Alkali metals

 ALKALI METALS BY B.T.SIR

The group IA elements like 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.

 Characteristic of alkali metals

Physical Properties of Alkali Metals

1.Physical state : 

• 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. 
• They are 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.

2.Atomic and Ionic Radii:

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.{alertWarning}

3.Hydration energy :

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.

Note:

# Hydration energy ∝ 

Hydration energy ∝ oxidation potential ∝ tendency to lose electron ∝ reducing properties # SIZE OF HYDRATED METALLIC CATION ∝

HYDRATION ENERGY ∝

ଵ

௠௢௕௜௟௜௧௬

# Mobility ∝ conductivity

# In aq. state Fr is the strongest reducing agent while in aq. state

Li is the strongest reducing agent.

4.Ionization Energy (Ionization enthalpy or

ionization potential):

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. This is due to

increase in the atomic radius on moving downward a group.eg.

Li>Na>K>Rb>Cs>Fr

5.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⁻

electropositive ion.

10.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.

11.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.

12.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.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

1.Reaction with Oxygen

The alkali metals tarnish in air due to the formation of an oxide,

hydroxide and carbonate 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 + O2 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.

NOTE ;

Normal oxide dissolves in water to give alkali.eg.

Na₂O + H₂O → 2NaOH

Peroxide reacts with water or dil.acid to give hydrogen peroxide.

Na2O2 + 2H₂O → 2NaOH + H2O2.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 (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.

 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.

3.Reaction with Water

 Alkali metals dissolve into water to give alkali and hydrogen gas

with the evolution of heat due to which it may catches fire also. It

also reacts with other compounds containing acidic hydrogen

atoms such as hydrogen halides (HX) and acetylene (C2H2) and

liberate H2 gas

2Na + H₂O → 2NaOH + H₂+ heat

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.

Except LiOH all hydroxides are thermally stable.

2LiOH → Li₂O + H₂O

On moving downward a group basic character of hydroxide

increases .

4.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 Fr

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 .

5.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.

6.Reaction with Sulphur and Phosphorus

Alkali metals react with sulphur and phosphorus on heating to form sulphides and phosphides respectively. 16 Na + S8 8Na2S Sodium sulphide 12Na + P4 4Na3P Sodium Phosphide 

7.Reaction with Mercury Alkali metals combine with mercury to form amalgams. The reactions is highly exothermic in nature Na + Hg → Na (Hg) Sodium amalgam 

8.Carbonates and Bicarbonates:all alkali metas form carbonate of the type M2CO3.(M = Li ,Na,K,Rb,Cs) Alkali metal carbonates except lithium carbonate are ionic, thermally stable, and water-soluble. These properties increase from lithium carbonate to carbonate. Li2CO3 → Li2O + CO2 All other carbonates fuses on heating. Bicarbonates, except lithium bicarbonate, are solid, watersoluble and on heating liberate carbon dioxide. 2NaHCO3 →Na2CO3 + Li2O + CO2 9.Sulphates: All alkali metals form sulphate of the type M2SO4.Sulphates except lithium are soluble in water. Sulphates can be reduced by carbon to sulphide. Forms double salts with trivalent metal sulphates (alum). 

10.Nitrates: They form nitrate of the type MNO3.They are soluble in water and on heating except lithium nitrate decomposes to nitrites. 2MNO3 → 2MNO2 + O2 2NaNO3 ∆→ 2NaNO2 + O2 Lithium nitrate on heating gives 4LiNO3 ∆→ 2Li2O + 4NO2 + O

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