Alkali Metals or Group IA Elements
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. 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.
2.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 selectrons. 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.
3.Hydration energy :
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.
Note: # Hydration energy ∝ 1/catonic size # Hydration energy ∝ oxidation potential ∝ tendency to lose electron ∝ reducing properties # SIZE OF HYDRATED METALLIC CATION ∝ HYDRATION ENERGY ∝ 1/mobility # 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⁻
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 ions as the magnitude of second ionization energy is very high.
6.Reducing Properties
The alkali metals have low values of reduction potential and therefore have a strong tendency to lose electrons and act as good reducing agents. The reducing character increases from sodium to caesium in solid state. However lithium is the strongest reducing agent in aq. medium.
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. But in aq. medium Li + ion has high hydration energy that favours the easy removal of electron and hence,Li is the strongest reducing agent in aq.medium.
7.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.
8.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 unusual increase in atomic size of potassium.
9.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.
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
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 .
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 (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
2Na + H₂O → 2NaOH + H₂ + heat
2Na + 2HCI → 2NaCI + H₂
2Na + 2HC ≡ CH → 2NaC≡ CH + H2 Sodium acetylide
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 where 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
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 + O2
Diagonal Relationship of Lithium and Magnesium
The main points of similarity:
- Both have almost similar electronegatives.
- Both Li and Mg are quite hard. They are harder and lighter than other elements in their respective groups.
- Both LiOH and Mg(OH)2 are weak bases.
- Both form ionic nitrides when heated in atmosphere of nitrogen. 6Li+N2
- The hydroxides of both lithium and magnesium decompose upon heating
- Both lithium and ma
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