the comparison of the elements of group 1[IA] i.e. alkali metals and group 17 [VIIA] i.e., halogens:

PropertyGroup 1[IA]Group 17 [VIIA]
ElementsLi, Na, K, Rb, Cs, FrF, Cl, Br, I, At
ValencyUnivalent (1 valence electron)Univalent (7 valence electrons)
NatureHighly reactive, highly electropositive, light soft metalsHighly reactive, highly electronegative, non - metals
ConductivityGood conductors of heat and electricityBad or non conductors of heat and electricity
Reducing / Oxidizing natureStrong reducing agentsStrong oxidizing agents
ElectronegativityLow electronegativityHigh electronegativity

#### Question 13

Explain the term:

(a) periodicity in properties of elements

(b) periodic properties

(c) periodicity of elements

(a) Periodicity in properties of elements — The phenomenon of occurrence of characteristic properties of elements at definite intervals in the modern periodic table when elements are arranged in increasing order of their atomic numbers is called Periodicity in properties of elements.

(b) Periodic properties — The properties which appear at regular intervals in the periodic table are called periodic properties. Periodic Properties are:

2. Ionization potential
3. Electron affinity
4. Electronegativity
5. Non-Metallic and Metallic character.
6. Density
7. Melting and boiling points.
8. Nature of oxides, oxy-acids, hydrides.

(c) Periodicity of elements — Gradual change in properties of elements when they are arranged in increasing order of atomic number in the periodic table is called the periodicity of elements.

#### Question 14

State the reasons for periodicity of elements in periods and groups.

Reasons for Periodicity in properties in periods and groups is as follows —

1. After definite intervals of atomic number, similar valence shell electronic configuration occurs i.e., same number of electrons are present in the outermost orbit of the elements.
2. Properties of elements depend upon the number and arrangement of electrons in various shells including valence shells.
3. In the same period or sub-group, increase or decrease in a particular property is due to gradual change in electronic configuration in the arranged elements.

#### Question 15

Explain the meaning of the following periodic properties:

(b) Ionization potential

(c) Electron affinity

(d) Electronegativity

(e) Non-metallic and metallic character.

(a) Atomic radius — It is distance between the center of the nucleus and outer most shell of the atom.

(b) Ionization potential [I.P.] — It is the amount of energy required to remove an electron from the outer most shell of an isolated gaseous atom.

(c) Electron affinity [E.A.] — It is the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.

(d) Electronegativity [E.N.] — It is the tendency of an atom to attract electrons to itself when combined in a compound.

(e) Non-metallic and Metallic character — In terms of electron loss or gain, an element is a:

• Non-metal — if it gains one or more electrons and,
• Metal — if it loses one or more electrons.

#### Question 16

State the factors which affect the atomic size of elements in a periodic table. In period 2 from left to right, state which element has the largest atomic size and which has the smallest, giving reasons.

Factors affecting the atomic size are:

1. Number of shells — An increase in the number of shells increases the size of an atom because the distance between the outermost shell and the nucleus increases.
2. Nuclear charge — An increase in nuclear charge decreases the size of the atom because the electrons are then attracted towards the nucleus with a greater force thereby bringing the outermost shell closer to the nucleus.

In period 2 from left to right, Li has the largest atomic size and F has the smallest atomic size. This is so because the nuclear charge i.e., the atomic number increases from left to right in the same period, thereby bringing the outermost shell closer to the nucleus.

#### Question 17

Explain the trend in atomic radii on moving down a group, with reference to the alkali metals in Group 1 [IA].

Atomic radii of an atom increases on moving down a group. As we move down a group, the number of shells increases and the nuclear charge also increases. But increase in the number of shells dominates over increase in nuclear charge therefore, overall atomic size increases.
For alkali metals in Group 1 [IA], Li which is at the top of the group has the smallest atomic radius whereas Cs at the bottom of the group has the largest atomic radius. The atomic radii of the elements has the following relation:

Li < Na < K < Rb < Cs

#### Question 18

State the factors which influence or affect the ionization potential of elements in a periodic table.

The factors which influence the ionization potential of an element are —

1. Atomic size — As atomic size increases, the nuclear attraction on the outer electrons decreases and outer electron is loosely held. Hence, ionization potential decreases.

2. Nuclear charge — As nuclear charge increases, the nuclear attraction on the outer most electron increases and the outer electron is more firmly held. Therefore, ionization potential increases.

#### Question 19

Explain the trend in general of ionization potential of elements :

(a) on moving from left to right across a period

(b) on moving down a group.

Give reasons for the change in the periodic trend in each case.

(a) On moving from left to right across a period, the ionization potential increases.
The ionization potential of element tends to increase across a period because the atomic size decreases due to an increase in nuclear charge and electrons in the outermost shell are more strongly held because of which greater energy is required to remove the electron.

(b) On moving down a group ionization potential decreases.
On moving down the group, atomic size, as well as, nuclear charge increases. Due to the increase in atomic size the ionization potential decreases and due to the increase in nuclear charge the ionization potential should increase However, the effect of increase in atomic size dominates over the effect of increase in nuclear charge. Hence, ionization potential decreases down the group.

#### Question 20

State the factors which affect :

(a) electron affinity

(b) electronegativity of elements in a periodic table.

(a) The Factors which affect the electron affinity are:

1. Atomic size — The smaller the atomic size, the greater the electron affinity, because a small atom takes up electrons more readily than a large atom since nucleus has greater attraction on electrons.
2. Nuclear charge — As the nuclear charge increases electron affinity increases because with an increase in nuclear charge the tendency of an atom to accept electrons increases.

(b) The factors affecting electronegativity are:

1. Atomic size — The greater the size of the atom, the lesser the electronegativity because a small atom will pull the shared pair of electrons more than a large atom.
2. Nuclear charge — As the nuclear charge increases electronegativity increases because an increase in nuclear charge increases the force of attraction on the electron.

#### Question 21

Explain the trend in general of (i) electron affinity (ii) electronegativity of elements:

(a) on moving from left to right across a period

(b) on moving down a group.

Give reasons for the change in each periodic trend.

1. Electron affinity
1. Across a period atomic size decreases and nuclear charge increases. Both these factors lead to an increase in the Electron affinity across a period from left to right as a small atom with more nuclear charge can attract electrons more easily.
2. On moving down a group both atomic size and nuclear charge increases but the increase in atomic size dominates over increase in nuclear charge hence Electron affinity decreases on moving down a group.
2. Electronegativity
1. Across a period atomic size decreases and nuclear charge increases. Both these factors lead to an increase in Electronegativity across a period from left to right as a small atom will pull the shared pair of electrons more than a large atom.
2. On moving down a group both atomic size and nuclear charge increases but the increase in atomic size dominates over increase in nuclear charge hence Electronegativity decreases on moving down a group.

#### Question 22

With reference to the alkali metals in Group 1 [IA] and the halogens in 17 [VIIA] explain the trend in ionization potential, electron affinity and electronegativity on moving down the groups in the periodic table.

Below table explains the trends in ionization potential, electron affinity and electronegativity on moving down the groups in the periodic table:

Alkali Metals
Group 1 [IA]
Halogens
Group 17 [VIIA]
Ionization
Potential
On moving down the group both atomic size and nuclear charge increases but atomic size increase dominates over nuclear charge increase hence Ionization Potential decreases down the group.
∴ Li > Na > K > Rb > Cs
Similarly for Halogens too Ionization Potential decreases down the group.
∴ F > Cl > Br > I
Electron AffinityOn moving down the group, electron affinity decreases as atomic size increase dominates over nuclear charge increase hence the atom attracts the electrons with a lesser force.
∴ Li > Na > K > Rb > Cs
Here also electron affinity decreases in general on moving down the group but Fluorine is an exception here as it has lesser electron affinity than Chlorine due to its electronic configuration.
∴ F < Cl > Br > I
ElectronegativitySimilar to electron affinity, electronegativity decreases on moving down the group.
∴ Li > Na > K > Rb > Cs
Electronegativity decreases on moving down the group.
∴ F > Cl > Br > I

#### Question 23

State the factors which affect the metallic and the non-metallic character of elements in a periodic table.

Factors which affect metallic and non-metallic character in a periodic table are as follows :

1. Atomic size — The greater the atomic size, the farther the outermost orbit, and thus lesser is the nuclear pull exerted on it. As a result, electrons can be removed more easily from the valence shell, this making the element more metallic and less non-metallic.

2. Nuclear charge — The greater the nuclear charge, the greater is the force exerted by the nucleus on the electron of the outermost orbit. This makes it difficult to remove the electron from the outermost orbit. Thus, metallic nature decreases and non-metallic nature increases.

#### Question 24

Explain the trends from metallic to non-metallic character of the different elements in the first three periods.

On moving across a period, nuclear pull increases due to increase in atomic number and thus the atomic size decreases. Hence, the elements cannot lose electrons easily. Therefore, metallic nature decreases across a period moving from left to right and non-metallic character increases.
The alkali metals ( group 1 [IA]) and alkaline earth metals (group 2 [IIA]), placed on the left side of the table are most metallic in nature and the halogens (group 16 [VIA]) placed on the right side of the table are most non-metallic in nature.

#### Question 25

Explain with reasons the trends in metallic and non-metallic character down a group.

Metallic/Non-Metallic character depends on atomic size and nuclear charge. Increase in atomic size increases the Metallic nature and increase in nuclear charge decreases the metallic nature. On moving down the group both atomic size and nuclear charge increases. But the atomic size increase dominates over nuclear charge increase. Hence, tendency to lose electrons increases. Thus, metallic character increases as one moves down a group and non-metallic character decreases.

#### Question 26

State how density and melting points of elements varies across a period and down a group.

(i) Density — Density of elements across a period increases gradually to maximum and then it decreases slightly. Down a group density of elements increases gradually.

(ii) Melting point — Across a period from left to right, melting point increase upto group 14(IV A) and then decreases. The melting point of metals decrease down the group. The melting point of non-metals increase going down a group

#### Question 27

State the general trend in periodicity in properties of oxides, hydroxides, oxy-acids and hydrides of compounds of elements across a period and down a group.

Below table shows the general trend in periodicity in properties of oxides, hydroxides, oxy-acids and hydrides of compounds of elements across a period and down a group:

Across a periodDown a group
OxidesVaries from strongly basic to strongly acidicVaries from acidic to basic
HydroxidesVaries from strongly basic to amphotericVaries from less basic to strongly basic
Oxy-acidsVaries from weak oxy-acids to strong oxy-acidsVaries from strong oxy-acids to weak oxy-acids
HydridesVaries from strongly basic to strongly acidicVaries from less acidic to more acidic

#### Question 28

State the relation between atomic number and atomic mass for light elements. State which elements are considered radioactive giving reasons.

Electronic configuration of lighter elements shows that the elements which have an even number of proton, for example, atomic numbers like $_2^4\text{He}$, $_6^{12}\text{C}$ etc., have their mass numbers twice the atomic numbers except for $_4^9\text{Be}$ and $_{18}^{40}\text{Ar}$.
Elements which have an odd number of protons like $_3^7\text{Li}$, $_5^{11}\text{B}$ etc., have their mass number twice the atomic numbers + 1 (A = 2Z + 1) except $_7^{14}\text{N}$ and $_1^1\text{H}$

Elements with neutron:proton ratio 1.5 and above are considered as radioactive as the nucleus of such elements becomes unstable causing their radioactive decay. For example, in Uranium $(_{92}^{235}\text{U})$:

No. of protons (p) = 92
No. of neutrons (n) = 235 - 92 = 143

n : p = $\dfrac{143}{92}$ = 1.5543

A neutron:proton ratio greater than 1.5 makes Uranium radioactive.

## State the following

#### Question 1

The alkali metal in period 2 and the halogen in period 3.

Lithium, chlorine.

#### Question 2

The noble gas having duplet arrangement of electrons.

Helium

#### Question 3

The noble gas having an electronic configuration 2, 8, 8.

Argon

#### Question 4

The number of electron shells in elements of period 3.

3
Reason — The number of shells present in an atom determines it's period. Hence, element of period 3 will have 3 shells.

#### Question 5

The valency of elements in group 1[IA].

One
Reason — Group number signifies the number of valence electrons of an element. Hence valency of group 1[1A] elements will be one.

#### Question 6

The metals present in period 3 and the non-metals present in period 2.

Metals present in period 3 are Na, Mg and Al.
Non-metals present in period 2 are C, N, O and F.

#### Question 7

The group whose elements have zero valency.

Group 18 (Inert gases group)
Reason — The group 18 elements have 2 or 8 valence electrons and are stable elements hence have zero valency.

#### Question 8

The non-metal in period 3 having a valency 1.

Chlorine

#### Question 9

The formula of the hydroxide of the element having electronic configuration 2, 8, 2.

Mg(OH)2

#### Question 10

The formula of the hydride of the halogen in period 3.

HCl

#### Question 11

The formula of the sulphite of the element in period-3, group 1 [IA].

Na2SO3

#### Question 12

The element in period-3 which does not form an oxide.

Argon
Reason — Electronic configuration of argon is 2, 8, 8 and so it's octet is complete and is stable and it does not form oxide.

#### Question 13

The bonding [i.e. electrovalent or covalent] of the oxide of the element in period-3 group 16 [VIA].

Covalent
Reason — As the combining atoms i.e., S and O are non-metals and they both have high electron affinity hence they share two pairs of valence electrons in order to complete their octet and attain a stable state and hence form covalent bond.

#### Question 14

The character of the hydroxide of the element in period-3 group 13 [IIIA].

Amphoteric
Reason — As we move from left to right in periodic table, the basic character decreases and acidic character increases. Group 13 [IIIA] lies in the middle of the periodic table, hence it has amphoteric nature.

#### Question 15

A light element in period-3 with a neutron/proton ratio around 1.

Sodium is a stable element having neutron/proton ratio around 1 (approximately).
Reason — No. of protons (p) = 11
No. of neutrons (n) = 23 - 11 = 12
n : p = $\dfrac{12}{11}$ = 1.09 ≈ 1

#### Question 16

The element with the least atomic size from carbon, nitrogen, boron and beryllium.

Nitrogen
Reason — As atomic size decreases across a period (from left to right) and nitrogen lies at the right side of the periodic table, hence it has least size.

#### Question 17

The element, from the elements Li, Na, K, having the least number of electron shells.

Li (Lithium)
Reason — As the number of shells increases as we move down a group and Li lies in period 2 and is at top among the given elements hence has the least number of shells.

#### Question 18

The element from the elements C, O, N, F, having the maximum nuclear charge.

F (Fluorine)
Reason — Nuclear charge increases across a period (from left to right) and fluorine lies at the right side of the periodic table hence it has maximum nuclear charge.

#### Question 19

The element from the elements Be and Mg having a lower nuclear charge.

Be (Beryllium)
Reason — Nuclear charge increases down a group and among Be and Mg, Be is on top of Mg hence Be will have lower nuclear charge.

#### Question 20

The element from the elements fluorine and neon having a higher electron affinity.

Fluorine
Reason — Electron affinity of neon is zero because it is an inert element hence fluorine has higher electron affinity amongst the two.

#### Question 21

The period and group to which the element 'X' with electronic configuration 2, 8, 8, 2 belongs.

Group 2, period 4
Reason — Period is determined by the number of shells of the element and group is determined by the number of valence electron hence, element X will belong to group 2 and period 4 because it has 4 shells and 2 valence electrons.

#### Question 22

The more electronegative element from the elements Ar, S, Cl of period-3.

Cl (Chlorine)
Reason — As electronegativity increases across a period (from left to right) and Ar is an inert element hence, Cl will have the highest value of electronegativity among the given elements.

#### Question 23

The element with the largest atomic size from the elements of period-1, 2 and 3.

Sodium
Reason — As atomic size decreases across a period (from left to right) and increases down a group hence sodium will have the largest atomic size from the elements of period-1, 2 and 3 because it is placed in period 3 and group 1 and is at the lower left side among the given elements.

#### Question 24

The element with the highest ionization potential from the elements of period 1, 2 and 3.

Helium
Reason — As ionization potential increases across a period (from left to right) and decreases down a group hence helium will have the highest ionization potential among the elements of period-1, 2 and 3 because it is placed in period 1 and group 18 and is at the top right side of the periodic table.

#### Question 25

The element from the elements Li, Na, K which has maximum metallic character.

K (Potassium)
Reason — As metallic character increases down a group hence K will have maximum metallic character among the given elements.

#### Question 26

The element with maximum non-metallic character from the elements of period-2.

F (Fluorine)
Reason — As non-metallic character increases across a period (from left to right) hence F will have maximum non-metallic character as it lies in group 17 in period-2

#### Question 27

The more non-metallic element from the elements S, P, Cl and Ar.

Cl (Chlorine)
Reason — As non-metallic character increases across a period (from left to right) and Ar is a noble gas hence Cl will have maximum non-metallic character as it lies on the right side of the periodic table.

#### Question 28

The more non-metallic element from the elements 'X' and 'Y' having electronic configuration 2, 8, 5 and 2, 8, 6 respectively.

'Y' with electronic configuration 2, 8, 6.
Reason — X will belong to group 15 and Y will belong to group 16 and as non-metallic increases across a period (i.e., from left to right) hence Y will be more non-metallic character than X.

#### Question 29

The periodic property which relates to the amount of energy required to remove an electron from the outermost shell of an isolated gaseous atom.

Ionization potential

#### Question 30

The periodic property which refers to the character of element, which loses electron/s when supplied with energy.

Metallic property.

## Fill in the blanks

#### Question 1

Periods are ............... [5, 6, 7] horizontal rows of elements in the periodic table and an element with three electron shells and two electrons in it's valence shell belongs to period ............... [6, 3, 1] and group ............... [3, 6, 2].

Periods are 7 horizontal rows of elements in the periodic table and an element with three electron shells and two electrons in it's valence shell belongs to period 3 and group 2.

#### Question 2

Across a period the valence electrons ..............., while down a subgroup they ............... [remain same / increase by 1]

Across a period the valence electrons increase by 1 while down a subgroup they remain same.

#### Question 3

Across a period, the electropositive character ..............., and down a group the electronegative character ............... [increases/decreases].

Across a period, the electropositive character decreases and down a group the electronegative character decreases.

#### Question 4

Elements at the extreme left of the modern periodic table are ............... reactive, while elements on the extreme right [group 18 (0)] are ............... reactive [least/un/most].

Elements at the extreme left of the modern periodic table are most reactive, while elements on the extreme right [group 18 (0)] are unreactive.

#### Question 5

Elements of group 1 [IA] are strong ............... [oxidizing/reducing] agents since they are electron ............... [acceptors/donors].

Elements of group 1 [IA] are strong reducing agents since they are electron donors.

#### Question 6

The element in group 17 [VIIA] which is a liquid at room temperature is ............... [F, Cl, Br, I].

The element in group 17 [VIIA] which is a liquid at room temperature is Br.

#### Question 7

Periodicity in properties is observed in elements after definite intervals due to similar ............... [electronic configuration, number of valence electrons, atomic numbers] of elements.

Periodicity in properties is observed in elements after definite intervals due to similar electronic configuration.

#### Question 8

Across a period the nature of oxides and hydrides varies from ............... to ............... [acidic/basic] while the strength of oxy-acids ............... [decreases/increases] from left to right.

Across a period the nature of oxides and hydrides varies from basic to acidic while the strength of oxy-acids increases from left to right.

#### Question 9

Nuclear charge of an atom is the ............... [negative/positive] charge on the nucleus of an atom, equivalent to the atomic ............... [number/mass] of an atom.

Nuclear charge of an atom is the positive charge on the nucleus of an atom, equivalent to the atomic number of an atom.

#### Question 10

Atomic size of neon is ............... [more/less] than the atomic size of fluorine.

Atomic size of neon is more than the atomic size of fluorine.

#### Question 11

Atomic size across a period ............... [increases/decreases] with increase in nuclear charge of the element.

Atomic size across a period decreases with increase in nuclear charge of the element.

#### Question 12

With increase in nuclear charge the nuclear attraction for outer electrons ............... [increases/decreases], hence ionization potential ............... [increases/decreases].

With increase in nuclear charge the nuclear attraction for outer electrons increases, hence ionization potential increases.

#### Question 13

Increase in nuclear charge of an atom ............... [decreases/increases] the tendency of the atom to lose electrons.

Increase in nuclear charge of an atom decreases the tendency of the atom to lose electrons.

#### Question 14

Elements with stable electronic configuration e.g. neon have an electron affinity value of ............... [1, 0, -1].

Elements with stable electronic configuration e.g. neon have an electron affinity value of 0 .

#### Question 15

An atom with a small atomic radii takes up electrons ............... [less/more] readily than an atom with a large radii.

An atom with a small atomic radii takes up electrons more readily than an atom with a large radii.

#### Question 16

If combining atoms of a compound have nearly similar electronegativities the bond between them is ............... [electrovalent/covalent].

If combining atoms of a compound have nearly similar electronegativities the bond between them is electrovalent

#### Question 17

Elements with low electronegativity are usually ............... [metallic/non-metallic].

Elements with low electronegativity are usually metallic.

#### Question 18

An atom is said to be a non-metal, if it ............... [gains/loses] one or more electrons.

An atom is said to be a non-metal, if it gains one or more electrons.

#### Question 19

Atoms with ............... [small/large] atomic radii and ............... [high/low] ionization potential tend to gain electrons.

Atoms with small atomic radii and high ionization potential tend to gain electrons.

#### Question 20

Element 'X' in period 3 has high electron affinity and electronegativity. It is likely to be a ............... [metal/non-metal].

Element 'X' in period 3 has high electron affinity and electronegativity. It is likely to be a non-metal.

#### Question 21

Element 'B' in period 2 is to the right of the element 'A'.
Element 'B' is likely to be ............... [more/less] non-metallic in character than element 'A'.

Element 'B' in period 2 is to the right of the element 'A'.
Element 'B' is likely to be more non-metallic in character than element 'A'.

#### Question 22

Element 'Z' in sub-group 2[IIA] is below element 'Y' in the same sub-group. The element 'Z' will be expected to have ............... [higher/lower] atomic size and ............... [more/less] metallic character than 'Y'

Element 'Z' in sub-group 2[IIA] is below element 'Y' in the same sub-group. The element 'Z' will be expected to have higher atomic size and more metallic character than 'Y'

#### Question 23

Argon in period 3 is likely to have a ............... [larger/smaller] atomic size than chlorine and it's electron affinity value would be ............... [greater/lesser/zero] compared to chlorine.

Argon in period 3 is likely to have a larger atomic size than chlorine and it's electron affinity value would be zero compared to chlorine.

#### Question 24

Across a period — Atomic size and metallic character ............... while I.P., E.A., E.N. and non-metallic character ............... and nuclear charge ............... [increases/decreases].

Down a group — Atomic size and metallic character ............... while I.P., E.A., E.N. and non-metallic character ............... and nuclear charge ............... [increases/decreases].

Across a period — Atomic size and metallic character decreases while I.P., E.A., E.N. and non-metallic character increases and nuclear charge increases.

Down a group — Atomic size and metallic character increases while I.P., E.A., E.N. and non-metallic character decreases and nuclear charge increases.

## Give reasons

#### Question 1

Give reasons for the following:

In the same period or subgroup a gradual change in a particular property may be seen.

In the same period or sub group a gradual change in particular property may be seen because physical and chemical properties are periodic function of their atomic number and as the elements are arranged in an increasing order of atomic number in the periodic table hence we see a gradual change.
In other words, it is due to the gradual change in electronic configuration in the arranged elements.

#### Question 2

Give reasons for the following:

Atomic size of group 18 [0 group] elements is more than the atomic size of group 17 [VIIA] elements.

The outer shell of group 18 [0 group] elements is completely filled. Due to this force of repulsion is maximum. The effect of nuclear pull over the valence shell electrons is not seen. Hence, Atomic size of group 18 [0 group] elements is more than group 17 [VIIA] elements.

#### Question 3

Give reasons for the following:

Ionization potential increases with increase in nuclear charge of the elements.

As nuclear charge increases, the nuclear attraction on the outer most electron increases and the outer electron is more firmly held. Therefore, ionization potential increases.

#### Question 4

Give reasons for the following:

Electron affinity of noble gas elements is zero.

Noble gas elements have completely filled outer-shell. Such electronic configurations are highly stable and as such noble gases find it difficult to accept electrons. Thus electron affinity of noble gas elements is zero.

#### Question 5

Give reasons for the following:

Phosphorus, sulphur and chlorine are electronegative elements of the periodic table.

Phosphorus, Sulphur and Chlorine are the rightmost elements of period 3 in the periodic table. Across a period from left to right electronegativity increases as nuclear charge increases and atomic size decreases. This makes Phosphorus, Sulphur and Chlorine electronegative elements of the periodic table.

#### Question 6

Give reasons for the following:

Sulphur is placed in group 16 [VIA], chlorine in group 17[VIIA] but argon in group 18 [0 group] of the periodic table.

Elements are arranged in increasing order of atomic number in the periodic table. The elements placed in group 16 have 6 electrons in the outer most shell, elements of group 17 have 7 valence electrons and the group 18 elements have 8 valence electrons.

Electronic configurations of Sulphur (S), Chlorine (Cl) and Argon (Ar) are:

S = 2, 8, 6
Cl = 2, 8, 7
Ar = 2, 8, 8

With 6 valence electrons S is placed in Group 16, Cl with 7 valence electrons is placed in Group 17 and Ar with 8 valence electrons is in Group 18.

#### Question 7

Give reasons for the following:

Fluorine is the most electronegative element of the periodic table.

Fluorine is placed in period 2 group 17(VIIA) in the modern periodic table i.e., it occupies the upper right hand corner of the periodic table. Along a period from left to right electronegativity increases and down a group from top to bottom it decreases. Hence, Fluorine being the topmost and rightmost element in its group and period is the most electronegative element of the periodic table.

#### Question 8

Give reasons for the following:

Atoms with large atomic radii and low ionization potential are more metallic in nature.

Metals have the tendency to lose one or more electrons i.e., they are are electropositive in nature. Atoms with large atomic radii and low ionization potential can easily lose one or more electrons because the nuclear pull on the outer electrons is less. Therefore, these atoms are more metallic in nature.

#### Question 9

Give reasons for the following:

A decrease in ionization potential of an element leads to a decrease in non-metallic character of the element.

When there is a decrease in ionization potential then the tendency to lose electron increases. This results in a decrease in non-metallic character and increase in metallic character.

#### Question 10

Give reasons for the following:

Atomic size decreases across a period but increases down a group of the periodic table.

Across a period from left to right, nuclear charge increases. This decreases the size of the atom because the electrons are then attracted towards the nucleus with a greater force thereby bringing the outermost shell closer to the nucleus.
Down a group, the number of shells and nuclear charge both increase. But increase in the number of shells dominate over increase in nuclear charge. This increase in the number of shells increases the size of an atom because the distance between the outermost shell and the nucleus increases.

## Unit Test Paper I — Periodic Table

#### Question 1

In period 2, element 'A' is to the right of element 'B'

1. The element 'A' would probably have a ............... [smaller/larger] atomic size than 'B'.
2. The element 'B' would probably have ............... [lower/higher] ionization potential than 'A'.
3. The element 'A' would have ............... [lesser/higher] electron affinity than 'B'.
4. Nuclear charge of element 'B' would be ............... [less/more] than element 'A'.
5. If an element 'C' had a low electronegativity and ionization potential it would have more tendency to ............... [gain/lose] electrons.

1. The element 'A' would probably have a smaller atomic size than 'B'
2. The element 'B' would probably have lower ionization potential than 'A'.
3. The element 'A' would have higher electron affinity than 'B'.
4. Nuclear charge of element 'B' would be less than element 'A'.
5. If an element 'C' had a low electronegativity and ionization potential it would have more tendency to lose electrons.

#### Question 2

With reference to period 3 of the periodic table — State :

1. The type of bonding of the element with electronic configuration 2, 8, 7.
2. The formula of the chloride of the element with electronic configuration 2, 8, 4.
3. The nature of the oxide of the alkaline earth metal in the period.
4. The number of electrons in the penultimate shell of the element with valency -1.
5. The electronic configuration of the element whose hydroxide is a weak base.

1. The element with electronic configuration 2, 8, 7 will form ionic bonds with metals and with non-metal it will form covalent bond.
2. The element with electronic configuration is 2, 8, 4 is Si and with chloride it forms SiCl4.
3. The alkaline earth metal in the period 3 is Mg and it's oxide is basic in nature (i.e. MgO is basic is nature).
4. The number of electrons in the penultimate shell of the element with valency -1 is 8.
5. The element is Magnesium and electronic configuration is (2, 8, 2)

#### Question 3

With reference to group 1 [IA] of the periodic table – fill in the blanks with the correct word:

The elements are ............... [light/heavy] ............... [metals/non metals] since their atomic size is ............... [large/small]. The energy binding the atoms is ............... [high/low] and hence the elements have ............... [high/low] melting points. The melting points of the elements ............... [increases/decreases] down the subgroup. The electropositive character ............... [increases/decreases] down the subgroup and the elements are strong ............... [reducing/oxidizing] agents. The element with electronic configuration 2, 8, 1 will have ............... [higher/lower] electron affinity and ............... [smaller/larger] atomic size than the element with electronic configuration 2, 1.

The elements are light metals since their atomic size is large. The energy binding the atoms is low and hence the elements have low melting points. The melting points of the elements decreases down the subgroup. The electropositive character increases down the subgroup and the elements are strong reducing agents. The element with electronic configuration 2, 8, 1 will have lower electron affinity and larger atomic size than the element with electronic configuration 2, 1.

#### Question 4

Match the elements in column 'X' with the correct group they belong from column 'Y'

X    Y
1: Element with atomic number 19A: Group 18 [0 group]
2: Element with electronic configuration 2B: Group 16 [VI A]
3: Element with a valency of -2C: Group 1 [IA]
4: Element 'P' which loses 3 electrons to form a cationD: Group 17 [VII A]
5: Element 'Q' in period-3 which has the highest electron affinityE: Group 13 [III A]

XY
1: Element with atomic number 19A: Group 1 [IA]
2: Element with electronic configuration 2B: Group 18 [0 group]
3: Element with a valency of -2C: Group 16 [VI A]
4: Element 'P' which loses 3 electrons to form a cationE: Group 13 [III A]
5: Element 'Q' in period-3 which has the highest electron affinityD: Group 17 [VII A]

#### Question 5.1

Give reasons for the following:

Occurrence of characteristic properties of elements takes place at definite intervals in the modern periodic table.

After definite intervals of atomic number, similar valence shell electronic configuration occurs. Properties of elements depend upon their electronic configuration. As in the modern periodic table, the elements are arranged in increasing order of their atomic numbers hence we observe the occurrence of characteristic properties of elements at definite intervals.

#### Question 5.2

Give reasons for the following:

Properties of elements are periodic functions of their atomic numbers and not atomic weights.

Atomic number of an element is equal to the number of protons (or electrons in case of a neutral atom). Physical and chemical properties of elements depend on the number of electrons and their arrangement. Thus, properties of elements are periodic function of their atomic numbers and not atomic weights.

#### Question 5.3

Give reasons for the following:

Atomic size of an element depends on the nuclear charge of that element.

Nuclear charge is the positive charge present in the nucleus of an atom and when there is an increase in nuclear charge then the atomic size decreases because the electrons are attracted towards the nucleus with a greater force thereby bringing the outer most shell closer to the nucleus.

#### Question 5.4

Give reasons for the following:

Down a group electronegativity should increase with increase in nuclear charge but it is seen that the electronegativity decreases.

On moving down a group both atomic size and nuclear charge increases but the increase in atomic size dominates over increase in nuclear charge hence Electronegativity decreases on moving down a group.

#### Question 5.5

Give reasons for the following:

If combining atoms have nearly similar electro­negativities the bond between them is covalent.

If two combining atoms have almost similar electronegativities then transference of electrons cannot take place and the atoms share their valence electron equally with each other hence covalent bond is formed by mutual sharing of electrons.

#### Question 6

Arrange the following elements as per the guidelines in brackets.

1. Na, Cl, Mg, P [in decreasing order of atomic size]
2. C, Li, F, N [in increasing order of electronegativity]
3. Cl, Al, Na, S [in increasing order of ionization potential]
4. Li, F, C, O [in increasing order of electron affinity]
5. Ar, He, Ne [in increasing order of number of electron shells]