S-Block Elements And Their Compounds
Group 1 of the periodic table consists of the elements : lithium, sodium, potassium, rubidium, caesium and francium .
The elements of Group 2 include beryllium, magnesium, calcium, strontium, barium and radium.
Ionization Enthalpy
- The ionization enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs.
- This is due to the effect of increasing size outweighs the increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge
Hydration Enthalpy:
The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
$\mathrm{Li}^{+} > \mathrm{Na}^{+} > \mathrm{K}^{+} > \mathrm{Rb}^{+} > \mathrm{Cs}^{+}$
has maximum degree of hydration and for this reasons lithium salts are mostly hydrated e.g., $\mathrm{LiCl} .2 \mathrm{H}_{2} \mathrm{O}$
Physical properties:
All the alkali metal are silvery white, soft and light metals. Because of the larger size, these element have low density. The melting and boiling point of the alkali metals are low indicating weak metallic bonding alkali metals and their salts impart characteristic colour to an oxidizing flame.
Density : All are light metals. The densities are low. Lithium, sodium and potassium are lighter than water, for this very reason they float on water. Density gradually increases in moving down from Li to Cs. Potassium is, however, lighter than sodium.
| Li | Na | K | Rb | Cs | |
|---|---|---|---|---|---|
| Density (g/mL) | 0.54 | 0.97 | 0.86 | 1.53 | 1.87 |
The reason for the low values is that these metals have high atomic volumes. The abnormal value of potassium is due to unusual increase in atomic size, i.e., atomic volume. Lithium is the lightest among all the alkali metals. $$ \mathrm{Li}<\mathrm{K}<\mathrm{Na}<\mathrm{Rb}<\mathrm{Cs} $$ $$ \text { Density increases } \longrightarrow $$ [Note: The density of potassium is less than sodium contrary to expectation. This is probably due to abnormal increase in atomic size or atomic volume from ${Na}$ to ${K}$. $]$
Melting and boiling points: The energy binding the atoms in the crystal lattices of these metals is relatively low on account of a single electron in the valency shell. Consequently, the metals have low melting and boiling points. These decrease in moving down from $\mathrm{Li}$ to $\mathrm{Cs}$ as the metallic bond strength decreases or cohesive force decreases.
| $\mathrm{Li}$ | $\mathrm{Na}$ | $\mathrm{K}$ | $\mathrm{Rb}$ | $\mathrm{Cs}$ | |
|---|---|---|---|---|---|
| Melting point (K) | 454 | 371 | 336 | 312 | 302 |
| Boiling point $(\mathrm{K})$ | 1615 | 1156 | 1032 | 961 | 944 |
Chemical Properties:
The alkali metal are highly reactive due to their larger size and low ionization enthalpy.
Reactivity towards air :
They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxide.
Reducing nature:
- The alkali metals, are strong reducing agents, lithium being the most and sodium the least powerful.
- The alkali metals and their salts impart characteristic colour to an oxidizing flame
Solution in liquid ammonia:
The alkali metals dissolve in liquid ammonia giving deep blue solution which are conducting in nature.
$ M+(x+y)NH_3 \longrightarrow [M(NH_3)_x]^+ + [e(NH_3)_y]^-$
The blue colour of the solution is due to the ammoniated electron and the solutions is paramagnetic.
$ M^+ (aq) + e^- + NH_3 (l) \xrightarrow{ \text{on standing }} MNH_2 (aq) + 1/2 H_2(g) $
In concentrated solution, the blue colour changes to bronze colour and becomes, diamagnetic.
ANOMALOUS PROPERTIES OF LITHIUM
(i) exceptionally small size of its atom and ion, and
(ii) high polarizing power (i.e., charge/ radius ratio ).
The similarity between lithium and magnesium is particularly striking and arises because of their similar size: atomic radii, $\mathrm{Li}=152 \mathrm{pm}, \mathrm{Mg}=160$ pm; ionic radii : $\mathrm{Li}^{+}=76 \mathrm{pm}, \mathrm{Mg}^{2+}=72 \mathrm{pm}$.
GROUP 2 ELEMENTS : ALKALINE EARTH METALS
The first element beryllium differs from the rest of the member and shows diagonal relationship to aluminium.
Hydration Enthalpies
Hydration enthalpies of alkaline earth metal ions. $\mathrm{Be}^{2+}>\mathrm{Mg}^{2+}>\mathrm{Ca}^{2+}>$ $\mathrm{Sr}^{2+}>\mathrm{Ba}^{2+}$. The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., $ MgCl_2 $ and $ CaCl_2 $ exist as $ MgCl_2.6 H_2O $ and $ CaCl_2.6H_2O $ while $NaCl$ and $KCl$ do not form such hydrates.
Physical Properties
The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. The melting and boiling point of these metals are higher due to smaller sizes. Because of the low ionization enthalpies they are strongly electropositive in nature. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence these elements do not impart any colour to the flame.
Calcium, strontium and barium impart characteristic colour to the flame.
Difference between Lithium and other Alkali Metals
Lithium nitrate when heated gives lithium oxide, $Li_2O$, whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$4LiNO_3 \rightarrow 2Li_2O + 4NO_2 + O_2$
$2NaNO_3 \rightarrow 2NaNO_2 + O_2$
Chemical Properties
Reactivity towards air and water :
Beryllium and magnesium are inert to oxygen and water. Magnesium is more electropositive and burns with dazzling brilliance in air to give $\mathrm{MgO}$ and $ Mg_3 N_2 $. Calcium, strontium and barium are readily attacked by air to form the oxide and nitride.
Reducing nature :
The alkaline earth metals are strong reducing agent. This is indicated by large negative value of their reduction potentials.
Solution in liquid ammonia:
The alkaline earth metals dissolve in liquid ammonia to give deep blue black solution forming ammoniated ions.
$ M+(x+y)NH_3 \longrightarrow [M(NH_3)_x]^{+2} + 2[e(NH_3)_y]^-$
From these solutions, the ammoniates, $ [ M(NH_3)_6 ]^{+2} $can be recovered.
General characterstics of compounds of the alkaline earth metals
The alkaline earth metals form compounds which are predominantly ionic but less ionic than the corresponding compounds of alkali metals due to increased nuclear charge and smaller size.
Oxides and Hydroxides:
- Except for BeO the alkaline earth metals burn in oxygen to form the monoxide MO
- BeO is amphoteric while oxides of other elements are ionic in nature
- All these oxides except BeO are basic in nature and react with water to form sparingly soluble hydroxides.
$$ MO + H_2O \rightarrow M(OH)_2 $$
- Except $Be(OH)_2$ all other hydroxides are ionic in nature and their basic nature increase with increase in atomic number.
Halides:
- EXcept beryllium halide all other halides of alkaline earth metals are ionic in nature. Beryllium halide are essentially covalent and soluble in organic solvents.
- Beryllium chloride has a chain structure in the solid state
- In the vapour phase $BeCl_2$ form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K.
- The tendency to form halide hydrates gradually decreases
Salts of Oxoacids
The alkaline earth metals also form salts of oxoacids. Some of these are :
Carbonates:
Carbonates of alkaline earth metals are insoluble in water and can be precipitated by addition of a sodium or ammonium carbonate solution to a solution of a soluble salt of these metals.
- The solubility of carbonates in water decreases as the atomic number of the metal ion increases.
- All the carbonates decompose on heating to give carbon dioxide and the oxide.
- Beryllium carbonate is unstable and can be kept only in the atmosphere of $CO_2$.
- The thermal stability increases with increasing cationic size.
Sulphates:
- The sulphates of the alkaline earth metals are all white solids and stable to heat.
- $BeSO_4$, and $MgSO_4$ are readily soluble in water.
- The solubility decreases from $CaSO_4$ to $BaSO_4$.
- The greater hydration enthalpies of $Be^{2+}$ and $Mg^{2+}$ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.
ANOMALOUS BEHAVIOUR OF BERYLLIUM
Beryllium the first member of the Group 2 metals, shows anomalous behaviour as compared to magnesium and rest of the members. Further, it shows diagonal relationship to aluminium.
Diagonal Relationship between Beryllium and Aluminium
The ionic radius of $\mathrm{Be}^{2+}$ is estimated to be $31 \mathrm{pm}$; the charge/ radius ratio is nearly the same as that of the $\mathrm{Al}^{3+}$ ion. Hence beryllium resembles aluminium in some ways.
Compounds of s-block elements
1. Sodium Oxide $ ( Na_2O ) $
2. Sodium peroxide $ ( Na_2 O_2 ) $
3. Sodium Hydroxide (NaOH)
4. Sodium Carbonate $ ( Na_2 CO_3 ) $
5. Quick Lime, Slaked Lime and Lime Water
$ Ca(OH)_2+2 NH_4 Cl \xrightarrow{\Delta} 2 NH_3+CaCl_2+2 H_2 O $
Quick Lime Uses:
- It is an important primary material for manufacturing cement and is the cheapest form of alkali.
- It is used in the manufacture of sodium carbonate from caustic soda.
- It is employed in the purification of sugar and in the manufacture of dye stuffs
Slaked lime Uses:
- It is used in the preparation of mortar, a building material.
- It is used in white wash due to its disinfectant nature.
- It is used in glass making, in tanning industry, for the preparation of bleaching powder and for purification of sugar.
Gypsum Uses:
- Gypsum is used as fertilizer.
- Gypsum prevents soil erosion, improves soil composition, helps the movement of water and air, and facilitates root growth.
- Gypsum balances micronutrients like zinc, iron etc.
- It is used as a component of cement for controlling the force at which concrete sets in.
- Due to its low thermal conductivity, it is used in the manufacturing of drywall or wallboards.
- Gypsum is used for making fireproof wall boards.
Calcium Carbonate Uses:
- It is used as a building material in the form of marble and in the manufacture of quick lime.
- Calcium carbonate along with magnesium carbonate is used as a flux in the extraction of metals such as iron.
- It is also used as an antacid, mild abrasive in tooth paste, a constituent of chewing gum, and a filler in cosmetics.
Calcium Sulphate (Plaster of Paris) Uses:
- The largest use of Plaster of Paris is in the building industry as well as plasters.
- It is used for immoblising the affected part of organ where there is a bone fracture or sprain.
- It is also employed in dentistry, in ornamental work and for making casts of statues and busts.
Cement:
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It is an important building material and obtained by combining a material rich in lime, CaO with other material such as clay which contains silica, $SiO_2$ along with the oxides of aluminium, iron and magnesium.
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The average composition of Portland cement is : CaO, 50-60%; $SiO_2$, 20-25%; $Al_2O_3$, 5-10%; MgO, 2-3%; $Fe_2O_3$, 1-2% and $SO_3$, 1-2%.
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For a good quality cement, the ratio of silica $(SiO_2)$ to alumina $(Al_2O_3)$ should be between 2.5 and 4 and the ratio of lime $(CaO)$ to the total of the oxides of silicon $(SiO_2)$ aluminium $(Al_2O_3)$ and iron $(Fe_2O_3)$ should be as close as possible to 2.
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The raw materials for the manufacture of cement are limestone and clay. When clay and lime are strongly heated together they fuse and react to form ‘cement clinker’. This clinker is mixed with 2-3% by weight of gypsum $(CaSO_4·2H_2O)$ to form cement. Thus important ingredients present in Portland cement are dicalcium silicate $(Ca_2SiO_4)$ 26%, tricalcium silicate $(Ca_3SiO_5$) 51% and tricalcium aluminate $(Ca_3Al_2O_6)$ 11%.
Setting of Cement:
When mixed with water, the setting of cement takes place to give a hard mass. This is due to the hydration of the molecules of the constituents and their rearrangement.
- Gypsum is added only to slow down the process of setting of the cement so that it gets sufficiently hardened. Uses: It is used in concrete and reinforced concrete, in plastering and in the construction of bridges,dams and buildings.
BIOLOGICAL IMPORTANCE OF MAGNESIUM AND CALCIUM
- All enzymes that utilise ATP in phosphate transfer require magnesium as the cofactor.
- The main pigment for the absorption of light in plants is chlorophyll which contains magnesium.
- About 99 % of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation.
- The calcium concentration in plasma is regulated at about 100 $mgL^{–1}$. It is maintained by two hormones: calcitonin and parathyroid hormone.
