An important philosophy for any relationship to last till eternity is unconditional love, which actually can be elaborated as unconditional giving. If one is willing to be happily stable in a relationship, he should give unconditional love to his counterpart, without a doubt!
Hey, wait! This is definitely not a philosophical lecture that we are proceeding with. In fact, the topic is Alkali Metals. These metals are just like those persons, who happily create relationships by giving their share of and are very much willing to give away their ‘last possessions’ for the sake of love and in order to make their relationship stable!
Sounds interesting! Let's delve deeper to find out more about this beautiful and important set of elements that constitute the 1st group members of the periodic table.
TABLE OF CONTENTS
Alkali metals comprise of Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs) and Francium (Fr). They belong to the Group 1 and s-block of the modern periodic table, occupying the leftmost side of the periodic table.
They are silvery-white coloured metals and are highly reactive. This high reactivity is because of the presence of a single electron in their outermost shell which they can easily lose in order to attain high stability.
Because they react vigorously in the presence of water to form soluble hydroxides known as alkalis, they are collectively known as alkali metals.
The atomic size of alkali metals increases as we move down the group. Every alkali metal has the largest radii than any other element in the corresponding period. Alkali metals readily lose an electron and form cation . The ionic radii of their respective cations also increase down the group.
|Atomic radius (pm)||152||186||227||248||265|
Alkali metals have a large size and volume, which results in a low density. So they are very soft and can be cut with a knife. Lithium, sodium and potassium are lighter than water. Potassium has the lowest density among alkali metals.
Exceptionally, potassium is less dense than sodium. We know that density is inversely proportional to the atomic volume and directly proportional to atomic mass. In the case of potassium, the increase in atomic size (volume) outweighs the pull of the core on the outer shell electron, so potassium is less dense than sodium.
|Density (g cm-3)||0.53||0.97||0.86||1.53||1.9|
The ionisation enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs. This is because the effect of increasing atomic radii outweighs the increasing nuclear charge, and the outermost electron is screened very well from the inner orbital electrons. Hence, it is easy to remove electron from the outermost shell.
|Ionisation energy (KJ mol-1)||520||496||419||403||376|
Hydration enthalpy is defined as the amount of the energy released when one mole of gaseous ions dissolves in a large amount of water i.e. infinitely diluted.
Hydration energy is released after attaining the stability gained because of the electrostatic attraction between water molecules and metal cations (ion-dipole interaction). The more the electrostatic attraction between the metal cation and the water molecules, the more the hydration energy.
|Hydration enthalpy (kj mol-1)||-506||-406||-330||-310||-276|
Li+ <Na+ <K+ <Rb+ <Cs+
Hydrated radius is defined as the radius of ion and closely bound water molecules (distance between the nucleus of a metal ion and the boundary of the water molecule). The ions with bigger size and low charge hold water molecules less tightly and thus have a smaller hydrated radius.
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.
Alkali metals are strong reducing agents. Reducing ability is, related to the ease of electron donation or lower ionization energy.
As ionisation energy decreases down the group, reducing property is expected to increase from Lithium to Cesium. In reality, reducing ability increases from Sodium to Rubidium and then again decreases down the group. Lithium has the highest reduction potential (-3.04V) and is the strongest reducing agent of of all elements.
|Reduction potential (v)||-3.04||-2.714||-2.925||-2.930||-2.927|
Reduction potential and reducing ability depend on the combined energy difference of three processes:
Being the smallest ion, the hydration enthalpy of Li+ is much higher than others and compensates for its higher ionization enthalpy. So, for the above-mentioned reasons, standard reduction potential varies as
Alkali metals form strong ionic bonds as they readily lose their valence electrons to form cation because of their low electronegativity and low electron gain enthalpy values.
The electronegativity values of alkali metals are small in magnitude and they decrease as we move down the group. The order being, Li < Na < K <Rb <Cs
The alkali metals and their salts impart a characteristic colour to an oxidising flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level.
When an excited electron returns to its ground state, radiation in the visible region of the spectrum is emitted. As a result, they give the flame a distinctive colour that reflects their emission or absorption spectrum, which can be used to identify them qualitatively.
Chemical Properties and Reactivity of Alkali Metals
Atomic number increase down the group for alkali metals moving from Li towards Cs. They are highly electropositive and form compounds that are ionic in nature.
The alkali metals are highly reactive due to their large size and low ionisation enthalpy. The reactivity of these metals increases down the group.
Reactions of Alkali Metals With Air (O2)
4Li (s) + O2(g) ⟶ 2Li2O(S) [Oxide]
4Na (s) + O2(g) ⟶ 2Na2O(s)[Oxide]
2Na (s) + O2(g) ⟶ 2Na2O2(s)[Peroxide]
M(S) + O2(g) ⟶ MO2(s)[Superoxide]
[Where, M = K,Rb,Cs
6Li (s) + N2(g) ⟶ 2Li3N(s)
Hydroxides of Alkali Metals
The oxides of the alkali metals react with moisture to form hydroxides and hydrogen gas. Alkali metal peroxides react with water to form metal hydroxides and hydrogen peroxide. Superoxides of alkali metals react with water to form hydroxides, and hydrogen peroxide and liberate oxygen gas.
Li2O(s) + 2H2O (l) ⟶ 2LiOH (aq) + H2(g) [Oxide + Water]
Na2O2 (s)+ 2H2O (l) ⟶ 2NaOH (aq)+ H2O2 (aq) [Peroxide + Water]
2KO2 (s)+ 2H2O (l) ⟶ 2KOH (aq)+ H2O2 (aq)+O2(g) [Superoxide + Water]
2RbO2 (s)+ 2H2O (l) ⟶ 2RbOH (aq)+ H2O2 (aq)+O2(g) [Superoxide + Water]
2CsO2 (s)+ 2H2O (l) ⟶ 2CsOH (aq)+ H2O2 (aq)+O2(g) [Superoxide + Water]
Reactions of Alkali Metals With Water
The alkali metals react with water to form metal hydroxides and release hydrogen gas.
2M(s) + 2H2O(l) ⟶ 2MOH(aq) + H2(g)
Reaction of the alkali metal with water is exothermic and the enthalpy of the reaction increases from Lithium to Caesium. Alkali metals float in the water during the reaction.
Example- 2Na(s) + 2H2O(l) ⟶ 2NaOH(aq) +H2(g)
Reactions of Alkali Metals With Dihydrogen
The alkali metals react with dihydrogen to form metal hydrides.
2M (s) +H2(g) ⟶ 2MH (s)
Reactions of Alkali Metals With Halogens
The alkali metals react vigorously with halogens to form ionic halides. An alkali metal loses one electron and a halogen accepts that electron, which results in the formation of an ionic salt .
2M (s)+X2(g)2MX (s)
KI (s) +I2(g)KI3 (∵ I-+I2I3-)
Carbonates and Bicarbonates
Alkali metal hydroxides on reacting with carbon dioxide forms metal carbonates. Also, carbonic acid reacts with metal hydroxides to form bicarbonates.
2NaOH (s)+CO2(g)⟶Na2CO3(s) +H2O(l)
NaOH (s)+H2CO3(aq)⟶NaHCO3(s) +H2O(l)
Alkali metal carbonates except LI2CO3 are ionic, water-soluble and thermally stable in nature.
Alkali Metal Sulphates and Nitrates
Alum, KAl(SO4)2.12H2O .
Reactions of Alkali Metals with Liquid Ammonia
Na + (x + y)NH3 [Na(NH3)x]+ ⟶ + [e(NH3)y]-
Ammoniated Cation Ammoniated electron
2e- + 2(NH3)y ⟶ [e-(NH3)y]2
NaNH2(s)+H2O(l)⟶NaOH(aq) + NH3(g)
Alkali metals are the most electropositive elements in the periodic table and highly reactive as well, so displacement by other metals and electrolysis is usually not applied directly.
They have high standard electrode potential values, which restrict reducing agents like carbon from reducing them directly.
Also, in the electrolysis of aqueous solutions, hydrogen ions get preferentially reduced to gaseous hydrogen before metal ions can get deposited at the cathode. Hence, sodium and potassium are obtained only by the electrolysis of the fused salts of sodium hydroxide and sodium chloride, which too in the form of amalgamation, which prevents them from further reactions. Alkali metals form alloys with themselves, other metals, and mercury amalgams.
Chemical precipitation, solvent extraction, and leaching techniques have all been used to extract Rubidium and Caesium.
Anomalous Properties of Lithium
Lithium differs in various chemical properties from the rest of the alkali metals owing to its smallest size, highest ionization energy and strongest polarising power in the group.
Also, Lithium has the strongest reducing character which can be attributed to its smaller atomic radius, larger solubility, and highest electrode potential. So those exceptional properties are:
2Li (s)+NH3(l) ⟶Li2NH(s)+H2(g)
Q.1. Which of the following metals are used to make photoelectric cells?
Alkali metals which less ionization enthalpy, Caesium is most widely used to make photoelectric cells as they can readily convert sunlight into electrical energy.
Q 2. Alkali metals when dissolved in ammonia behave as good reducing agents due to the:
1. Solvated cation
2. Solvated unpaired electron
3. Liberation of H2gas
4. Both A and B
The reduction is the gain of electrons and reducing agents are the one which reduces others and itself gets oxidized. So, ammoniated solutions of alkali metals are good reducing agents due to the presence of free yet solvated unpaired electrons.
Na + (x + y)NH3⟶[Na(NH3 )x]++ [e(NH3)y]-
Ammoniated Cation Ammoniated electron
Q 3. The correct order of ionic radii among alkali metals is:
1. Li+ ˂ Na+ ˂ K+ ˂ Rb+ ˂ Cs+
2. Na+ ˂ Li+ ˂ K+ ˂ Rb+ ˂ Cs+
3. Na+ ˂ Li+ ˂ Rb+ ˂ K+ ˂ Cs+
4. Cs+ ˂ Li+ ˂ K+ ˂ Rb+ ˂ Na+
Alkali metals readily lose an electron and become cationic. The ionic radii of their respective cations also increase down the group. Increasing order of ionic radius therefore will be Li+ ˂ Na+ ˂ K+ ˂ Rb+ ˂ Cs+.
Q 4. Which among the following alkali metals is capable of reacting with atmospheric nitrogen?
LI+ ion is the smallest among alkali metal ions and this size is quite compatible with that of nitride ion N3-. Hence, lattice energy released during the bond formation is high, which stabilises the molecule. Thus, only Lithium can react with atmospheric nitrogen among alkali metals.
Q 1. Are alkali metals paramagnetic or diamagnetic?
Answer: Alkali metals are paramagnetic in nature as their outer electronic configuration is ns1. Hence due to one unpaired electron in their outermost shell, they are paramagnetic. But their salts, where they exist as monovalent ions like in Li+, Na+, K+ etc, they have fully filled outer s-orbital (ns2) and hence they are diamagnetic.
Q 2. Which of the following alkali metals gives hydrated halide salts?
Answer: Due to the small size of LI+, infact smallest among all alkali metals, it is quite able to polarise water molecules unlike other alkali metal ions, and hence it forms a hydrated salt, i.e., LiCl.2H2O.
Q 3. Why solubility of LiF and CsI is less as compared to other alkali metal halides?
Answer: As the sizes of LI+ and F- are very small, the lattice energy of LiF is very high. Thus, due to this, the solubility of LiF decreases.
As the sizes of Cs+ and I- are very big, the hydration energy of CsI is very low. Thus, due to this, the solubility of CsI decreases.
Q 4. What is so special about alkali metals?
Answer: Alkali metals are highly reactive at normal temperature and pressure conditions and readily loose one valence electron to form monovalent cations. Thus they always appear in a combined state.
They appear silvery and are quite lusturous. Although they are solids but they are also called soft metals as they can be cut using a plastic knife as well.
|Sodium Hydroxide||Sodium Carbonate|