Periodic Table — Periodic Properties & Variation of Properties

Select the correct 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

Answer

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.

Identify the term :

(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.

Answer

(i) Electronegativity.

(ii) Valence Electrons.

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.

Answer

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

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

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

(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?

Answer

(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 4^th period.

Fill in the blanks :

(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.

Answer

(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.

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)

Answer

Ionization Potential

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

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

(ii) A metal of valency 1.

(iii) A gaseous element with valency 2.

(iv) An element in Group 2.

(v) A rare gas.

Answer

(i) 15

(ii) 19

(iii) 8

(iv) 4

(v) 2

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).

Answer

(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.

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.

Answer

Electron affinity

Give reasons :

(i) Inert gases do not form ions.

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

Answer

(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.

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 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.

Answer

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.

The most electronegative element from the following elements is :

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

Answer

Chlorine

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

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

Answer

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

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.

Answer

Electronegativity

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).

Answer

(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.

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 :

(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.

Answer

(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.

The element with highest ionization potential is :

1. Hydrogen
2. Caesium
3. Radon
4. Helium

Answer

Helium

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

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

Answer

Electronegativity

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.

Answer

(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.

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

Answer

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

In the Periodic Table, elements of Period 3 are arranged in the increasing order of ionization potential as:

1. B, N, CI, Ar
2. Mg, Si, S, Ar
3. Ar, Si, S, Mg
4. Si, Ar, Cl, Mg

Answer

Mg, Si, S, Ar

Reason — Ionisation potential increases across a period from left to right. Argon in period 3 has maximum ionisation potential, since its outermost shell is completely filled. Hence, the increasing order of ionisation potential of period 3 is - Mg < Si < S < Ar

The tendency of an atom to attract shared pair of electrons to itself when forming a chemical bond is known as

1. Electron affinity
2. Electronegativity
3. lonization potential
4. Nuclear charge

Answer

Electronegativity

Reason — Electronegativity is the tendency of an atom to attract shared pair of electrons to itself when forming a chemical bond.

Elements A & B have electronic configuration : 8 & 13 respectively. The chemical formula formed between A & B will be:

1. AB
2. B₃A₃
3. A₂B₃
4. B₂A₃

Answer

B₂A₃

Reason —

Electronic configuration of Element A : 8 — 2, 6
Electronic configuration of Element B : 13 — 2, 8, 3

Valency of element A : 2
Valency of element B : 3

Formula of the compound formed by the combination of A and B:

$\overset{\phantom{3}{3}}{\text{B}} \space {\swarrow}\mathllap{\searrow} \space \overset{2}{\text{A}} \Rightarrow \underset{\phantom{2}{2}}{\text{B}} \space {\swarrow}\mathllap{\searrow} \underset{\phantom{3}{3}}{\text{A}} \\[1em] \text{B}_2\text{A}_3$

∴ Formula of the compound is B₂A₃

Alkaline earth metals have the same:

1. Number of valence electrons
2. Number of shells
3. Metallic property
4. lonization potential

Answer

Number of valence electrons

Reason — All group 2 alkali earth metals are the elements having 2 valence electrons in their outermost electrons. Whereas, number of shells and metallic property increases down the group and ionisation potential decreases down the group.

The electronic configuration of four elements is:

(i) W: 2,6
(ii) X: 2,8
(iii) Y: 2,8,1
(iv) Z: 2,8,7

Which pair of atoms will form a covalent compound.

1. Two atoms of W
2. Two atoms of X
3. An atom of W & an atom of X
4. An atom of Y & an atom of Z

Answer

Two atoms of W

Reason — The electronic configuration of four elements :

(i) W: 2,6 — Atomic number 8
(ii) X: 2, 8 — Atomic number 10
(iii) Y: 2,8,1 — Atomic number 11
(iv) Z: 2,8,7 — Atomic number 17

Covalent bonding occurs between non-metals by sharing electrons. W needs 2 more electrons to complete its octet. So, two W atoms can share electrons with each other, forming a covalent bond
Hence, two atoms of W will form covalent bond.
However, X has completely filled outer shell, so it is not involved in covalent bond formation and bond between Y and Z is ionic.

Element with an atomic number 19 will:

1. Accept an electron & get oxidized
2. Accept an electron & get reduced
3. Lose an electron & get oxidized
4. Lose an electron & get reduced

Answer

Lose an electron & get oxidized

Reason — The element with atomic number 19 has the electronic configuration is 2, 8, 8, 1. It has 1 valence electron. It tends to lose its 1 valence electron to achieve a stable noble gas configuration. The loss of electron is defined as oxidation.

Element with low ionization potential likely to be:

1. Metal
2. Metalloid
3. Non-metal
4. Inert gas

Answer

Metal

Reason — The amount of energy required to remove a loosely bound electron from the outermost shell of an atom is known as ionisation potential. Metals have loosely held outer electrons and tend to lose electrons easily. Hence, they have low ionisation potential.

The non-metallic properties of elements from left to right in a Periodic Table:

1. Increases
2. Decreases
3. Remains same
4. First increases & then decreases

Answer

Increases

Reason — As we move from left to right across a period in the periodic table the atomic size decreases, nuclear charge increases, tendency to gain electrons increases. These changes cause the non-metallic character to increase.

The electronic arrangements of six atoms, A to F is shown below.

A: 2,5    B: 2    C: 2,6    D: 2,8,6    E: 2,8,8    F: 2,8,3

The atoms from the same group of the periodic table:

1. D & E
2. C & D
3. E & F
4. C & E

Answer

C & D

Reason — The electronic configuration of six atoms is given below :

Atom C (2,6) and D (2,8,6) both have 6 valence electrons. So, they both belong to Group 16.

The electronic arrangements of six atoms, A to F is shown below.

A: 2,5    B: 2    C: 2,6    D: 2,8,6    E: 2,8,8    F: 2,8,3

Two noble gases:

1. A & B
2. E & F
3. B & E
4. D & E

Answer

B & E

Reason — The electronic configuration of six atoms is given below :

Atom B (2) and E (2, 8, 8) both have completely filled outermost electron shells. Their electronic configurations reflect stable duplet and octet arrangements. Hence, B and E are nobel gases.

The electronic arrangements of six atoms, A to F is shown below.

A: 2,5    B: 2    C: 2,6    D: 2,8,6    E: 2,8,8    F: 2,8,3

The atom which is the most electronegative:

1. A
2. B
3. C
4. F

Answer

C

Reason — The electronic configuration of six atoms is given below :

Electronegativity increases with increasing number of valence electrons. A has 5 valence electrons and C has 6. Hence, atom C has the highest electronegativity. Atom D has also has 6 valence electrons, but they are in the 3rd shell so less attraction due to increased distance. Hence, D is less electronegative than C

The electronic arrangements of six atoms, A to F is shown below.

A: 2,5    B: 2    C: 2,6    D: 2,8,6    E: 2,8,8    F: 2,8,3

The atom which has the highest ionization potential:

1. A
2. B
3. E
4. F

Answer

B

Reason — The electronic configuration of six atoms is given below :

Element B is helium, it has a completely filled outer shell, so helium has the highest ionisation potential. Element E also has completely filled outer shell, but it has 3 shells, so electrons are further away, hence it has lower ionisation energy than B.

An element in period 3, whose electron affinity is zero:

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

Answer

Argon

Reason — Argon belongs to period 3. It is an inert element and its octet is complete hence, it's electron affinity is zero.

An element with the largest atomic radius is:

1. Carbon
2. Nitrogen
3. Lithium
4. Beryllium

Answer

Lithium

Reason — The given elements belong to the second period and Lithium is the first element of second period. Atomic size decreases from left to right in a period hence, Lithium will have the largest atomic radius.

State the term - The energy released when an atom in the gaseous state accepts an electron to form an anion.

Answer

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

Arrange the following:

(i) N, Be, O, C (in increasing order of non-metallic character)

(ii) P. Si, F. Be (in decreasing order of valence electrons)

Answer

(i) Be < C < N < O

(ii) F > P > Si > Be

Reason

(i) Across a period, left to right in a periodic table, the non-metallic character increases.

(ii) The valence electrons decreases from right to left in the periodic table.

Give a reason for — Ionisation potential decreases down a group.

Answer

On moving down the group, atomic size, as well as, nuclear charge increases. However, the effect of increase in atomic size dominates over the effect of increase in nuclear charge. Hence, ionization potential decreases down the group.

In the 2nd period Neon has maximum Ionization Potential because

1. It has unstable electronic configuration
2. It easily accepts electrons
3. It easily loses electrons.
4. The outermost shell is completely filled

Answer

The outer most shell is completely filled.

Reason — The outer most shell is completely filled making it extremely stable and requiring more energy to remove an electron. Therefore, Neon has maximum Ionization Potential in the 2nd period.

Electron Affinity is maximum in :

1. Mg
2. Ar
3. Li
4. Br

Answer

Br

Reason — Electron affinity tends to be highest for elements on the right side of the periodic table. Therefore, Bromine (Br) is likely to have the highest electron affinity.

State the term : The tendency of an atom to pull a shared pair of electrons towards itself in a compound.

Answer

Electronegativity

Give reasons for — Inert gases do not form ions.

Answer

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

Arrange the following :

(a) Carbon, Fluorine, Beryllium (in decreasing order of atomic size).

(b) Potassium, Lithium, Sodium (in increasing order of ionization potential).

Answer

(a) Beryllium > Carbon > Fluorine

(b) Potassium < Sodium < Lithium

Identify the following:

(a) An element in period 1 which can be placed in both group 1 & group 17 of the Periodic Table.

(b) The element having electronic configuration 2, 8, 6.

(c) The most electronegative element of period 3.

Answer

(a) Hydrogen

(b) Sulphur

(c) Chlorine

The table shows, the electronic configuration of the atoms A, B, C & D. Answer the following:

(a) The formula of the compound formed between

(i) A and B
(ii) D and C

(b) Which element's exhibits catenation ?

Answer

(a) The formula of the compounds formed are :

(i) AB is the formula of the compound formed between A and B.
Reason — A has 2 electrons in its outermost shell. B has 6 electrons in its outermost shell. To form a compound between A and B, A will donate its 2 outermost electrons to B. Hence, the formula of the compound is AB as both have 2 as valency.

(ii) DC₄ is the formula of the compound formed between D and C.
Reason — D has 4 electrons in its outermost shell. C has 7 electrons in its outermost shell. Therefore, their valency is 4 and 1 respectively. So, 4 atoms of C will share their one electron each with one atom of D to form a compound. Hence, the formula of the compound is DC₄.

(b) D will exhibit catenation.
Reason — Electronic configuration of D (2, 4) suggests that it is Carbon. Carbon exhibits catenation to the maximum extent due to tetra-covalency of carbon and greater strength of carbon-carbon bond.

Assertion (A): Noble gases Ne, Ar have electron affinity zero.

Reason (R): Greater the value of electron affinity, the more electronegative is the element.

1. Both A & R are true and R is the correct explanation of A.
2. Both A & R are true but R is not the correct explanation of A.
3. A is true but R is false
4. A is false but R is true.

Answer

Both A & R are true but R is not the correct explanation of A.

Reason — 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. Hence, the assertion (A) is true.
Elements with high electron affinity tend to strongly attract electrons, and are therefore usually more electronegative. So, greater the value of electron affinity, the more electronegative is the element. Hence, the reason (R) is true but it doesn't explain why nobel gases have zero electron affinity.

Assertion (A): Elements with high electronegativity are usually non-metals.

Reason (R): Non-metallic atoms have high ionisation potential & tend to gain electrons.

1. Both A & R are true and R is the correct explanation of A.
2. Both A & R are true but R is not the correct explanation of A.
3. A is true but R is false
4. A is false but R is true.

Answer

Both A & R are true and R is the correct explanation of A.

Reason — Electronegativity is the tendency of an atom to attract electrons in a chemical bond. Non-metals readily attract electrons and have high electronegativity. Hence, the assertion (A) is true.
Non-metals don’t lose electrons easily, and they tend to gain electrons so they have high ionization energy and are highly electronegative. Hence, the reason (R) is true and it explains why elements have high electronegativity.

The oxide of the metallic element with three valence electrons, in its third shell is:

1. Basic
2. Weakly acidic
3. Amphoteric
4. Strongly basic

Answer

Amphoteric

Reason — Three valence electrons in the third shell means the element has the configuration: 2,8,3. This corresponds to atomic number 13, which is Aluminum (Al).
Aluminum forms Al₂O₃ (aluminum oxide). This oxide is amphoteric, it can act as both an acid and a base.

'ACl₂' is the formula of an electrovalent solid. 'A' is likely to be in the same group in period 3, as the element given below present in period 2.

1. Be
2. Li
3. C
4. B

Answer

Be

Reason — A is in period 3. The formula ACl₂ suggests that A has a valency of +2, so, A is a metal that forms A²⁺ ions.
In Period 3, the element with valency 2 is Magnesium (Mg), it forms MgCl₂, an ionic solid.

∴ Be (Beryllium) is in Group 2, period 2, the same group as Mg. Be and Mg both form +2 ions, and their chlorides have the formula ACl₂.

Metallic element 'X' in period 3, having valency 2, reacts with a non-metallic element 'Y' of the same period also having valency 2. The number of non-metallic atoms in the compound formed between 'X' & 'Y' is:

1. 3
2. 1
3. 4
4. 2

Answer

1

Reason —

Element X : A metal in Period 3 with valency 2 will be present in group 2, so the element will be magnesium (Mg)

Element Y : A non-metal in Period 3 with valency 2 will be present in group 16, so the element will be sulphur (S).

Magnesium (Mg) with valency of 2 will form Mg²⁺ by the loss of 2 electrons.
Sulphur (S) with valency of 2 will form S^2- by the gain of 2 electrons.

Mg²⁺ and S^2- combine to form MgS.

∴ The number of non-metallic atoms in the compound formed between 'X' & 'Y' is 1

State the element with the lowest electron affinity from the elements whose electronic configurations are below:

1. 2, 8, 8
2. 2, 7
3. 2, 8, 6
4. 2, 8, 7

Answer

2, 8, 8

Reason — The electronic configuration 2, 8, 8 is of argon, a nobel gas. 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.

Non-metallic elements have:

1. Large atomic size
2. Low electron affinity
3. High ionisation potential
4. Greater tendency to lose electrons.

Answer

High ionisation potential

Reason — Non-metals possess relatively small atomic radii and their valence electrons are strongly attracted by the nucleus. Consequently, a large amount of energy is required to remove an electron, giving them high ionization energy.

The electronic configuration of the element which has the largest atomic size is:

1. 2, 8, 5
2. 2, 8, 1
3. 2, 8, 7
4. 2, 1

Answer

2, 8, 1

Reason — Atomic radius increases down a group because each successive period adds a new electron shell, and it decreases from left to right across a period owing to increasing nuclear charge.

• Configurations 2, 1 and 2, 8, 1 belong to Group 1.
• Element with electronic configuration (2, 8, 1) is one period below lithium, so it possesses an additional electron shell and therefore a larger atomic radius.
• The remaining configurations (2, 8, 5) and (2, 8, 7) lie further to the right in Period 3, where atomic size is smaller.

Hence, the element with configuration 2, 8, 1 has the largest atomic size among the choices given.

Rewrite the incorrect statements only by rectifying the incorrect word/s in the statement.

(a) Periodic table is based on the fundamental property atomic weight.

(b) Elements in the periodic table are arranged in seven horizontal rows called 'groups', which ends with an element having one electron in outermost shell.

(c) Period numbers i.e., 1, 2, 3 signifies the no. of electron shells of an element e.g. period 3 element 'Mg' has 2 shells.

(d) Across a period from left to right the no. of electron shells increases by one & the valence electron remains same.

(e) Group numbers i.e., 1, 2, 3 signifies the number of electron shells of an element.

(f) Down a sub-group, the valence electrons & the number of electron shells remains same.

Answer

The corrected statements are given below:

(a) The periodic table is based on the fundamental property atomic number.

(b) Elements in the periodic table are arranged in Seven horizontal rows called periods, which ends with an element having a full outermost shell.

(c) Period numbers i.e., 1, 2, 3 signify the number of electron shells of an element e.g. period 3 element 'Mg' has 3 shells.

(d) Across a period from left to right, the no. of electron shells remains same and the number of valence electrons increases.

(e) Group numbers i.e., 1, 2, 3 signify the number of valence electrons of an element.

(f) Down a sub-group, the valence electrons remain the same & the number of electron shells increases.

Correct the order of elements in period 1 to 3 of the periodic table & state their correct electronic configurations.

Answer

Give reason — Noble gases are considered, unreactive elements.

Answer

Noble gases are largely unreactive because their outermost electron shells are completely filled (helium has a stable duplet of two electrons, while the other noble gases have an octet of eight electrons). With a full valence shell, they have no strong tendency to gain, lose, or share electrons, so they rarely take part in chemical reactions.

Give reason — In a sub-group the chemical properties of elements remain similar.

Answer

Elements belonging to the same group have the same number of valence (outermost) electrons, so their outermost electronic configurations are alike. As chemical properties depend mainly on this outer electronic configuration, the elements in a group exhibit similar chemical properties.

Give reason — Halogens in group 17 are strong oxidising agents.

Answer

Halogens have seven valence electrons and therefore need only one additional electron to attain a stable octet. Their very high electron affinity and electronegativity make them readily accept an electron from other substances, so they act as powerful oxidising agents.

Give reason — Atomic radii is considered a periodic property, exhibiting periodicity.

Answer

When the elements are arranged in order of increasing atomic number, the atomic radius shows a regular, repeating pattern: it generally decreases from left to right across a period (because of increasing nuclear charge) and increases down a group (due to the addition of new electron shells). Because this variation repeats in each successive period, atomic radius is classified as a periodic property.

Complete the definitions of the periodic properties given below :

Answer

(a) Atomic size — The distance between the centre of the nucleus of an atom and its outermost shell.

(b) Ionisation potential — The amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom.

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

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

(e) Metallic character — Metallic character of an element is exhibited when an atom of the element loses one or more electrons when supplied with energy.

Complete the table of trends in periodic properties by -

(a) Filling in the blanks with 'decreases' or 'increases' across a period & down a group.

(b) The reasons for the 'decrease' or 'increase' based on the criteria given.

Answer

(a) Periodic properties

(b) Reasons for trends in periodic property

Atomic Size

1. Across a period — Decreases
   Reason — The atomic size decreases across a period because of increase in nuclear charge. The number of shells remain same across a period.
2. Down a group — Increases
   Reason — Atomic size increases due to increase in number of shells as new shells are added with increasing atomic number. Nuclear charge also increases but number of shells dominates over nuclear charge.

Ionisation potential

1. Across a period — Increases
   Reason — The atomic radii decreases because the electrons are added to same shell. The nuclear charge increases. Hence the outer electrons are more firmly held.
2. Down a group — Decreases
   Reason — Ionisation potential decreases because atomic radii increases as number of shells increases, outermost electrons can be removed easily. Nuclear charge also increases but number of shells dominates over nuclear charge.

Electron affinity

1. Across a period — Increases
   Reason — As nuclear charge increases and atomic size decreases, the nucleus exerts a stronger pull on an added electron. Atoms are closer to achieving a full octet, so they release more energy when gaining an electron.
2. Down a group — Decreases
   Reason — Atomic size increases, so the added electron experiences less nuclear attraction due to its greater distance and inner electrons reduce the attraction between the nucleus and outer electrons. Therefore, less energy is released.

Electronegativity

1. Across a period — Increases
   Reason — Nuclear charge increases while atomic radius decreases, making it easier for atoms to attract shared electrons in a bond.
2. Down a group — Decreases
   Reason — Even though nuclear charge increases, the atomic radius becomes larger, and the outer electrons are farther from the nucleus. This reduces the nucleus's ability to attract bonding electrons.

Metallic character

1. Across a period — Decreases
   Reason — Atomic radius decreases and nuclear charge increases, so atoms hold electrons more tightly and do not lose them easily reducing metallic behavior.
2. Down a group — Increases
   Reason — Atomic radius increases, and outer electrons are farther from the nucleus and held less tightly. This makes it easier for atoms to lose electrons, which is a typical property of metals.

State the following

(i) Physical properties of elements which exhibit periodicity

(ii) Nature of hydroxides of elements which exhibit periodicity

Answer

(i) Physical properties of elements which exhibit periodicity are :

1. Density
2. Melting point
3. Boiling point

Across a period — Density and melting point of elements increases gradually and then slightly decreases thereafter

Down a group — Density increases gradually, melting point and boiling point of elements decreases gradually

(ii) Nature of hydroxides of elements which exhibit periodicity is given below

1. Across a period — The nature of hydroxides changes from strongly basic to amphoteric.
2. Down a group — The nature of hydroxides changes from less basic to strongly basic, due to increasing atomic size and metallic character.

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

Answer

Lithium, chlorine.

The noble gas having duplet arrangement of electrons.

Answer

Helium

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

Answer

Argon

The number of electron shells in elements of period 3.

Answer

3

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

The valency of elements in group 1[IA].

Answer

One

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

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

Answer

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

The group whose elements have zero valency.

Answer

Group 18 (Inert gases group)

Reason — The group 18 elements have 2 or 8 valence electrons and are stable elements hence have zero valency.

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

Answer

Chlorine

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

Answer

Mg(OH)₂

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

Answer

HCl

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

Answer

Na₂SO₃

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

Answer

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.

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

Answer

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.

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

Answer

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.

A light element in period 3 Which is a stable metal and a light metal.

Answer

Magnesium (Mg)

Reason — Magnesium is a light element in Period 3 that is considered a stable and light metal because it has a low atomic mass, moderate reactivity, and forms a protective oxide layer that prevents rapid corrosion, making it chemically stable.

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

Answer

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.

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

Answer

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.

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

Answer

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.

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

Answer

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.

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

Answer

Fluorine

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

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

Answer

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.

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

Answer

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.

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

Answer

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.

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

Answer

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.

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

Answer

K (Potassium)

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

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].

Answer

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.

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

Answer

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

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

Answer

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

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].

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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.

Answer

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.

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.

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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

Answer

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

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

Answer

Elements with low electronegativity are usually metallic.

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

Answer

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

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

Answer

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

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

Answer

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

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'.

Answer

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'.

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'

Answer

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'

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.

Answer

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.

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].

Answer

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 for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

Electron affinity of noble gas elements is zero.

Answer

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.

Give reasons for the following:

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

Answer

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.

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.

Answer

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.

Give reasons for the following:

Fluorine is the most electronegative element of the periodic table.

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

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.

Answer

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.

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.

Answer

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)

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.

Answer

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.

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

Answer

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

Give reasons for the following:

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

Answer

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.

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]

Answer

1. Na > Mg > P > Cl
2. Li < C < N < F
3. Na < Al < S < Cl
4. Li < C < O < F
5. He < Ne < Ar