The S-block

elements are those in which the last electron enters the outermost s-orbital. As the sorbital can accommodate only two electrons, two groups ( 1 & 2) belong to the s-block of the Periodic Table.

Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr) are called alkali metals since their oxides and hydroxides are soluble in water and form strong alkalies. Francium is, however, radioactive.

General electronic configuration

The general electronic configuration of alkali metals is [noble gas] $n s^{1}$ where $n=2-7$. Due to the presence of one unpaired electron, all alkali metals are paramagnetic but their salts are diamagnetic.

Abundance

Sodium and potassium are abundant and lithium, rubidium and caesium have much lower abundances. Francium is highly radioactive, its longest lived isotope ${ }^{223} \mathrm{Fr}$ has a halflife of only 21 minutes.

Atomic radii

The atomic radii of alkali metals are the largest (after inert gases) in their respective periods. Both atomic and ionic radii increase down the group from $\mathrm{Li}$ to $\mathrm{Cs}$ as the number of inner filled shells and screening effect increases.

Enthalpy of hydration

The enthalpy of hydration of alkali metal ions decreases with increase is ionic radii, i.e., $\mathrm{Li}^{+}>\mathrm{Na}^{+}>\mathrm{K}^{+}>\mathrm{Cs}^{+}$. In other words, degree of hydration decreases down the group from $\mathrm{Li}^{+}$to $\mathrm{Cs}^{+}$. Consequently, radii of hydrated ions decreases in the order : $\left.\left.\left.\left.\mathrm{Li}^{+}\right\rangle \mathrm{Na}^{+}\right\rangle \mathrm{K}^{+}\right\rangle \mathrm{Rb}^{+}\right\rangle \mathrm{Cs}^{+}$. Conversely, the ionic mobility decreases in the opposite order, i.e., $\mathrm{Li}^{+}<\mathrm{Na}^{+}<\mathrm{K}^{+}<\mathrm{Rb}^{+}<\mathrm{Cs}^{+}$.

Further since $\mathrm{Li}$ has the maximum degree of hydration, therefore, several $\mathrm{Li}$ salts are hydrated, i.e. $\mathrm{LiClO} _{4} \cdot 3 \mathrm{H} _{2} \mathrm{O}, \mathrm{Li} _{2} \mathrm{CO} _{3} \cdot 3 \mathrm{H} _{2} \mathrm{O}, \mathrm{LiCl} ; 2 \mathrm{H} _{2} \mathrm{O}$ etc.

Ionization enthalpy

The first ionization enthalpy $\left(\Delta _{i} H _{1}\right)$ of alkali metals is quite low and decreases further down the group due to increasing atomic size and shielding effect. Further, due to inert gas configuration of the unipositive ions the second ionization enthalpies $\left(\Delta _{i} \mathrm{H} _{2}\right)$ of alkali metals are very high and decrease down the group from Li to Cs.

Electropositive character

All alkali metals are strongly electropositive due to their low ionization enthalpies. Further, the electropositive character increases down the group as the ionization enthalpies decrease. Thus, Fr is the most metallic and electropositive element in the periodic table. Further, due to the presence of a single electron outside the noble gas core, all the alkali metals have a strong tendency to lose this electron to acquire the electronic configuration of the nearest noble gas.

Consequently, all the alkali metals uniformly show an oxidation state of +1 .

Metallic character

Alkali metals have low ionization energies, they have high tendency to lose valence electrons. The metallic character increases down the group because ionization energy decreases down the group.

Melting and boiling points

The melting and boiling points of alkali metals are very low and decrease with increase in atomic number.

As alkali metals have large size, the intermetallic bonds in them are quite weak. Hence, they have low melting and boiling points which decrease down the group with the increase in the atomic size.

Photoelectric effect

Due to low ionization enthalpies, all alkali metals except Li show photoelectric effect. Due to strong photoelectric effect, K and Cs are commonly used in photoelectric cells.

Nature of bonding

Due to low ionization enthalpies, alkali metals form ionic compounds. Further, the ionic character increases down the group due to decreasing ionization enthalpies. Lithium, however, because of its high ionization enthalpy forms covalent compounds, i.e., alkyllithium ( $\mathrm{R}-\mathrm{Li}$ ), aryllithium (Ar-Li).

Flame colouration

Due to low ionization enthalpies, all the alkali metals and their salts impart characteristic colours to the flame. For example.

$ \begin{array}{llllll} \text{Metal} & \text{Li} & \mathrm{Na} & \text{K} & \mathbf{Rb} & \text{Cs} \\ \text{Colour} & \text{Crimson red} & \text{Golden yellow} & \text{Pink violet} & \text{Red violet} & \text{Sky blue} \\ \lambda n \mathrm{~m} & (670-8 \mathrm{~nm}) & (589-2 \mathrm{~nm}) & (766-5 \mathrm{nm}) & (780-0 \mathrm{~nm}) & (455-5 \mathrm{nm}) \\ \end{array} $

On heating an alkali metal or its salts especially chloride due to its more volatile nature in a flame, the electrons are excited easily to higher energy levels because of absorption of energy. When these excited electrons return to their ground states, they emit extra energy in the visible region thereby imparting a characteristic colour to the flame.

Reactivity and electrode potential

Reducing character

Due to large negative electrode potentials, alkali metals are strong reducing agents. The reducing power increases from Na to $\mathrm{Cs}$ but $\mathrm{Li}$ is the strongest reducing agent.

Electrode potential is a measure of the tendency of an element to lose electrons in the aqueous solution. More negative is the electrode potential, higher is the tendency of the element to lose electrons and hence stronger is the reducing agent.

Lithium is the strongest reducing agent in the aqueous solution.

Electrode potential depends upon

i. Sublimation enthalpy

ii. Ionization enthalpy

iii. Enthalpy of hydration i. $\mathrm{Li}(\mathrm{s}) \xrightarrow{\text { sublimation enthalpy }} \mathrm{Li}(\mathrm{g})$

ii. $\quad \mathrm{Li}(\mathrm{g}) \xrightarrow{\text { Ionization Enthalpy }} \mathrm{Li}^{+}(\mathrm{g})+\mathrm{e}^{-}$

iii. $\mathrm{Li}^{+}(\mathrm{g})+\mathrm{aq} \longrightarrow \mathrm{Li}^{+}(\mathrm{aq})+$ enthalpy of hydration

The sublimation enthalpies of alkali metals are almost similar. Since, Lithium has the smallest ionic size among alkali metals, its enthalpy of hydration is the highest. Althrough ionization enthalpy of Lithium is the highest among alkali metals, it is more than compensated by the large hydration enthalpy.

Action of water

All alkali metals react with water evolving $\mathrm{H} _{2}$ and forming their corresponding metal hydroxides,

$$ 2 \mathrm{M}+2 \mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{MOH}+\mathrm{H} _{2} $$

Lithium is the least reactive while the reactivity of other alkali metals increases down the group as the electropositive character of the metal increases :

$$ \mathrm{Li}<\mathrm{Na}<\mathrm{K}<\mathrm{Rb}<\mathrm{Cs} $$

Basic Character

Alkali metal hydroxides are strong bases and their solubility and basic character increases down the group from Li to Cs as the ionization enthalpy of the metal decreases or the electropositive character of the metal increases. Thus, basicity increases in the order:

$$ \mathrm{LiOH}<\mathrm{NaOH}<\mathrm{KOH}<\mathrm{RbOH}<\mathrm{CsOH} $$

Action of oxygen

Lithium forms monoxide

$4 \mathrm{Li}+\mathrm{O} _{2} \xrightarrow{\Delta} 2 \mathrm{Li} _{2} \mathrm{O}$

Sodium forms peroxide

$2 \mathrm{Na}+\mathrm{O} _{2} \xrightarrow{\Delta} \mathrm{Na} _{2} \mathrm{O} _{2}$

$\mathrm{K}$, $\mathrm{Rb}$ and $\mathrm{Cs}$ react with oxygen form superoxides

$$ \mathrm{K}+\mathrm{O} _{2} \longrightarrow \mathrm{KO} _{2} $$

Increasing stability of peroxide or superoxide as the size of the metal cation increases is due to the stabilization of larger anions by larger cations through higher lattice enthalpies.

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 an oxidizing agent in organic chemistry.

Action of air and moisture

Alkali metals get tarnished when exposed to air and moisture due to the formation of oxides, then hydroxides and finally carbonates. Because of these reactions, alkali metals are stored in inert hydrocarbon solvents like petroleum ether and kerosene oil which prevent them from coming in contact with air and moisture.

Action of hydrogen

All alkali metals on heating with hydrogen form ionic hydrides of the general formula $\mathrm{M}^{+} \mathrm{H}^{-}$. The reactivity of alkali metals towards hydrogen decreases down the group from Li to Cs due to the reason that lattice enthalpies and hence the stability of these hydrides decreases down the group as the size of the metal cation increases, i.e., stability decreases in the order:

$\mathrm{LiH}>\mathrm{NaH}>\mathrm{KH}>\mathrm{RbH}>\mathrm{CsH}$

Action of halogens

Alkali metals combine directly with halogens to form metal halides (MX). With the exception of lithium halides, all other alkali metal halides are ionic.

The reactivity of alkali metals towards a particular halogen increases in the order. $\mathrm{Li}<\mathrm{Na}<\mathrm{K}<\mathrm{Rb}<\mathrm{Cs}$ while that of halogens towards a particular alkali metal decreases in the order: $\mathrm{F} _{2}>\mathrm{Cl} _{2}>\mathrm{Br} _{2}>\mathrm{I} _{2}$.

(a) Solubility of alkali metal halides.

With the exception of LiF, all other lithium halides are covalent. Being covalent, LiCl, LiBr and Lil are insoluble in water but are soluble in organic solvents like pyridine, benzene, alcohols and ethers.

(b) Melting points and boiling points.

(i) For the same alkali metal, the melting points decrease in the order: fluorides > chlorides > bromides > iodides due to a corresponding decrease in their lattice enthalpies. Thus, the m.p. of sodium halides decrease in the order : $\mathrm{NaF}(1261 \mathrm{~K})>$ $\mathrm{NaCl}(1084 \mathrm{~K})>\mathrm{NaBr}(1028 \mathrm{~K})>\mathrm{Nal}(944 \mathrm{~K})$.

(ii) For the same halide ion, lithium halides being covalent have lower melting points than their corresponding sodium halides. Thereafter, the melting points decrease due to a corresponding decrease in lattice enthalpies as the size of the metal increases. Thus, the m.p. of alkali metal chlorides decrease in the order: $\mathrm{NaCl}$ ( 1084 K) > KCl (1039 K) > RbCl (988 K) > CSCl (925 K) > LiCl (887 K).

Solubility in liquid ammonia

All the alkali metals dissolve in liquid ammonia to give blue solutions due to the presence of ammoniated (solvated) electrons in the solution.

$$ \mathrm{M}+(\mathrm{x}+\mathrm{y}) \mathrm{NH} _{3} \longrightarrow \mathrm{M}^{+}\left(\mathrm{NH} _{3}\right) _{\mathrm{x}}+\mathrm{e}^{-}\left(\mathrm{NH} _{3}\right) _{\mathrm{y}} $$

When ordinary light falls on these ammoniated electrons, they get excited to higher energy levels by absorbing energy corresponding to red region of visible light. As a result, transmitted light is blue which imparts blue colour to the solution.

Dilute solutions of alkali metals in liquid ammonia are dark blue in colour but as the concentration increases above $3 \mathrm{M}$, the colour changes to copper bronze and solution acquires metallic lustre due to formation of metal ion clusters.

Blue coloured solutions are paramagnetic due to presence of large unpaired electrons but bronze solutions are diamagnetic due to formation of metal ion clusters. These solutions behave as strong reducing agents.

Salts of oxo-acids

Oxo acids are those in which the acidic proton is a hydroxyl group with an oxo group attached to the same ex-carbonic acid $\left(\mathrm{H} _{2} \mathrm{CO} _{3}\right)$. The alkali metals form salts with all oxo acids. They are generally soluble in water and thermally stable.

Their carbonates and in most cases their hydrogen carbonates are $\left(\mathrm{M} _{2} \mathrm{CO} _{3} \& \mathrm{MHCO} _{3}\right)$ are stable to heat.

Stability of carbonates.

The stability of carbonates towards heat increases down the group as the basic character of the alkali metal hydroxides increases down the group. $\mathrm{Li} _{2} \mathrm{CO} _{3}$ is, however, less stable and decomposes on heating

$$ \mathrm{Li} _{2} \mathrm{CO} _{3} \xrightarrow{\Delta} \mathrm{Li} _{2} \mathrm{O}+\mathrm{CO} _{2} $$

Lithium differs from other alkali metals because of its (i) small size, (ii) high polarizing power, (ii) high ionization enthalpy and (iv) absence of $d$-orbitals in the valence shell.

Lithium forms bicarbonate in solution while all other alkali metals form solid bicarbonates. All the bicarbonates on gentle heating undergo decomposition to form carbonates with the evolution of $\mathrm{CO} _{2}$

$$ 2 \mathrm{MHCO} _{3} \xrightarrow{\Delta} \mathrm{M} _{2} \mathrm{CO} _{3}+\mathrm{CO} _{2}+\mathrm{H} _{2} \mathrm{O} $$

All the carbonates and bicarbonates are soluble in water and their solubility increases down the group. Lithium differs from other alkali metals

(i) Li is hard. Its m.p. and b.p. are higher

(ii) Li is least reactive but the strongest reducing agent

(iii) LiCl is deliquescent and crystallises as a hydrate

(iv) $\mathrm{LiHCO} _{3}$ is not obtained in the solid form while all other elements form solid bicarbonates

(v) Lidoes not form ethynide on reaction with ethyne.

(vi) $\mathrm{LiNO} _{3}$ when heated gives $\mathrm{L} _{2} \mathrm{O}$ white others decompose to give the corresponding nitrite $2 \mathrm{NaNO} _{3} \longrightarrow 2 \mathrm{NaNO} _{2}+\mathrm{O} _{2}$ $4 \mathrm{LiNO} _{3} \longrightarrow 2 \mathrm{Li} _{2} \mathrm{O}+4 \mathrm{NO} _{2}+\mathrm{O} _{2}$

Lithium shows diagonal relationship with magnesium mainly due to the similarity in sizes of their atoms and ions.

1. Both react with nitrogen to form ionic nitride

$6 \mathrm{Li}+\mathrm{N} _{2} \xrightarrow{\Delta} 2 \mathrm{Li} _{3} \mathrm{~N} \quad ; \quad 3 \mathrm{Mg}+\mathrm{N} _{2} \xrightarrow{\Delta} \mathrm{Mg} _{3} \mathrm{~N} _{2}$

2. Carbonates of these metals decompose on heating to the corresponding oxide with the evolution of $\mathrm{CO} _{2}$.

$\mathrm{Li} _{2} \mathrm{CO} _{3} \xrightarrow{\Delta} \mathrm{Li} _{2} \mathrm{O}+\mathrm{CO} _{2} \quad ; \quad \mathrm{MgCO} _{3} \xrightarrow{\Delta} \mathrm{MgO}+\mathrm{CO} _{2}$

3. Hydroxides of both Liand $\mathrm{Mg}$ decompose on heating

$2 \mathrm{LiOH} \xrightarrow{\Delta} \mathrm{Li}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \quad ; \quad \mathrm{Mg}(\mathrm{OH})_2 \xrightarrow{\Delta} \mathrm{MgO}+\mathrm{H}_2 \mathrm{O}$

4. Both Li and magnesium nitrates decompose on heating to produce $\mathrm{NO} _{2}$

$\begin{aligned} & 4 \mathrm{LiNO} _{3} \xrightarrow{\Delta} 2 \mathrm{Li} _{2} \mathrm{O}+4 \mathrm{NO} _{2}+\mathrm{O} _{2} \\ & 2 \mathrm{Mg}\left(\mathrm{NO} _{3}\right) _{2} \xrightarrow[\Delta]{\Delta} 2 \mathrm{MgO}+4 \mathrm{NO} _{2}+\mathrm{O} _{2} \end{aligned}$

5. $\mathrm{LiCl}$ and $\mathrm{MgCl} _{2}$ are deliquescent and crystallise from aqueous solution as hydrates, $\mathrm{LiCl} _{2} .2 \mathrm{H} _{2} \mathrm{O}$ and $\mathrm{Mg} \mathrm{Cl} _{2} .6 \mathrm{H} _{2} \mathrm{O}$.

Extraction of alkali metals

Alkali metals cannot be isolated by electrolysis of aqueous solution of their salts since hydrogen is liberated at the cathode instead of alkali metal because the standard electrode potentials of alkali metals are much lower ( $\mathrm{L} i=-3.04 \mathrm{~V}, \mathrm{Na}=-2.71 \mathrm{~V}$ ) than that of $\mathrm{H} _{2} \mathrm{O}\left(-0.82 \mathrm{~V}\right.$ ). As a result, $\mathrm{H} _{2}$ is produced in preference to alkali metals.

Compounds of Sodium

Sodium carbonate (washing soda), $\mathrm{Na} _{2} \mathrm{CO} _{3} \cdot 10 \mathrm{H} _{2} \mathrm{O}$ is prepared by Solvay ammonia process. In this process, $\mathrm{CO} _{2}$ is passed through brine (sodium chloride) solution saturated with ammonia, when sodium bicarbonate being sparingly soluble separates out. It is filtered and then ignited to get sodium carbonate (called soda ash). $\mathrm{NaCl}+\mathrm{H} _{2} \mathrm{O}+\mathrm{CO} _{2}+\mathrm{NH} _{3} \longrightarrow \mathrm{NaHCO} _{3}+\mathrm{NH} _{4} \mathrm{Cl}$

$2 \mathrm{NaHCO} _{3} \xrightarrow{\Delta} \mathrm{Na} _{2} \mathrm{CO} _{3}+\mathrm{CO} _{2}+\mathrm{H} _{2} \mathrm{O}$

Potassium carbonate cannot be prepared by this process because $\mathrm{KHCO} _{3}$ is fairly soluble in water. As a result, when $\mathrm{CO} _{2}$ is passed through ammoniated brine, $\mathrm{KHCO} _{3}$ does not separate out.

Baking soda is $\mathrm{NaHCO} _{3}$ and is obtained by saturating a solution of $\mathrm{Na} _{2} \mathrm{CO} _{3}$ with $\mathrm{CO} _{2}$ when $\mathrm{NaHCO} _{3}$ being less soluble separates out.

$$ \mathrm{Na} _{2} \mathrm{CO} _{3}+\mathrm{CO} _{2}+\mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{NaHCO} _{3} $$

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.

Crude $\mathrm{NaCl}$ is generally obtained by crystallisation of brine solution. $\mathrm{NaCl}$ melts at $1081 \mathrm{~K}$. It has a solubility of $36.0 \mathrm{~g}$ in $100 \mathrm{~g}$ of water at $273 \mathrm{~K}$. It is used as a common salt for domestic purpose, for preparation of $\mathrm{Na} _{2} \mathrm{O} _{2}, \mathrm{NaOH}$ and $\mathrm{Na} _{2} \mathrm{CO} _{3}$.

Sodium hydroxide (Caustic soda)

NaOH is prepared by the electrolysis of NaCl by Castner - Kellner cell. Arine solution is electrolyzed using a mercury cathode and a carbon anode

$$ \begin{aligned} & \text { Cathode } \mathrm{Na}^{+}+\mathrm{e}^{-} \longrightarrow \text { Hg.Na-amalgam } \\ & \text { Anode } \mathrm{Cl}^{-} \longrightarrow 1 / 2 \mathrm{Cl} _{2}+\mathrm{e}^{-} \\ & 2 \mathrm{Na} \text {-amalgam }+2 \mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{NaOH}+2 \mathrm{Hg}+2 \mathrm{H} _{2} \end{aligned} $$

It is a while translucent solid, melts at 591K. It is readily soluble in water to give a strong alkaline solution.

It is used in manufacture of soap, paper, artificial silk and a number of chemicals, in petroleum refining, in purification of bauxite, in textile industry in preparation of pure fats and oils and as laboratory reagent.

Sodium hydrogen carbonate (Baking Soda) $\mathrm{NaHCO} _{3}$

It is known as baking soda as it decomposes on heating to generate bubbles of $\mathrm{CO} _{2}$ leaving holes in cakes or pastries and make them light and fluffy.

It is made by saturating a solution of $\mathrm{Na} _{2} \mathrm{CO} _{3}$ with $\mathrm{CO} _{2}$. The white powder of $\mathrm{NaHCO} _{3}$ being less soluble separates out.

$$ \mathrm{Na} _{2} \mathrm{CO} _{3}+\mathrm{H} _{2} \mathrm{O}+\mathrm{CO} _{2} \longrightarrow 2 \mathrm{NaHCO} _{3} $$

It is used as a mild antiseptic for skin infection and is used in fire extinguishes.

Alkaline Earth Metals (Group 2 Elements)

Members of the group.

Group 2 of the periodic table consists of six elements : beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr) and radium (Ra). Radium is, however, radioactive. These elements are also called alkaline earth elements because the oxides of $\mathrm{Mg}, \mathrm{Ca}, \mathrm{Sr}$ and $\mathrm{Ba}$ are found in earth’s crust and are alkaline in mixture.

General electronic configuration.

The general electronic configuration of alkaline earth metals is [noble gas] $n s^{2}$ where $n=2-7$.

Atomic and ionic radii. The first and second ionization enthalpies of alkaline earth metals are fairly low; the second ionization enthalpy $\left(\Delta _{i} \mathrm{H} _{2}\right)$ being almost double of the first ionization enthalpy $\left(\Delta _{i} \mathrm{H} _{1}\right)$.

$ \begin{array}{lccc} \text{Element} & \text{IE1 } (\mathbf{K J ~ m o l}^{-1}) & \text{IE2 } (\mathbf{K J ~ m o l}^{-1}) \\ \mathrm{Na} \text{ (group 1)} & 496 & 4562 \\ \mathrm{Mg} \text{ (group 2)} & 737 & 1450 \\ \end{array} $

The second electron in case of alkali metals (e.g. Na) is to be removed from a cation (unipositive ion) which has already acquired the stable noble gas configuration whereas in case of alkaline earth metals (e.g. Mg), the second electron is to be removed from a cation (unipositive ion) which is yet to acquire the stable noble gas configuration. Therefore, removal of second electron in case of alkaline earth metals requires much less energy than that in case of alkali metals.

Hydration enthalpy.

Like alkali metal ions the hydration enthalpies of alkaline earth metal ions decrease as the size of metal ion increases down the group ie. $\mathrm{Be}^{2+}>\mathrm{Mg}^{2+}>\mathrm{Ca}^{2+}>\mathrm{Sr}^{2+}>\mathrm{Ba}^{2+}$. Due to smaller size of alkaline earth metal ions, the hydration enthalpies of alkaline earth metal ions are larger than alkali metal ions.

The compounds of alkaline earth metal ions are more extensively hydrated than those of alkali metal ions. $\mathrm{MgCl} _{2}$ exists as $\mathrm{Mg} \mathrm{Cl} _{2}$. $6 \mathrm{H} _{2} \mathrm{O}$.

Oxidation state.

All alkaline earth metals uniformly show an oxidation state of +2 both in the solid state or in the aqueous solutions. In the solid state, the greater lattice enthalpy of the $\mathrm{M}^{2+}$ ions (relative to $\mathrm{M}^{+}$ ions) more than compensates the higher second ionization enthalpy while in the aqueous solution, it is the greater hydration enthalpy of the $\mathrm{M}^{2+}$ ions (relative to $\mathrm{M}^{+}$ions) which more than compensates the higher second ionization enthalpy.

Electropositive character.

Alkaline earth metals are strongly electropositive since they have a strong tendency to lose both the valence electrons to form $\mathrm{M}^{2+}$ ions having noble gas configuration. However due to small size and higher ionization enthalpies $\left(\Delta _{i} \mathrm{H} _{1}, \Delta _{i} \mathrm{H} _{2}\right)$ they are less electropositive than alkali metals. Further, the electropositive character increases down the group (i.e., $\mathrm{Be}<\mathrm{Mg}<\mathrm{Ca}<\mathrm{Sr}<\mathrm{Ba}$ ) as the ionization enthalpies decrease.

Melting and boiling points. The melting points and boiling points of alkaline earth metals are low but higher than those of alkali metals. However, down the group, the melting points and boiling points do not show any regular trend.

Density. The alkaline earth metals are denser than the alkali metals due to smaller size and better packing of the atoms in the crystal lattice.

Nature of bonds formed. They form ionic compounds because they have low ionization enthalpy. Their compounds are less ionic because their ionization enthalpies are higher than those of corresponding alkali metals. Be forms covalent compounds because of its small size and high ionization enthalpy.

Flame colouration. Like alkali metal salts, alkaline earth metals and their salts also impart a characteristic colour to the flame.

Calcium $\qquad $ Strontium $\qquad $ Barium $\qquad $ Radium

Brick red $\qquad $ Crimson $\qquad $ Apple green $\quad $ Crimson

Be and $\mathrm{Mg}$ being smaller in size, their electrons are strongly held by the nucleus. As a result, they need huge amounts of energy for excitation of electrons to higher energy levels which is not available in the Bunsen flame. So they do not impart any colour to the flame.

Reducing character.

The reducing character of alkaline earth metals increases down the group as the ionization enthalpy of the metals decreases or the electrode potential becomes more and more negative down the group.

Chemical Reactivity.

Since ionization enthalpies of alkaline earth metals are higher and their corresponding alkali metals, so alkaline earth metals are less reactive than alkali metals.

Since ionization enthalpies of alkaline earth metals are higher and their electrode potentials are less negative than the corresponding alkali metals so alkaline earth metals are weaker reducing agents than alkali metals.

Action of air.

Alkaline earth metals on heating in air form a mixture of oxides and nitrides.

$$ 2 \mathrm{M}+\mathrm{O} _{2} \xrightarrow{\Delta} 2 \mathrm{MO} ; \quad 3 \mathrm{M}+\mathrm{N} _{2} \xrightarrow{\Delta} 3 \mathrm{M} _{3} \mathrm{~N} _{2} $$

Action of water.

The reactivity of alkaline earth metals towards water increases down the group from $\mathrm{Mg}$ to $\mathrm{Ba}$, i.e., $\mathrm{Mg}<\mathrm{Ca}<\mathrm{Sr}<\mathrm{Ba}$. Be does not react with hot water or even steam, $\mathrm{Mg}$ reacts with hot water but $\mathrm{Ca}$, $\mathrm{Sr}$ and $\mathrm{Ba}$ react vigorously even with cold water evolving $\mathrm{H} _{2}$ gas.

$$ \mathrm{M}+2 \mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{M}(\mathrm{OH}) _{2}+\mathrm{H} _{2} \text { where } \mathrm{M}=\mathrm{Mg}, \mathrm{Ca}, \text { Sr or } \mathrm{Ba} $$

$\mathrm{Be}(\mathrm{OH}) _{2}$ is amphoteric while all other alkaline earth metal hydroxides are basic in nature; their solubility in water and basicity increases down the group from $\mathrm{Mg}(\mathrm{OH}) _{2}$ to $\mathrm{Ba}(\mathrm{OH}) _{2}$.

Action of hydrogen.

All the alkaline earth metals except Be combine with hydrogen directly on heating to form metal hydrides, $\mathrm{MH} _{2}$. The stabilities of these hydrides decreases down the group from $\mathrm{Be}$ to $\mathrm{Ba}$ because their lattice enthalpies decrease down the group as the size of the metal cation increases. Both $\mathrm{BeH} _{2}$ and $\mathrm{MgH} _{2}$ are covalent, electron deficient molecules. They have polymeric structures involving hydrogen bridges while all other hydrides $\left(\mathrm{CaH} _{2}, \mathrm{SrH} _{2}\right.$ and $\left.\mathrm{BaH} _{2}\right)$ are ionic and liberate $\mathrm{H} _{2}$ on treatment with $\mathrm{H} _{2} \mathrm{O}$.

$$ \mathrm{BeCl} _{2}+\mathrm{LiAlH} _{4} \longrightarrow 2 \mathrm{BeH} _{2}+\mathrm{LiCl}+\mathrm{AlCl} _{3} $$

Action of halogens.

When heated with halogens $\left(\mathrm{F} _{2}, \mathrm{Cl} _{2}, \mathrm{Br} _{2}\right.$ and $\left.\mathrm{I} _{2}\right)$. All alkaline earth metals form halides of the general formula, $\mathrm{MX} _{2}$.

(i) Due to small size and high polarizing power of $\mathrm{Be}^{2+}$ ions, beryllium halides are covalent, soluble in organic solvents, hygroscopic and fume in air due to hydrolysis. The halides of other alkaline earth metals are fairly ionic and their ionic character increases as the size of the metal increases.

(ii) Being electron deficient, $\mathrm{BeCl} _{2}$ has a polymeric structure (with chlorine bridges) in the solid state but exists as a dimer in the vapour state and as a monomer at $1200 \mathrm{~K}$.

(iii) $\mathrm{BeF} _{2}$ is very soluble in water due to the high hydration enthalpy of $\mathrm{Be}^{2+}$ ions. The fluorides of other alkaline earth metals are insoluble in water.

Formation of sulphates.

Alkaline earth metals form sulphates of the type $\mathrm{MSO} _{4}$. These are formed by action of $\mathrm{H} _{2} \mathrm{SO} _{4}$ on corresponding oxides, hydroxides or carbonates.

(i) The solubility of sulphates in water decreases down the group from Be to Ba. This is mainly due to the reason that as the size of the cation increases, the enthalpy of hydration decreases while the lattice enthalpy remains about the same. Thus, $\mathrm{BeSO} _{4}$ and $\mathrm{MgSO} _{4}$ are highly soluble, $\mathrm{CaSO} _{4}$ is sparingly soluble while the sulphates of $\mathrm{Sr}$, Ba and Ra are virtually insoluble.

(ii) These sulphates decompose on heating forming the corresponding oxides and $\mathrm{SO} _{3}$. The temperature of decomposition of these sulphates, however, increases as the electropositive character of the metal or the basicity of the metal hydroxide increases down the group.

$\mathrm{M} _{2} \mathrm{SO} _{4} \xrightarrow{\Delta} \mathrm{M} _{2} \mathrm{O}+\mathrm{SO} _{3}$

Nitrates

The nitrates are made by dissolution of the carbonates in dilute nitric acid.

All of them decompose on heating to give the oxide like lithium nitrate.

$2 \mathrm{M}\left(\mathrm{NO} _{3}\right) _{2} \longrightarrow 2 \mathrm{MO}+4 \mathrm{NO} _{2}+\mathrm{O} _{2}$

Solutions in liquid ammonia

They dessolve is liquid ammonia to give deep blue black solutions forming ammoniated Cations and ammoniated electrons.

$$ \mathrm{M}+(\mathrm{x}+\mathrm{y}) \mathrm{NH} _{3} \longrightarrow\left[\mathrm{M}\left(\mathrm{NH} _{3}\right) _{\mathrm{x}}\right]^{2+}+2\left[\mathrm{e}\left(\mathrm{NH} _{3}\right) _{\mathrm{y}}\right]^{-} $$

Carbonates and bicarbonates.

i. The bicarbonates of the alkaline earth metals do not exist in the solid state but are known only in solution. On heating, these bicarbonates decompose forming carbonates with the evolution of $\mathrm{CO} _{2}$.

ii. The solubilities of the carbonates decrease as we move down the group from $\mathrm{Be}$ to $\mathrm{Ba}$, i.e., $\mathrm{BeCO} _{3}>\mathrm{MgCO} _{3}>\mathrm{CaCO} _{3}>\mathrm{SrCO} _{3}>\mathrm{BaCO} _{3}$. This is mainly due to the reason that as the size of the cation increases, the lattice enthalpy of their carbonates remains almost unchanged (like that of sulphates) but the enthalpies of hydration of the cations decrease.

iii. The carbonates of all alkaline earth metals decompose on heating to form the corresponding metal oxide and $\mathrm{CO} _{2}$. The temperature of decomposition of these carbonates, however, increases down the group as the electropositive character of the metal or the basicity of metal hydroxide increases from $\mathrm{Be}(\mathrm{OH}) _{2}$ to $\mathrm{Ba}(\mathrm{OH}) _{2}$.

Beryllium differs from rest of the elements of its family because of

(i) exceptionally small atomic and ionic size, (ii) high ionization enthalpy and (iii) absence of $\mathrm{d}$-orbitals in its valence shell.

Beryllium shows diagonal relationship with aluminium because both these elements have same electronegativity and same polarizing power.

1. Both $\mathrm{BeCl} _{2}$ and $\mathrm{AlCl} _{3}$ act as strong Lewis acids and are used as Friedel crafts catalysts.

2. Both $\mathrm{BeCl} _{2}$ and $\mathrm{AICl} _{3}$ have chlorine bridged structures in the vapour phase.

3. Both the metals dissolve in strong alkalies to from soluble complexes: beryllates $\left[\mathrm{Be}(\mathrm{OH}) _{4}\right]^{2-}$ and aluminates $\left[\mathrm{Al}(\mathrm{OH}) _{4}^{-}\right]$

4. The oxides and hydroxides of both $\mathrm{Be}$ and $\mathrm{Al}$ are amphoteric and dissolve in sodium hydroxide solution as well as in hydrochloric acid.

$$ \begin{aligned} \mathrm{BeO}+2 \mathrm{HCl} \longrightarrow \mathrm{BeCl} _{2}+\mathrm{H} _{2} \mathrm{O} ; \mathrm{BeO}+2 \mathrm{NaOH} \longrightarrow \mathrm{H}^{2} \longrightarrow & \begin{array}{l} \mathrm{Na} _{2} \mathrm{BeO} _{2}+\mathrm{H} _{2} \mathrm{O} \\ \text { sod. berrylate } \end{array} \\ \mathrm{Al} _{2} \mathrm{O} _{3}+6 \mathrm{HCl} \longrightarrow 2 \mathrm{AlCl} _{3}+\mathrm{H} _{2} \mathrm{O} ; \mathrm{Al} _{2} \mathrm{O} _{3}+2 \mathrm{NaOH} \longrightarrow & 2 \mathrm{NaAlO} _{2}+\mathrm{H} _{2} \mathrm{O} \end{aligned} $$

Compounds of Ca

Calcium oxide (CaO)

Calcium oxide (Ca0) is also called quick lime. It is obtained when limestone is heated to $1070-2070 \mathrm{~K}$. On adding water, quick lime gives a hissing sound and forms calcium hydroxide known as slaked lime while the filtered and clear solution is called lime water. Chemically both are calcium hydroxide.

$$ \mathrm{CaCO} _{3} \xrightarrow{\Delta} \mathrm{CaO}+\mathrm{CO} _{2} $$

Reactions of Ca0

$\mathrm{CaO}+ \mathrm{SiO}_{2} \xrightarrow{\Delta} \mathrm{CaSiO} _{3}$

$\qquad \qquad \qquad $ calcium silicate

$\mathrm{CaO}+2 \mathrm{H}_{2}O \longrightarrow \mathrm{Ca} (OH) _{2}$

$\quad \qquad \qquad \qquad $ slaked lime

$\mathrm{CaO}+\mathrm{CO} _{2} \longrightarrow \mathrm{CaCO} _{3}$

$\qquad \qquad \qquad $ calcium carbonate

$\mathrm{CaO}+2 \mathrm{HCl} \longrightarrow \mathrm{CaCl} _{2}+\mathrm{H} _{2} \mathrm{O}$

The suspension of slaked lime in water is called milk of lime while the filtered and clear solution is known as lime water.

Reactions of $\mathrm{Ca}(\mathrm{OH}) _{2}$

1. $\mathrm{Ca}(\mathrm{OH}) _{2} \xrightarrow{>700K} \mathrm{O00K} \mathrm{CaO}+\mathrm{H} _{2} \mathrm{O}$

2. $2 \mathrm{Ca}(\mathrm{OH}) _{2}+2 \mathrm{Cl} _{2} \longrightarrow \mathrm{CaCl} _{2}+\mathrm{Ca}(\mathrm{OCl}) _{2}+2 \mathrm{H} _{2} \mathrm{O}$

3. When carbon dioxide is passed through lime water, it turns lime water, milky due to the formation of insoluble carbonate.

$\mathrm{Ca}(\mathrm{OH}) _{2}+\mathrm{CO} _{2} \rightarrow \mathrm{CaCO} _{3} \downarrow+\mathrm{H} _{2} \mathrm{O}$

$ \qquad \quad \qquad \qquad \qquad \text { (milkiness) }$

On passing excess of $\mathrm{CO} _{2}$, the precipitate of calcum carbonate dissolve to form soluble calcium bicarbonate and hence milkiness disappears. $\mathrm{CaCO} _{3}+\mathrm{CO} _{2}+\mathrm{H} _{2} \mathrm{O} \longrightarrow \mathrm{Ca}\left(\mathrm{HCO} _{3}\right) _{2}$

If this clear solution of calcium bicarbonate is heated, the solution again turns milky due to the decomposition of calcium bicarbonate back to calcium carbonate.

$\mathrm{Ca}\left(\mathrm{HCO} _{3}\right) _{2} \xrightarrow{\Delta} \mathrm{CaCO} _{3} \downarrow+\mathrm{CO} _{2}+\mathrm{H} _{2} \mathrm{O}$

4. $\mathrm{Ca}(\mathrm{OH}) _{2}+2 \mathrm{HCl} \longrightarrow \mathrm{CaCl} _{2}+\mathrm{H} _{2} \mathrm{O}$

5. $\mathrm{Ca}(\mathrm{OH}) _{2}+\mathrm{SO} _{3} \longrightarrow \mathrm{CaSO} _{4}+\mathrm{H} _{2} \mathrm{O}$

Plaster of Paris

Plaster of Paris is $\mathrm{CaSO} _{4} \cdot 1 / 2 \mathrm{H} _{2} \mathrm{O}$. It is prepared by heating gypsum to $393 \mathrm{~K}$. On mixing with $1 / 3$ of its weight of water, it sets into a hard mass of gypsum $\left(\mathrm{CaSO} _{4} \cdot 2 \mathrm{H} _{2} \mathrm{O}\right)$ with slight expansion $(1 %)$ in volume. it is used for making casts for statues and busts, as decorative material and in dentistry. It is also used in surgical bandages used for plastering broken or fractured bones of the body and for making black board chalks. When heated above $393 \mathrm{~K}$, it loses its water of crystallization and forms anhydrous calcium sulphate. It is called dead burnt plaster since it loses the property of setting with water.

$\underset{\text { gypsum }}{2 \mathrm{CaSO} _{4}} \cdot 2 \mathrm{H} _{2} \mathrm{O} \xrightarrow{393 \mathrm{~K}} \underset{\text { plaster of Paris }}{2 \mathrm{CaSO} _{4} .1 / 2 \mathrm{H} _{2} \mathrm{O}+3 \mathrm{H} _{2} \mathrm{O}}$

$2 \mathrm{CaSO} _{4} \cdot 1 / 2 \mathrm{H} _{2} \mathrm{O} \longrightarrow 393 \mathrm{~K} 2 \mathrm{CaSO} _{4}+\mathrm{H} _{2} \mathrm{O}$

Plaster of paris

Anhyd. calcium sulphate

[Dead burnt plaster]

Cement

Cement is essentially a finely powdered mixture of calcium silicates and aluminates along with small quantities of gypsum which sets into a hard stone like mass when treated with water. The average composition of Portland cement is :

$\mathrm{CaO}=50-60 %, \mathrm{MgO}=2-3 %, \mathrm{SiO} _{2}=20-25 %, \mathrm{Fe} _{2} \mathrm{O} _{3}=1-2 %, \mathrm{Al} _{2} \mathrm{O} _{3}=5-10 %$ and $\mathrm{SO} _{3}=1-2 %$

The important constituents 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 $\mathrm{Ca} _{3} \mathrm{Al} _{2} \mathrm{O} _{6}=11 %$.

When water is added to cement, an exothermic reaction occurs. During this process, cement reacts with water to form a gelatinous mass which slowly sets into a hard mass having three-dimensional networkstructure involving $-\mathrm{Si}-0-\mathrm{SI}-$ and $-\mathrm{Si}-0-\mathrm{Al}-$ chains.

Biological importance of S-block elements

A typical $70 \mathrm{Kg}$ contains $90 \mathrm{~g}$ of $\mathrm{Na}$ and $170 \mathrm{~g}$ of $\mathrm{K}$. $\mathrm{Na}^{+}$ions are found in blood plasma, they participate in transmission of nerve signals. The $\mathrm{K}^{+}$ions are the most abundant cations within cell fluids, they activate many enzymes, participate in oxidation of glucose to produce ATP.

A Na-K pump operates across the cell membrane.

An adult body contains about $25 \mathrm{~g}$ of $\mathrm{Mg}$ and $1200 \mathrm{~g}$ of Ca about $99 %$ of body calcium is present in bones and teeth. It also plays an important roles in neuromuscular function inter neuronal transmission, cell membrane integrity and blood coagulation.

Solved Problems

1. Among the following alkali metals, the correct order of increasing atomic radius is

1. $\mathrm{Sr}<\mathrm{Ca}<\mathrm{Ba}<\mathrm{Ra}$

2. $\mathrm{Sr}<\mathrm{Ca}<\mathrm{Ra}<\mathrm{Ba}$

3. $\mathrm{Ca}<\mathrm{Ba}<\mathrm{Sr}<\mathrm{Ra}$

4. $\mathrm{Ca}<\mathrm{Sr}<\mathrm{Ba}<\mathrm{Ra}$

2. What is the best description of the change that occurs when $\mathrm{Na} _{2} \mathrm{O}$ is dissolved in water?

1. Oxide ion donates a pair of electrons

2. Oxidation number of oxygen increases

3. Oxidation number of sodium decreases

4. Oxide ion accepts sharing in a pair of electrons

3. The product obtained on heating $\mathrm{LiNO} _{3}$ will be:

1. $\mathrm{Li} _{3} \mathrm{~N}+\mathrm{O} _{2}$

2. $\mathrm{Li} _{2} \mathrm{O}+\mathrm{NO}+\mathrm{O} _{2}$

3. $\mathrm{LiNO} _{2}+\mathrm{O} _{2}$

4. $\mathrm{Li} _{2} \mathrm{O}+\mathrm{NO} _{2}+\mathrm{O} _{2}$

4. The hydration energies of Group II ions decreases as one goes down in the group from $\mathrm{Be}^{2+}$ to $\mathrm{Ba}^{2+}$ Hence

1. the solubilities of their sulphates, hydroxides and fluorides decrease on descending the group

2. the solubilities of the their sulphates and hydroxides decrease, while that of fluorides increase, on descending the group

3. the solubilities of their sulphates decrease, while that of hydroxides and fluorides increase, on descending the group

4. the solubilities of their sulphates and fluorides decrease, while that of hydroxides increase, on descending the group

Practice Questions

1. The sequence of ionic mobility in aqueous solution is

(a) $\mathrm{Rb}^{+}>\mathrm{K}^{+}>\mathrm{Cs}^{+}>\mathrm{Na}^{+}$

(b) $\mathrm{Na}^{+}>\mathrm{K}^{+}>\mathrm{Rb}^{+}>\mathrm{Cs}^{+}$

(c) $\mathrm{K}^{+}>\mathrm{Na}^{+}>\mathrm{Rb}^{+}>\mathrm{Cs}^{+}$

(d) $\mathrm{Cs}^{+}>\mathrm{Rb}^{+}>\mathrm{K}^{+}>\mathrm{Na}^{+}$

2. Alkali metals have negative reduction potential and hence they behave as

(a) oxidizing agents

(b) Lewis bases

(c) reducing agents

(d) electrolytes

3. Which of the following statements is false regarding alkali metals?

(a) Alkali metals are soft and can be cut with the help of knife

(b) Alkali metals do not occur in free state in nature

(c) Alkali metals are highly electropositive

(d) Alkali metal hydrides are covalent in character

4. In the Solvay process of manufacture of sodium carbonate, the by-product is

(a) $\mathrm{NH} _{4} \mathrm{Cl}$

(b) $\mathrm{NaHCO} _{3}$

(c) $\mathrm{CaCl} _{2}$

(d) $\mathrm{CO} _{2}$

5. The set representing the correct order of ionic radii is

(a) $\mathrm{Na}^{+}>\mathrm{Li}^{+}>\mathrm{Mg}^{2+}>\mathrm{Be}^{2+}$

(b) $\mathrm{Li}^{+}>\mathrm{Na}^{+}>\mathrm{Mg}^{2+}>\mathrm{Be}^{2+}$

(c) $\mathrm{Mg}^{2+}>\mathrm{Be}^{2+}>\mathrm{Li}^{+}>\mathrm{Na}^{+}$

(d) $\mathrm{Li}^{+}>\mathrm{Be}^{2+}>\mathrm{Na}^{+}>\mathrm{Mg}^{2+}$

6. Which pair of the following chlorides do not impart colour to the flame?

(a) $\mathrm{BeCl} _{2}$ and $\mathrm{SrCl} _{2}$

(b) $\mathrm{BeCl} _{2}$ and $\mathrm{MgCl} _{2}$

(c) $\mathrm{CaCl} _{2}$ and $\mathrm{BaCl} _{2}$

(d) $\mathrm{BaCl} _{2}$ and $\mathrm{SrCl} _{2}$

7. Which of the following oxides is most acidic in nature?

(a) $\mathrm{Be} 0$

(b) $\mathrm{MgO}$

(c) $\mathrm{CaO}$

(d) $\mathrm{Ba} 0$

8. The solubility in water of sulphates down the $\mathrm{Be}$ group is : $\mathrm{Be}>\mathrm{Mg}>\mathrm{Ca}>\mathrm{Sr}>\mathrm{Ba}$. This is due to

(a) increase in melting point

(b) increase in ionization energy

(c) decreasing lattice energy

(d) decreasing hydration enthalpy

9. Choose the incorrect statement in the following:

(a) $\mathrm{BeO}$ is almost insoluble but $\mathrm{BeSO} _{4}$ is soluble in water.

(b) $\mathrm{BaO}$ is soluble but $\mathrm{BaSO} _{4}$ is insoluble in water.

(c) Lil is more soluble than $\mathrm{Kl}$ in ethanol

(d) Both Liand Mg from solid hydrogen carbonates.

10. Which of the following is the weakest base?

(a) $\mathrm{Ca}(\mathrm{OH}) _{2}$

(b) $\mathrm{KOH}$

(c) $\mathrm{LiOH}$

(d) $\mathrm{Sr}(\mathrm{OH}) _{2}$