Solutions-Class-11-Science-Chemistry-Chapter-7-Modern Periodic Table-Maharashtra Board

• Li, B, Be and N belong to the same 2^nd
• As we move across a period from left to right in the periodic table, the effective nuclear charge increases steadily, and therefore, electronegativity increases.
• Hence, the elements Li, B, Be and N have the electronegativities 1.0, 2.0, 1.5, and 3.0, respectively on the Pauling scale.

• Cl, I and Br belong to group 17 (halogen group) in the periodic table.
• As we move down the group from top to bottom in the periodic table, a new shell gets added to the atom of the elements.
• As a result, the effective nuclear charge decreases due to an increase in the atomic size as well as an increased shielding effect.
• Therefore, the valence electrons experience less attractive force from the nucleus and are held less tightly resulting in the increased atomic radius.
• Thus, their atomic radii increase in the following order down the group.

Cl (99 pm) < Br (114 pm) < I (133 pm)

Hence, the atomic radii of Cl, I, and Br are 99, 133, and 114 pm, respectively.

• Both, Na⁺ and F^- are isoelectronic having 10 electrons each.
• Size of isoelectronic species decreases as nuclear charge increases.
• Na⁺ with nuclear charge (+11) which is greater than F^- (+9).
• Hence, Na⁺ has lower ionic radius (98 pm) than that of F^- (133 pm.)

• Al, Si and P belong to Third period. Aluminium has tendency to lose electrons and form cations, it shows all metallic properties hence, aluminium is a metal.
• Phosphorus has tendency to gain electrons and form anions, it shows nonmetallic properties, hence phosphorus is a nonmetal.
• Silicon has the properties of both metals and nonmetals, hence it is a metalloid.

Cu : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

Zn : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰

Since, Zn has completely filled s and d orbitals, it does not form coloured ions. Cu has incomplete 4s subshell hence it forms coloured ions.

Ge is in period 4 and group 14. Therefore, its outermost electronic configuration is 4s² 4p². In the fourth period 3d subshell is filled before 4p.

Hence, the electronic configuration is

₃₂Ge : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p².

Po is in period 6 and group 16. Therefore, its outermost electronic configuration is 6s² 6p⁴.

In sixth period, 4f and 5d subshells are filled up before 6p. Hence, electronic configuration of Po is

₈₄Po : 1s² 2s²2p⁶ 3s²3p⁶ 4s²3d¹⁰4p⁶ 5s²4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁴.

Cu is in Period 4 and group 11. Therefore, its outermost electronic configuration is 4s².

In fourth period, 3d subshell is filled up and outer electronic configuration should be 3d⁹ 4s².

Half-filled or completely filled subshell is more stable. Hence, the electronic configuration of Cu is

₂₉Cu : 1s² 2s²2p⁶ 3s²3p⁶ 3d¹⁰ 4s¹

• Li has one electron in valence shell while F has 7 electrons. Therefore, F has a tendency to accept electron to complete its octet.
• Li has larger atomic radius than F. Therefore, Li has less ionization enthalpy than F.

Screening effect : The inner shell electrons in an atom screen or shield the outermost valence electrons from the nuclear attraction. This effect is called screening effect or shielding effect.

The magnitude of screening effect depends upon the number of inner electrons. Higher the number of inner electrons, greater is the value of screening effect. The screening constant is represented by 'o' (sigma).

Example :

Across a period, the screening effect due to inner electrons remains the same as electrons are added to the same shell. Down the group, the screening effect due to inner electrons increases as a new valence shell is added.

e.g. Potassium (₁₉K) has electronic configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

K has 4 shells and thus, the valence shell electrons are effectively shielded by the electrons present in the inner three shells. As a result of this, valence shell electron (4s¹) in K experiences a much less effective nuclear charge and can be easily removed.

First electron is removed from a gaseous neutral atom while second electron is removed from positively charged gaseous cation. Hence the second ionization enthalpy (IE₂) is higher than the first ionization enthalpy (IE₁).

In general, if IE₁, IE₂ and IE₃ are the first, second and third ionization enthalpies respectively then,

IE₃ > IE₂ > IE₁.

The elements belonging to the same group have similar electronic configuration of the valence shell. These valence electrons are involved during a chemical reaction. Hence, they have similar chemical properties.

• Electronegativity is the ability of an atom in a compound to attract shared electrons of a covalent bond to itself. Electron gain enthalpy is the change in the enthalpy accompanying the addition of an electron to gaseous neutral atom forming a gaseous anion.
• Electronegativity of any given element is not constant, but depends upon the other element to which is bound in a compound. Electron gain enthalpy of an element is constant.
• Electronegativity cannot be measured experimentally. It is obtained from numerical scale of electronegativity developed mainly by Linus Pauling. Electron gain enthalpy is a measurable

(d) K > Mg > Al > B

(c) principal quantum number

(d) bottom

(b) halogens

(a) 17

(b) 3, 12

(b) Mg²⁺ and Ne

(a) 13, 31

(a) 1s² : Period first and group 18.

(b) 1s² 2s² 2p⁶ : Period 2nd and group 18.

(a) Fe²⁺

(b) Ca²⁺

(a) Li⁺ (b) F^-

(a) Sulphur should have smaller ionization energy then oxygen.

(b) Lithium (2, 1) has only one valence electron, hence forms +1 ion by the loss of 1 electron. Berylium (2, 2) has two valence electrons and form +2 ions by the loss of 2 electrons.

Ionic radius is the distance between neighbouring cations and anions in ionic crystals.

The ability of a covalently bonded atom to attract the shared electrons towards itself is called electronegativity. (EN).

• Metals react with oxygen to form corresponding oxides which are basic in nature. The oxides of reactive metals when dissolved in water give base. For example, sodium form sodium oxide (Na₂O) which form a base (NaOH) with water.
• The nonmetals react with O₂ to form acidic oxides. The oxides of nonmetals form acids when dissolved in water. For example, oxide of reactive nonmetals like chlorine form Cl₂O₇ which forms strong acid (HClO₄) with water.
• The oxides of the elements in the center of the main group elements form amphoteric oxides of metals (Al₂O₃) or neutral or weakly acidic oxides of nonmetal. For example, CO and NO are neutral whereas CO₂ is weakly acidic.

Valence electrons are the electrons present in the outermost shell of an atom.

Ionization enthalpy : The amount of energy required to remove an electron from the isolated gaseous atom in its ground state is called ionization enthalpy.

X_(g) → X_(g)⁺ + e^- Δ_iH  : [Ionization enthalpy is always positive as energy has to be supplied to remove outer electron from an atom.]

Factors on which ionisation enthalpy depends are :

(a) Radius of an atom

(b) Nuclear charge

(c) The screening effect of inner electrons

(d) Nature of electronic configuration.

Ionization enthalpy in the periods :

• The first ionization enthalpy for the elements increases from left to right.
• Across the period as the atomic number increases, the electrons are added successively to the same principal shell.
• The nuclear charge increases from left to right in the periods.
• Hence the added electrons poorly shield each other from the nucleus.
• As a result effective nuclear charge increases and causes the decrease in the atomic radii.

Hence, along a period ionization enthalpy increases. In each period, the last inert element has the highest ionization enthalpy, and alkali metals have the lowest value of first ionization enthalpy.

Ionization enthalpy in a group :

• Along a group from top to bottom, first ionization enthalpy decreases.
• As the atomic number increases down the group, the valence electron enters a new principal shell, increasing the atomic radius.
• As the number of inner electrons increases the screening or shielding effect also increases.
• Hence, the force of nuclear attraction on valence electrons decreases and it becomes easier to remove the outermost electrons.
• Therefore, along the group from top to bottom ionization enthalpy decreases, with the increase in atomic number.

(1) In a period : Atomic radius decreases in a period from left to right (upto group 17). This is because as we move across a period, atomic number increases from left to right, nuclear charge increases but electrons enter the same shell. Thus, the screening effect of the inner core remains the same, on the other hand the effective nuclear charge goes on increasing. As a result the valence electrons are tightly held by the nucleus and atomic size (radius) decreases.

For example, in 2nd Period, it decreases from Li(152 pm) to F(64 pm).

(2) In a group : Down the group, as the atomic number increase, the effective nuclear charge increases and shielding effect increases down a group. The valence electrons are held by weaker attrctive forces, thus, atomic radius increases.

For example, atomic radius increases. For group 17 elements, the value increase from 1st member F (64 pm) to the last member At (140 pm).

Alkali metals have large volume and only one electron in the outermost shell (ns¹). Therefore, less energy is required to remove the single unpaired electron from the valence shell. Hence alkali metals have low ionization energy.

• Across a period, the screening effect is the same, and the effective nuclear charge increases.
• As a result, the outer electron is held more tightly, and hence, the ionization enthalpy increases across a period.
• Inert gases are present on the extreme right of the periodic table i.e., in group 18. Also, inert gases have stable electronic configurations i.e., complete octet or duplet.
• Due to this, they are extremely stable and it is very difficult to remove electrons from their valence shell.

Hence, inert gases have exceptionally high ionization enthalpies.

• Fluorine (F) and chlorine (Cl) belong to the same group (17).
• Down the group, electron gain enthalpy becomes less negative hence F should have higher negative electron gain enthalpy than Cl. But actually it is reverse.
• This is due to the smaller atomic size and higher electron density in F than in Cl.
• Hence, adding an electron to the 2p-orbital in F leads to a greater repulsion than adding an electron to the larger 3p-orbital in Cl.
• Therefore F has less negative electron gain enthalpy than Cl.

F_(g) + e^- → F^- -328 kJ mol^-1

Cl(g) + e^- → Cl^- -349 kJ mol^-1

Noble gases (last element in each period) are inert elements, as their outermost shell is completely filled and the octet is completed. Thus, noble gases (group 18) have larger atomic radii than preceding halogens (group 17).

(a) Li₂O is the most basic oxide.

(b) CO₂ would be the most acidic oxide B₂O₃ is an acidic oxide.

(c) Al₂O₃ is an amphoteric oxide.