Periodic Table, Periodic Properties and Variations of Properties

(a) The elements with atomic number 9 (fluorine) & 35 (bromine) have physical properties resembling chlorine because they belong to the same group.

(b) The element with atomic number 9 (fluorine) is more electronegative than chlorine because on moving down the group electronegativity decreases.

It is given that,

X = 500 kJ mol^-1

Y = 375 kJ mol^-1

Their relative position in a group is X & Y as ionisation enthalpy decreases on moving down the group.

Their relative position in a period is Y & X as ionisation enthalpy increases on moving across a period.

(a) These elements belong to the 2^nd period.

(b) The missing element is Nitrogen. It belongs to the Group 15 and should be placed as shown below:
    Li Be B C N O F

(c) Beryllium < Nitrogen < Fluorine.

(d) Fluorine.

(a) The element belongs to 3^rd period.
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 3^rd period as it has three shells.

(b) 6 [∵ electronic configuration is 2, 8, 6]

(c) Z is a non metal.

(d) H₂Z

Given, M is a metal

(i) 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.

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

(iii) M is a metal which belongs to group 1 whereas oxygen is a non metal with higher electronegativity compared to a M.

(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.

1. J is the most electronegative element.

2. B is the most reactive element of group IA or 1.

3. K is the element from period 3 with smallest atomic size.

4. Ne the highest ionization potential.

5. 5 valence electrons present in G.

6. C is the element from group 2 which has least ionization energy.

7. L is the noble gas of the fourth period.

8. ionic bond is formed between A and H.
   The molecular formula is :
   2A + H ⟶ A₂H

1

Reason — The elements placed in group 1 of the periodic table are known as alkali metals (except hydrogen) as they form strong alkalis with water.

They are metallic in nature.

Reason — Halogens are non-metallic in nature.

Ionization potential increases from left to right across a period.

Reason — On moving from left to right across a period, the atomic size decreases due to increase in the nuclear charge, and thus, more energy is required to remove the electrons. Hence, ionization potential increases from left to right across a period.

Argon

Reason — Argon is the noble gas in period 3 having an electron affinity of zero.

7

Reason — Halogens have 7 valence electrons.

Fluorine

Reason — Electron affinity increases from left to right in a period. Out of the given options, Lithium, Carbon and Fluorine are period 2 elements in that order. Hence, Fluorine has the highest electron affinity among period 2 elements.

Atomic radius decreases and nuclear charge increases

Reason — On moving from left to right across a period, the atomic radius decreases due to increase in the nuclear charge, and thus, more energy is required to remove the electrons. Hence, ionization potential increases from left to right across a period.

3 shells and 2 valence electrons

Reason — The number of shells present in an atom determines its period and valence electrons determine its group.

Lithium

Reason — Lithium, Boron, Carbon, Fluorine is the order of given elements from left to right in period 2. As electronegativity increases from left to right in a period, hence, Lithium has the least electronegativity in the given elements.

17

Reason — Element with atomic number 19 has an electronic configuration of 2, 8, 8, 1. It is a metal with one valence electron. Element with atomic number 17 has an electronic configuration of 2, 8, 7. It is a non-metal. Hence, element with atomic number 19 will donate its one valence electron to element with atomic number 17. Both will achieve stable octet. Hence, these two elements are most likely to combine chemically.

(i) increases.
Reason — On moving across a period from left to right, nuclear pull increases due to increase in atomic number and thus atomic size decreases. Hence, elements cannot lose electrons easily. Therefore, non-metallic character increases.

(ii) increases.
Reason — On moving across a period from left to right, nuclear charge increases due to an increase in atomic number. Hence, electronegativity increases from left to right in a period.

(iii) increases.
Reason — On moving from left to right across a period, the atomic radius decreases due to increase in the nuclear charge, and thus, more energy is required to remove the electrons. Hence, ionization potential increases from left to right across a period.

(iv) decreases.
Reason — On moving across a period from left to right, nuclear pull increases due to increase in atomic number and thus atomic size decreases.

ionic radius will increase

Reason — On going down in the halogen group, number of shells increases, hence, ionic radius also increases. Reactivity, electronegativity and ionisation potential decreases while going down in the halogen group.

Energy released when an electron is added to an isolated atom in the gaseous state.

Reason — The amount of energy released while converting a neutral gaseous isolated atom into a negatively charged gaseous ion (anion) by the addition of electron is called Electron Affinity (E.A.).

Reactivity increases in alkali metals but decreases in halogens.

Reason — Atomic number increases while going down the group in alkali metals and halogens. The tendency of losing electrons increases down the group. Reactivity of metals depends on tendency to lose electrons, thus reactivity of alkali metals increases while going down the group. Reactivity of non-metals depends on tendency to gain electrons, thus reactivity of halogens decreases while going down the group.

Sr, Ca, Mg, Be

Reason — Ionisation energy decreases on moving down the group.

Both Q and R

Reason — Element X with an atomic number of 13 is Aluminium (Al), which is a metal. Element Y with an atomic number of 17 is Chlorine (Cl), which is a non-metal.
X being a metal tends to lose electrons to achieve a stable electron configuration and Y being a non-metal tends to gain electrons to achieve stable electron configuration. Hence, statements Q and R are true whereas statement P is false.

A is true but R is false.

Explanation — The second period includes the elements from lithium (Z=3) to neon (Z=10): Li, Be, B, C, N, O, F, Ne → 8 elements. Hence assertion is true. Reason given is incorrect. In the first period, electrons are being added to the K electronic configuration.

The number of available electronic configuration or orbital for elements of first period is 1.

The number of elements present in the first period is 2, which is not four times the number of atomic orbitals available. Hence the (R) statement is false.

A is false but R is true.

Explanation — In Dobereiner’s triads, the atomic mass of the middle element is approximately the average of the atomic masses of the other two. Hence, Assertion (A) is false.
Dobereiner grouped elements into triads because they had similar chemical properties. Hence, Reason (R) is true.

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

Explanation — Assertion (A) is true because smaller atoms attract electrons more strongly. Reason (R) correctly defines electronegativity but does not explain why smaller atoms have higher electronegativity.

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

Explanation — Hydrogen is placed in Group I with alkali metals because, like them, it has one electron in its outermost shell and forms (H⁺) ions by losing this electron. Although hydrogen can also gain an electron to form the hydride ion (H⁻) similar to halogens. Hence, Both Assertion (A) and Reason (R) are true but Reason (R) is not the correct explanation of Assertion (A).

A is false but R is true.

Explanation — Atomic size decreases across a period (from left to right) in the periodic table. This is because, although electrons are being added, they are added to the same shell, and the increasing effective nuclear charge pulls the electrons closer to the nucleus. Hence, Assertion (A) is false.
As atomic number increases, the number of protons increases, which increases the effective nuclear charge. Hence, Reason (R) is true.

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

Explanation — Elements in the same vertical column i.e., a group in the periodic table, have the same number of valence electrons. Since chemical properties of elements depend upon the number of valence electrons, thus elements in the same group have similar properties. Hence, Reason (R) correctly explains the Assertion (A).

Modern periodic table is based on Henry Moseley's modern periodic law which states that "the physical and chemical properties of elements are the periodic functions of their atomic number".

(i) In order of increasing radii:

(a) Cl < Cl^-.
Reason — Anion is greater in size than its parent atom.

(b) Mg²⁺ < Mg⁺ < Mg.
Reason — Cation is smaller in size than its parent atom.

(c) O < N < P.
Reason — Between Oxygen (O) and Nitrogen (N), atomic size of Nitrogen (N) is more as it is to the left of Oxygen (O) in period 2. Between Nitrogen (N) and Phosphorus (P), Phosphorus (P) has larger atomic size as it comes after Nitrogen (N) in group 5A.

(ii) In order of increasing ionisation energy:

(a) Na < P < Cl
Reason — Ionisation energy increases when moving across a period from left to right.

(b) O < F < Ne
Reason — Ionisation energy increases when moving across a period from left to right.

(c) Ar < Ne < He
Reason — Ionisation energy decreases when moving down the group.

The electronic configuration of the given elements are:

A(8) = (2, 6 ) — A forms anion.

B(7) = (2, 5 ) — B forms anion.

C(11) = (2, 8, 1) — C forms cation.

D(12) = (2, 8, 2) — D forms cation.

E(13) = (2, 8, 3) — E forms cation.

F(9) = (2, 7 ) — F forms anion.

(a) The oxidising 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 oxidising power depends on tendency to gain electrons and non-metals are good oxidising agents hence oxidising power of elements increases across a period.

(b) The ionisation 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.

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

(a) The element below sodium in the same group would be expected to have a lower (lower/higher) electro-negativity than sodium and the element above chlorine would be expected to have a higher (lower/higher) ionization potential than chlorine.

(b) On moving down a group, the number of valence electrons remains the same (remains the same/increases/ decreases).

(c) Metals are good reducing agent (oxidising agent/ reducing agent) because they are electron donors (acceptors/donors).

(d) Down the group, electron affinity decreases [increases, decreases, remains same].

(e) Electronegativity across the period increases [increases/ decreases].

(f) Non-metallic character down the group decreases [increases/ decreases].

(g) In a period, increase in electron affinity increases reduction (oxidation/reduction).

(h) On descending a group, decrease (increase/decrease) in ionisation potential as well as electron affinity decreases (increases/decreases) oxidising capacity.

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

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

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

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

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 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.

(a) Na > Mg > Si > S > Cl.

(b) Li < Na < K < Rb < Cs.

(c) K < Na < Si < S < Cl.

(d) I < Br < F* < Cl.
Flourine (F) in group 17 is an exception. It has lower electron affinity than Chlorine (Cl).

(e) Li > Na > K > Rb > Cs.

(f) Pb < Zn < Ca < K.

(g) H > Li > Na > K.

We know that electronegativity increases from left to right in a period. Hence, the order of the elements in the periodic table is:

Na, Mg, Al, Si, P, S, Cl.

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

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

2^nd Group.
Reason — Since A loses 2 electrons to form an anion, A²⁺ so its valency is 2. Hence, it belongs to the 2^nd group.

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

The Modern Periodic Law is stated as, "the physical and chemical properties of elements are the periodic functions of their atomic number."

Henry Moseley.

Groups - 18

Period - 07

Horizontal rows are known as Periods,

Vertical columns are known as Groups.

Periodicity is observed due to similar electronic configuration.

(i) As we move across a period the number of shells remain same & number of valence electron increases by one.

(ii) As we go from top to bottom in a group the number of shells increases one by one but the number of electron remain the same.

(i) Lithium (Li) and Sodium (Na).

(ii) Beryllium (Be) and Magnesium (Mg).

(iii) Fluorine (F) and Chlorine (Cl).

(iv) Helium (He) and Neon (Ne).

Valency is equal to the number of electrons an atom can donate or accept or share. Hence, it depends on the number of electrons in the outermost shell (i.e. valence shell).

Group 1 (IA) elements have one electron in their outermost shell. So they donate this one electron to become stable hence their valency is one. On the other hand, Group 17 (VIIA) elements have seven electrons in their outermost shell. So they accept one electron to attain stable electronic configuration hence their valency is also one.

(i) Elements in the same group have the same valency.

(ii) Valency depends upon the number of valence electrons in an atom.

(iii) Copper and Zinc are transition elements.

(iv) Noble gases are placed at the extreme right of the periodic table.

(i) The properties that reappear at regular intervals, or in which there is gradual variation (i.e. increase or decrease) at regular intervals, are called 'periodic properties' and the phenomenon is known as the periodicity of elements.

(ii) The third period elements are typical elements. They are Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl) .

(iii) Electrons revolve around the nucleus in certain definite circular paths called orbits.

Beryllium and Magnesium will show the similar chemical reactions as calcium. Since these elements belong to the same group 2 and also have same number of valence electrons as calcium.

1. Metals — Lithium (Li), Beryllium (Be), Sodium (Na), Magnesium (Mg), Aluminium (Al), Potassium (K), Calcium (Ca).
2. Metalloids — Boron (B), Silicon (Si).
3. Non-Metals — Hydrogen (H), Helium (He), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne), Phosphorus (P), Sulphur (S), Chlorine (Cl), Argon (Ar).

(i) Similar properties of Fluorine, Chlorine and Bromine are:

1. They have same number of valence electrons (7).
2. They have non-metallic character.
3. They all try to gain one electron to obtain stable electronic configuration.
4. They all have same valency, which is equal to 1.

(ii) Halogens is the common name of this group elements.

Main characteristics of the last element in each period of the Periodic Table are:

1. They have valency equal to zero.
2. They react under very special condition otherwise they don't react.
3. They have stable gas electronic configuration.

General name of such elements is Inert Gases.

Number of valence electrons determines which element will be the first and which element will be the last in a period.

(i) The number of valence electron successively increases by one as we move from left to right in the third period.

(ii) Valency firstly increase from 1 to 4 from element sodium (Na) to silicon (Si) and then reduces to zero from element phosphorus (P) till chlorine (Cl) as we move from left to right in the third period.

(i) Inert Gas or Noble Gas Elements.

(ii) Representative Element.

(iii) Transition Element.

(iv) Halogens.

(v) Alkaline Earth Metals.

(i) Atomic Number is 20.

(ii) It belongs to group 2 and fourth period.

(iii) It is a metal.

(iv) The name assigned to this group is IIA.

(v) Valency is 2.

Electronic configuration of S is : 2, 8, 6

(ii) 16th group and 3rd period.

(iii) Valency of S = 8-6 = 2.

(iv) Non-metal.

(v) It is an oxidising agent.

(vi) Formula with Hydrogen = H₂S

(a) Alkali metal in period 3 = Sodium

Halogen in period 2 = Fluorine

(b) Argon (Ar).

(c) Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne) are non-metals in period 2.

Sodium (Na), Magnesium (Mg) and Aluminium (Al)are metals in period 3.

(d) Silicon (Si).

(e) Argon (Ar).

(f) Be will have lower nuclear charge then Mg.

(i) Group Number of T is 1.

(ii) Period of T is 4.

(iii) Their is 1 valence electron present in atom T.

(iv) Valency of T is 1.

(v) It is a Metal.

(i) A metal of valency one — Atomic number 19 (Potassium).

(ii) A solid non-metal of period 3 — Atomic number 15 (Phosphorus).

(iii) A rare gas — Atomic number 2 (Helium).

(iv) A gaseous element with valency 2 — Atomic number 8 (Oxygen).

(v) An element of group 2 — Atomic number 4 (Beryllium).

Atomic Size is the distance between the centre of the nucleus of an atom and its outermost shell.

Unit — Angstrom (Å) and Picometer (pm).

(i) On moving down the group the atomic size increases as the number of valence shell increases.

(ii) On moving across the period the size of an atom decreases from left to right. This is 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.

Second Period — F < O < N < C < B < Be < Li.

Third Period — Cl < S < P < Si < Al < Mg < Na.

(i) The size of neon is bigger compared to fluorine because the outer shell of neon have a complete octet. They have maximum number of electrons in their outermost orbit thus the electronic repulsions are maximum. So, the effect of nuclear pull over the valence shell electrons is not observed. Hence, the size of neon is greater then fluorine.

(ii) Magnesium is placed to the right of Sodium in period 3. As the atomic size decreases on moving from left to right across the period, hence, atomic size of Sodium is greater than Magnesium.

(a) An atom.

Reason — An atom is greater in size than a cation because cation is formed by the loss of electron(s), hence proton(s) are more than electron(s) in a cation. So electrons are strongly attracted by the nucleus and are pulled inward. Hence, the size decreases.

(b) An anion.

Reason — An anion is greater in size than atom because anion is formed by the gain of electron(s). Thus, the number of electron(s) are more than proton(s). The effective positive charge in the nucleus is less, so less inward pull is experienced. Hence, the size increases.

(c) Fe²⁺.

Reason — Fe²⁺ is greater in size than Fe³⁺ because Fe²⁺ has one more electron than Fe³⁺. So, nuclear pull is more in Fe³⁺ than Fe²⁺, thereby bringing the outermost shell closer to the nucleus. Hence, size decreases.

(d) Oxygen.

Reason — Oxygen is greater in size than fluorine. The position of oxygen in periodic table is before fluorine so, as we move from left to right in a period the outermost shell with more number of electrons is brought closer to the nucleus, decreasing the size of the atom.

Potassium (K) has the maximum metallic character as potassium belongs to the group 4 whereas sodium belongs to group 3 & lithium belongs to group 2, On moving down the group the metallic character increases. So, among Sodium (Na), Lithium (Li), and Potassium (K), Potassium (K) has the maximum metallic character.

(i) F < O < N < C < B < Be < Li .

(ii) Cl > S > P > Si > Al > Mg > Na.

(i) On moving from left to right in the third period, the chemical reactivity of elements first decreases until silicon as the tendency to lose electrons deceases and then starts to increase from phosphorus to chlorine, as the tendency to gain electrons increases.

(ii) Chemical reactivity in metals increases on going down the group, but in non-metals it decreases on going down the group.
(a) As group IA (1) contains metals, hence, chemical reactivity increases on going down the group IA (1).
(b) As group VIIA (17) contains non-metals, hence, chemical reactivity decreases on going down the group VIIA (17).

The metal is Aluminium. Its
Atomic Number = 13,
Valency = 3.

Valency of O is 2 so valency of M is 3. Therefore, M belongs to 13^th group. It is given that M belongs to 3^rd period, hence, it has atomic number 13.

(i) 7.

(ii) Chlorine (Cl).

(iii) Halogens Family.

(iv) XY₃ [AlCl₃ (Aluminium Chloride)].

(i) They belong to the same group as they have same number of electrons in their valence shell.

(ii) As m.p. decrease on going down the group, and metallic character increases on going down the group. Therefore,

B < C < A.

Potassium.

Reason — Sodium (Na), Magnesium (Mg) and Aluminium (Al) belong to period 3 whereas Potassium (K) belongs to period 4. As atomic size increases on moving down the group hence, Potassium has the largest atomic radius of the given elements.

I^-
Reason — Cl, Br and I belong to the same group (group 17) and between the three, I is the bottom most. As atomic size increases down the group hence, between Cl, Br and I, I has the largest atomic size. Between I and I^-, atomic size of I^- is larger as anion is larger than the parent atom. Hence, I^- has the largest atomic size in the given elements.

(i) Barium forms ions more readily due to large atomic size and low ionisation energy. Among the given elements Barium has the largest atomic size because as we move down the group atomic size increases. With increase in atomic size, the ionisation energy decreases.

(ii) The common feature in their electronic configuration are:

1. There are two electrons in their valence shell.
2. They lose two electrons to gain stable gas electronic configuration.

For $_{19}^{39}\text{K}$:

No. of Protons = 19

No. of Neutrons = 39 - 19 = 20

Electronic Configuration = 2, 8, 8, 1

Position in periodic table = 4^th period (as it has 4 shells) & 1^st group (as it has one valence electron).

For $_{15}^{31}\text{P}$:

No. of Protons = 15

No. of Neutrons = 31 - 15 = 16

Electronic Configuration = 2, 8, 5

Position in periodic table = 3^rd period (as it has 3 shells) & 15^th group (as it has five valence electrons).

(i) Neon (Ne)

(ii) Aluminium (Al)

(iii) Phosphorus (P)

(iv) Calcium (Ca)

(v) Carbon (C)

(i) Metal.

(ii) Barium is more reactive than calcium.

(iii) Valency = 2.

(iv) Barium Phosphate Ba₃(PO₄)₂.

(v) Barium is smaller in size then Caesium.

Z < Y < X.

Atomic size increases as we move down the group and also the atomic no. increases on moving down the group.

(i) Incorrect — A few groups contain metalloids and inert gases also.

(ii) Incorrect — as we move down the group the number of electrons increases.

(iii) Incorrect — Non-metallic character increases across a period with increase in atomic number.

(iv) Incorrect — Chemical reactivity in metals increases on going down the group, but in non-metals decreases on going down the group.

No. of elements in:

Period 1 = 2

Hydrogen (H) , Helium (He).

Period 2 = 8

Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon(Ne).

Period 3 = 8

Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl), Argon (Ar).

Common feature of the electronic configuration of the elements at the end of Period 2 and Period 3 are:

1. They have complete octet (8 electrons in valence shell).
2. They are stable so they don't react.

Non-metallic , metallic.

Sodium (Na).

(i) Atomic Number.

(ii) Period, Non-metallic.

(iii) More.

(iv) Number of outer electrons.

(i) Cl^- has one more electron then Cl atom. Anion is formed by the gain of electron(s). Thus, the number of electron(s) are more than proton(s). The effective positive charge in the nucleus is less, so less inward pull is experienced. Hence, the size increases.

(ii) Argon is an inert gas and its outermost shell is complete. They have maximum number of electrons in their outermost orbit. Thus, the electronic repulsions are maximum. The effect of nuclear pull over the valence electrons is not seen. Hence, the size of the argon is greater then chlorine.

(iii) The chemical reactivity of non-metals decreases on going down the group as it depends upon the tendency to gain electrons, which decreases down the group. As fluorine is present at the top of the group it is more reactive than chlorine

(iv) Inert gases do not form ion due to their stable electronic configuration, inert gases find it difficult to accept or lose electrons hence they do not form ion.

The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion is called Ionisation Potential (I.P.) or Ionisation Energy (I.E.).

M (g) + I.E. ⟶ M⁺ (g) + e^-

where, M is any element

Unit : electronvolts per atom (eV/atom)

SI Unit : kilojoule per mole (kJ/mole).

(a) The greater the atomic size the lesser the force of attraction. Since the electrons of the outermost shell lie further away from the nucleus, it makes their removal easier. Thus, Ionisation Potential decreases with increase in atomic size.

(b) The greater the nuclear charge, greater is the attraction for the electrons of the outermost shell. Therefore, the electrons in the outermost shell are more firmly held because of which greater energy is required to remove the electron(s). Thus, Ionisation Potential increases with increase in nuclear pull.

(a) On moving left to right across a period, the ionisation energy tends to increase as the atomic size decreases.

(b) On moving down a group the ionisation energy decrease as the atomic size increases.

Helium has the highest I.E. and Sodium has the lowest I.E. in the first three period.

Second Period : Li < Be < B < C < N < O < F < Ne.

Third Period : Na < Mg < Al < Si < P < S < Cl < Ar.

Helium

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

Electron affinity — The amount of energy released while converting a neutral gaseous isolated atom into a negatively charged gaseous ion (anion) by the addition of electron is called Electron Affinity (E.A.)

Unit : Electron volts per atom (eV/atom) or kJ mol^-1. It is represented by negative sign.

Second Period : Li < B < C < O < F .

Ne , N , Be do not follow the trend.

Inert gases have zero or very low electron affinity because their outer shells are completely filled, making it difficult for them to gain electrons. Neon is an inert gas, so its electron affinity is nearly zero. Therefore, fluorine has a much higher electron affinity than neon.

Increases , Decreases.

Electronegativity — The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is called its electronegativity.

Electronegativity is a dimensionless property since it is only a tendency. As such, it has no unit.

Lithium.

Reason — The given elements belong to period 2. Among these elements, Lithium is the left most and Fluorine is the right most. As electronegativity increases from left to right in a period, hence, Lithium is the least Electronegative.

Chlorine.

Reason — The given elements belong to period 3. Among these elements, Chlorine is the right most. As electronegativity increases from left to right in a period, hence, Chlorine is the most Electronegative.

(a) Metals lose electrons whereas non-metals gain electrons to attain stable electronic configuration. Group 1 elements are placed on the left in the periodic table whereas Group 17 elements are towards right. Hence, the atomic size of Group 1 elements is greater when compared to the corresponding Group 17 elements as atomic size decreases from left to right across a period due to increase in nuclear charge. Due to this Group 1 elements can easily lose electrons making them strong metals. It is difficult for Group 17 elements to lose electrons due to stronger nuclear pull but it helps in attracting electrons from other elements making them strong non-metals.

(b) On moving across a period from left to right, nuclear pull increases due to the increase in atomic number and thus the atomic size decreases. Hence, elements cannot lose electrons easily. Therefore, the metallic character decreases across a period.
On moving down the group, the atomic size increases and the nuclear charge also increases. The effect of an increased atomic size is greater as compared to the increased nuclear charge. Therefore, tendency to lose electron increases and elements can lose electrons easily. Thus, metallic character of elements increases in moving down a group.

(c) The atomic size of halogens is very small. The smaller the atomic size, the greater the electron affinity, because the effective attractive force between the nucleus and the valence electrons is greater in smaller atoms, and so the electrons are strongly held with the atom.

(d) Greater the tendency to lose electrons, stronger is the reducing power. On moving down the group, the atomic size and nuclear charge increases but the effect of increased atomic size is greater. Therefore, the tendency to lose electrons increases which increases the reducing power of elements on moving down in a group.
On moving across the period, nuclear pull increases due to increase in atomic number and thus the atomic size decreases. Hence, elements cannot lose electrons easily. Therefore, reducing power of elements decreases on moving across a period from left to right.

(e) Size of atoms progressively becomes smaller when we move from Sodium (Na) to Chlorine (Cl) in the third period of the periodic table because on moving from left to right in a period nuclear pull increases because of increase in atomic number and so the atomic size decreases.

(a) Ionisation Energy (I.E.)

(b) Metallic Character.

(c) Electronegativity.

(a) The most electronegative elements are G(9) & O(17).

(b) G(9) & O(17) are halogens.

(c) A(3) & I(11) are alkali metals.

(d) D(6) & L(14) are elements with valency 4.

(e) I(11) have least Ionisation Energy.

(f) O(17) have least atomic size in period 3.

(a) Thallium has the most metallic character as metallic character increases on moving down the group.

(b) Boron has the highest electronegativity because the electronegativity decrease on moving down the group.

(c) There are 3 electrons in the outer shell of thallium because all the elements in a group have same number of valence electrons.

(d) BCl₃ (Boron Chloride).
Boron needs 3 electrons while each chlorine atom can donate one electron. So 3 chlorine atoms donate 1 electron each to react with 1 Boron atom forming BCl₃ (Boron Chloride).

(e) The elements in the group to the right of this boron group are less metallic because on moving left to right in a period metallic character decreases as atomic size decreases the electrons are strongly bound with the atom so it becomes more difficult to lose an electron.