Periodic Table

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.

(i) Electronegativity.

(ii) Valence Electrons.

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

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

(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.

(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.

Ionization Potential

(i) 15

(ii) 19

(iii) 8

(iv) 4

(v) 2

(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.

Electron affinity

(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.

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.

Chlorine

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

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

Electronegativity

(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.

(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.

Helium

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

Electronegativity

(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.

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

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

Electronegativity

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

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:

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

∴ Formula of the compound is B₂A₃

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.

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.

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.

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.

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.

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.

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.

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

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.

Argon

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

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.

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

(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.

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.

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.

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.

Electronegativity

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

(a) Beryllium > Carbon > Fluorine

(b) Potassium < Sodium < Lithium

(a) Hydrogen

(b) Sulphur

(c) Chlorine

(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.

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.

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.

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.

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.

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

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.

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.

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.

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.

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.

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.

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.

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.

(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.

(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.

(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.

Lithium, chlorine.

Helium

Argon

3

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

One

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

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

Group 18 (Inert gases group)

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

Chlorine

Mg(OH)₂

HCl

Na₂SO₃

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.

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.

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.

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.

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.

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.

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.

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.

Fluorine

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

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.

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.

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.

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.

K (Potassium)

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

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 increase by 1 while down a subgroup they remain same.

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 most reactive, while elements on the extreme right [group 18 (0)] are unreactive.

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 Br.

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 basic to acidic while the strength of oxy-acids increases from left to right.

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 than the atomic size of fluorine.

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, hence ionization potential increases.

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 0 .

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 covalent

Elements with low electronegativity are usually metallic.

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

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 non-metal.

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 atomic size and more metallic character than 'Y'

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 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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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)

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.

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.

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.

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.

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.

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.

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