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What are the common physical and chemical features of alkali metals ?

Physical properties

(1) Metallic character: Alkali metals are highly electropositive in nature and hence, they are typical metals. The metallic character is due to the low values of ionisation energies and consequently, they have tendency to lose the valence electrons.

(2) Low ionization energy: The first ionization energies of alkali metals are quite low as compared to the elements of the other groups belonging to the same period. The reason is that atoms of alkali metals are of large sizes. Therefore, the outermost electron is far away from the nucleus and can be easily removed. Within the group, ionization energies of alkali metals decrease as we move down the group.

(3) Low melting and boiling points: The melting and boiling points of alkali metals are very low because of large size of their atoms due to which inter- particle forces are very weak in them. The melting and boiling points decrease on going down the group (Lithium to Cesium) as the charge density decreases because of the increase in size of the mono-valent cation.

(4) Low electro- negativity values: Alkali metals have low values of electro- negativity. They have very little tendency to attract the shared pair of electrons towards themselves. The electro negativity values of alkali metals decrease as we move down in the group from Li to Cs e.g. Li(1.9) K(0.8), Cs(0.7).

(5) Soft metals: All the alkali metals are soft and can be cut with the help of a knife. Softness of alkali metal is due to weak metallic bonding in them as the result of large size of the atoms. As we move down the group, metallic bonding weakens and therefore, softness increases

(6) Density: Alkali metals have low density due to the large size of metals atoms.

(7) Oxidation state: The alkali metals exhibit oxidation state of +1 in their compounds and strongly electropositive in character. The electropositive character increases from lithium down to caesium in the group.

(8) Characteristic flame colour: Alkali metals impart characteristic colour to the flame, e.g. Li(red), Na(yellow), K(lilac), Rb (violet) and Cs(blue).

Chemical Properties

(1) Decomposition of water: The alkali metals decompose water at the ordinary temperature giving out hydrogen.

2Li + 2H2O 2LiOH + H2 (slow)

2Na + 2H2O 2NaOH + H2 (quick)

2K + 2H2O 2KOH + H2 (more vigorously)

(2) Reaction with oxygen: The alkali metals readily burn in oxygen or air to form their oxides, M2O.

4Li + O2®  2 Li2O

4Na + O2 2Na2O

(3) Combination with halogens: The akali metals burn in halogens (F, Cl. Br, I) forming their halides (MX).

2Na +Cl2 2NaCl

2K + Br2 2KBr

(4) Combination with hydrogen: Alkali metals combines with hydrogen to give white crystalline hydrides (MH). The ease with which the alkali metals form these hydrides decrease from Li to Cs, because the electropositive character of the metals increases in the same order. The hydrides are ionic in character, because hydrogen is present as H- in them. Thus these are represented and M+ H-.

2M + H2 2M+ H-[M = Li, Na, K, Rb, Cs]

(5) Combination with Sulphur: Alkali metals react with sulphur upon heating to the corresponding sulphides.

16Na + S8
8 Na2S
                   Sodium sulphide

(6) Solubility in liquid ammonia: Alkali metals dissolve in liquid ammonia to produce blue coloured solutions which conduct electricity.

(7) Reaction with acids: Alkali metals react with strong and weak acids (organic and inorganic) to give their respective salts and hydrogen gas is liberated.

M + H+ M+ + H2(g)
         Metal salt

(8) With Phosphorous: The alkali metals combine with phosphorous forming metals phosphide, M3P

12Na + P4
             Sodium phosphide

12K + P4
4K3 P
              Sodium phosphide
12K+ P4
            Potassium phosphide


Discuss the general characteristics and gradation in properties of alkaline earth metals.

Electronic configuration: Their general electronic configuration is ns2. They have 2 electrons in their outermost shells.

Oxidation state: They can lose two electrons to acquire nearest noble gas configuration; therefore their oxidation state is +2.

Atomic and ionic sizes: The atomic and ionic size goes on increasing down the group due to decrease in effective nuclear charge. Their atomic and ionic radii are smaller than the corresponding alkali metals in corresponding periods due to increased effective nuclear charge in their atoms.

Physical appearance: These metals in general are silvery white, lustrous and relatively soft but harder than the alkali metal. Be and Mg appear to be greyish.

Ionization enthalpies: The first ionization enthalpies of alkaline earth metals are higher than alkali metals because of smaller atomic size and higher effective nuclear charge. Second ionization enthalpy is smaller than that of alkali metals because alkali metals acquire noble gas configuration after losing one electron. Ionization enthalpies of alkaline earth elements decrease down the group due to increase in atomic size.

Melting and boiling points: They have fairly higher melting and boiling points than corresponding alkali metals. The trend is not systematic mg has lowest melting point.

Electrical and thermal conductivity: They are good conductors of heat and electricity.

Flame colour: Be and Mg does not impart any colour to the flame because electrons are strongly bound and therefore, cannot be excited by flame.

Ca, imparts brick red, Sr imparts crimson red, Ba gives apple green colour because electrons get excited to higher energy levels and when they come back to lower energy level, energy is radiated in form of visible light.

Density: Alkaline earth elements are denser than corresponding alkali metal elements which is due to greater strength of metallic bond and more closely packed crystal structure. However, the variation in density is irregular. Ca has lowest density.

Metallic character: They are less electropositive than alkali metals due to higher ionization enthalpies.

Metallic character increases down the group due to decrease in ionization enthalpy.

Reactivity with air and water: The alkaline earth metals are less reactive than alkali metals. Be and Mg are kinetically inert of O2 and H2O due to formation of oxide layer on its surface. Be does not react with water or steam even at red hot and does not get oxidized in air below 873 K.

Powdered Be burns on ignition to form BeO and Be3N2

2Be + O2
2 BeO

3Be + N2 Be3N2

Mg is more electropositive and burns in air with dazzling light forming MgO and Mg3N2.

2Mg +O2 2MgO

3Mg + M2 Mg3N2

Ca, Sr and Ba readily react with oxygen to give oxides. Calcium forms oxide whereas Sr and Ba form peroxide. They react with nitrogen to form nitrides.

2Ca + O2 2CaO

3Cs +N2 Ca3N2

Sr + O2 SrO2

Ba + O2 BaO2

Mg reacts with hot water, Ca, Ba, Sr, react with cold water vigorously.

Mg + H2O (hot) MgO + H2

Ca + 2H2O Ca(OH)2+ H2

Sr + 2H2O Sr(OH)2+ H2

Ba + 2H2O Ba(OH)2+ H2

Reaction with halogens: Group 2 elements react with halogens at increased temperature to from halides.

Be + Cl2 BeCl2

Mg + Cl2 MgCl2

Ca + Cl2 CaCl2

Action with acids: The alkaline earth metals readily react with acids to from salts and liberate H2 gas.

Be + 2HCl BeCl2 + H2

Mg + 2HCl MgCl2 + H2

Mg + 2HNO3(5%) Mg (NO3)2 + H2

Mg + H2SO4(dil) MgSO4 + H2

Reaction with NH3 : Alkaline earth metals dissolve in liquid NH3 to give deep blue black solutions due to solvated electron.

Reducing power: They are weaker reducing agents than alkali metals. Reducing power goes on increasing down the group due to decrease in standard reduction potential.

Be has less negative value of reduction potential due to high enthalpy of atomization and ionization inspite of higher hydration enthalpies.

Reaction with H2: All metals combine with H2 to from hydrides except Be.

Ca+ H2 CaH2 (Hydrolith)
                    Calcium hydride

Mg + H2 MgH2

BeH2 is prepared by reduction of BeCl2 with LiAlH4.

2BeCl2 + LiAlH4 2BeH2+LiCl + AlCl3

They are ionic hydrides except BeH2 which is covalent.

CaH2 + H2O Ca2+(aq) +2OH-(aq) + 2H2(g)

They are basic in nature.


Why are alkali metals not found in nature ?

Alkali metals are highly reactive due to low ionisation enthalpy and strong electropositive character. They do not occur in free or native state and are always combined with other elements. As a result, alkali metals are not found in nature.


Find out the oxidation state of sodium in Na2O2.


        x        -1

     Na2      O2

2x + 2(-1) = 0

or          x = +1.


Explain why is sodium less reactive than potassium.

This is mainly due to low ionisation enthalpy of sodium as compared to potassium. With the result, the potassium is more electropositive and stronger reducing agent than sodium. Thus, sodium is less reactive than potassium.


Compare the alkali metals and alkaline earth metals with respect to (i) ionisation enthalpy (ii) basicity of oxides and (iii) solubility of hydroxides.



Alkali metals

Alkaline metals

i) Ionization energy

Ionization energy is low, but only +1 state is stable.

Ionization energy is very high and +2 state is stable.

ii) Basicity of the oxides Oxides of alkali metals are highly basic, the basic strength increases from Li2O to Cs2O. Beryllium oxide (BeO) is amphoteric but oxides of other alkaline earth metals are basic. The basic strength increases from MgO to BaO.
iii) Solubility of hydroxides Hydroxides of alkali metals are highly soluble in water. Hydroxides of alkaline earth metals are less soluble in water.


In what ways lithium shows similarities to magnesium in its chemical behaviour?

(1) Atomic radius of lithium is 155 pm. It is comparable to that of magnesium (160pm).

(2) The ionic sizes of Li+ (r= 60pm) and Mg2+ (r= 65pm) are nerly equal.

(3) Lithium reacts with nitrogen to give lithium nitride Li3N. Magnesium also reacts with nitrogen to give magnesium nitride, Mg3N2.

(4) Lithium forms only monoxide, Li2O. magnesium also prefers to form only the monoxide, MgO.

(5) Lithium hydroxide, alike magnesium hydroxide, is sparingly soluble in water and behaves as a weak base.

(6) Lithium chloride, like magnesium chloride, separates out from aqueous solution as hydrated crystal, e.g. LiCl.2H2O and MgCl2.6H2O.

(7) Lithium chloride, like magnesium chloride is deliquescent.


Explain why alkali and alkaline earth metals cannot be obtained by chemical reduction methods?

Alkali and alkaline earth metals are very strong reducing agents. So their oxides or halides cannot be reduced by any other elements/ compounds chemically.


Why are potassium and caesium, rather than lithium used in photoelectric cells?

In a photoelectric cell, electrons are knocked out of the surface of a metal under the influence of striking photons carrying the required energy (threshold energy E0 or more). Now, lithium and sodium due to their small atomic sizes have high ionization energies and photoelectric effect does not take place in the visible region of light. However, both potassium and cesium have comparatively bigger sizes and the electrons can be knocked out more easily. They, therefore, exhibit photoelectric effect in visible light.


When an alkali metal dissolves in liquid ammonia, the solution can acquire different colours. Explain the reasons for this type of colour change.

The alkali metals dissolve in liquid ammonia to give blue to bronze colour.

The colour is attributed to the presence of solvated electrons e.g. [e(NH3)4]- in dilute solution.

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

However in concentrated solutions, the ammoniated metal ions are bound by the free unpaired electrons which give bronze colour.

The blue solutions are paramagnetic, whereas the bronze coloured solutions are diamagnetic.


Beryllium and magnesium do not give colour to flame whereas other alkaline earth metals do so. Why ?

Chlorides of alkaline earth metals, except that of Be and Mg, produce characteristic colour to flame due to easy excitation of electrons to higher energy levels. Beryllium and magnesium atoms, due to their small size, bond their electrons more strongly, i.e., their ionisation energies are high. Hence they possess high excitation energy and are not excited by the energy of the flame to higher energy state with the result no colour is produced by the flame.


Discuss the various reactions that occur in the Solvay process.

When carbon dioxide is passed into a concentrated solution of brine saturated with ammonia, ammonium bicarbonate is produced.


CO2 + H2O H2CO3

                        Ammonium bicarbonate

The ammonium bicarbonate then reacts with common salt forming sodium bicarbonate.

NH4CO3 + NaCl NaHCO3 + NH4Cl
                             Sodium bicarbonate

Sodium bicarbonate being slightly soluble (in presence of sodium ions) gets precipitated. The precipitated sodium bicarbonate is removed by filtration and changed into sodium carbonate by heating.

2NaHCO3 Na2CO3 + H2O +CO2

the mother liquor remaining after the precipitation of sodium bicarbonate contains ammonium chloride. This is then heated by steam with milk of lime to regenerate ammonia which can be as one of the raw materials.

2NH4Cl + Ca(OH)2
CaCl2 + 2H2O + 2NH3

Lime is obtained by heating lime stone

CaO + CO2.


Potassium carbonate cannot be prepared by Solvay process. Why ?

Potassium carbonate cannot be prepared by Solvay process because potassium bicarbonate is too soluble to be precipitated by the addition of ammonium bicarbonate to a saturated solution of potassium chloride.


Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher temperature?

Strong polarizing action of small Li+ ion distorts the electron cloud on the nearby oxygen atom of the large CO32- ion to such and extent that the C-O bond gets weakened and Li-O bond becomes stronger which ultimately leads to the decomposition of lithium carbonate to oxide and carbon dioxide.

Li2O + CO2

Replacement of the larger carbonate ion by smaller oxide ions leads to increase lattice energy and thus favours the decomposition of Li2CO3.


Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals. (a) Nitrates (b) Carbonates (c) Sulphates.


Alkali metals


Thermal stability

1. Nitrates

Soluble in organic solvents due to their ionic nature.

When heated give their oxides.

4LiNO3 2 Li2O + 4NO2 +O2

2NaNO3 2NaNO2+ O2

2. Carbonates They are generally soluble in water. The carbonates are generally highly stable to heat and thus thermally stable.
3. Sulphates Sparingly soluble in H2O. Potassium and sodium slats are readily soluble. Conform to thermal stability.


Alkaline earth metals


Thermal stability

1. Nitrates


Dissolve in water sparingly soluble and possess crystalline six molecules of H2O.

They decompose upon heating giving oxide.

2M(NO3)2 2MO + 4NO2+O2

2. Carbonates Sparingly soluble in water. Thermally unstable which increases with increasing cationic size.
3. Sulphates Readily soluble in H2O due to their greater hydration energy. Thermally stable.



Starting with sodium chloride how would you proceed to prepare (i) sodium metal (ii) sodium hydroxide (iii) sodium peroxide (iv) sodium carbonate

(1)Na is prepared from NaCl by the following method:

Electrolysis : NaCl Na+ +Cl-

At cathode: Na++e- Na (sodium)

At anode:     Cl- Cl +e-

                   Cl+Cl Cl2

(2) Sodium hydroxide is prepared by carrying out the electrolysis of the aqueous solution of sodium chloride either in Nelson's cell or Castner -Kellner cell.

Preparation of sodium peroxide: Sodium chloride is first converted in sodium by electrolytic reduction. The metals are then heated with excess of oxygen at about 573K in an atmosphere free from moisture and carbon dioxide to form sodium peroxide.

2Na + O2
                sodium peroxide

Sodium carbonate is prepared by the following processes:

2Na + 2H2O 2NaOH + H2

2NaOH + CO2 Na2CO3 + H2O
                    (sodium carbonate)


What happens when (i) magnesium is burnt in air (ii) quick lime is heated with silica (iii) chlorine reacts with slaked lime (iv) calcium nitrate is heated ?

(i) Magnesium burns in air with dazzling white light.
5Mg + O2+ N2 2MgO + Mg3N
                             (in air)

(ii) CaO + SiO2 CaSiO3
                                                        Calcium silicate

(iii) Ca(OH)2 + Cl2 CaOCl2 + H2O
                             Bleaching powder

(iv) 2Ca(NO3)2 2CaO + 4NO2+O2.


Describe two important uses of each of the following : (i) caustic soda (ii) sodium carbonate (iii) quicklime.

(i) Caustic soda

(a) It is used in petroleum refining and purification of bauxite.

(b) It is used in the manufacture of soap, Paper, artificial silk and a number of chemicals.

(ii) Sodium carbonate

(a) It is used in the manufacture of soap, glass, paper, borax, caustic soda etc.

(iii) Quick lime

(a) It is used in the manufacture of sodium carbonate from caustic soda.

(b) It is used in the manufacture of dye stuffs.

(c) It is used in the manufacture of sodium carbonate from caustic soda.


Draw the structure of (i) BeCl2 (vapour) (ii) BeCl2 (solid).

(i) Structure of BeCl2 in vapour phase:

(ii)Structure of BeCl2 in solid phase:


The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.

The hydroxides and carbonates of sodium and potassium are easily soluble in water, because they have low lattice energies owing to their smaller charge density and larger atomic size.

Polar water molecules can easily break the electrostatic force of attraction in their ionic crystal lattice. Moreover, these metals have heats of hydration.

On the other hand hydroxides of Mg and Ca are sparingly soluble in water, because they have higher lattice energies owing to their large charge density and smaller atomic size than sodium and potassium.

Polar molecules can thus break only partially the electrostatic forces of attraction in their ionic crystal lattices. Moreover, these metals possess low heat of hydration.


Describe the importance of the following : (i) limestone (ii) cement (iii) plaster of paris.

(i) Importance of limestone

(a) It is used in the manufacture of quick lime.

(b) It is used as a building material in the form of marble.

(c) It is used in the manufacture of quick lime.

(d) It is also used as a raw material for the manufacture of sodium carbonate in solvay process.

(e) It is used as a constituent of toothpaste.

(ii) Importance of Cement

(a) It has become so important next to iron and steel it can be called as a commodity of national necessiry for any country.

(b) It is used in concrete and reinforced concrete.

(c) It is used in plastering and in the construction of bridges, dams and buildings.

(iii) Importance of Plaster of Paris

It is used for producing moulds for pottery, ceramics etc.

(b) It is used for making statues, models and other decorative materials.

(c) It is used in surgical bandages known as plasters for setting broken and fractured bones in the body.


Why are lithium salts commonly hydrated and those of the other alkali ions usually anhydrous?

Since the hydration tendency depends upon charge to radius ratio (q/r); the hydration tendency decreases from Li+ to Cs+. Thus Li+ in aqueous soluiton is very strongly hydrated. Thus lithium salts are commonly hydrated and those of the other alkali ions usually anhydrous.


Why is LiF almost insoluble in water whereas LiCl soluble not only in water but also in acetone?

The low solubility of LiF is due to its high lattice enthalpy. LiCl has much higher solubility in water. This is due to the small size of Li+ ion and much higher hydration energy.


Explain the significance of sodium, potassium, magnesium and calcium in biological fluids.

Significance of sodium and potassium in biological fluids: K+ and Na+ cations are present in the red blood cells. The ratio of K+ to Na+ ions in mammals such as human beings, rabbits, rats and horses is 7:1. In cats and dogs the ratio is 1:15. In order to establish this ratio called concentration gradient in the cell, work has to be done. Biologists have suggested different mechanisms involving sodium pump and potassium pump for this purpose. The cation gradients control the development and the functioning of the nerve cells. During the state of rest the potential of the nerve cells is linked to the K+ ion concentration across the membrane. During the activation of the nerve cells, a chemical called acetyl choline is released near its end plate and membrane potential is discharged. This discharge is transmitted through the length of the nerve cell by an electric pulse. Thus it is clear that Na+ and K+ ions have a significant role in biological fluids.

Significance of magnesium and calcium in biological fluids:

(i) Mg2+ ions are present in chlorophyll which is the green colouring matter present in plants. Chlorophyll absorbs light from sun and carries the process of photosynthesis in plants.

(ii) Ca2+ ions occur as phosphates in the bones of both human beings and animals. These ions also play an important role in the miscle contraction. The malnutrition in children is mainly due to the deficiency of the Ca2+ion.

(iii) Both Mg2+ and Ca2+ ions catalyze the formation of pyrophosphate linkages which control the various biological systems. The pyrophosphates undergo hydrolysis with the release of energy. This process is controlled by Ca2+ ions.


What happens when

(i) sodium metal is dropped in water ?

(ii) sodium metal is heated in free supply of air ?

(iii) sodium peroxide dissolves in water ?

(i) 2Na(s) +2H2O (l) 2NaOH(aq) + H2(g).

(ii)2Na(s) + O2(g)
Na2O2(s) (excess from air sodium peroxide free from CO2)

(iii) 2Na2O2(s) +2H2O(l) 4NaOH (aq) + O2(g).


(b) Lithium is the only alkali metal to form a nitride directly.
E0 for M2+ (aq) + 2e– _ M(s) (where M = Ca, Sr or Ba) is nearly constant.


(a) This shows that the enthalpy of ionization is increasing of Cs+ to Na +ion it means that the ionising strength increases from Na+ to Cs+

(b) Lithium is the only alkali metal which reacts directly with nitrogen in air to form the nitride, Li3N

6Li + N2 2Li3N
               (Lithium nitride)

This is due to its small size and greater polarizing power. Its reducing power and E0 are high.

From the electrode potential (E0) of alkali metals which measuring the reducing power represents another change from Ms to M+(aq). It depends upon three parameters: (a) sublimation, (b) ionization and (c) hydration enthalpies.

The equation
M2+ +2e-

shows that the standard electrode potential of Ca, Sr and Ra are always the same.


State as to why

(i) a solution of Na2CO3 is alkaline ?

(ii) alkali metals are prepared by electrolysis of their fused chlorides?

(iii) sodium is found to be more useful than potassium ?


(i) The solution of sodium carbonate is alkaline in nature. Because when sodium carbonate is treated with water, it gets hydrolysed to form alkaline

CO32- + H2O HCO3- + OH-

(ii) The alkali metals are very reactive and are strong reducing agents so they cannot be extracted by usual methods. This is due to the following reasons:

(a) They are strong reducing agents.

(b) They cannot be extracted by electrolysis of their aqueous solutions. Because the metals formed will react immediately react with water forming their hydroxides. Therefore these metals are generally isolated by the electrolysis if their fused metals chlorides.

(iii) The sodium metals are more useful than potassium due to the following reasons:

(a) Industrial uses of sodium metal reflects its strong reducing power, about 60% of world production of sodium are used to make tetraethyl lead PbEt4 for gasoline antiknocks.

(b) Sodium is also used for the preparation of sodium compounds such as peroxide, amide and sodium cyanide.

(c) It is also used in dye industry.

(d) It is also used for detecting the presence of nitrogen, sulphur and halogens in organic compounds. It is widely used as sodium amalgam as a reductant.

(e) Liquid sodium is used in nuclear reactors as a medium for heat exchange. So sodium is a more useful metal than potassium.


Write balanced equations for reactions between

(a) Na2O2 and water

(b) KO2 and water

(c) Na2O and CO2

(a) 2Na2O2+2H2O 4NaOH +O2

(b) 2KO2+ 2H2O 2KOH + H2O2+O2

(c) Na2O +CO2 Na2CO3


How would you explain the following observations?

(i) BeO is almost insoluble but BeSO4 in soluble in water,

(ii) BaO is soluble but BaSO4 is insoluble in water,


(i) The solubility depends upon two factors:
(a) Lattice energy, (b) hydration energy.
BeO oxide is insoluble in water because its lattice energy is large than its hydration energy. On the other hand BeSO4 are readily soluble because the hydration energy of Be2+ion are higher which over come the lattice energy factor and therefore BeSO4 readily dissolves in water.

(ii) Barium oxide (BaO) is soluble in water because its hydration energy is greater than its lattice energy. On the other hand the lattice energy of BaSO4 is very high due to its bivalent charges so that hydration energy released is not sufficient to overcome the lattice energy and to break the bond and it remains insoluble.


Which of the alkali metal is having least melting point ?
(a) Na (b) K (c) Rb (d) Cs

(d) Cs.


Which one of the alkaline earth metal carbonates is thermally the most stable ?
(a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO

(d) BaCO3.

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