What are Alkali Metals: Extraction of Alkali Metals, Uses, Anomalous Behaviour of Lithium, Physical and Chemical Properties

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

• What are Alkali Metals?
• Electronic Configuration of Alkali Metals
• Trends in Physical Properties of Alkali Metals
• Chemical Properties of Alkali Metals
• Extraction of Alkali Metals
• Anomalous Behaviour of Lithium
• Uses of Alkali Metal
• Practice Problems
• Frequently Asked Questions-FAQs

What are Alkali Metals?

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. 

Electronic Configuration of Alkali Metals

• Alkali metals have one electron in their valence shell.
• The outer electronic configuration of alkali metals is ns^-1. For example, the electronic configuration of lithium is given as 1s²s¹.
• They tend to lose the outer shell electron to form cations with a charge of +1 (monovalent ions).
• So their oxidation state is +1.
• The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. Hence they are never found in the free state in nature.

Trends in Physical Properties of Alkali Metals

Atomic Radii

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.

• Increasing order of Atomic and Ionic Radius: Li < Na < K <Rb <Cs and Li⁺ <Na⁺ <K⁺ <Rb⁺ <Cs⁺

Density of Alkali Metals

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 atomic volume and directly proportional to atomic. 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.

Ionisation Enthalpy and Electropositive Metallic Character

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 electrons from the outermost shell. 

• Alkali metals shall donate the single valence electron to get a noble gas configuration. Thus they are all univalent electropositive metals.
• Order of Ionization Enthalpy is: Li < Na < K <Rb <Cs

Hydration Enthalpy and Solubility of Alkali Metal Ions

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. 

• The smaller the ion, the higher the charge density, the stronger the electrostatic attraction between water molecules and metal cations, and the higher the hydration enthalpy.
• Hence, on moving down the group, the hydration enthalpy of alkali metal ions decreases with the increase in ionic size as charge density decreases. 

• Thus, the correct order of the hydration energy of alkali metals is given as follows:

          Li⁺ <Na⁺ <K⁺ <Rb⁺ <Cs⁺

• Li⁺  has the maximum degree of hydration or is most soluble and for this reason, lithium salts are mostly hydrated. Example: LiCl. 2H₂0.
• It can be said that the higher the hydration enthalpy, the more will be the solubility of metal ions in water.
• Hydration and solubility decrease with increasing size so that Cs⁺ ion is the least water-soluble alkali metal ion.
• Smaller ions have a higher charge density and can be hydrated by more water molecules. This releases a higher enthalpy of hydration and makes the hydrated ions more stable.

Hydrated Ionic Radii & Order of Conductance in Alkali Metal Ions

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. 

• As the size of the cation decreases, the charge density of the cation increases and the degree of hydration increases. 
• As the degree of hydration increases, the hydrated ionic radius increases and the ionic mobility of the metal ions decreases. 
• As the ionic mobility of the metal ions decreases, the molar conductance and the conductivity of the solution decrease.
• The correct order of the ionic radius of alkali metals is given as:  Li⁺ <Na⁺ <K⁺ <Rb⁺ <Cs⁺
• The correct order of the hydrated radii of alkali metal ions is given as: Li⁺ <Na⁺ <K⁺ <Rb⁺ <Cs⁺
• The correct order of the conductance for the alkali metals is given as: Li⁺ <Na⁺ <K⁺ <Rb⁺ <Cs⁺

Melting and Boiling Point of Alkali Metals Electronegativity

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.

• As we go down the group, the boiling and melting points decrease continuously because of increasing atomic size i.e., the distance between the nucleus of one atom and the outer electrons of another atom of the element increases which leads to a decrease in the attraction between the nucleus and the electrons.
• Example- Melting and Boiling point of Li is 105°  and 1342 ° C respectively. Whereas Melting and Boiling point of  K is  63.38 ° C and 759 ° C respectively.

Reduction Potential

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 (-3.04 V) and is the strongest reducing agent of all elements.

Reduction potential and reducing ability depend on the combined energy difference of three processes:

• Sublimation of the atom
• Ionization of the metal ion
• Hydration of the ions with water

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

 

Nature of Bond

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.

Electronegativity

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

Flame Test for Alkali Metals

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) 

• The alkali metals tarnish in dry air due to the formation of their oxides. Lithium forms monoxide, sodium forms oxide as well as peroxide, and the other metals form oxide, peroxide as well as superoxides.

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

• They are normally kept in kerosene oil because of their high reactivity to air and water.
• Lithium shows exceptional behaviour in reacting directly with the nitrogen of air to form lithium nitride.

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.

Example:

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.

• The density of Sodium and potassium is lower than in water. In heavier alkali metal, reaction enthalpy is high such that the metal gets melted and raises to the surface. Hence, the reaction with water becomes kinetically faster, highly exothermic, and explosive leading to fire from lithium to caesium.

Example- 2Na(s) + 2H₂O(l) ⟶ 2NaOH(aq) +H₂(g)

• The melting point decreases down the group, and the reaction with water becomes more and more vigorous because as the metal is melted, the surface area exposed to water becomes more. Due to this, the reaction is kinetically faster. 

Reactions of Alkali Metals With Dihydrogen

The alkali metals react with dihydrogen to form metal hydrides. 

• All the alkali metal hydrides are ionic solids with high melting points. 

2M (s) +H₂(g) ⟶ 2MH (s)

• As we move down the group, the stability of the hydride decreases.
• The metal hydrides react with water to give metal hydroxides (MOH) and release H₂gas.
  
  LiH (s) +H₂O (l)⟶ LiOH (aq) + H₂(g) 

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)

• Halides of potassium, rubidium, and caesium have the property of combining with extra halogen atoms forming polyhalides.

KI (s) +I₂(g)KI₃ (∵ I^-+I₂I₃^-)

• Due to the small size of ions, Lithium halides are somewhat covalent. It is because of the high polarisation capability of lithium ions (the distortion of the electron cloud of the anion by the cation is known as polarisation). 

Carbonates and Bicarbonates

Alkali metal hydroxides on reacting with carbon dioxide form 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.

• Bicarbonates, except for LIHCO₃ (liquid state), are solid, water-soluble and on heating liberate carbon dioxide.

Alkali Metal Sulphates and Nitrates

• Except for LI₂SO₄, all alkali metal sulphates are soluble. 
• Alkali metal Sulphates can be reduced from carbon to sulphide. 

Na₂SO₄(aq)+2C(s)⟶Na₂S(aq)+2CO₂(g)

• Also alkali metals form double salts with trivalent metal sulphates, for example- 

Alum, KAl(SO₄)₂.12H₂O .

• Nitrates of alkali metals are soluble in water and on heating except lithium nitrate decomposes to nitrites.

Reactions of Alkali Metals with Liquid Ammonia

• Alkali metals dissolve in liquid ammonia, The solubility of alkali metals in liquid ammonia is due to the formation of ammoniated cation and ammoniated electrons. Example:

Na + (x + y)NH₃ [Na(NH₃)_x]⁺ ⟶ +    [e(NH₃)_y]^-

                                Ammoniated Cation           Ammoniated electron

• Point to note here, is liquid ammonia is ammonia gas liquified under pressure and liquor ammonia is an aqueous solution of ammonia or ammonia dissolved in water. So they both are different.
• The overall solution is deep blue in colour. It is also conducting, reducing and paramagnetic in nature. 
• Blue colour is due to the empty cages in solvated ammonia electrons which have different energy levels and due to different interactions with the partially charged side of NH₃ electronic transitions between these cages take place, and energy is absorbed in the visible region and the blue colour is observed. 
• It is paramagnetic due to the presence of unpaired electrons present in the solution in the form of ammoniated electrons.
• It is due to the presence of ammoniated cation and ammoniated electrons that the solution is conducting in nature.
• On increasing the concentration of sodium metal, the concentrated metal-ammonia solution gets a metallic bronze colour and becomes diamagnetic due to the pairing of e^- .

2e^- + 2(NH₃)_y ⟶ [e^-(NH₃)_y]₂

• On keeping this solution standing for long, the colour fades due to formation of amide after liberating hydrogen. Also, hot alkali metals react with dry ammonia gas to form amide and liberate hydrogen. Amide on hydrolysis produces sodium hydroxide.

2Na (s)+2NH₃(g)⟶2NaNH₂(s)+H₂(g)

NaNH₂(s)+H₂O(l)⟶NaOH(aq) + NH₃(g)

Extraction of Alkali Metals      

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, 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:

• Lithium halides are covalent in nature owing to their small ionic radii of Li⁺ ions and high polarising power.
• It is harder than other alkali metals.
• Lithium reacts slowly with oxygen to form a normal oxide that does not get tarnished quickly.
• Lithium reacts very slowly with water.
• LiOH is less basic. Only lithium hydroxide is able to decompose back into its oxide and water.
• Li₂CO₃ is less stable due to its covalent nature and decomposes into oxide and carbon dioxide.
• Only Lithium reacts with atmospheric nitrogen to form nitride (Li₃N₂).
• LiNO₃ decomposes into nitrogen dioxide (NO₂), oxygen (O₂) and oxide (Li₂O), while the other nitrates of alkali metals yield nitrites and oxygen.

• Lithium forms an imide while other alkalis form amide with liquid ammonia.

2Li (s)+NH₃(l) ⟶Li₂NH(s)+H₂(g)

2Na (s)+2NH₃(l)⟶2NaNH₂(s)+H₂(g)

• Lithium salts are less soluble compared to other alkali metal salts.

Uses of Alkali Metals

• Lithium is used to create heat-resistant ceramics, glasses, rechargeable batteries for cell phones, and alloys for use in aeroplanes.
• Sodium salts are used in the food industry, soap-making, pharmaceutical industries etc.
• Liquid sodium works as a coolant in nuclear reactors.
• Potassium chloride, potash is used as fertilizer, insecticide, for water-retention, enhancing crop yield and many more.
• An isotope of Rubidium (RB-82) is used in medical treatments in myocardial perfusion.
• Rubidium is used in laser cooling, engines of space crafts, optical glasses, atomic clocks etc.
• Caesium-134 is used in the nuclear industry.
• Caesium-137 is used in the treatment of cancer.
• It is also used as a standard in analytical methods such as spectrophotometry. 

Practice Problems

Q.1. Which of the following metals are used to make photoelectric cells?

1. Caesium
2. Lithium
3. Magnesium
4. Calcium

Answer: (A)

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

Answer: (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⁺

Answer: (A)

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?

1. Fr
2. Cs
3. Li
4. Na

Answer: (C)

 LI⁺ ion is the smallest among alkali metal ions and this size is quite compatible with that of nitride ion N^3-. 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. 

Frequently Asked Questions-FAQs

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 is the 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. So, 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 lose one valence electron to form monovalent cations. Thus they always appear in a combined state.

They appear silvery and are quite lusturous. Though they are solids, they are also called soft metals as they can even be cut using a plastic knife.