Alkali Metals

The group I comprising Li, Na, K, Rb, Cs & Fr are commonly called alkali metals. These are called alkali metals because hydroxides of these metals are strong alkali. For example NaOH and KOH

Francium is radioactive and has a very short life (half life of 21 minutes), therefore very little is known about it.

Physical Properties of Alkali Metals

The general electronic configuration of alkali metals may be represented by [noble gas]  where n = 2 to 7











Atomic Number







Electronic Configuration







Atomic Mass







Metallic radius (pm)







Ionic radius (M+/pm)







Ionization enthalpy

(kJ mol–1)













Electro negativity

(Pauling Scale)







Density/g cm–3 (at 298K)







Melting point/K







Boiling point/K







E°(V) at 298K for

M+(aq) + e→ M(s)







Occurrence in







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

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.

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.

Atomic and Ionic radii


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

Li> Na> K+ > Rb+ > Cs+              

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

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.

Ionization Energy (Ionization enthalpy) of alkali metals


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.

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 → M+ + e-         


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  ions as the magnitude of second ionization energy is very high.

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.


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.  

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.


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.

Density of Alkali Metals


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.

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.


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.

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).

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.

Flame colouration

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









Crimson Red

Golden Yellow

Pale Violet


Sky Blue







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.




Flame Test for Lithium

Flame Test for Sodium

Flame Test for Potassium

Refer to the follwoing video for flame test

Chemical Properties of Alkali Metals

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

  • Low value of first ionization energy

  • Large size

  • Low heat of atomization

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 + O2 → 2M2O (Where M = Li, Na, K, Rb, Cs)

When heated with excess of air, lithium forms normal oxide,Li2O ; sodium forms peroxide, Na2O2, whereas potassium rubidium and caesium form superoxides having general formula MO2

4Li O2 → 2Li2O ( Lithium oxide)                     

2Na + O\overset{575 K}{\rightarrow} Na2O2 ( Sodium peroxide)      

K + O2 → KO2 ( Potassium Superoxide)       

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, Li+  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, O2-2. On the other hand,  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 O2-2 . The larger K+ , Rb+ , and Cs+ ions have still weaker positive field around them which cannot prevent even peroxide O2-2 ion,  to combine with another oxygen atom to form superoxide O-2 .  

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 + H2 \overset{\Delta }{\rightarrow}(M = Li, Na, K, Rb or Cs)
Some important features of hydrides are
  • 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.

  • These hydrides react with water to form corresponding hydroxides and hydrogen gas.

 LiH + H2O → LiOH + H2    

NaH + H2O → NaOH + H2                                          

  • These hydrides are strong reducing agents and their reducing    nature increases down the group.

  • Alkali metals also form complex hydrides such as LiAIH4 and NaBH4 which are good reducing agents.

  • All these hydrides react with proton donors such as water, alcohols, gaseous ammonia and alkynes liberating H2 gas.   

NaH(S) + ROH (l) \overset{\Delta }{\rightarrow} RONa (S) + H2 (g)           

NaH (S) + NH3 (g) \overset{\Delta }{\rightarrow} NaNH2 (S) + H2 (g)                       

2KH (S) + HC ≡ CH (g) \overset{\Delta }{\rightarrow} KC ≡ CK (S) + 2H2 (g)                      

  • 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.

Reaction of sodium with waterReaction 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 + H2O → 2NaOH + H2

2Na + 2HCI → 2NaCI + H2              

2Na + 2HC ≡ CH → 2NaC≡ CH + H2               
                                Sodium acetylide               

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.

Refer to the following video for reaction of alkali metals with water

Reaction with Halogens

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

2M + X2 → 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 F2 to I2 .

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 + ( x + y ) NH3   → M+  ( NH3 )x + e-(NH3)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 + 2NH3 → 2MNH2 + H2     
                      (Mital amide)          

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.

Reaction with Sulphur and Phosphorus

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

16 Na + S8 \overset{\Delta }{\rightarrow} 8Na2S
                  Sodium sulphide          
12Na + P4 \overset{\Delta }{\rightarrow} 4Na3P              
                   Sodium Phosphide              

Reaction with Mercury

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

Na + Hg → Na (Hg)
              Sodium amalgam              

Question 1: On prolonged exposure to air, sodium finally changes to

a. Na2CO3

b. Na2O

c. NaOH

d. NaHCO3

Question 2: Which reaction below is explosive?

a. Sodium with water

b. Magnesium with water

c. Beryllium with water

d. Caesium with water

Question 3: Group 1 elements are also known as

a. Alkali metals

b. Alkali earth emtals

c. Halogens

d. Chalogens

Question 4:  Which of the following elements have highets first ionization enthalpy?

a. Na

b. Li

c. K

d. Rb









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