S-Block

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Intro of group-1 & group-2

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Anomalous behaviour of first elements.

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In these anomalous properties there is ample the second element of the next group. Thus Lithium shows similarities to magnesium and beryllium shows similarities to aluminium in many of their properties, and this similarity is known as diagonal relationship.

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Cause of Diagonal relationship of first elements.

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Electronic configuration of alkali metals

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All the alkali metals have one valence electron, ns1 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 M+ ions. Hence they are never found in free state in nature.

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Atomic and ionic radii of alkali metals

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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 (M+) 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 Li to Cs.

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Basic trends of periodic table

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Ionization enthalpy of alkali metals

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The ionization enthalpies of the alkali metals are considerably low and decrease down the group from 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.

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Hydration enthalpy of alkali metals

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

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

Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., LiCl· 2H2O

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Physical properties of alkali metals

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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.

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Flame colouration

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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. Alkali metals can therefore, be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy.

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Photoelectric effect of alkali metals

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These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron. Alkali metals (except Li) exhibit photoelectric effect. This property makes caesium and potassium useful as electrodes in photoelectric cells.

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Reactivity of alkali metals towards air

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The alkali metals are highly reactive due to their large size and low ionization enthalpy. The reactivity of these metals increases down the group.

The alkali metals tarnish in dry air due to the formation of their oxides which in turn react withmoisture to form hydroxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O2– ion is stable only in the presence of large cations such as K, Rb, Cs.

4 Li + O2 -----> 2Li2O (oxide)

2 Na + O2 -------> Na2O2 (peroxide)

M + O2 ---------> MO2 (superoxide)

(M = K, Rb, Cs)

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, Li3N as well. Because of their high reactivity towards air and water,alkali metals are normally kept in kerosene oil.

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Reactivity of alkali metals towards water

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The alkali metals react with water to form hydroxide and dihydrogen.

2M + 2H2O → 2M+ + 2OH- + H2

(M = an alkali metal)

This tendency of forming Hydroxide with water depends upon their electrode potential (E0) and this electrode potential is a measure of the tendency of an element to lose electron in the aqueous solution. Thus, more negative is the electrode potential, higher is the tendency of the element to lose electron and hence stronger is the reducing agent, with decrease in negative value of electrode potential, the reaction with water become more vigorous.

It may be noted that although lithium has most negative E0 value, its reaction with water is less vigorous than that of sodium which has the least negative E0 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.

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Reactivity of alkali metals towards dihydrogen

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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.

2M + H2 ------>2(M+) (H−)

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Reactivity of alkali metal towards halogens

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The alkali metals readily react vigorously with halogens to form ionic halides, M+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 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.

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Reducing nature of alkali metals

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The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The standard electrode potential (E0) which measures the reducing power represents the overall change :

2M(s)→ M(g) [sublimation enthalpy]

M(g) → M (g) + e [ionization enthalpy]

M (g) + H2O →M (aq) [hydration enthalpy]

With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative E0 value and its high reducing power.

Higher the negative value of E0 value, more reducing will be the ionic species.

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Solutions of alkali metals in liquid Ammonia

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The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature.

M + (x + y)NH3 → [M(NH3)x ]+ + [e (NH3)y]-

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.

M+ (am) + e- + NH3(I) → MNH2 + ½H2 (g)

(where ‘am’ denotes solution in ammonia.)

In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic. Blue solutions are paramagnetic.

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Uses of alkali metals

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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 Na/Pb alloy needed to make PbEt4 and PbMe4. These organo lead 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 and also acts an an absorbent of carbon dioxide. It is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising photoelectric cells.

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Oxides and hydroxides of alkali metals.

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On combustion in excess of air, lithium forms mainly the oxide, Li2O (plus some peroxide Li2O2), sodium forms the peroxide, Na2O2 (and some superoxide NaO2) whilst potassium, rubidium and caesium form the superoxides, MO2. Under appropriate conditions pure compounds M2O, M2O2 and MO2 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. 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. 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. The higher oxides generally act as good oxidizing agent, Sodium Peroxide is widely used as an oxidising agent in inorganic chemistry. The hydroxides of alkali metals are strong bases. The basic character of alkali metal hydroxides increases in going down the group.

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Halides of alkali metals

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The alkali metal halides, MX, (X=F,Cl,Br,I) are all ionic, high melting, colourless crystalline solids. Reactivity of alkali metals with particular halogen increases from Lithium to caesium. On the other hand, reactivity of halogens with particular alkali metal and decreases from fluoride to iodide. 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 Δf H0 values for fluorides become less negative as we go down the group, whilst the reverse is true for Δf H0 for chlorides, bromides and iodides. For a given metal Δf H0 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 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; LiCl is soluble in pyridine also.

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Salt of oxo acids of alkali metals

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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,H2CO3 (OC(OH)2; sulphuric acid, H2SO4(O2S(OH)2).

The alkali metals form salts with all the oxo-acids. They are generally soluble in water and thermally stable. Their carbonates (M2CO3) and in most cases the hydrogen carbonates (MHCO3) also are highly stable to heat. As the electropositive character increases down the group, the stability of the carbonates and hydorgen carbonates increases. Lithium carbonate is not so stable to heat; lithium being very small in size polarises a large (CO3)2– ion leading to the formation of more stable Li2O and CO2. Its hydrogen carbonate does not exist as a solid.

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ANOMALOUS PROPERTIES OFLITHIUM

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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.

Points of Difference betweenLithium 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, Li2O and the nitride, Li3N unlike other alkali metals.

(iii) LiCl is deliquescent and crystallises as a hydrate, LiCl.2H2O whereas other alkali metal chlorides do not form hydrates.

(iv) Lithium hydrogen carbonate is not obtained in the solid form while all other elements form solid hydrogen carbonates.

(v) Lithium unlike other alkali metals forms no ethynide on reaction with ethyne.

(vi) Lithium nitrate when heated gives lithium oxide, Li2O, whereas other alkali metal nitrates decompose to give the corresponding nitrite.

(vii) LiF and Li2O are comparatively much less soluble in water than the corresponding compounds of other alkali metals.

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Points of Similarities betweenLithium and Magnesium

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The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes :

atomic radii,

Li = 152 pm,

Mg = 160 pm;

ionic radii :

Li+ = 76 pm,

Mg2+= 72 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,Li3N and Mg3N2, by direct combination with nitrogen.

(iii) The oxides, Li2O and 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 CO2. Solid hydrogen carbonates are not formed by lithium and magnesium.

(v) Both LiCl and MgCl2 are soluble in ethanol.

(vi) Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates, LiCl·2H2O and MgCl2·8H2O.

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Some trends of physical properties of alkali metals

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Front 25

Some important points about group 1

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