THE s-BLOCK ELEMENTS
UNIT 10
THE s-BLOCK ELEMENTS
“The first element of alkali and alkaline earth metals differs in many respects from the other members of the group”
The s-block elements of the Periodic Table are those in which the last electron enters the outermost s-orbital. As the s-orbital can accommodate only two electrons, two groups ( 1 & 2) belong to the s-block of the Periodic Table. Group 1 of the Periodic Table consists of the elements: lithium, sodium, potassium, rubidium, caesium and francium. They are collectively known as the alkali metals. These are so called because they form hydroxides on reaction with water which are strongly alkaline in nature. The elements of Group 2 include beryllium, magnesium, calcium, strontium, barium and radium. These elements with the exception of beryllium are commonly known as the alkaline earth metals. These are so called because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust[^0].
Among the alkali metals sodium and potassium are abundant and lithium, rubidium and caesium have much lower abundances (Table 10.1). Francium is highly radioactive; its longest-lived isotope ${ }^{223} \mathrm{Fr}$ has a half-life of only 21 minutes. Of the alkaline earth metals calcium and magnesium rank fifth and sixth in abundance respectively in the earth’s crust. Strontium and barium have much lower abundances. Beryllium is rare and radium is the rarest of all comprising only $10^{-10}$ per cent of igneous rocks $^{\dagger}$ (Table 10.2, page 299).
The general electronic configuration of s-block elements is [noble gas] $n s^{1}$ for alkali metals and [noble gas] $n s^{2}$ for alkaline earth metals.[^1] from magma (molten rock) that has cooled and hardened.
Lithium and beryllium, the first elements of Group 1 and Group 2 respectively exhibit some properties which are different from those of the other members of the respective group. In these anomalous properties they resemble the second element of the following group. Thus, lithium shows similarities to magnesium and beryllium to aluminium in many of their properties. This type of diagonal similarity is commonly referred to as diagonal relationship in the periodic table. The diagonal relationship is due to the similarity in ionic sizes and /or charge/radius ratio of the elements. Monovalent sodium and potassium ions and divalent magnesium and calcium ions are found in large proportions in biological fluids. These ions perform important biological functions such as maintenance of ion balance and nerve impulse conduction.
10.1 GROUP 1 ELEMENTS: ALKALI METALS
The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number. The atomic, physical and chemical properties of alkali metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron, $n s^{1}$ (Table 10.1) outside the noble gas core. The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. They readily lose electron to give monovalent $\mathrm{M}^{+}$ions. Hence they are never found in free state in nature.
| Element | Symbol | Electronic configuration |
|---|---|---|
| Lithium | $\mathrm{Li}$ | $1 s^{2} 2 s^{1}$ |
| Sodium | $\mathrm{Na}$ | $1 \mathrm{~s}^{2} 2 s^{2} 2 p^{6} 3 s^{1}$ |
| Potassium | $\mathrm{K}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{1}$ |
| Rubidium | $\mathrm{Rb}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 5 s^{1}$ |
| Caesium | $\mathrm{Cs}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2}$ $4 p^{6} 4 d^{10} 5 s^{2} 5 p^{6} 6 s^{1}$ or $[\mathrm{Xe}] 6 s^{1}$ |
| Francium | $\mathrm{Fr}$ | $[\mathrm{Rn}] 7 s^{1}$ |
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes in a particular period of the periodic table. With increase in atomic number, the atom becomes larger. The monovalent ions $\left(\mathrm{M}^{+}\right)$are smaller than the parent atom. The atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size while going from $\mathrm{Li}$ to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals are considerably low and decrease down the group from $\mathrm{Li}$ to Cs. This is because the effect of increasing size outweighs the increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
$\mathrm{Li}^{+}>\mathrm{Na}^{+}>\mathrm{K}^{+}>\mathrm{Rb}^{+}>\mathrm{Cs}^{+}$
$\mathrm{Li}^{+}$has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., $\mathrm{LiCl} \cdot 2 \mathrm{H_2} \mathrm{O}$
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and light metals. Because of the large size, these elements have low density which increases down the group from Li to Cs. However, potassium is lighter than sodium. The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them. The alkali metals and their salts impart characteristic colour to an oxidizing 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 as given below:
| Metal | Li | $\mathbf{N a}$ | $\mathbf{K}$ | $\mathbf{R b}$ | $\mathbf{C s}$ |
|---|---|---|---|---|---|
| Colour | Crimson red |
Yellow | Violet | Red violet |
Blue |
| $\lambda / \mathrm{nm}$ | 670.8 | 589.2 | 766.5 | 780.0 | 455.5 |
Alkali metals can therefore, be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy. These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron.
Table 10.1 Atomic and Physical Properties of the Alkali Metals
| Property | Lithium Li |
Sodium $\mathbf{N a}$ |
Potassium $\mathbf{K}$ |
Rubidium Rb |
Caesium Cs |
Francium Fr |
|---|---|---|---|---|---|---|
| Atomic number | 3 | 11 | 19 | 37 | 55 | 87 |
| Atomic mass $\left(\mathrm{g} \mathrm{mol}^{-1}\right)$ | 6.94 | 22.99 | 39.10 | 85.47 | 132.91 | $(223)$ |
| Electronic configuration |
$[\mathrm{He}] 2 s^{1}$ | $[\mathrm{Ne}] 3 \mathrm{~s}^{1}$ | $[\mathrm{Ar}] 4 \mathrm{~s}^{1}$ | $[\mathrm{Kr}] 5 \mathrm{~s}^{1}$ | $[\mathrm{Xe}] 6 s^{1}$ | $[\mathrm{Rn}] 7 \mathrm{~s}^{1}$ |
| Ionization enthalpy $/ \mathrm{kJ} \mathrm{mol}^{-1}$ |
520 | 496 | 419 | 403 | 376 | $\sim 375$ |
| Hydration enthalpy $/ \mathrm{kJ} \mathrm{mol}^{-1}$ |
-506 | -406 | -330 | -310 | -276 | - |
| Metallic radius / pm |
152 | 186 | 227 | 248 | 265 | - |
| Ionic radius $\mathrm{M}^{+} / \mathrm{pm}$ |
76 | 102 | 138 | 152 | 167 | $(180)$ |
| m.p. / K | 454 | 371 | 336 | 312 | 302 | - |
| b.p / K | 1615 | 1156 | 1032 | 961 | 944 | - |
| Density $/ \mathrm{g} \mathrm{cm}^{-3}$ | 0.53 | 0.97 | 0.86 | 1.53 | 1.90 | - |
| Standard potentials $\mathrm{E}^{\ominus} / \mathrm{V}$ for $\left(\mathrm{M}^{+} / \mathrm{M}\right)$ |
-3.04 | -2.714 | -2.925 | -2.930 | -2.927 | - |
| Occurrence in lithosphere $^{\dagger}$ |
$18^{*}$ | $2.27^{* *}$ | $1.84^{* *}$ | $78-12^{*}$ | $2-6^{*}$ | $\sim 10^{-18 *}$ |
*ppm (part per million), ** percentage by weight; $\dagger$ Lithosphere: The Earth’s outer layer: its crust and part of the upper mantle
This property makes caesium and potassium useful as electrodes in photoelectric cells.
10.1.6 Chemical Properties
The alkali metals are highly reactive due to their large size and low ionization enthalpy. The reactivity of these metals increases down the group.
(i) Reactivity towards air: The alkali metals tarnish in dry air due to the formation of their oxides which in turn react with moisture to form hydroxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide $\mathrm{O_2}^{-}$ion is stable only in the presence of large cations such as $\mathrm{K}, \mathrm{Rb}$, $\mathrm{Cs}$.
$$ 4 \mathrm{Li}+\mathrm{O_2} \rightarrow 2 \mathrm{Li_2} \mathrm{O} \text { (oxide) } $$
$$ \begin{aligned} & 2 \mathrm{Na}+\mathrm{O_2} \rightarrow \mathrm{Na_2} \mathrm{O_2} \text { (peroxide) } \\ & \mathrm{M}+\mathrm{O_2} \rightarrow \mathrm{MO_2} \text { (superoxide) } \\ & (\mathrm{M}=\mathrm{K}, \mathrm{Rb}, \mathrm{Cs}) \end{aligned} $$
In all these oxides the oxidation state of the alkali metal is +1 . Lithium shows exceptional behaviour in reacting directly with nitrogen of air to form the nitride, $\mathrm{Li_3} \mathrm{~N}$ as well. Because of their high reactivity towards air and water, alkali metals are normally kept in kerosene oil.
Problem 10.1
What is the oxidation state of $\mathrm{K}$ in $\mathrm{KO_2}$ ?
Solution
The superoxide species is represented as $\mathrm{O_2}^{-}$; since the compound is neutral, therefore, the oxidation state of potassium is +1 . (ii) Reactivity towards water: The alkali metals react with water to form hydroxide and dihydrogen.
$$ \begin{array}{r} 2 \mathrm{M}+2 \mathrm{H_2} \mathrm{O} \rightarrow 2 \mathrm{M}^{+}+2 \mathrm{OH}^{-}+\mathrm{H_2} \\ (\mathrm{M}=\text { an alkali metal }) \end{array} $$
It may be noted that although lithium has most negative $\mathrm{E}^{\ominus}$ value (Table 10.1), its reaction with water is less vigorous than that of sodium which has the least negative $\mathrm{E}^{\ominus}$ value among the alkali metals. This behaviour of lithium is attributed to its small size and very high hydration energy. Other metals of the group react explosively with water.
They also react with proton donors such as alcohol, gaseous ammonia and alkynes.
(iii) Reactivity towards dihydrogen: The alkali metals react with dihydrogen at about 673K (lithium at 1073K) to form hydrides. All the alkali metal hydrides are ionic solids with high melting points.
$2 \mathrm{M}+\mathrm{H_2} \rightarrow 2 \mathrm{M}^{+} \mathrm{H}^{-}$
(iv) Reactivity towards halogens : The alkali metals readily react vigorously with halogens to form ionic halides, $\mathrm{M}^{+} \mathrm{X}^{-}$. However, lithium halides are somewhat covalent. It is because of the high polarisation capability of lithium ion (The distortion of electron cloud of the anion by the cation is called polarisation). The $\mathrm{Li}^{+}$ion is very small in size and has high tendency to distort electron cloud around the negative halide ion. Since anion with large size can be easily distorted, among halides, lithium iodide is the most covalent in nature.
(v) Reducing nature: The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful (Table 10.1). The standard electrode potential $\left(\mathrm{E}^{\ominus}\right)$ which measures the reducing power represents the overall change :
$\mathrm{M}(\mathrm{s}) \rightarrow \mathrm{M}(\mathrm{g}) \quad$ sublimationenthalpy
$\mathrm{M}(\mathrm{g}) \rightarrow \mathrm{M}^{+}(\mathrm{g})+\mathrm{e}^{-} \quad$ ionization enthalpy
$\mathrm{M}^{+}(\mathrm{g})+\mathrm{H_2} \mathrm{O} \rightarrow \mathrm{M}^{+}$(aq) hydrationenthalpy
With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative $\mathrm{E}^{\ominus}$ value and its high reducing power.
Problem 10.2
The $\mathrm{E}^{\ominus}$ for $\mathrm{Cl_2} / \mathrm{Cl}^{-}$is +1.36 , for $\mathrm{I_2} / \mathrm{I}^{-}$is +0.53 , for $\mathrm{Ag}^{+} / \mathrm{Ag}$ is $+0.79, \mathrm{Na}^{+} / \mathrm{Na}$ is -2.71 and for $\mathrm{Li}^{+} / \mathrm{Li}$ is -3.04 . Arrange the following ionic species in decreasing order of reducing strength:
$\mathrm{I}^{-}, \mathrm{Ag}, \mathrm{Cl}^{-}, \mathrm{Li}, \mathrm{Na}$
Solution
The order is $\mathrm{Li}>\mathrm{Na}>\mathrm{I}^{-}>\mathrm{Ag}>\mathrm{Cl}^{-}$
(vi) Solutions in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature.
$\mathrm{M}+(\mathrm{x}+\mathrm{y}) \mathrm{NH_3} \rightarrow \left[\mathrm{M} \left(\mathrm{NH_3} \right)_x \right]^{+}+ \left[\mathrm{e} \left(\mathrm{NH_3} \right)_y \right]^{-}$ The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution. The solutions are paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide.
$\mathrm{M}^{+}{ _(\mathrm{am})}+\mathrm{e}^{-}+\mathrm{NH_3}(\mathrm{l}) \rightarrow \mathrm{MNH_2(\mathrm{am})}+1 / 2 \mathrm{H_2}(\mathrm{~g})$
(where ‘am’ denotes solution in ammonia.) In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.
10.1.7 Uses
Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ bearings for motor engines, with aluminium to make aircraft parts, and with magnesium to make armour plates. It is used in thermonuclear reactions. Lithium is also used to make electrochemical cells. Sodium is used to make a $\mathrm{Na} / \mathrm{Pb}$ alloy needed to make $\mathrm{PbEt_4}$ and $\mathrm{PbMe_4}$. These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Potassium has a vital role in biological systems. Potassium chloride is used as a fertilizer. Potassium hydroxide is used in the manufacture of soft soap. It is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising photoelectric cells.
10.2 GENERAL CHARACTERISTICS OF THE COMPOUNDS OF THE ALKALI METALS
All the common compounds of the alkali metals are generally ionic in nature. General characteristics of some of their compounds are discussed here.
10.2.1 Oxides and Hydroxides
On combustion in excess of air, lithium forms mainly the oxide, $\mathrm{Li_2} \mathrm{O}$ (plus some peroxide $\mathrm{Li_2} \mathrm{O_2}$ ), sodium forms the peroxide, $\mathrm{Na_2} \mathrm{O_2}$ (and some superoxide $\mathrm{NaO_2}$ ) whilst potassium, rubidium and caesium form the superoxides, $\mathrm{MO_2}$. Under appropriate conditions pure compounds $\mathrm{M_2} \mathrm{O}, \mathrm{M_2} \mathrm{O_2}$ and $\mathrm{MO_2}$ may be prepared. The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice energy effects. These oxides are easily hydrolysed by water to form the hydroxides according to the following reactions :
$$ \begin{aligned} & \mathrm{M_2} \mathrm{O}+\mathrm{H_2} \mathrm{O} \rightarrow 2 \mathrm{M}^{+}+2 \mathrm{OH}^{-} \\ & \mathrm{M_2} \mathrm{O_2}+2 \mathrm{H_2} \mathrm{O} \rightarrow 2 \mathrm{M}^{+}+2 \mathrm{OH}^{-}+\mathrm{H_2} \mathrm{O_2} \\ & 2 \mathrm{MO_2}+2 \mathrm{H_2} \mathrm{O} \rightarrow 2 \mathrm{M}^{+}+2 \mathrm{OH}^{-}+\mathrm{H_2} \mathrm{O_2}+\mathrm{O_2} \end{aligned} $$
The oxides and the peroxides are colourless when pure, but the superoxides are yellow or orange in colour. The superoxides are also paramagnetic. Sodium peroxide is widely used as an oxidising agent in inorganic chemistry.
Problem 10.3
Why is $\mathrm{KO_2}$ paramagnetic?
Solution
The superoxide $\mathrm{O_2}^{-}$is paramagnetic because of one unpaired electron in $\pi^{*} 2 p$ molecular orbital. The hydroxides which are obtained by the reaction of the oxides with water are all white crystalline solids. The alkali metal hydroxides are the strongest of all bases and dissolve freely in water with evolution of much heat on account of intense hydration.
10.2.2 Halides
The alkali metal halides, $\mathrm{MX},(\mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I})$ are all high melting, colourless crystalline solids. They can be prepared by the reaction of the appropriate oxide, hydroxide or carbonate with aqueous hydrohalic acid (HX). All of these halides have high negative enthalpies of formation; the $\Delta_{f} H^{\ominus}$ values for fluorides become less negative as we go down the group, whilst the reverse is true for $\Delta_{f} H^{\ominus}$ for chlorides, bromides and iodides. For a given metal $\Delta_{f} H^{\ominus}$ always becomes less negative from fluoride to iodide.
The melting and boiling points always follow the trend: fluoride $>$ chloride $>$ bromide $>$ iodide. All these halides are soluble in water. The low solubility of $\mathrm{LiF}$ in water is due to its high lattice enthalpy whereas the low solubility of CsI is due to smaller hydration enthalpy of its two ions. Other halides of lithium are soluble in ethanol, acetone and ethylacetate; $\mathrm{LiCl}$ is soluble in pyridine also.
10.2.3 Salts of Oxo-Acids
Oxo-acids are those in which the acidic proton is on a hydroxyl group with an oxo group attached to the same atom e.g., carbonic acid, $\mathrm{H_2} \mathrm{CO_3} \left(\mathrm{OC}(\mathrm{OH})_2 \right.$; sulphuric acid, $\mathrm{H_2} \mathrm{SO_4}$ $ \left(\mathrm{O_2} \mathrm{~S}(\mathrm{OH})_2 \right)$. The alkali metals form salts with all the oxo-acids. They are generally soluble in water and thermally stable. Their carbonates $ \left(\mathrm{M_2} \mathrm{CO_3} \right)$ and in most cases the hydrogencarbonates $ \left(\mathrm{MHCO_3} \right)$ also are highly stable to heat. As the electropositive character increases down the group, the stability of the carbonates and hydorgencarbonates increases. Lithium carbonate is not so stable to heat; lithium being very small in size polarises a large $\mathrm{CO_3}^{2-}$ ion leading to the formation of more stable $\mathrm{Li_2} \mathrm{O}$ and $\mathrm{CO_2}$. Its hydrogencarbonate does not exist as a solid.
10.3 ANOMALOUS PROPERTIES OF LITHIUM
The anomalous behaviour of lithium is due to the : (i) exceptionally small size of its atom and ion, and (ii) high polarising power (i.e., charge/ radius ratio). As a result, there is increased covalent character of lithium compounds which is responsible for their solubility in organic solvents. Further, lithium shows diagonal relationship to magnesium which has been discussed subsequently.
10.3.1 Points of Difference between Lithium and other Alkali Metals
(i) Lithium is much harder. Its m.p. and b.p. are higher than the other alkali metals.
(ii) Lithium is least reactive but the strongest reducing agent among all the alkali metals. On combustion in air it forms mainly monoxide, $\mathrm{Li_2} \mathrm{O}$ and the nitride, $\mathrm{Li_3} \mathrm{~N}$ unlike other alkali metals.
(iii) $\mathrm{LiCl}$ is deliquescent and crystallises as a hydrate, $\mathrm{LiCl} .2 \mathrm{H_2} \mathrm{O}$ whereas other alkali metal chlorides do not form hydrates.
(iv) Lithium hydrogencarbonate is not obtained in the solid form while all other elements form solid hydrogencarbonates.
(v) Lithium unlike other alkali metals forms no ethynide on reaction with ethyne.
(vi) Lithium nitrate when heated gives lithium oxide, $\mathrm{Li_2} \mathrm{O}$, whereas other alkali metal nitrates decompose to give the corresponding nitrite.
$$ \begin{aligned} & 4 \mathrm{LiNO_3} \rightarrow 2 \mathrm{Li_2} \mathrm{O}+4 \mathrm{NO_2}+\mathrm{O_2} \\ & 2 \mathrm{NaNO_3} \rightarrow 2 \mathrm{NaNO_2}+\mathrm{O_2} \end{aligned} $$
(vii) $\mathrm{LiF}$ and $\mathrm{Li_2} \mathrm{O}$ are comparatively much less soluble in water than the corresponding compounds of other alkali metals.
10.3.2 Points of Similarities between Lithium and Magnesium
The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes : atomic radii, $\mathrm{Li}=152 \mathrm{pm}$, $\mathrm{Mg}=160 \mathrm{pm}$; ionic radii : $\mathrm{Li}^{+}=76 \mathrm{pm}$, $\mathrm{Mg}^{2+}=72 \mathrm{pm}$. The main points of similarity are:
(i) Both lithium and magnesium are harder and lighter than other elements in the respective groups.
(ii) Lithium and magnesium react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitride, $\mathrm{Li_3} \mathrm{~N}$ and $\mathrm{Mg_3} \mathrm{~N_2}$, by direct combination with nitrogen.
(iii) The oxides, $\mathrm{Li_2} \mathrm{O}$ and $\mathrm{MgO}$ do not combine with excess oxygen to give any superoxide.
(iv) The carbonates of lithium and magnesium decompose easily on heating to form the oxides and $\mathrm{CO_2}$. Solid hydrogencarbonates are not formed by lithium and magnesium.
(v) Both $\mathrm{LiCl}$ and $\mathrm{MgCl_2}$ are soluble in ethanol.
(vi) Both $\mathrm{LiCl}$ and $\mathrm{MgCl_2}$ are deliquescent and crystallise from aqueous solution as hydrates, $\mathrm{LiCl} \cdot 2 \mathrm{H_2} \mathrm{O}$ and $\mathrm{MgCl_2} \cdot 8 \mathrm{H_2} \mathrm{O}$.
10.4 SOME IMPORTANT COMPOUNDS OF SODIUM
Industrially important compounds of sodium include sodium carbonate, sodium hydroxide, sodium chloride and sodium bicarbonate. The large scale production of these compounds and their uses are described below:
Sodium Carbonate (Washing Soda), $\mathrm{Na_2} \mathrm{CO_3} \cdot \mathbf{1 0 H_ 2 \mathrm { O }}$
Sodium carbonate is generally prepared by Solvay Process. In this process, advantage is taken of the low solubility of sodium hydrogencarbonate whereby it gets precipitated in the reaction of sodium chloride with ammonium hydrogencarbonate. The latter is prepared by passing $\mathrm{CO_2}$ to a concentrated solution of sodium chloride saturated with ammonia, where ammonium carbonate followed by ammonium hydrogencarbonate are formed. The equations for the complete process may be written as :
$$ \begin{aligned} & 2 \mathrm{NH_3}+\mathrm{H_2} \mathrm{O}+\mathrm{CO_2} \rightarrow \left(\mathrm{NH_4} \right)_2 \mathrm{CO_3} \\ & \left(\mathrm{NH_4} \right)_2 \mathrm{CO_3}+\mathrm{H_2} \mathrm{O}+\mathrm{CO_2} \rightarrow 2 \mathrm{NH_4} \mathrm{HCO_3} \\ & \mathrm{NH_4} \mathrm{HCO_3}+\mathrm{NaCl} \rightarrow \mathrm{NH_4} \mathrm{Cl}+\mathrm{NaHCO_3} \\ \end{aligned} $$
Sodium hydrogencarbonate crystal separates. These are heated to give sodium carbonate.
$2 \mathrm{NaHCO_3} \rightarrow \mathrm{Na_2} \mathrm{CO_3}+\mathrm{CO_2}+\mathrm{H_2} \mathrm{O}$
In this process $\mathrm{NH_3}$ is recovered when the solution containing $\mathrm{NH_4} \mathrm{Cl}$ is treated with $\mathrm{Ca}(\mathrm{OH})_{2}$. Calcium chloride is obtained as a by-product.
$$ 2 \mathrm{NH_4} \mathrm{Cl}+\mathrm{Ca}(\mathrm{OH})_{2} \rightarrow 2 \mathrm{NH_3}+\mathrm{CaCl_2}+\mathrm{H_2} \mathrm{O} $$
It may be mentioned here that Solvay process cannot be extended to the manufacture of potassium carbonate because potassium hydrogencarbonate is too soluble to be precipitated by the addition of ammonium hydrogencarbonate to a saturated solution of potassium chloride.
Properties : Sodium carbonate is a white crystalline solid which exists as a decahydrate, $\mathrm{Na_2} \mathrm{CO_3} \cdot 10 \mathrm{H_2} \mathrm{O}$. This is also called washing soda. It is readily soluble in water. On heating, the decahydrate loses its water of crystallisation to form monohydrate. Above 373K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash.
$$ \begin{aligned} & \mathrm{Na_2} \mathrm{CO_3} + 10 \mathrm{H_2} \mathrm{O} \xrightarrow{375 \mathrm{~K}} \mathrm{Na_2} \mathrm{CO_3} + \mathrm{H_2} \mathrm{O}+9 \mathrm{H_2} \mathrm{O} \\ & \mathrm{Na_2} \mathrm{CO_3} + \mathrm{H_2} \mathrm{O} \xrightarrow{>373 \mathrm{~K}} \mathrm{Na_2} \mathrm{CO_3}+\mathrm{H_2} \mathrm{O} \end{aligned} $$
Carbonate part of sodium carbonate gets hydrolysed by water to form an alkaline solution.
$$ \mathrm{CO_3}^{2-}+\mathrm{H_2} \mathrm{O} \rightarrow \mathrm{HCO_3}^{-}+\mathrm{OH}^{-} $$
Uses:
(i) It is used in water softening, laundering and cleaning.
(ii) It is used in the manufacture of glass, soap, borax and caustic soda.
(iii) It is used in paper, paints and textile industries.
(iv) It is an important laboratory reagent both in qualitative and quantitative analysis.
Sodium Chloride, $\mathrm{NaCl}$
The most abundant source of sodium chloride is sea water which contains 2.7 to $2.9 %$ by mass of the salt. In tropical countries like India, common salt is generally obtained by evaporation of sea water. Approximately 50 lakh tons of salt are produced annually in India by solar evaporation. Crude sodium chloride, generally obtained by crystallisation of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Calcium chloride, $\mathrm{CaCl_2}$, and magnesium chloride, $\mathrm{MgCl_2}$ are impurities because they are deliquescent (absorb moisture easily from the atmosphere). To obtain pure sodium chloride, the crude salt is dissolved in minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium chloride separate out. Calcium and magnesium chloride, being more soluble than sodium chloride, remain in solution.
Sodium chloride melts at $1081 \mathrm{~K}$. It has a solubility of $36.0 \mathrm{~g}$ in $100 \mathrm{~g}$ of water at $273 \mathrm{~K}$. The solubility does not increase appreciably with increase in temperature.
Uses :
(i) It is used as a common salt or table salt for domestic purpose.
(ii) It is used for the preparation of $\mathrm{Na_2} \mathrm{O_2}$, $\mathrm{NaOH}$ and $\mathrm{Na_2} \mathrm{CO_3}$.
Sodium Hydroxide (Caustic Soda), $\mathrm{NaOH}$
Sodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in Castner-Kellner cell. A brine solution is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode.
Cathode: $\mathrm{Na}^{+}+\mathrm{e}^{-} \xrightarrow{\mathrm{Hg}} \mathrm{Na}-$ amalgam
Anode : $\mathrm{Cl}^{-} \rightarrow \frac{1}{2} \mathrm{Cl_2}+\mathrm{e}^{-}$
The amalgam is treated with water to give sodium hydroxide and hydrogen gas.
$2 \mathrm{Na}$-amalgam $+2 \mathrm{H_2} \mathrm{O} \rightarrow 2 \mathrm{NaOH}+2 \mathrm{Hg}+\mathrm{H_2}$
Sodium hydroxide is a white, translucent solid. It melts at $591 \mathrm{~K}$. It is readily soluble in water to give a strong alkaline solution. Crystals of sodium hydroxide are deliquescent. The sodium hydroxide solution at the surface reacts with the $\mathrm{CO_2}$ in the atmosphere to form $\mathrm{Na_2} \mathrm{CO_3}$.
Uses: It is used in (i) the manufacture of soap, paper, artificial silk and a number of chemicals, (ii) in petroleum refining, (iii) in the purification of bauxite, (iv) in the textile industries for mercerising cotton fabrics, (v) for the preparation of pure fats and oils, and (vi) as a laboratory reagent.
Sodium Hydrogencarbonate (Baking Soda), $\mathrm{NaHCO_3}$
Sodium hydrogencarbonate is known as baking soda because it decomposes on heating to generate bubbles of carbon dioxide (leaving holes in cakes or pastries and making them light and fluffy).
Sodium hydrogencarbonate is made by saturating a solution of sodium carbonate with carbon dioxide. The white crystalline powder of sodium hydrogencarbonate, being less soluble, gets separated out.
$\mathrm{Na_2} \mathrm{CO_3} + \mathrm{H_2} \mathrm{O} + \mathrm{CO_2} \rightarrow 2 \mathrm{NaHCO_3}$
Sodium hydrogencarbonate is a mild antiseptic for skin infections. It is used in fire extinguishers.
10.5 BIOLOGICAL IMPORTANCE OF SODIUM AND POTASSIUM
A typical $70 \mathrm{~kg}$ man contains about $90 \mathrm{~g}$ of $\mathrm{Na}$ and $170 \mathrm{~g}$ of $\mathrm{K}$ compared with only $5 \mathrm{~g}$ of iron and $0.06 \mathrm{~g}$ of copper.
Sodium ions are found primarily on the outside of cells, being located in blood plasma and in the interstitial fluid which surrounds the cells. These ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the transport of sugars and amino acids into cells. Sodium and potassium, although so similar chemically, differ quantitatively in their ability to penetrate cell membranes, in their transport mechanisms and in their efficiency to activate enzymes. Thus, potassium ions are the most abundant cations within cell fluids, where they activate many enzymes, participate in the oxidation of glucose to produce ATP and, with sodium, are responsible for the transmission of nerve signals.
There is a very considerable variation in the concentration of sodium and potassium ions found on the opposite sides of cell membranes. As a typical example, in blood plasma, sodium is present to the extent of $143 \mathrm{mmolL}^{-1}$, whereas the potassium level is only $5 \mathrm{mmolL}^{-1}$ within the red blood cells. These concentrations change to $10 \mathrm{mmolL}^{-1}\left(\mathrm{Na}^{+}\right)$and $105 \mathrm{mmolL}^{-1}\left(\mathrm{~K}^{+}\right)$. These ionic gradients demonstrate that a discriminatory mechanism, called the sodium-potassium pump, operates across the cell membranes which consumes more than one-third of the ATP used by a resting animal and about $15 \mathrm{~kg}$ per $24 \mathrm{~h}$ in a resting human.
10.6 GROUP 2 ELEMENTS : ALKALINE EARTH METALS
The group 2 elements comprise beryllium, magnesium, calcium, strontium, barium and radium. They follow alkali metals in the periodic table. These (except beryllium) are known as alkaline earth metals. The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium. The atomic and physical properties of the alkaline earth metals are shown in Table 10.2.
10.6.1 Electronic Configuration
These elements have two electrons in the $s$-orbital of the valence shell (Table 10.2). Their general electronic configuration may be represented as [noble gas] $n s^{2}$. Like alkali metals, the compounds of these elements are also predominantly ionic.
| Element | Symbol | Electronic configuration |
|---|---|---|
| Beryllium | $\mathrm{Be}$ | $1 s^{2} 2 s^{2}$ |
| Magnesium | $\mathrm{Mg}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2}$ |
| Calcium | $\mathrm{Ca}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2}$ |
| Strontium | $\mathrm{Sr}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 5 s^{2}$ |
| Barium | $\mathrm{Ba}$ | $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 4 d^{10} 5 s^{2} 5 p^{6} 6 s^{2}$ or $[\mathrm{Xe}] 6 s^{2}$ |
| Radium | Ra | $[\mathrm{Rn}] 7 s^{2}$ |
10.6.2 Atomic and Ionic Radii
The atomic and ionic radii of the alkaline earth metals are smaller than those of the
Table 10.2 Atomic and Physical Properties of the Alkaline Earth Metals
| Property | Beryllium Be |
Magnesium Mg |
Calcium Ca |
Strontium $\mathbf{S r}$ |
Barium Ba |
Radium Ra |
|---|---|---|---|---|---|---|
| Atomic number | 4 | 12 | 20 | 38 | 56 | 88 |
| Atomic mass $\left(\mathrm{g} \mathrm{mol}^{-1}\right)$ | 9.01 | 24.31 | 40.08 | 87.62 | 137.33 | 226.03 |
| Electronic configuration |
$[\mathrm{He}] 2 \mathrm{~s}^{2}$ | $[\mathrm{Ne}] 3 \mathrm{~s}^{2}$ | $[\mathrm{Ar}] 4 \mathrm{~s}^{2}$ | $[\mathrm{Kr}] 5 \mathrm{~s}^{2}$ | $[\mathrm{Xe}] 6 \mathrm{~s}^{2}$ | $[\mathrm{Rn}] 7 \mathrm{~s}^{2}$ |
| Ionization enthalpy (I) / kJ mol |
899 | 737 | 590 | 549 | 503 | 509 |
| Ionization enthalpy (II) $/ \mathrm{kJ} \mathrm{mol}^{-1}$ |
1757 | 1450 | 1145 | 1064 | 965 | 979 |
| Hydration enthalpy (kJ/mol) |
-2494 | -1921 | -1577 | -1443 | -1305 | - |
| Metallic radius / pm |
111 | 160 | 197 | 215 | 222 | - |
| Ionic radius $\mathrm{M}^{2+} / \mathrm{pm}$ |
31 | 72 | 100 | 118 | 135 | 148 |
| m.p. / K | 1560 | 924 | 1124 | 1062 | 1002 | 973 |
| b.p / K | 2745 | 1363 | 1767 | 1655 | 2078 | $(1973)$ |
| Density / $\mathrm{g} \mathrm{cm}^{-3}$ | 1.84 | 1.74 | 1.55 | 2.63 | 3.59 | $(5.5)$ |
| Standard potential $E^{\ominus} / V$ for $\left(M^{2+} / M\right)$ |
-1.97 | -2.36 | -2.84 | -2.89 | -2.92 | -2.92 |
| Occurrence in lithosphere |
$2 *$ | $2.76^{* *}$ | $4.6^{* *}$ | $384^{*}$ | 390 * | $10^{-6 *}$ |
*ppm (part per million); ** percentage by weight
corresponding alkali metals in the same periods. This is due to the increased nuclear charge in these elements. Within the group, the atomic and ionic radii increase with increase in atomic number.
10.6.3 Ionization Enthalpies
The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the atomic size increases down the group, their ionization enthalpy decreases (Table 10.2). The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals. This is due to their small size as compared to the corresponding alkali metals. It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.
10.6.4 Hydration Enthalpies
Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.
$$ \mathrm{Be}^{2+}>\mathrm{Mg}^{2+}>\mathrm{Ca}^{2+}>\mathrm{Sr}^{2+}>\mathrm{Ba}^{2+} $$
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., $\mathrm{MgCl_2}$ and $\mathrm{CaCl_2}$ exist as $\mathrm{MgCl_2} \cdot 6 \mathrm{H_2} \mathrm{O}$ and $\mathrm{CaCl_2} \cdot 6 \mathrm{H_2} \mathrm{O}$ while $\mathrm{NaCl}$ and $\mathrm{KCl}$ do not form such hydrates.
10.6.5 Physical Properties
The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish. The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The trend is, however, not systematic. Because of the low ionisation enthalpies, they are strongly electropositive in nature. The electropositive character increases down the group from $\mathrm{Be}$ to $\mathrm{Ba}$. Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame. In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. 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. The flame test for $\mathrm{Ca}, \mathrm{Sr}$ and $\mathrm{Ba}$ is helpful in their detection in qualitative analysis and estimation by flame photometry. The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.
10.6.6 Chemical Properties
The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.
(i) Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. However, powdered beryllium burns brilliantly on ignition in air to give $\mathrm{BeO}$ and $\mathrm{Be_3} \mathrm{~N_2}$. Magnesium is more electropositive and burns with dazzling brilliance in air to give $\mathrm{MgO}$ and $\mathrm{Mg_3} \mathrm{~N_2}$. Calcium, strontium and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold to form hydroxides.
(ii) Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.
$$ \mathrm{M}+\mathrm{X_2} \rightarrow \mathrm{MX_2}(\mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, 1) $$
Thermal decomposition of $\left(\mathrm{NH_4}\right)_{2} \mathrm{BeF_4}$ is the best route for the preparation of $\mathrm{BeF_2}$, and $\mathrm{BeCl_2}$ is conveniently made from the oxide.
$$\mathrm{BeO} + \mathrm{C} + \mathrm{Cl_2} \xrightarrow{\text{600 - 800 K}} \mathrm{BeCl_2} + \mathrm{CO}$$
(iii) Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides, $\mathrm{MH_2}$.
$\mathrm{BeH_2}$, however, can be prepared by the reaction of $\mathrm{BeCl_2}$ with $\mathrm{LiAlH_4}$.
$2 \mathrm{BeCl_2}+\mathrm{LiAlH_4} \rightarrow 2 \mathrm{BeH_2}+\mathrm{LiCl}+\mathrm{AlCl_3}$
(iv) Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen.
$\mathrm{M}+2 \mathrm{HCl} \rightarrow \mathrm{MCl_2}+\mathrm{H_2}$
(v) Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials (Table 10.2). However their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative value compared to other alkaline earth metals. However, its reducing nature is due to large hydration energy associated with the small size of $\mathrm{Be}^{2+}$ ion and relatively large value of the atomization enthalpy of the metal.
(vi) Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.
$\mathrm{M}+(\mathrm{x}+\mathrm{y}) \mathrm{NH_3} \rightarrow \left[\mathrm{M} \left(\mathrm{NH_3} \right)_x\right]^{2+}+2 \left[\mathrm{e}\left(\mathrm{NH_3} \right)_Y \right]^{-}$
From these solutions, the ammoniates, $\left[\mathrm{M}\left(\mathrm{NH_3}\right)_{6}\right]^{2+}$ can be recovered.
10.6.7 Uses
Beryllium is used in the manufacture of alloys. Copper-beryllium alloys are used in the preparation of high strength springs. Metallic beryllium is used for making windows of X-ray tubes. Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Magnesium carbonate is an ingredient of toothpaste. Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes. Radium salts are used in radiotherapy, for example, in the treatment of cancer.
10.7 GENERAL CHARACTERISTICS OF COMPOUNDS OF THE ALKALINE EARTH METALS
The dipositive oxidation state $\left(\mathrm{M}^{2+}\right)$ is the predominant valence of Group 2 elements. The alkaline earth metals form compounds which are predominantly ionic but less ionic than the corresponding compounds of alkali metals. This is due to increased nuclear charge and smaller size. The oxides and other compounds of beryllium and magnesium are more covalent than those formed by the heavier and large sized members ( $\mathrm{Ca}, \mathrm{Sr}, \mathrm{Ba})$. The general characteristics of some of the compounds of alkali earth metals are described below.
(i) Oxides and Hydroxides: The alkaline earth metals burn in oxygen to form the monoxide, MO which, except for BeO, have rock-salt structure. The BeO is essentially covalent in nature. The enthalpies of formation of these oxides are quite high and consequently they are very stable to heat. $\mathrm{BeO}$ is amphoteric while oxides of other elements are ionic in nature. All these oxides except $\mathrm{BeO}$ are basic in nature and react with water to form sparingly soluble hydroxides.
$$ \mathrm{MO}+\mathrm{H_2} \mathrm{O} \rightarrow \mathrm{M}(\mathrm{OH})_{2} $$
The solubility, thermal stability and the basic character of these hydroxides increase with increasing atomic number from $\mathrm{Mg}(\mathrm{OH})_2$ to $\mathrm{Ba}(\mathrm{OH})_2$. The alkaline earth metal hydroxides are, however, less basic and less stable than alkali metal hydroxides. Beryllium hydroxide is amphoteric in nature as it reacts with acid and alkali both.
$$ \begin{aligned} & \mathrm{Be}(\mathrm{OH})_2+2 \mathrm{OH}^{-} \rightarrow \underset{\text { Beryllate ion }}{\left[\mathrm{Be}(\mathrm{OH})_4 \right]^{2-}} \\ & \mathrm{Be}(\mathrm{OH})_2+2 \mathrm{HCl}+2 \mathrm{H_2} \mathrm{O} \rightarrow \left[\mathrm{Be}(\mathrm{OH})_4\right] \mathrm{Cl_2} \end{aligned} $$
(ii) Halides: Except for beryllium halides, all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown below:
In the vapour phase $\mathrm{BeCl_2}$ tends to form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of $1200 \mathrm{~K}$. The tendency to form halide hydrates gradually decreases (for example, $\mathrm{MgCl_2} \cdot 8 \mathrm{H_2} \mathrm{O}, \mathrm{CaCl_2} \cdot 6 \mathrm{H_2} \mathrm{O}, \mathrm{SrCl_2} \cdot 6 \mathrm{H_2} \mathrm{O}$ and $\mathrm{BaCl_2} \cdot 2 \mathrm{H_2} \mathrm{O}$ ) down the group. The dehydration of hydrated chlorides, bromides and iodides of $\mathrm{Ca}, \mathrm{Sr}$ and $\mathrm{Ba}$ can be achieved on heating; however, the corresponding hydrated halides of Be and Mg on heating suffer hydrolysis. The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.
(iii) 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 $\mathrm{CO_2}$. The thermal stability increases with increasing cationic size.
Sulphates: The sulphates of the alkaline earth metals are all white solids and stable to heat. $\mathrm{BeSO_4}$, and $\mathrm{MgSO_4}$ are readily soluble in water; the solubility decreases from $\mathrm{CaSO_4}$ to $\mathrm{BaSO_4}$. The greater hydration enthalpies of $\mathrm{Be}^{2+}$ and $\mathrm{Mg}^{2+}$ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.
Nitrates: The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six molecules of water, whereas barium nitrate crystallises as the anhydrous salt. This again shows a decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy. All of them decompose on heating to give the oxide like lithium nitrate.
$$ \begin{aligned} & 2 \mathrm{M}\left(\mathrm{NO_3}\right)_{2} \rightarrow 2 \mathrm{MO}+4 \mathrm{NO_2}+\mathrm{O_2} \\ & (\mathrm{M}=\mathrm{Be}, \mathrm{Mg}, \mathrm{Ca}, \mathrm{Sr}, \mathrm{Ba}) \end{aligned} $$
Problem 10.4
Why does the solubility of alkaline earth metal hydroxides in water increase down the group?
Solution
Among alkaline earth metal hydroxides, the anion being common the cationic radius will influence the lattice enthalpy. Since lattice enthalpy decreases much more than the hydration enthalpy with increasing ionic size, the solubility increases as we go down the group.
Problem 10.5
Why does the solubility of alkaline earth metal carbonates and sulphates in water decrease down the group?
Solution
The size of anions being much larger compared to cations, the lattice enthalpy will remain almost constant within a particular group. Since the hydration enthalpies decrease down the group, solubility will decrease as found for alkaline earth metal carbonates and sulphates.
10.8 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 which is discussed subsequently.
(i) Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of high ionisation enthalpy and small size it forms compounds which are largely covalent and get easily hydrolysed.
(ii) Beryllium does not exhibit coordination number more than four as in its valence shell there are only four orbitals. The remaining members of the group can have a coordination number of six by making use of $d$-orbitals. (iii) The oxide and hydroxide of beryllium, unlike the hydroxides of other elements in the group, are amphoteric in nature.
10.8.1 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. Some of the similarities are:
(i) Like aluminium, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.
(ii) Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion, $\left[\mathrm{Be}(\mathrm{OH})_4\right]^{2-}$ just as aluminium hydroxide gives aluminate ion, $\left[\mathrm{Al}(\mathrm{OH})_4\right]^{-}$.
(iii) The chlorides of both beryllium and aluminium have $\mathrm{Cl}^{-}$bridged chloride structure in vapour phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel Craft catalysts.
(iv) Beryllium and aluminium ions have strong tendency to form complexes, $\mathrm{BeF_4}{ }^{2-}, \mathrm{AlF_6}{ }^{3-}$.
10.9 SOME IMPORTANT COMPOUNDS OF CALCIUM
Important compounds of calcium are calcium oxide, calcium hydroxide, calcium sulphate, calcium carbonate and cement. These are industrially important compounds. The large scale preparation of these compounds and their uses are described below.
Calcium Oxide or Quick Lime, $\mathrm{CaO}$
It is prepared on a commercial scale by heating limestone $\left(\mathrm{CaCO_3}\right)$ in a rotary kiln at 1070-1270 K.
$$\mathrm{CaCO_3} \xrightarrow{\text{ heat }} \mathrm{CaO} + \mathrm{CO_2}$$
The carbon dioxide is removed as soon as it is produced to enable the reaction to proceed to completion.
Calcium oxide is a white amorphous solid. It has a melting point of $2870 \mathrm{~K}$. On exposure to atmosphere, it absorbs moisture and carbon dioxide.
$$ \begin{aligned} & \mathrm{CaO}+\mathrm{H_2} \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_{2} \\ & \mathrm{CaO}+\mathrm{CO_2} \rightarrow \mathrm{CaCO_3} \end{aligned} $$
The addition of limited amount of water breaks the lump of lime. This process is called slaking of lime. Quick lime slaked with soda gives solid sodalime. Being a basic oxide, it combines with acidic oxides at high temperature.
$$ \begin{aligned} & \mathrm{CaO}+\mathrm{SiO_2} \rightarrow \mathrm{CaSiO_3} \\ & 6 \mathrm{CaO}+\mathrm{P_4} \mathrm{O_10} \rightarrow 2 \mathrm{Ca_3}\left(\mathrm{PO_4}\right)_{2} \end{aligned} $$
Uses:
(i) It is an important primary material for manufacturing cement and is the cheapest form of alkali.
(ii) It is used in the manufacture of sodium carbonate from caustic soda.
(iii) It is employed in the purification of sugar and in the manufacture of dye stuffs.
Calcium Hydroxide (Slaked lime), $\mathrm{Ca}(\mathrm{OH})_{2}$ Calcium hydroxide is prepared by adding water to quick lime, $\mathrm{CaO}$.
It is a white amorphous powder. It is sparingly soluble in water. The aqueous solution is known as lime water and a suspension of slaked lime in water is known as milk of lime.
When carbon dioxide is passed through lime water it turns milky due to the formation of calcium carbonate.
$$ \mathrm{Ca}(\mathrm{OH})_{2}+\mathrm{CO_2} \rightarrow \mathrm{CaCO_3}+\mathrm{H_2} \mathrm{O} $$
On passing excess of carbon dioxide, the precipitate dissolves to form calcium hydrogencarbonate.
$$ \mathrm{CaCO_3}+\mathrm{CO_2}+\mathrm{H_2} \mathrm{O} \rightarrow \mathrm{Ca}\left(\mathrm{HCO_3}\right)_{2} $$
Milk of lime reacts with chlorine to form hypochlorite, a constituent of bleaching powder.
$$ 2 \mathrm{Ca}(\mathrm{OH})_2+2 \mathrm{Cl_2} \rightarrow \mathrm{CaCl_2}+\underset{\substack{\text { Bleaching powder }}}{\mathrm{Ca}(\mathrm{OCl})_2}+2 \mathrm{H_2} \mathrm{O} $$
Uses:
(i) It is used in the preparation of mortar, a building material. (ii) It is used in white wash due to its disinfectant nature.
(iii) It is used in glass making, in tanning industry, for the preparation of bleaching powder and for purification of sugar.
Calcium Carbonate, $\mathrm{CaCO_3}$
Calcium carbonate occurs in nature in several forms like limestone, chalk, marble etc. It can be prepared by passing carbon dioxide through slaked lime or by the addition of sodium carbonate to calcium chloride.
$\mathrm{Ca}(\mathrm{OH})_{2}+\mathrm{CO_2} \rightarrow \mathrm{CaCO_3}+\mathrm{H_2} \mathrm{O}$
$\mathrm{CaCl_2}+\mathrm{Na_2} \mathrm{CO_3} \rightarrow \mathrm{CaCO_3}+2 \mathrm{NaCl}$
Excess of carbon dioxide should be avoided since this leads to the formation of water soluble calcium hydrogencarbonate.
Calcium carbonate is a white fluffy powder. It is almost insoluble in water. When heated to $1200 \mathrm{~K}$, it decomposes to evolve carbon dioxide.
$$ \mathrm{CaCO_3} \xrightarrow{1200 \mathrm{~K}} \mathrm{CaO}+\mathrm{CO_2} $$
It reacts with dilute acid to liberate carbon dioxide.
$\mathrm{CaCO_3}+2 \mathrm{HCl} \rightarrow \mathrm{CaCl_2}+\mathrm{H_2} \mathrm{O}+\mathrm{CO_2}$
$\mathrm{CaCO_3}+\mathrm{H_2} \mathrm{SO_4} \rightarrow \mathrm{CaSO_4}+\mathrm{H_2} \mathrm{O}+\mathrm{CO_2}$
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. Specially precipitated $\mathrm{CaCO_3}$ is extensively used in the manufacture of high quality paper. 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), $\mathrm{CaSO_4} \cdot \mathrm{H_2} \mathrm{O}$
It is a hemihydrate of calcium sulphate. It is obtained when gypsum, $\mathrm{CaSO_4} \cdot 2 \mathrm{H_2} \mathrm{O}$, is heated to $393 \mathrm{~K}$.
$2\left(\mathrm{CaSO_4} \cdot 2 \mathrm{H_2} \mathrm{O}\right) \rightarrow 2\left(\mathrm{CaSO_4}\right) \cdot \mathrm{H_2} \mathrm{O}+3 \mathrm{H_2} \mathrm{O}$
Above $393 \mathrm{~K}$, no water of crystallisation is left and anhydrous calcium sulphate, $\mathrm{CaSO_4}$ is formed. This is known as ‘dead burnt plaster’.
It has a remarkable property of setting with water. On mixing with an adequate quantity of water it forms a plastic mass that gets into a hard solid in 5 to 15 minutes.
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: Cement is an important building material. It was first introduced in England in 1824 by Joseph Aspdin. It is also called Portland cement because it resembles with the natural limestone quarried in the Isle of Portland, England.
Cement is a product obtained by combining a material rich in lime, $\mathrm{CaO}$ with other material such as clay which contains silica, $\mathrm{SiO_2}$ along with the oxides of aluminium, iron and magnesium. The average composition of Portland cement is : $\mathrm{CaO}, 50-$ $60 % ; \mathrm{SiO_2}, 20-25 % ; \mathrm{Al_2} \mathrm{O_3}, 5-10 % ; \mathrm{MgO}, 2-$ $3 % ; \mathrm{Fe_2} \mathrm{O_3}, 1-2 %$ and $\mathrm{SO_3}, 1-2 %$. For a good quality cement, the ratio of silica $\left(\mathrm{SiO_2}\right)$ to alumina $\left(\mathrm{Al_2} \mathrm{O_3}\right)$ should be between 2.5 and 4 and the ratio of lime $(\mathrm{CaO})$ to the total of the oxides of silicon $\left(\mathrm{SiO_2}\right)$ aluminium $\left(\mathrm{Al_2} \mathrm{O_3}\right)$ and iron $\left(\mathrm{Fe_2} \mathrm{O_3}\right)$ should be as close as possible to 2 .
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 $\left(\mathrm{CaSO_4} \cdot 2 \mathrm{H_2} \mathrm{O}\right)$ to form cement. Thus important ingredients present in Portland cement are dicalcium silicate $\left(\mathrm{Ca_2} \mathrm{SiO_4}\right) 26 %$, tricalcium silicate $\left(\mathrm{Ca_3} \mathrm{SiO_5}\right) 51 %$ and tricalcium aluminate $\left(\mathrm{Ca_3} \mathrm{Al_2} \mathrm{O_6}\right) 11 %$.
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. The purpose of adding gypsum is only to slow down the process of setting of the cement so that it gets sufficiently hardened.
Uses: Cement has become a commodity of national necessity for any country next to iron and steel. It is used in concrete and reinforced concrete, in plastering and in the construction of bridges, dams and buildings.
10.10 BIOLOGICAL IMPORTANCE OF MAGNESIUM AND CALCIUM
An adult body contains about $25 \mathrm{~g}$ of $\mathrm{Mg}$ and $1200 \mathrm{~g}$ of Ca compared with only $5 \mathrm{~g}$ of iron and $0.06 \mathrm{~g}$ of copper. The daily requirement in the human body has been estimated to be $200-300 \mathrm{mg}$.
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 $100 \mathrm{mgL}^{-1}$. It is maintained by two hormones: calcitonin and parathyroid hormone. Do you know that bone is not an inert and unchanging substance but is continuously being solubilised and redeposited to the extent of $400 \mathrm{mg}$ per day in man? All this calcium passes through the plasma.
SUMMARY
The s-Block of the periodic table constitutes Group1 (alkali metals) and Group 2 (alkaline earth metals). They are so called because their oxides and hydroxides are alkaline in nature. The alkali metals are characterised by one s-electron and the alkaline earth metals by two $s$-electrons in the valence shell of their atoms. These are highly reactive metals forming monopositive $\left(\mathbf{M}^{+}\right)$and dipositve $\left(\mathbf{M}^{2+}\right)$ ions respectively.
There is a regular trend in the physical and chemical properties of the alkali metal with increasing atomic numbers. The atomic and ionic sizes increase and the ionization enthalpies decrease systematically down the group. Somewhat similar trends are observed among the properties of the alkaline earth metals.
The first element in each of these groups, lithium in Group 1 and beryllium in Group 2 shows similarities in properties to the second member of the next group. Such similarities are termed as the ‘diagonal relationship’ in the periodic table. As such these elements are anomalous as far as their group characteristics are concerned.
The alkali metals are silvery white, soft and low melting. They are highly reactive. The compounds of alkali metals are predominantly ionic. Their oxides and hydroxides are soluble in water forming strong alkalies. Important compounds of sodium includes sodium carbonate, sodium chloride, sodium hydroxide and sodium hydrogencarbonate. Sodium hydroxide is manufactured by Castner-Kellner process and sodium carbonate by Solvay process.
The chemistry of alkaline earth metals is very much like that of the alkali metals. However, some differences arise because of reduced atomic and ionic sizes and increased cationic charges in case of alkaline earth metals. Their oxides and hydroxides are less basic than the alkali metal oxides and hydroxides. Industrially important compounds of calcium include calcium oxide (lime), calcium hydroxide (slaked lime), calcium sulphate (Plaster of Paris), calcium carbonate (limestone) and cement. Portland cement is an important constructional material. It is manufactured by heating a pulverised mixture of limestone and clay in a rotary kiln. The clinker thus obtained is mixed with some gypsum $(2-3 %)$ to give a fine powder of cement. All these substances find variety of uses in different areas.
Monovalent sodium and potassium ions and divalent magnesium and calcium ions are found in large proportions in biological fluids. These ions perform important biological functions such as maintenance of ion balance and nerve impulse conduction.
EXERCISES
10.1 What are the common physical and chemical features of alkali metals ?
Answer Physical properties of alkali metalsare as follows. (1) They are quite soft and can be cut easily. Sodium metal can be easily cut using a knife. (2) They are light coloured and are mostly silvery white in appearance. (3) They have low density because of the large atomic sizes. The density increases down the group from Li to Cs. The only exceptionto this isK, which has lower density than $Na$. (4) The metallic bonding present in alkali metals is quite weak. Therefore, they have low melting and boiling points. (5) Alkali metals and their salts impart a characteristic colour to flames. This is because the heat from the flame excites the electron present in the outermost orbital to a high energy level. When this excited electron reverts back to the ground state, it emits excess energy as radiation that falls in the visible region. (6) They also display photoelectric effect. When metals such as Cs and K are irradiated with light, they lose electrons. Chemical properties of alkali metals Alkali metals are highly reactive due to their low ionization enthalpy. As we move down the group, the reactivity increases. (1) They react with water to form respective oxides or hydroxides. As we move down the group, the reaction becomes more and more spontaneous. (2) They react with water to form their respective hydroxides and dihydrogens. The general reaction for the same is given as $
2 M+2 H_2 O \longrightarrow 2 M^{+}+2 OH^{\ominus}+H_2
$ (3) They react with dihydrogen to form metal hydrides. These hydrides are ionic solids and have high melting points. $
2 M+H_2 \longrightarrow 2 M^{+} H^{-}
$ (4) Almost all alkali metals, except Li, react directly with halogens to form ionic halides. $
\begin{aligned}
& 2 M+Cl_2 \longrightarrow 2 MCl \\
& (M=Li, K, Rb, Cs)
\end{aligned}
$ Since Li+ion is very small in size, it can easily distort the electron cloud around the negative halide ion. Therefore, lithium halides are covalent in nature. (5) They are strong reducing agents. The reducing power of alkali metals increases on moving down the group. However, lithium is an exception. It is the strongest reducing agent among the alkali metals. It is because of its high hydration energy. (6) They dissolve in liquid ammonia to form deep blue coloured solutions. These solutions are conducting in nature. $
M+(x+y) NH_3 \longrightarrow[M(NH_3) _{x}]^{+}+[M(NH_3) _{y}]^{-}
$ The ammoniated electrons cause the blue colour of the solution. These solutions are paramagnetic and if allowed to stand for some time, then they liberate hydrogen. This results in the formation of amides. $
M _{(a m)}^{+}+e^{-}+NH _{3(l)} \longrightarrow MNH _{(a m)}+\frac{1}{2} H _{2(g)}
$ In a highly concentrated solution, the blue colour changes to bronze and the solution becomes diamagnetic.Show Answer
Answer General characteristics of alkaline earth metals are as follows. (i) The general electronic configuration of alkaline earth metals is [noble gas] $n s^{2}$. (ii) These metals lose two electrons to acquire the nearest noble gas configuration. Therefore, their oxidation state is +2 . (iii)These metals have atomic and ionic radii smaller than that of alkali metals. Also, when moved down the group, the effective nuclear charge decreases and this causes an increase in their atomic radii and ionic radii. (iv)Since the alkaline earth metals have large size, their ionization enthalpies are found to be fairly low. However, their first ionization enthalpies are higher than the corresponding group 1 metals. (v) These metals are lustrous and silvery white in appearance. They are relatively less soft as compared to alkali metals. (vi)Atoms of alkaline earth metals are smaller than that of alkali metals. Also, they have two valence electrons forming stronger metallic bonds. These two factors cause alkaline earth metals to have high melting and boiling points as compared to alkali metals. (vii) They are highly electropositive in nature. This is due to their low ionization enthalpies. Also, the electropositive character increases on moving down the group from $Be$ to $Ba$. (viii) $Ca, Sr$, and $Ba$ impart characteristic colours to flames. Ca â€" Brick red Sr â€" Crimson red Ba –“Apple green In Be and Mg, the electrons are too strongly bound to be excited. Hence, these do not impart any colour to the flame. The alkaline earth metals are less reactive than alkali metals and their reactivity increases on moving down the group. Chemical properties of alkaline earth metals are as follows. (i) Reaction with air and water: $Be$ and $Mg$ are almost inert to air and water because of the formation of oxide layer on their surface. (a) Powdered Be burns in air to form $BeO$ and $Be_3 N_2$. (b) $Mg$, being more electropositive, burns in air with a dazzling sparkle to form $MgO$ and $Mg_3 N_2$. (c) $Ca, Sr$, and $Ba$ react readily with air to form respective oxides and nitrides. (d) $Ca, Ba$, and $Sr$ react vigorously even with cold water. (ii) Alkaline earth metals react with halogens at high temperatures to form halides. $
M+X_2 \longrightarrow MX_2(X=F, Cl, Br, I)
$ (iii) All the alkaline earth metals, except Be, react with hydrogen to form hydrides. (iv) They react readily with acids to form salts and liberate hydrogen gas. $
M+2 HCl \longrightarrow MCl_2+H _{2(g)} \uparrow
$ (v) They are strong reducing agents. However, their reducing power is less than that of alkali metals. As we move down the group, the reducing power increases. (vi) Similar to alkali metals, the alkaline earth metals also dissolve in liquid ammonia to give deep blue coloured solutions. $
M+(x-y) NH_3 \longrightarrow[M(NH_3) _{x}]^{+2}+2[e(NH_3) _{y}]^{-}
$Show Answer
Answer Alkali metals include lithium, sodium, potassium, rubidium, cesium, and francium. These metals have only one electron in their valence shell, which they lose easily, owing to their low ionization energies. Therefore, alkali metals are highly reactive and are not found in nature in their elemental state.Show Answer
Answer Let the oxidation state of $Na$ be $x$. The oxidation state of oxygen, in case of peroxides, is â $\hat{\text{ " }}$. Therefore,
$2(x)+2(-1)=0$ $2 x-2=0$ $2 x=2$ $x=+1$ Therefore, the oxidation sate of sodium is +1 .Show Answer
Answer In alkali metals, on moving down the group, the atomic size increases and the effective nuclear charge decreases. Because of these factors, the outermost electron in potassium can be lost easily as compared to sodium. Hence, potassium is more reactive than sodium.Show Answer
AnswerShow Answer
Alkali metals
Alkaline earth metals
(i)
Ionization enthalpy:
These have lowest ionization enthalpies
in respective periods. This is because of
their large atomic sizes. Also, they lose
their only valence electron easily as
they attain stable noble gas
configuration after losing it.(i)
Ionization enthalpy:
Alkaline earth metals have smaller atomic size and
higher effective nuclear charge as compared to alkali
metals. This causes their first ionization enthalpies
to be higher than that of alkali metals. However,
their second ionization enthalpy is less than the
corresponding alkali metals. This is because alkali
metals, after losing one electron, acquires noble gas
configuration, which is very stable.
(ii)
Basicity of oxides:
The oxides of alkali metals are very
basic in nature. This happens due to the
highly electropositive nature of alkali
metals, which makes these oxides
highly ionic. Hence, they readily
dissociate in water to give hydroxide
ions.Basicity of oxides:
The oxides of alkaline earth metals are quite basic
but not as basic as those of alkali metals. This is
because alkaline earth metals are less electropositive
than alkali metals.
(iii)
Solubility of hydroxides:
The hydroxides of alkali metals are
more soluble than those of alkaline
earth metals.The hydroxides of alkaline earth metals are less
soluble than those of alkali metals. This is due to the
high lattice energies of alkaline earth metals. Their
higher charge densities (as compared to alkali
metals) account for higher lattice energies.
Answer Similarities between lithium and magnesium are as follows. (i) Both $Li$ and $Mg$ react slowly with cold water. (ii) The oxides of both $Li$ and $Mg$ are much less soluble in water and their hydroxides decompose at high temperature. $
\begin{aligned}
& 2 LiOH \xrightarrow{\text{ heat }} Li_2 O+H_2 O \\
& Mg(OH)_2 \xrightarrow{\text{ heat }} MgO+H_2 O
\end{aligned}
$ (iii) Both $Li$ and $Mg$ react with $N_2$ to form nitrides. $6 Li+N_2 \xrightarrow{\text{ heat }} 2 Li_3 N$ $3 Mg+N_2 \xrightarrow{\text{ heat }} Mg_3 N_2$ (iv) Neither Li nor Mg form peroxides or superoxides. (v) The carbonates of both are covalent in nature. Also, these decompose on heating. $
\begin{aligned}
& Li_2 CO_3 \xrightarrow{\text{ heat }} Li_2 O+CO_2 \\
& MgCO_3 \xrightarrow{\text{ heat }} MgO+CO_2
\end{aligned}
$ (vi) Li and $Mg$ do not form solid bicarbonates. (vii) Both $LiCl$ and $MgCl_2$ are soluble in ethanol owing to their covalent nature. (viii) Both $LiCl$ and $MgCl_2$ are deliquescent in nature. They crystallize from aqueous solutions as hydrates, for example, $LiCl \cdot 2 H_2 O$ and $MgCl_2 \cdot 8 H_2 O$.Show Answer
Answer In the process of chemical reduction, oxides of metals are reduced using a stronger reducing agent. Alkali metals and alkaline earth metals are among the strongest reducing agents and the reducing agents that are stronger than them are not available. Therefore, they cannot be obtained by chemical reduction of their oxides.Show Answer
Answer All the three, lithium, potassium, and cesium, are alkali metals. Still, $K$ and $Cs$ are used in the photoelectric cell and not Li. This is because as compared to Cs and K, Li is smaller in size and therefore, requires high energy to lose an electron. While on the other hand, $K$ and $Cs$ have low ionization energy. Hence, they can easily lose electrons. This property of $K$ and $Cs$ is utilized in photoelectric cells.Show Answer
Answer When an alkali metal is dissolved in liquid ammonia, it results in the formation of a deep blue coloured solution. $
M+(x+y) NH_3 \longrightarrow M^{+}(NH_3) _{x}+e^{-1}(NH_3) _{y}
$ The ammoniated electrons absorb energy corresponding to red region of visible light. Therefore, the transmitted light is blue in colour. At a higher concentration ( $3 M$ ), clusters of metal ions are formed. This causes the solution to attain a copper–bronze colour and a characteristic metallic lustre.Show Answer
Answer When an alkaline earth metal is heated, the valence electrons get excited to a higher energy level. When this excited electron comes back to its lower energy level, it radiates energy, which belongs to the visible region. Hence, the colour is observed. In Be and Mg, the electrons are strongly bound. The energy required to excite these electrons is very high. Therefore, when the electron reverts back to its original position, the energy released does not fall in the visible region. Hence, no colour in the flame is seen.Show Answer
Answer Solvay process is used to prepare sodium carbonate. When carbon dioxide gas is bubbled through a brine solution saturated with ammonia, sodium hydrogen carbonate is formed. This sodium hydrogen carbonate is then converted to sodium carbonate. Step 1: Brine solution is saturated with ammonia. $2 NH_3+H_2 O+CO_2 \longrightarrow(NH_4)_2 CO_3$ This ammoniated brine is filtered to remove any impurity. Step 2: Carbon dioxide is reacted with this ammoniated brine to result in the formation of insoluble sodium hydrogen carbonate. $NH+H_2 O+CO_2 \tilde{A}$ câ $€$ ’ $NH_4 HCO_3 NaCl+NH_4 HCO_3 \tilde{A}$ câ $€$ ’ $NaHCO_3+NH_4 Cl$ Step 3: The solution containing crystals of $NaHCO_3$ is filtered to obtain $NaHCO_3$. Step 4: $NaHCO_3$ is heated strongly to convert it into $NaHCO_3$. $2 NaHCO_3 \longrightarrow Na_2 CO_3+CO_2+H_2 O$ Step 5: To recover ammonia, the filtrate (after removing $NaHCO_3$ ) is mixed with $Ca(OH)_2$ and heated. $Ca(OH)_2+2 NH_4 Cl$ Ã The overall reaction taking place in Solvay process is $2 NaCl+CaCO_3 \longrightarrow Na_2 CO_3+CaCl_2$Show Answer
Answer Solvay process cannot be used to prepare potassium carbonate. This is because unlike sodium bicarbonate, potassium bicarbonate is fairly soluble in water and does not precipitate out.Show Answer
Answer As we move down the alkali metal group, the electropositive character increases. This causes an increase in the stability of alkali carbonates. However, lithium carbonate is not so stable to heat. This is because lithium carbonate is
covalent. Lithium ion, being very small in size, polarizes a large carbonate ion, leading to the formation of more stable lithium oxide. $
Li_2 CO_3 \xrightarrow{\Delta} Li_2 O+CO_2
$ Therefore, lithium carbonate decomposes at a low temperature while a stable sodium carbonate decomposes at a high temperature.Show Answer
Answer (i) Nitrates Thermal stability Nitrates of alkali metals, except $LiNO_3$, decompose on strong heating to form nitrites. $
2 KNO _{3(s)} \longrightarrow 2 KNO _{2(s)}+O _{2(g)}
$ $LiNO_3$, on decomposition, gives oxide. $
2 LiNO _{3(s)} \xrightarrow{\Delta} Li_2 O _{(s)}+2 NO _{2(g)}+O _{2(g)}
$ Similar to lithium nitrate, alkaline earth metal nitrates also decompose to give oxides. $
2 Ca(NO_3) _{(s)} \xrightarrow{\Delta} 2 CaO _{(s)}+4 NO _{2(g)}+O _{2(g)}
$ As we move down group 1 and group 2, the thermal stability of nitrate increases. Solubility Nitrates of both group 1 and group 2 metals are soluble in water. (ii) Carbonates Thermal stability The carbonates of alkali metals are stable towards heat. However, carbonate of lithium, when heated, decomposes to form lithium oxide. The carbonates of alkaline earth metals also decompose on heating to form oxide and carbon dioxide. $
\begin{aligned}
& Na_2 CO_3 \xrightarrow{\Delta} \text{ No effect } \\
& Li_2 CO_3 \xrightarrow{\Delta} Li_2 O+CO_2 \\
& MgCO_3 \xrightarrow{\Delta} MgO+CO_2
\end{aligned}
$ Solubility Carbonates of alkali metals are soluble in water with the exception of $Li_2 CO_3$. Also, the solubility increases as we move down the group. Carbonates of alkaline earth metals are insoluble in water. (iii) Sulphates Thermal stability Sulphates of both group 1 and group 2 metals are stable towards heat. Solubility Sulphates of alkali metals are soluble in water. However, sulphates of alkaline earth metals show varied trends. $BeSO_4$ Fairly soluble $MgSO_4$ Soluble $CaSO_4$ Sparingly soluble $SrSO_4$ Insoluble $BaSO_4$ Insoluble In other words, while moving down the alkaline earth metals, the solubility of their sulphates decreases.Show Answer
Answer (a) Sodium can be extracted from sodium chloride by Downs process. This process involves the electrolysis of fused $NaCl(40 %)$ and $CaCl_2(60 %)$ at a temperature of $1123 K$ in Downs cell. Steel is the cathode and a block of graphite acts as the anode. Metallic $Na$ and $Ca$ are formed at cathode. Molten sodium is taken out of the cell and collected over kerosene. $NaCl$ $\xrightarrow{\text{ Electrolysis }} Na^{+}+Cl^{-}$ Molten At Cathode: $Na^{+}+e^{-} \longrightarrow Na$ At Anode: $Cl^{-}+e^{-} \longrightarrow Cl$ $
Cl+Cl \longrightarrow Cl_2
$ (ii) Sodium hydroxide can be prepared by the electrolysis of sodium chloride. This is called CastnerâE"Kellner process. In this process, the brine solution is electrolysed using a carbon anode and a mercury cathode. The sodium metal, which is discharged at cathode, combines with mercury to form an amalgam. Cathode: $Na^{+}+e^{-} \xrightarrow{Hg} Na$ - amalgam Anode: $Cl^{-} \longrightarrow \frac{1}{2} Cl_2+e^{-}$ (iii) Sodium peroxide First, $NaCl$ is electrolysed to result in the formation of $Na$ metal (Downs process). This sodium metal is then heated on aluminium trays in air (free of $CO_2$ ) to form its peroxide. $2 Na+O _{2 \text{ (air) }} \longrightarrow Na_2 O_2$ (iv) Sodium carbonate is prepared by Solvay process. Sodium hydrogen carbonate is precipitated in a reaction of sodium chloride and ammonium hydrogen carbonate. $
\begin{aligned}
& 2 NH_3+H_2 O+CO_2 \longrightarrow(NH_4)_2 CO_3 \\
& (NH_4)_2 CO_3+H_2 O+CO_2 \longrightarrow 2 NH_4 HCO_3 \\
& NH_4 HCO_3+NaCl \longrightarrow NH_4 Cl+NaHCO_3
\end{aligned}
$ These sodium hydrogen carbonate crystals are heated to give sodium carbonate. $2 NaHCO_3 \longrightarrow Na_2 CO_3+CO_2+H_2 O$Show Answer
Answer (i) Magnesium burns in air with a dazzling light to form $MgO$ and $Mg_3 N_2$. $
\begin{aligned}
& 2 Mg+O_2 \xrightarrow{\text{ Buming }} 2 MgO \\
& 3 Mg+N_2 \xrightarrow{\text{ Burning }} Mg_3 N_2
\end{aligned}
$ (ii) Quick lime ( $CaO)$ combines with silica $(SiO_2)$ to form slag. $
CaO+SiO_2 \xrightarrow{\text{ heat }} CaSiO_3
$ (iii) When chloride is added to slaked lime, it gives bleaching powder. $Ca(OH)_2+Cl_2 \xrightarrow{\Delta} CaOCl_2+H_2 O$ (iv) Calcium nitrate, on heating, decomposes to give calcium oxide. $
2 Ca(NO_3) _{2(s)} \xrightarrow{\Delta} 2 CaO _{(s)}+4 NO _{2(g)}+O _{2(g)}
$Show Answer
Bleaching
powder
Answer (i) Uses of caustic soda (a) It is used in soap industry. (b) It is used as a reagent in laboratory. (ii) Uses of sodium carbonate (a) It is generally used in glass and soap industry. (b) It is used as a water softener. (iii) Uses of quick lime (a) It is used as a starting material for obtaining slaked lime. (b) It is used in the manufacture of glass and cement.Show Answer
Answer (a) Structure of $BeCl_2$ (solid) $BeCl_2$ exists as a polymer in condensed (solid) phase. In the vapour state, $BeCl_2$ exists as a monomer with a linear structure.Show Answer
Answer The atomic size of sodium and potassium is larger than that of magnesium and calcium. Thus, the lattice energies of carbonates and hydroxides formed by calcium and magnesium are much more than those of sodium and potassium. Hence, carbonates and hydroxides of sodium and potassium dissolve readily in water whereas those of calcium and magnesium are only sparingly soluble.Show Answer
Answer (i) Chemically, limestone is $CaCO_3$. Importance of limestone (a) It is used in the preparation of lime and cement. (b) It is used as a flux during the smelting of iron ores. (ii) Chemically, cement is a mixture of calcium silicate and calcium aluminate. Importance of cement (a) It is used in plastering and in construction of bridges. (b) It is used in concrete. (iii) Chemically, plaster of Paris is $2 CaSO_4 \cdot H_2 O$. Importance of plaster of Paris (a) It is used in surgical bandages. (b) It is also used for making casts and moulds.Show Answer
Answer Lithium is the smallest in size among the alkali metals. Hence, $Li^{+}$ion can polarize water molecules more easily than other alkali metals. As a result, water molecules get attached to lithium salts as water of crystallization. Hence, lithium salts such as trihydrated lithium chloride $(LiCl .3 H_2 O)$ are commonly hydrated. As the size of the ions increases, their polarizing power decreases. Hence, other alkali metal ions usually form anhydrous salts.Show Answer
Answer LiF is insoluble in water. On the contrary, $LiCl$ is soluble not only in water, but also in acetone. This is mainly because of the greater ionic character of LiF as compared to LiCl. The solubility of a compound in water depends on the balance between lattice energy and hydration energy. Since fluoride ion is much smaller in size than chloride ion, the lattice energy of $LiF$ is greater than that of $LiCl$. Also, there is not much difference between the hydration energies of fluoride ion and chloride ion. Thus, the net energy change during the dissolution of $LiCl$ in water is more exothermic than that during the dissolution of LiF in water. Hence, low lattice energy and greater covalent character are the factors making $LiCl$ soluble not only in water, but also in acetone.Show Answer
Answer Importance of sodium, potassium, magnesium, and calcium in biological fluids: (i) Sodium (Na): Sodium ions are found primarily in the blood plasma. They are also found in the interstitial fluids surrounding the cells. (a) Sodium ions help in the transmission of nerve signals. (b) They help in regulating the flow of water across the cell membranes. (c) They also help in transporting sugars and amino acids into the cells. (ii) Potassium (K): Potassium ions are found in the highest quantity within the cell fluids. (a) $K$ ions help in activating many enzymes. (b) They also participate in oxidising glucose to produce ATP. (c) They also help in transmitting nerve signals. (iii) Magnesium (Mg) and calcium (Ca): Magnesium and calcium are referred to as macro-minerals. This term indicates their higher abundance in the human body system. (a) Mghelps in relaxing nerves and muscles. (b) $Mg$ helps in building and strengthening bones. (c) Mg maintains normal blood circulation in the human body system. (d) Ca helps in the coagulation of blood (e) Ca also helps in maintaining homeostasis.Show Answer
(i) sodium metal is dropped in water ?
(ii) sodium metal is heated in free supply of air ?
(iii) sodium peroxide dissolves in water ?
Answer (i) When $Na$ metal is dropped in water, it reacts violently to form sodium hydroxide and hydrogen gas. The chemical equation involved in the reaction is: $
2 Na _{(s)}+2 H_2 O _{(f)} \longrightarrow 2 NaOH _{(a q)}+H _{2(g)}
$ (ii) On being heated in air, sodium reacts vigorously with oxygen to form sodium peroxide. The chemical equation involved in the reaction is: $
2 Na _{(s)}+O _{2(s)} \longrightarrow Na_2 O _{2(s)}
$ (iii) When sodium peroxide is dissolved in water, it is readily hydrolysed to form sodium hydroxide and water. The chemical equation involved in the reaction is: $
Na_2 O _{2(s)}+2 H_2 O _{(l)} \longrightarrow 2 NaOH _{(a q)}+H_2 O _{2(a q)}
$Show Answer
(a) The mobilities of the alkali metal ions in aqueous solution are $\mathrm{Li}^{+}<\mathrm{Na}^{+}<\mathrm{K}^{+}$ $<\mathrm{Rb}^{+}<\mathrm{Cs}^{+}$
(b) Lithium is the only alkali metal to form a nitride directly.
(c) $\mathrm{E}^{\ominus}$ for $\mathrm{M}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{M}(\mathrm{s})$ (where $\mathrm{M}=\mathrm{Ca}$, $\mathrm{Sr}$ or $\mathrm{Ba}$ ) is nearly constant.
Answer (a) On moving down the alkali group, the ionic and atomic sizes of the metals increase. The given alkali metal ions can be arranged in the increasing order of their ionic sizes as: $Li^{+}<Na^{+}<K^{+}<Rb^{+}<Cs^{+}$ Smaller the size of an ion, the more highly is it hydrated. Since $Li^{+}$is the smallest, it gets heavily hydrated in an aqueous solution. On the other hand, $Cs^{+}$is the largest and so it is the least hydrated. The given alkali metal ions can be arranged in the decreasing order of their hydrations as: $Li^{+}>Na^{+}>K^{+}>Rb^{+}>Cs^{+}$ Greater the mass of a hydrated ion, the lower is its ionic mobility. Therefore, hydrated $Li^{+}$is the least mobile and hydrated $Cs^{+}$is the most mobile. Thus, the given alkali metal ions can be arranged in the increasing order of their mobilities as: $Li^{+}<Na^{+}<K^{+}<Rb^{+}<Cs^{+}$ (b) Unlike the other elements of group 1, Li reacts directly with nitrogen to form lithium nitride. This is because $Li^{+}$is very small in size and so its size is the most compatible with the $N^{3}$ ion. Hence, the lattice energy released is very high. This energy also overcomes the high amount of energy required for the formation of the $N^{3-}$ ion. (c) Electrode potential $(E^{\circ})$ of any $M^{2+} / M$ electrode depends upon three factors: (i) Ionisation enthalpy (ii) Enthalpy of hydration (iii) Enthalpy of vaporisation The combined effect of these factors is approximately the same for $Ca, Sr$, and $Ba$. Hence, their electrode potentials are nearly constant.Show Answer
(a) a solution of $\mathrm{Na_2} \mathrm{CO_3}$ is alkaline?
(b) alkali metals are prepared by electrolysis of their fused chlorides ?
(c) sodium is found to be more useful than potassium ?
Answer (a) When sodium carbonate is added to water, it hydrolyses to give sodium bicarbonate and sodium hydroxide (a strong base). As a result, the solution becomes alkaline. $
Na_2 CO_3+H_2 O \longrightarrow NaHCO_3+NaOH
$ (b) It is not possible to prepare alkali metals by the chemical reduction of their oxides as they themselves are very strong reducing agents. They cannot be prepared by displacement reactions either (wherein one element is displaced by another). This is because these elements are highly electropositive. Neither can electrolysis of aqueous solutions be used to extract these elements. This is because the liberated metals react with water. Hence, to overcome these difficulties, alkali metals are usually prepared by the electrolysis of their fused chlorides.
(c) Blood plasma and the interstitial fluids surrounding the cells are the regions where sodium ions are primarily found. Potassium ions are located within the cell fluids. Sodium ions are involved in the transmission of nerve signals, in regulating the flow of water across the cell membranes, and in transporting sugars and amino acids into the cells. Hence, sodium is found to be more useful than potassium.Show Answer
(a) $\mathrm{Na_2} \mathrm{O_2}$ and water
(b) $\mathrm{KO_2}$ and water
(c) $\mathrm{Na_2} \mathrm{O}$ and $\mathrm{CO_2}$.
Answer (a) The balanced chemical equation for the reaction between $Na_2 O_2$ and water is: $
2 Na_2 O _{2(s)}+2 H_2 O _{(l)} \longrightarrow 4 NaOH _{(a q)}+O _{2(a q)}
$ (b) The balanced chemical equation for the reaction between $KO_2$ and water is: $
2 KO _{2(s)}+2 H_2 O _{(j)} \longrightarrow 2 KOH _{(a q)}+H_2 O _{2(a q)}+O _{2(g)}
$ (c) The balanced chemical equation for the reaction between $Na_2 O$ and $CO_2$ is: $
Na_2 O _{(s)}+CO _{2(g)} \longrightarrow Na_2 CO_3
$Show Answer
(i) $\mathrm{BeO}$ is almost insoluble but $\mathrm{BeSO_4}$ is soluble in water,
(ii) $\mathrm{BaO}$ is soluble but $\mathrm{BaSO_4}$ is insoluble in water,
(iii) LiI is more soluble than KI in ethanol.
Answer (i) $BeO$ is almost insoluble in water and $BeSO_4$ is soluble in water. $Be^{2+}$ is a small cation with a high polarising power and $O^{22 E}$ is a small anion. The size compatibility of $Be^{2+}$ and $O^{25 \epsilon}$ is high. Therefore, the lattice energy released during their formation is also very high. When $BeO$ is dissolved in water, the hydration energy of its ions is not sufficient to overcome the high lattice energy. Therefore, $BeO$ is insoluble in water. On the other hand, $SO_4^{2-}$ ion is a large anion. Hence, $Be^{2+}$ can easily polarise $SO_4^{2-}$ ions, making $BeSO_4$ unstable. Thus, the lattice energy of $BeSO_4$ is not very high and so it is soluble in water. (ii) $BaO$ is soluble in water, but $BaSO_4$ is not. $Ba^{2+}$ is a large cation and $O^{2 a+}$ is a small anion. The size compatibility of $Ba^{2+}$ and $O^{2 \bar{{}\epsilon} \text{ is }}$ not high. As a result, $BaO$ is unstable. The lattice energy released during its formation is also not very
large. It can easily be overcome by the hydration energy of the ions. Therefore, $BaO$ is soluble in water. In $BaSO_4$, $Ba^{2+}$ and $SO_4^{2-}$ are both large-sized. The lattice energy released is high. Hence, it is not soluble in water. (iii) Lil is more soluble than $KI$ in ethanol. As a result of its small size, the lithium ion has a higher polarising power than the potassium ion. It polarises the electron cloud of the iodide ion to a much greater extent than the potassium ion. This causes a greater covalent character in Lil than in KI. Hence, Lil is more soluble in ethanol.Show Answer
(a) $\mathrm{Na}$
(b) $\mathrm{K}$
(c) $\mathrm{Rb}$
(d) $\mathrm{Cs}$
Answer Atomic size increases as we move down the alkali group. As a result, the binding energies of their atoms in the crystal lattice decrease. Also, the strength of metallic bonds decreases on moving down a group in the periodic table. This causes a decrease in the melting point. Among the given metals, Cs is the largest and has the least melting point.Show Answer
(b) $\mathrm{Na}$
(c) $\mathrm{K}$
(d) $\mathrm{Cs}$
Answer Smaller the size of an ion, the more highly is it hydrated. Among the given alkali metals, Li is the smallest in size. Also, it has the highest charge density and highest polarising power. Hence, it attracts water molecules more strongly than the other alkali metals. As a result, it forms hydrated salts such as $LiCl .2 H_2 O$. The other alkali metals are larger than $Li$ and have weaker charge densities. Hence, they usually do not form hydrated salts.Show Answer
(b) $\mathrm{CaCO_3}$
(c) $\mathrm{SrCO_3}$
(d) $\mathrm{BaCO_3}$
Show Answer
Answer
Thermal stability increases with the increase in the size of the cation present in the carbonate. The increasing order of the cationic size of the given alkaline earth metals is
$Mg<Ca<Sr<Ba$
Hence, the increasing order of the thermal stability of the given alkaline earth metal carbonates is $MgCO_3<CaCO_3<SrCO_3<BaCO_3$
