Group 1 Elements: Alkali Metals

INTRODUCTION

The elements in the long form of the periodic table has been divided into four blocks, namely

s, p, d & f blocks. The elements of group I & II receive their last electron in s-orbital. So they are called as s – block elements.

The metals Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr) which have one electron in their outermost shell belongs to group I and are called alkali metals as they react with water to form hydroxides which are strong bases or alkalies.

The elements of group II are Berryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), barium (Ba) and radium (Ra) which have two electrons in their outermost shell. All these elements are also metallic in nature and are commonly known as alkaline earth metals with the exception of beryllium. Because of their low density, alkali metals and alkaline earth metals are called lighter metals. Both alkali and alkaline earth metals are highly reactive and hence do not occur in free state but found in combined state. Whereas alkali metals mostly occur as halides, oxides, silicates, borates and nitrates, alkaline earth metals mainly occur as silicates, carbonates, sulphates and phosphates. Some alkali & alkaline earth metals occur abundantly in nature. Calcium is the fifth, magnesium is the sixth, sodium is seventh and potassium is eight barium is the fourteenth and strontium is the fifteenth most abundant element by weight in the earth's crust. Sodium and magnesium are also present in sea water brine wells and few salt lakes.

Anamolous behavour of first element

The first element of a group differs considerably from the rest of the elements of the same group. This anomolous behaviour is due to

(i) Smaller size of their atoms

(ii) Their higher ionization energies

(iii) Their higher electronegativites

(iv) Absence of vacant d – orbitals in their valence shell

(v) High polarizing power of its cation.

Thus Li differes from the rest of alkali metals (Na, K, Rb & Cs) and Be differs from rest of the alkaline earth metals (Mg, Ca, Sr & Ba)

Diagonal relationships

On moving diagonally some members show similar properties with the members of next higher group which is particularly seen in the elements of second and third periods of the periodic table. However the similarities shown are far less pronounced than the similarities with in a group.

The main reasons for the diagonal relationship are

(i) Similarity in electropositive character

The electropositive character decreases along a period, but increases down a group. Hence on moving diagonally the two opposing trend partially cancels out. As a result diagonally related elements have similar electropositive character and hence exhibit similar properties.

(ii) Similarity in polarizing power

On moving along a period from left to right, the charge on the ions increases while ionic size decreases, hence polarizing power increases. On moving down the group, the ionic size increases and hence polarizing power decreases. On moving diagonally, these two trends partially cancel out. As a result thus diagonally related elements have same polarizing power and thus exhibit similar properties.

(iii) Similarity in atomic or ionic radii

The atomic and ionic radii decrease across a period and increase in a group. Evidently on moving diagonally, the two trends partially cancel out. As a result, diagonally related elements have similar atomic & ionic radii and hence have similar properties.

Alkali Metals

The group I comprising Li, Na, K, Rb, Cs & Fr are commonly called alkali metals. Francium is radioactive and has a very short life (half life of 21 minutes), therefore very little is known about it.

Electronic Configuration

The general electronic configuration of alkali metals may be represented by [noble gas] $n{s}^1$ where n = 2 to 7

Table – I: [Some Physical properties of group 1 elements (alkali metals)]

* ppm (parts per million) ** percentage by weight.

Illustration 1. Why are Group 1 elements called alkali metals?

Solution: The Group 1 elements are called alkali metals because they form water soluble hydroxides.

PHYSICAL PROPERTIES

(i) All the alkali elements are silvery white solid. These are soft in nature and can be cut with the help of knife except the lithium. When freshly cut, they have a bright lusture which quickly fades due to surface oxidation. These are highly malleable and ductile. The silvery luster of alkali metals is due to the presence of highly mobile electrons of the metallic lattice. There being only a single electron per atom, the metallic bonding is not so strong. As the result, the metals are soft in nature. However, the softness increases with increase in atomic number due to continuous decrease in metallic bond strength on account of an increase in atomic size.

(ii) Atomic and Ionic radii

The atoms of alkali metals have the largest size in their respective periods. The atomic radii increase on moving down the group among the alkali metals.

Reason

On moving down the group a new shell is progressively added. Although, the nuclear charge also increases down the group but the effect of addition of new shells is more predominant due to increasing screening effect of inner filled shell on the valence s-electrons. Hence the atomic size increases in a group.

Alkali metals change into positively charged ions by losing their valence electron. The size of cation is smaller than parent atom of alkali metals. However, within the group the ionic radii increase with increases in atomic number.

The alkali metal ions get extensively hydrated in aqueous solutions. Smaller the ion more is the extent or degree of hydration. Thus, the ionic radii in aqueous solution follow the order

$L{i}^{+}>N{a}^{+}>K^{+}>R{b}^{+}>C{s}^{+}$

The charge density on Li⁺ is higher in comparison to other alkali metals due to which it is extensively hydrated.

(iii) Ionization Energy (Ionization enthalpy)

The first ionization energy of the alkali metals are the lowest as compared to the elements in the other group. The ionization energy of alkali metals decreases down the group.

Reason

The size of alkali metals is largest in their respective period. So the outermost electron experiences less force of attraction from the nucleus and hence can be easily removed.

The value of ionization energy decreases down the group because the size of metal increases due to the addition of new shell along with increase in the magnitude of screening effect.

(iv) Oxidation State

The alkali metals show +1 oxidation state. The alkali metals can easily loose their valence electron and change into uni-positive ions

$M\to M^{+}+e^{-}$

Reason

Due to low ionization energy, the alkali metals can easily lose their valence electron and gain stable noble gas configuration. But the alkali metals cannot form $M^{2+}$ ions as the magnitude of second ionization energy is very high.

(v) Reducing Properties

The alkali metals have low values of reduction potential (as shown in table-I) and therefore have a strong tendency to lose electrons and act as good reducing agents. The reducing character increases from sodium to caesium. However lithium is the strongest reducing agent.

Reason

The alkali metals have low value of ionization energy which decreases down the group and so can easily lose their valence electron and thus act as good reducing agents.

(vi) Melting and Boiling Points

The melting and boiling points of alkali metals are very low because the intermetallic bonds in them are quite weak. And this decreases with increase in atomic number with increases in atomic size.

(vii) Density

The densities of alkali metals are quite low as compared to other metals. Li, Na and K are even lighter than water. The density increases from Li to Cs.

Reason

Due to their large size, the atoms of alkali metals are less closely packed. Consequently have low density. On going down the group, both the atomic size and atomic mass increase but the increase in atomic mass compensates the bigger atomic size. As a result, the density of alkali metals increases from Li to Cs. Potassium is however lighter than sodium. It is probably due to an unusal increase in atomic size of potassium.

(viii) Nature of bond formed

All the alkali metals form ionic (electrovalent) compounds. The ionic character increases from Li to Cs because the alkali metals have low value of ionization energies which decreases down the group and hence tendency to give electron increases to form electropositive ion.

(ix) Conductivity

The alkali metals are good conductors of heat and electricity. This is due to the presence of loosely held valence electrons which are free to move throughout the metal structure.

(x) Photoelectric Effect

Alkali metals (except Li) exhibit photoelectric effect (A phenomenon of emission of electrons from the surface of metal when light falls on them). The ability to exhibit photoelectric effect is due to low value of ionization energy of alkali metals. Li does not emit photoelectrons due to high value of ionization energy.

(xi) Flame colouration

The alkali metals and their salts impart a characteristic colour to flame

On heating an alkali metal or its salt (especially chlorides 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 electrons return to their ground states, they emit extra energy in form of radiations which fall in the visible region thereby imparting a characteristic colour to the flame.

Illustration 2. What is the most reactive alkali metal and why?

Solution: The most reactive alkali metal is cesium due to its lowest first ionization enthalpy and lowest electronegativity.

CHEMICAL PROPERTIES

The alkali metals are highly reactive metals and the reactivity increases down the group. The reactivity is due to-

(a) low value of first ionization energy

(b) large size

(c) low heat of atomization

(i) Reaction with Oxygen

The alkali metals tarnish in air due to the formation of an oxide or hydroxide on the surface. Alkali metals when burnt in air form different kinds of oxides. For example the alkali metals on reaction with limited quantity of oxygen form normal oxides of formula, M₂O

$4M+O_2\to 2M_{2}O$ (Where M = Li, Na, K, Rb, Cs)

When heated with excess of air, lithium forms normal oxide, $L{i}_{2}O$; sodium forms peroxide,$N{a}_2{O}_2$, whereas potassium rubidium and caesium form superoxides having general formula $M{O}_2$

$4Li+O_2\to 2L{i}_{2}O\left(Lithiumoxide\right)$

$2Na+O_2\stackrel{575K}{\to }N{a}_2{O}_2\left(Sodiumperoxide\right)$

$K+O_2\to K{O}_2\left(PotassiumSuperoxide\right)$

Thus the reactivity of alkali metals with oxygen increases down the group. Further, the increasing stability of peroxide or superoxide, as the size of the metal ion increases is due to the stabilization of larger anions by larger cation through higher lattice energies.

Due to small size, $L{i}^{+}$ has a strong positive field around it which attracts the negative charge so strongly that it does not permit the oxide anion $O^{-2}$ to combine with another oxygen to form peroxide ion,$O_2^{-2}$. On the other hand, $N{a}^{+}$ ion because of its large size than Li⁺ ion has comparatively weaker positive field around it which cannot prevent $O^{-2}$ ion to combine with another oxygen to form peroxide ion$O_2^{-2}$. The larger $K_{}^{+}$, $R{b}_{}^{+}$, and$C{s}_{}^{+}$ ions have still weaker positive field around them which cannot prevent even peroxide ion, $O_2^{2-}$ to combine with another oxygen atom to form superoxide $O_2^{-}$.

(ii) Reaction with Hydrogen

Alkali metals react with dry hydrogen at about 673K to form colourless crystalline hydrides. All the alkali metal hydrides are ionic solids with high melting points.

$2M+H_2\stackrel{\Delta }{\to }2MH$ (M = Li, Na, K, Rb or Cs)

Some important features of hydrides are

(a) The stability of hydrides decrease from Li to Cs. It is because of the fact that M-H bond becomes weaker due to increase in the size of alkali metals down the group.

(b) These hydrides react with water to form corresponding hydroxides and hydrogen gas.

$LiH+H_{2}O\to LiOH+H_2$

$NaH+H_{2}O\to NaOH+H_2$

(c) These hydrides are strong reducing agents and their reducing nature increases down the group.

Alkali metals also form complex hydrides such as $LiAl{H}_4$ and $NaB{H}_4$ which are good reducing agents.

(d) All these hydrides react with proton donors such as water, alcohols, gaseous ammonia and alkynes liberating $H_2$ gas.

$NaH\left(s\right)+ROH\left(\ell \right)\stackrel{\Delta }{\to }RONa\left(s\right)+H_2\left(g\right)$

$NaH\left(s\right)+N{H}_3\left(g\right)\stackrel{\Delta }{\to }NaN{H}_2\left(s\right)+H_2\left(g\right)$

$2KH\left(s\right)+HC\equiv CH\left(g\right)\stackrel{\Delta }{\to }KC\equiv CK\left(s\right)+2H_2\left(g\right)$

The order of reactivity of the alkali metals towards hydrogen decreases as we move down the group from Li to Cs. This is due to the reason that the lattice energies of these hydrides decreases progressively as the size of the metal cation increases and thus the stability of these hydrides decreases from LiH to CsH.

(iii) Reaction with water

The alkali metals are known to have large negative reduction potential values. As a result they can act as better reducing agents as compared to hydrogen. Hence, alkali metals react with water and other compounds containing acidic hydrogen atoms such as hydrogen halides (HX) and acetylene (C₂H₂) and liberate H₂ gas

$2Na+H_{2}O\to 2NaOH+H_2$

$2Na+2HCl\to 2NaCl+H_2$

$2Na+2HC\equiv CH\to \underset{Sodiumacetylide}{2NaC\equiv CH+H_2}$

The reaction becomes more and more violent as we move down the group. Thus, Lithium reacts gently, sodium melts on the surface of water and the molten metal moves around vigorously and may sometimes catch fire. Potassium melts and always catches fire and so are Rb and Cs.

(iv) Reaction with halogens

Alkali metals react vigorously with halogens to form metal halides of general formula MX, which are ionic crystalline solids.

$2M+X_2\to 2MX$

M = Li, Na, K, Rb or Cs and

X = F, Cl, Br or I

Reactivity of alkali metals with particular halogens increases from Li to Cs. On the other hand, reactivity of halogens decreases from $F_2$ to $I_2$.

(v) Solubility in liquid ammonia

All alkali metal dissolve in liquid ammonia giving deep blue solutions which are conducting in nature. These solutions contain ammoniated cations and ammoniated electrons as shown below:

$M+\left(x+y\right)N{H}_3\to M^{+}{\left(N{H}_3\right)}_x+e^{-1}{\left(N{H}_3\right)}_y$

The blue colour of the solution is considered to be due to ammoniated electrons which absorb energy corresponding to red region of the visible light for the their excitation to higher energy levels. The transmitted light is blue which imparts blue colour to the solutions. The electrical conductivity of the solution is due to both ammoniated cations and ammoniated electrons. The blue solution on standing slowly liberates hydrogen resulting in formation of amide :

$2M+2N{H}_3\to \underset{\left(Metalamide\right)}{2MN{H}_2}+H_2$

At concentrations above 3M, the solutions of alkali metals in liquid ammonia are copper-bronze coloured. These solutions contains clusters of metal ions and hence possess metallic lusture. The blue coloured solutions are paramagnetic due to presence of large number of unpaired electrons, but bronze solutions are diamagnetic due to formation of electron clusters in which ammoniated electrons with opposite spin group together

These solutions are stronger reducing agents than hydrogen and hence will react with water to liberate hydrogen.

(vi) Reaction with sulphur and phosphorus

Alkali metals react with sulphur and phosphorus on heating to form sulphides and phosphides respectively.

$16Na+S_8\stackrel{\Delta }{\to }\underset{Sodiumsulphide}{8N{a}_{2}S}$

$12Na+P_4\stackrel{\Delta }{\to }\underset{SodiumPhosphide}{4N{a}_{3}P}$

(vii) Reaction with Mercury

Alkali metals combine with mercury to form amalgams. The reactions is highly exothermic in nature

$Na+Hg\to \underset{Sodiumama\lg am}{Na\left(Hg\right)}$

Illustration 3. What happens when (give chemical equation only):

(i) sodium is exposed to moist air

(ii) sodium reacts with water

Solution: (i) $2Na+\frac{1}{2}{O}_2+H_{2}O+C{O}_2\to N{a}_{2}C{O}_3.H_{2}O$

(ii) $2Na+2H_{2}O\to 2NaOH+H_2+Heat$

General characteristic of the compounds of the alkali metals

(i) Oxides and Hydroxides

All the alkali metals, their oxides, peroxides and superoxides readily dissolve in water to produce corresponding hydroxides which are strong alkalies eg

$2Na+2H_{2}O\to 2NaOH+H_2$

$N{a}_{2}O+2H_{2}O\to 2NaOH$

$N{a}_2{O}_2+2H_{2}O\to 2NaOH+H_2{O}_2$

$2K{O}_2+2H_{2}O\to 2KOH+H_2{O}_2+O_2$

Thus peroxides and superoxides also act as oxidizing agents since they react with $H_{2}O$ forming H₂O₂ and $O_2$ respectively. The hydroxides of all the alkali metals are white crystalline solids. They are strongest of all base and readily dissolve in water with the evolution of much heat.

(a) Basic Strength

The basic strength of these hydroxides increases as we move down the group Li to Cs. The hydroxides of alkali metals behave as strong bases due to their low ionization energies which decrease down the group. The decrease in ionization energies leads to weakening of the bond between metal and hydroxide ion and M – O bond in M – O – H can easily break giving $M^{+}$ and $O{H}^{-}$. This results in the increased concentration of hydroxyl ions in the solution i.e increased basic characters.

(b) Solubility and stability

All these hydroxides are highly soluble in water and thermally stable except lithium hydroxide.

$2LiOH\stackrel{\Delta }{\to }L{i}_{2}O+H_{2}O$

(c) Formation of salts with acids:

Alkali metals hydroxides being strongly basic react with all acids forming salts.

$NaOH+HCl\to \mathrm{NaCl}+H_{2}O$

$2NaOH+H_{2}S{O}_4\to N{a}_{2}S{O}_4+2H_{2}O$

Illustration 4: Name the alkali metal which forms superoxides when heated in excess of air and why?

Solution: Potassium forms superoxides when heated in excess of air. This is due to the stabilization of large size cation by large size anion.

\begin{align}K+{{O}_{2}}\xrightarrow[{}]{{}}K{{O}_{2}}\\\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\text{potassium}\,\,\text{superoxide} \\ \end{align}

Illustration 5. Why lithium forms only lithium oxide and not peroxide or superoxides?

Solution: Due to the small size of lithium, it has a strong positive field around it. On combination with the oxide anion (O^2–), the positive field of lithium ion restricts the spread of negative charge towards another oxygen atom and thus prevents the formation of higher oxides.

(ii) Halides

The alkali metals combine directly with halogens under appropriate conditions forming halides of general formula MX. These halides can also be prepared by the action of aqueous halogen acids (HX) on metals oxides, hydroxides or carbonate.

$M_{2}O+2HX\to 2MX+H_{2}O$

$MOH+HX\to \mathrm{MX}+H_{2}O$

$M_{2}C{O}_3+2HX\to 2MX+C{O}_2+H_{2}O$ (M = Li, Na, K, Rb or Cs)

(X = F, Cl, Br or I)

All these halides are colourless, high melting crystalline solids having high negative enthalpies of formation.

Table – II [Standard enthalpies of formation in (kJ/mol^-1)]

The value decreases in the order: Fluoride > Chloride > bromides > Iodide

Thus fluorides are the most stable while iodides are the least stable.

The trends in melting points, boiling points and solubility of alkali metals halides can be understood in terms of polarization effects, lattice energy and hydration of ions.

(a) Polarization effects

Comparison of ionic and covalent character of alkali metal halides.

When a cation approaches an anion, the electron cloud of the anion is attracted towards the cation and hence gets distorted. This effect is called polarization. The power of the cation to polarize the

anion is called its polarizing power and the tendency of the anion to get polarized is called its polarizability. The greater the polarization produced more is the concentration of the electrons between the two atoms thereby decreasing the ionic character or increasing the covalent character. The covalent character of any compound in general depends upon the following factors.

(i) Size of the cations

Smaller the cation greater is its polarizing power and hence larger is the covalent character. The covalent character decreases as size of cation increases.

LiCl > NaCl > KCl > RbCl > CsCl

Thus LiCl is more covalent than KCl.

(ii) Size of the anion

Larger the anion, greater is its polarizability. This explains the covalent character of lithium halides is in order

LiI > LiBr > LiCl > LiF

(iii) Charge of the ion

Greater the charge on the cation greater is its polarizing power and hence larger is the covalent character. The covalent character of some halides increase in the order

$N{a}^{+}C{l}^{-}<M{g}^{+2}C{l}_2<A{l}^{+3}C{l}_3$

Similarly greater the charge on the anion, more easily it gets polarized thereby imparting more covalent character to the compound formed eg covalent character increase in the order

$NaCl<N{a}_{2}S{O}_4<N{a}_{3}P{O}_4$

Thus the covalent character decreases as the charge of the anion decrease.

(iv) Electronic configuration of the cation

If two cations have the same charge and size, the one with pseudo noble gas configuration i.e having 18 electrons in the outermost shell has greater polarizing power than a cation with noble gas configuration i.e having 8 electrons. For example CuCl is more covalent than NaCl.

(b) Lattice Energies

lattice energy is defined as the amount of energy required to separate one mole of solid ionic compound into its gaseous ions. Evidently greater the lattice energy, higher is the melting point of the alkali metals halide and lower is its solubility in water

Table – III

(c) Hydration Energy

It is the amount of energy released when one mole of gaseous ions combine with water to form hydrated ions.

$M^{+}\left(g\right)+aq\to M^{+}\left(aq\right)+hydrationenergy$

$X^{-}\left(g\right)+aq\to X^{-}\left(aq\right)+hydrationenergy$

Higher the hydration energy of the ions greater is the solubility of the compound in water.

Further the extent of hydration depends upon the size of the ions. Smaller the size of the ion, more highly it is hydrated and hence greater is its hydrated ionic radius and less is its ionic mobility (Conductance).

From above arguments, the melting point and solubility in water or organic solvent of alkali metal halides can be explained

(i) A delicate balance between lattice enthalpy and hydration enthalpy determines the ultimate solubility of a compound in water. For eg. Low solubility of LiF (0.27 g/100 g$H_{2}O$) is due to its high lattice energy $\left(-1005KJmo{l}^{-1}\right)$whereas the low solubility of CsI (44g/100g $H_{2}O$) is due to smaller hydration energy of the two ions (-670 KJ/mol)

(ii) The solubility of the most of alkali metal halides except those of fluorides decreases on descending the group since the decrease in hydration energy is more than the corresponding decrease in the lattice energy.

(iii) Due to small size and high electronegativity, lithium halides except LiF are predominatantly covalent and hence are soluble in covalent solvents such as alcohol, acetone, ethyl acetate, LiCl is also soluble in pyridine. In contrast NaCl being ionic is insoluble in organic solvents.

(iv) Due to high hydration energy of $L{i}^{+}$ ion, Lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice energy.

(v) For the same alkali metal the melting point decreases in the order

fluoride > chloride > bromide > iodide

because for the same alkali metal ion, the lattice energies decreases as the size of the halide ion increases.

(vi) for the same halide ion, the melting point of lithium halides are lower than those of the corresponding sodium halides and thereafter they decrease as we move down the group from Na to Cs.

The low melting point of LiCl (887 K) as compared to NaCl is probably because LiCl is covalent in nature and NaCl is ionic.

Illustration 6. Why are alkali metal halides soluble in water?

Solution: Alkali metal halides are soluble in water due to their high ionic character and low lattice energy.

Salts of oxoacids

Since the alkali metals are highly electropositive, therefore their hydroxides are very strong bases and hence they form salts with all oxoacids $\left(H_{2}C{O}_3,H_{3}P{O}_4,H_{2}S{O}_4,HN{O}_3,HN{O}_{2}etc\right)$. They are generally soluble in water and stable towards heat. The carbonates $\left(M_{2}C{O}_3\right)$ of alkali metals are remarkably stable upto 1273 K, above which they first melt and then eventually decompose to form oxides. $L{i}_{2}C{O}_3$, however is considerably less stable and decomposes readily.

$L{i}_{2}C{O}_3\stackrel{\Delta }{\to }L{i}_{2}O+C{O}_2$

This is presumably due to large size difference between Li⁺ and $C{O}_3^{2-}$ which makes the crystal lattice unstable.

Being strongly basic, alkali metals also form solid bicarbonates. No other metals forms solid bicarbonates though$N{H}_{4}HC{O}_3$ also exists as a solid. Lithium, however does not form solid bicarbonate though it does exist in solution. All the bicarbonate on gentle heating undergo decomposition to form carbonates with the evolution of $C{O}_2$. .

$2MHC{O}_3\stackrel{\Delta }{\to }{M}_{2}C{O}_3+C{O}_2+H_{2}O$

All the carbonates and bicarbonates are soluble in water and their solubilities increase rapidly on descending the group. This is due to the reason that lattice energies decrease more rapidly than their hydration energies on moving down the group.

Illustration 7. Complete and balance the following:

(i) $LiN{O}_3\stackrel{Heat}{\to }$

(ii) $NaN{O}_3\stackrel{Heat}{\to }$

Solution: (i) $4LiN{O}_3\stackrel{\Delta }{\to }2L{i}_{2}O+4N{O}_2+O_2$

(ii) $2NaN{O}_3\stackrel{\Delta }{\to }2NaN{O}_2+O_2$

Anomolous Behaviour of Lithium and its Diagonal Relationship with Magnesium

The properties of lithium are quite different from the properties of other alkali metals. On the other hand, it shows greater resemblance with magnesium, which is diagonally opposite element of it. The main reasons for the anomalous behaviour of lithium as compared to other alkali metals are

(i) The extremely small size of lithium atom and its ion.

(ii) Greater polarizing power of lithium ion$\left(L{i}^{+}\right)$, due to its small size which result in the covalent character in its compounds.

(iii) Least electropositive character and highest ionization energy as compared to other alkali metals.

(iv) Non availability of vacant d-orbitals in the valence shell.

The reason for resemblance of properties of lithium with magnesium is that these two elements have almost same polarizing power.

The following points illustrate the anomalous properties of lithium and its diagonal relationship with magnesium:

(a) The melting point and boiling point of lithium are comparatively high.

(b) Lithium is much harder than the other alkali metals. Magnesium is also hard metal.

(c) Lithium reacts with oxygen least readily to form normal oxide whereas other alkali metals form peroxides and superoxides.

(d) $LiOHlikeMg{\left(OH\right)}_2$ is weak base. Hydroxides of other alkali metals are strong bases.

(e) Due to their appreciable covalent nature, the halides and alkyls of lithum and magnesium are soluble in organic solvents.

(f) Unlike elements of group 1 but like magnesium. Lithium forms nitride with nitrogen.

$6Li+N_2\to 2L{i}_{3}N$

(g) LiCl is deliquescent and crystallizes as a hydrate, $LiCl.2H_{2}O$. Other alkali metals do not form hydrates. $MgC{l}_2$ also forms hydrate, $MgC{l}_2.8H_{2}O$.

(h) Unlike other alkali metals lithium reacts directly with carbon to form an ionic carbide. Magnesium also forms a similar carbide.

(i) The carbonates, hydroxides and nitrates of lithium as well as magnesium decompose on heating.

$L{i}_{2}C{O}_3\to L{i}_{2}O+C{O}_2$

$MgC{O}_3\to MgO+C{O}_2$

$2LiOH\to L{i}_{2}O+H_{2}O$

$Mg{\left(OH\right)}_2\to MgO+H_{2}O$

$4LiN{O}_3\to 2L{i}_{2}O+4N{O}_2+O_2$

$2Mg{\left(N{O}_3\right)}_2\to 2Mg+4N{O}_2+O_2$

The corresponding salts of other alkali metals are stable towards heat.

(j) Lithium nitrate, on heating, decomposes to give lithium oxide, $L{i}_{2}O$ whereas other alkali metals nitrates decomposes to give the corresponding nitrite.

$4LiN{O}_3\to 2L{i}_{2}O+4N{O}_2+O_2$

$2NaN{O}_3\to 2NaN{O}_2+O_2$

$2KN{O}_3\to 2KN{O}_2+O_2$

(k) $L{i}_{2}C{O}_3,LiOH,LiFandL{i}_{3}P{O}_4$ are the only alkali metal salts which are insoluble in water. The corresponding magnesium compounds are also insoluble in water.

(l) Hydrogen carbonates of both lithium and magnesium can not be isolated in solid state. Hydrogen carbonates of other alkali metals can be isolated in solid state.

Illustration 8. Name the chief factor responsible for the anomalous behaviour of lithium.

Solution: The chief factors responsible for the anomalous behaviour of lithium are:

(i) its very small size,

(ii) high electronegativity,

(iii) high ionization enthalpy and

(iv) absence of vacant d-atomic orbital in the valence shell

Difficulties encountered during extraction of alkali metals

Alkali metals, can not be extracted from their ores by the usual methods of extraction of metals because of the following difficulties:

(i) Alkali metals are strong reducing agents and hence can not be extracted by reduction of their oxides or chlorides.

(ii) Alkali metals being highly electropositive can not be displaced from the aqueous solutions of their salts by other metals.

(iii) Alkali metals can not be isolated by electrolysis of the aqueous solution of their salts since hydrogen is liberated at the cathode instead of the alkali metal because the discharge potentials of alkali metals are much higher than that of the hydrogen. However, by using mercury as cathode, the alkali metals can be deposited at the cathode but the alkali metals so deposited readily combines with mercury to form an amalgam from which its recovery is very difficult.

Therefore in view of above difficulties, only successful method is the electrolysis of their molten (fused) salts usually chlorides.