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Block Elements ….. class 11 chemistry docx, Essays (high school) of Chemistry

Chemistry

Class 11 chemistry block elements pdf

Typology: Essays (high school)

2017/2018

Uploaded on 10/30/2023

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Class – 11 Unit – 10 S-block elements
• General Electronic Configuration of s-Block Elements
For alkali metals [noble gas] ns1
For alkaline earth metals [noble gas] ns2
• Group 1 Elements: Alkali metals
Electronic Configuration, ns1, where n represents the valence shell.
These elements are called alkali metals because they readily dissolve in water to form
soluble hydroxides, which are strongly alkaline in nature.
• Atomic and Ionic Radii
Atomic and ionic radii of alkali metals increase on moving down the group i.e., they
increase in size going from Li to Cs. Alkali metals form monovalent cations by losing
one valence electron. Thus cationic radius is less as compared to the parent atom.
• Ionization Enthalpy
The ionization enthalpies of the alkali metals are generally low and decrease down
the group from Li to Cs.
Reason: Since alkali metals possess large atomic sizes as a result of which the
valence s-electron (ns1) can be easily removed. These values decrease down the
group because of decrease in the magnitude of the force of attraction with the
nucleus on account of increased atomic radii and screening effect.
• Hydration Enthalpy
Smaller the size of the ion, more is its tendency to get hydrated hence more is the
hydration enthalpy.
Hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
Li+.> Na+.> K+.> Rb+.> Cs+
• Physical Properties
(i) All the alkali metals are silvery white, soft and light metals.
(ii) They have generally low density which increases down the group.
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Class – 11 Unit – 10 S-block elements

- General Electronic Configuration of s-Block Elements For alkali metals [noble gas] ns^1 For alkaline earth metals [noble gas] ns^2 - Group 1 Elements: Alkali metals Electronic Configuration, ns^1 , where n represents the valence shell. These elements are called alkali metals because they readily dissolve in water to form soluble hydroxides, which are strongly alkaline in nature.

• Atomic and Ionic Radii

Atomic and ionic radii of alkali metals increase on moving down the group i.e., they

increase in size going from Li to Cs. Alkali metals form monovalent cations by losing

one valence electron. Thus cationic radius is less as compared to the parent atom.

• Ionization Enthalpy

The ionization enthalpies of the alkali metals are generally low and decrease down

the group from Li to Cs.

Reason: Since alkali metals possess large atomic sizes as a result of which the

valence s-electron (ns^1 ) can be easily removed. These values decrease down the

group because of decrease in the magnitude of the force of attraction with the

nucleus on account of increased atomic radii and screening effect.

• Hydration Enthalpy

Smaller the size of the ion, more is its tendency to get hydrated hence more is the

hydration enthalpy.

Hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.

Li+^ > Na+^ > K+^ > Rb+^ > Cs+

• Physical Properties

(i) All the alkali metals are silvery white, soft and light metals.

(ii) They have generally low density which increases down the group.

(iii) They impart colour to an oxidising flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region.

- Chemical Properties of Alkali Metals (i) Reaction with air: When exposed to air surface of the alkali metals get tarnished due the formation of oxides and hydroxides. Alkali metals combine with oxygen upon heating to form different oxides depending upon their nature. (ii) Reaction with water: Alkali metals react with water to form hydroxide and dihydrogen (iii) Reaction with hydrogen: The alkali metals combine with hydrogen at about 673 K (lithium at 1073 K) to form hydrides. 2M + H 2 ————-> 2M+ The ionic character of hydrides increases from Li to Cs. (iv) Reaction with halogens: Alkali metals combine with halogens directly to form metal halides. 2M + X 2 ————–> 2MX They have high melting and boiling points. Order of reactivity of M: (v) Reducing nature: The alkali metals are strong reducing agents. In aqueous solution it has been observed that the reducing character of alkali metals follows the sequence Na < K < Rb < Cs < Li, Li is the strongest while sodium is least powerful reducing agent. This can be explained in terms of electrode potentials (E°). Since the electrode potential of Li is the lowest. Thus Li is the strongest reducing agent.

(vi) Solubility in liquid ammonia:

The alkali metals dissolve in liquid ammonia to give deep blue solution. The solution

is conducting in nature.

M+ (x + y) NH 3 ———-> [M (NH 3 ) X]+^ + [e (NH3) y]–

When light falls on the ammoniated electrons, they absorb energy corresponding to red colour and the light which emits from it has blue colour. In concentrated solution colour changes from blue to bronze. The blue solutions are paramagnetic while the concentrated solutions are diamagnetic.

  • Uses of Alkali Metals Uses of Lithium

Due to their small size in comparison to alkali metals first ionization enthalpies of alkaline earth metals is higher than that of alkali metals.

- Hydration Enthalpies The hydration enthalpies of alkaline earth metal ions are larger than those of the alkali metals. Thus alkaline earth metals have more tendency to become hydrate. The hydration enthalpies decreases down the group since the cationic size increases. Be2+^ > Mg2+^ > Ca2+^ > Sr2+^ > Ba2+ Metallic character: They have strong metallic bonds as compared to the alkali metals in the same period. This is due to the smaller kernel size of alkaline earth metal and two valence electrons present in the outermost shell. **- Chemical Properties

  1. Reaction with oxygen.** Beryllium and magnesium are kinetically inert to oxygen because of the formation of a thin film of oxide on their surface. Reactivity towards oxygen increases as going down the group. 2. Reaction with water. Since these metals are less electropositive than alkali metals, they are less reactive towards water. Magnesium reacts with boiling water or steam. Rest of the members reacts even with cold water. Mg + 2H 2 0 ——-> Mg(OH) 2 + H 2 Ca + 2H 2 0 ————> Ca(OH) 2 + H 2 3. Reaction with halogens. They combine with the halogens at appropriate temperature to form corresponding halides MX 2. M + X 2 ——–> MX 2 (X = F, Cl, Br, I) Thermal decomposition of (NH 4 ) 2 BeF 4 is used for the preparation of BeF 2. 4. Reaction with hydrogen. These metals except Be combine with hydrogen directly upon heating to form metal hydrides. - General Characteristics of Compounds of Alkaline Earth Metals Oxides and Hydroxides (i) The alkaline earth metals bum in oxygen to form MO (monoxide). (ii) These oxides are very stable to heat. (iii) BeO is amphoteric in nature while oxides of other elements are ionic. (iv) Except BeO, they are basic in nature and react with water to form sparingly soluble hydroxides. MO + H 2 O ———-> M(OH) 2 (v) Hydroxides of alkaline earth metals are less stable and less basic than alkali metal hydroxides. (vi) Beryllium hydroxide is amphoteric in nature. Halides The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX 2. They can also be prepared by the action of halogen acids (HX) on metals, metal oxides, metal hydroxides. M + 2HX ——-> MX 2 + H 2 MO + 2HX ——> MX 2 + H 20 M (OH) 2 + 2HX —–> MX 2 + 2H 20 (i) Except beryllium halides, all other halides of alkaline earth metals are ionic in nature. (ii) Except BeCl 2 and MgCl 2 other chloride of alkaline earth metals impart characteristic colours to flame.

(iii) The tendency to form halide hydrates decreases down the group.

For example, (MgCl 2 – 8 H 2 0, CaCl 2 – 6 H 2 0, SrCl 2 – 6 H 2 0, BaCl 2 – 2 H 2 O)

(iv) BeCl 2 has a chain structure in the solid phase as shown below.

In vapour phase the compound exist as a dimer which decomposes at about 1000K to give monomer in which Be atom is in sp hybridisation state. Sulphates (i) The sulphates of alkaline earth metals are white solids and quite stable to heat. (ii) BeS0 4 and MgS0 4 are readily soluble in water. Solubility decreases from BeS0 4 to BaS0 4. Reason. Due to greater hydration enthalpies of Be2+^ ions and Mg2+^ ions they overcome the lattice enthalpy factor. Their sulphates are soluble in water. Carbonates Carbonates of alkaline earth metals are thermally unstable and decompose on heating. Uses: (i) In the manufacture of cement, sodium carbonate, calcium carbide etc. (ii) Used in the purification of sugar. (iii) In the manufacture of dye stuffs. iii) Calcium Carbonate or Limestone (CaC0 3 ) Preparation: Calcium carbonate occurs in nature in different forms like limestone, marble, chalk etc. It can be prepared by passing C0 2 through slaked lime in limited amount. Ca(OH) 2 + C0 2 ———> CaC0 3 + H 20 It can also prepared by the reaction of a solution of sodium carbonate with calcium chloride. CaCl 2 + Na 2 C0 3 ————> CaC0 3 + 2NaCl Uses: (i) In the manufacturing of Quick Lime. (ii) With MgC0 3 used as flux in the extraction of metals. (iii) Used as an antacid. (iv) In the manufacture of high quality paper. (iv) Calcium Sulphate (Plaster of Paris) CaS0 4 -1/2H 20 Preparation: It is obtained when gypsum CaS0 4 – 2 H 2 0 is heated to 393 K