Chemical & Physical Properties of Alkali Metals | Inorganic Chemistry PDF Download

Alkali metals belong to the s-block elements occupying the leftmost side of the periodic table. 

• Alkali metals readily lose electrons, making them count among the most reactive elements on Earth. 
• In general ‘alkali’ refers to the basic or alkaline nature of their metal hydroxides. The compounds are called alkali metals because when they react with water they usually form alkalies which are nothing but strong bases that can easily neutralize acids. 

• Alkali elements are Lithium(Li), Sodium(Na), Potassium (K), Rubidium (Ru), Cesium (Cs) and Francium (Fr) occupying successive periods from first to seven. Francium is a radioactive element with a very low half-life.
• The main reason why hydrogen (H) is not considered an alkali metal is that it is mostly found as a gas when the temperature and pressure are normal. Hydrogen can show properties or transform into an alkali metal when it is exposed to extremely high pressure. 

Electronic Configuration of Alkali Metals

• Alkali metals have one electron in their valence shell.
• The electronic configuration is given by ns₁. For example, the electronic configuration of lithium is given by 1ns₁ 2ns₁.
• They tend to lose the outer shell electron to form cations with charge +1 (monovalent ions).
• This makes them the most electropositive elements and due to the same reason, they are not found in the pure state.

Trends in Physical Properties of Alkali Metals

1. Atomic and Ionic Radii

• Atomic and ionic radii of elements increase, regularly down the column. Also, every alkali metal has the largest radii of any other element in the corresponding period. 
• Alkali metals readily lose an electron and become cationic. The cationic radius is smaller than the neutral atom. 
• The relative ionic radii also increase down the column.
  Increasing order of Atomic and Ionic Radius: Li ˂ Na ˂ K ˂ Rb ˂ Cs and Li+ ˂ Na+ ˂ K+ ˂ Rb+ ˂ Cs+

2. Density

• All are light metals. The densities of metals Li, Na, and K Are lighter than water.
• Density gradually increases in moving down from Li to Cs. 
• Potassium is, however, lighter than sodium. 
• Sequence of densities: Li < Na < K < Rb < Cs

3. Structures of the Metals, Hardness & Cohesive Energy 

• At normal temperatures, all the Group 1 metals adopt a body-centred cubic type of lattice with a coordination number of 8. However, at very low temperatures lithium forms a hexagonal close-packed structure with a coordination number of 12.
• The cohesive energy is the force holding the atoms or ions together in the solid.
• The atoms become larger on descending the group from lithium to caesium, so the bonds are weaker, the cohesive energy decreases, and the softness of the metals increases.
• The crystal structures of Li₂O, Na₂O, K₂O, and Rb₂O are anti-fluorite structures, and Cs₂O has an anti-CdCl2 layer structure.

4. Melting and boiling points

• The cohesive energy decreases down the group, and the melting points decrease correspondingly.
• The energy binding the atoms in the crystal lattices of these metals is relatively low on account of a single electron in the valency shell. 
• Consequently, the metals have low melting and boiling points. These decrease in moving down from Li to Cs as the metallic bond strength decreases or cohesive force decreases.

5. Ionisation Energies & Electropositive Character 

• Due to their large size, the outermost electron is far from the nucleus and can easily be removed. 
• Their ionization potentials are relatively low. Thus, the metals have a great tendency to lose the ns¹ electron to change into M ⁺ ions.
• These metals are highly electropositive. As the ionization potential decreases from Li to Cs, the electropositive character increases, i.e. metallic character increases. 

• The reactivity of these metals increases from Li to Cs.

• The ns¹ electron is so loosely held that even the low-energy photons (light) can eject this electron from the metal surface. This property is termed a photoelectronic effect. K and Cs are used in photoelectric cells which are sensitive to blue light.

6. Oxidation States 

The alkali metals can lose their ns¹ electron quite easily to form a univalent positive ion, M⁺. 

• The ion has a stable configuration of inert gas. These metals are univalent and show electrovalency, i.e., form electrovalent compounds.
• Since the electron configuration of M + ions is similar to those of inert gases, these ions have no unpaired electrons and consequently are colourless and diamagnetic. 

7. Hydration of Ions, Hydrated Radii, and Hydration Energy

The salts of alkali metals are ionic and soluble in water. The solubility is due to the test that cations get hydrated by water molecules M⁺ + aq   —→  [M(aq)]⁺  

• The smaller the cation, the greater the degree of its hydration of M ⁺ ions & decrease from Li to Cs⁺. 
• Consequently, the radii of the hydrated ion decrease from Li⁺ to Cs⁺
• The ionic conductance of these hydrated ions increases from [Li(aq)] ⁺ to [Cs(aq)]⁺.
• Hydration of ions is an exothermic process.
• The energy released when one gram mole of an ion is dissolved in water to get it hydrated is called hydration energy. Since the degree of hydration decreases from Li ⁺ to Cs⁺, the hydration energy of alkali metal ions also decreases from Li⁺ to Cs⁺.
• Some water molecules touch the metal ions and bond to them, forming a complex. These water molecules constitute the primary shell of water. Thus Li⁺ is tetrahedrally surrounded by four water molecules sp₃ hybridization. 
•  With the heavier ions, particularly Rb⁺ and Cs⁺, the number of water molecules increases to six. VSEPR theory predicts an octahedral structure (d₂sp₃ hybridization). 
•  A secondary layer of water molecules further hydrates the ions, though these are only held by weak ion-dipole attractive forces. 

8. Conductivity

The alkali metals are good conductors of heat and electricity. This is due to the presence of loosely held valency electrons which are free to move throughout the metal structure.

9. Heat of Atomization 

• The heat of atomization decreases from Li to Cs.
•  This is due to the decrease in the metallic bond strength from Li to Cs

10. Flame Colors & Spectra

• The alkali metals and their salts impart a characteristic colour to a flame.
• The colour arises from electronic transitions in short-lived species which are formed

• The reason for flame coloration is that the energy of the flame causes an excitation of the outermost electrons which on return to their original position give out the energy so absorbed in the visible region. 
• The energy released is minimal in the case of Li ⁺ and increases from Li ⁺ to Cs⁺. Thus the frequency of the light emitted increases by the formula E = hν. The frequency of light in lithium is minimum which corresponds to the red region of the spectra.
• Compounds of Group 1 metals are typically white, except those where the anion is coloured. For example sodium chromate Na₂(CrO₄) (yellow), potassium dichromate K₂[Cr₂O₇] (orange), and potassium permanganate K[MnO₄] (deep purple).

11. Reduction Potential

The substances that can donate electrons are reducing agents. Reducing ability is, related to the ease of electron donation or lower ionization energy.

As ionization energy decreases down the column, reducing property is expected to increase from Lithium to Cesium. While reducing ability increases from Sodium to Cesium, Lithium has the highest reduction potential (-3.04V) 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 to the metal ion,
• Hydration of the ion with water.

Lithium, being the smallest ion, its hydration enthalpy is much higher than others and compensates more than its higher ionization enthalpy: ENa ˂ EK ˂ ERb ˂ ECs ˂ RLi

Chemical Properties of Alkali Metals

1. Action with Air

• On exposure to moist air, all alkali metals except lithium tarnish quickly. The effect of the atmosphere increases from Li to Cs. These are, therefore always kept under kerosene oil to protect them from the air. 

• Lithium, when heated in air, combines with nitrogen to form nitride, it is due to the diagonal relationship with magnesium.

2. Action with Water

• Alkali metals decompose water with the evolution of hydrogen gas.

3. Action with Hydrogen 

• The alkali metals combine directly with hydrogen to form crystalline hydrides of the formula M. H. These hydrides are ionic and contain the hydrides ion. H^–
  2M +H₂ ——→ 2MH
• The ionic character of the bonds in these hydrides increases from LiH to CsH and the stability decreases in the same order. 
• They are powerful reducing agents, especially at high temperatures.
  SiCl₄ + 4NaH ——→ SiH₄ + 4NaCl 

4. Action with Oxygen (Oxides and Hydroxides)

• Down the group affinity towards oxygen increases.
• The metals all burn in the air to form oxides, though the product varies depending on the metal. lithium forms the monoxide Li₂O (and some peroxide Li₂O₂), sodium forms the peroxide Na₂O₂ and some monoxide Na₂O), and the others form superoxides of the type MO₂. 

• The normal oxides ‘M₂O’ react with water to form hydroxides.  
  M₂O + H₂O ——→ 2MOH
• The basic nature of the oxides (M₂O) increases gradually on moving down in the group. The hydroxides (MOH) are colourless, strong alkaline, and corrosive compounds. 
• These are soluble in water and dissolve with the evolution of heat. The hydroxides are thermally stable except LiOH. 
• The relative strength of the hydroxides increases from LiOH to CsOH.
  CsOH > RbOH > KOH > NaOH > LiOH

5. Action with Halogens 

• The alkali metals directly react with halogens forming the halides of the type MX. 2M + X₂ → 2MX
• Except for certain lithium haloids, the metal halides are ionic compounds (M+X–). 
• The halides are crystalline and have high melting and boiling points. The fused halides are good conductors of electricity and are used for the preparation of alkali metals. 
• All halides except LiF dissolve in water.
• The alkali metal halides react with the halogens and interhalogen compounds forming ionic polyhalide compounds.

KI + I₂ → K[I₃],                 KBr + ICI → K[BrICI],                 KF + BrF₃ → K[BrF₄]

6. Nature of Oxysalts

• Alkali metals readily react with oxyacids forming corresponding salts with the evolution of hydrogen. 
• Lithium salt behaves abnormally due to the polarizing power of Li+ ions (small size) and lattice energy effects.

7. Nature of Carbonates and Bicarbonates

• As the electropositive character increases from Li to Cs, the solubility of the carbonates increases in the same order
  Cs₂CO₃ > Rb₂CO₃ > K₂CO₃ > Na₂CO₃ > Li₂CO₃
• Li₂CO₃ decomposes on heating and is insoluble in water. The aqueous solution of carbonates is alkaline. This is due to hydrolysis as carbonates are salts of strong bases and weak acid (H₂CO₃) carbonic acid).
  ⇆  
• The bicarbonates, MHCO₃, of the alkali metals, except lithium, are known in solid-state. The bicarbonates are soluble in water. On heating, bicarbonates decompose into carbonates with the evolution of CO₂.
  2MHCO₃ ——→ +M₂CO₃ + H₂O + CO₂
• The abnormal behaviour of Li₂CO₃ towards heat can be explained in the following manner. The Li ⁺ ion exerts a strong polarizing action and distorts the electron cloud of the nearby atom the large CO₃^2- ion. 
• This results in the weakening of the C—O bond and the strengthening of the Li—O bond. This ultimately facilitates the decomposition of Li₂CO₃ into Li₂O and CO₂. 
• The lattice energy of Li₂O is higher than the lattice energy of carbonate. This also favours the decomposition of Li₂CO₃.
• Lithium due to its less electropositive nature, does not form solid bicarbonate and LiHCO ₃ exists in solutions only.

8. Nature of nitrates 

Nitrates of the type, MNO₃, are known. These are colourless, soluble in water, and electrovalent in nature. The nitrates do not undergo hydrolysis. Except for LiNO₃, the other nitrates decompose into nitrites and oxygen.
2MNO₃ ——→ 2MNO₂ + O₂
Lithium nitrate decomposes to oxide on heating, it is due to the diagonal relationship with magnesium.

9. Nature of Sulphates 

Sulphates of the type M₂SO₄ are known. Except for Li₂SO₄, other sulphates are soluble in water. The sulfates when fused with carbon form sulphides.
M₂SO₄ + 4C ——→ M₂S + 4CO
The sulphates of alkali metals form double salts with the sulphates of the trivalent metals like Fe, Al, Cr, etc. The double sulphates crystallize with a large number of water molecules potash alum K2SO₄, Al₂(SO₄)₃.24H₂O consists of 24 water molecules. Sulphate of lithium is not known to form alum.

10. Action of Liquid Ammonia 

The alkali metals dissolve in liquid ammonia without the evolution of hydrogen. The colour of the dilute solutions is blue. The metal atom loses electrons and it combines with the ammonia molecule.
M → M+ (in liquid ammonia) + e (ammoniated)
M⁺(x + y) NH₃ → [M(NH₃)x]⁺ + e(NH₃)y solvated electron

On heating its blue colour changes to bronze. It is an ammoniated electron that is responsible for colour.
The solutions are good conductors of electricity and have strong reducing properties. The solutions are paramagnetic.
When dry ammonia is passed over hot metal, amides are formed.

The amides are decomposed by cold water with the evolution of NH₃ MNH₂ + 2NH₃ ——→ MOH + NH₃
Recent studies proved the existence of Li(NH₃)₄, a golden yellow solid.

11. Formation of Alloys 

The alkali metals form alloys amongst themselves and with other metals. These combine with mercury readily forming amalgams.

12. Complex Formation

Alkali metals have a very little tendency to form complexes, Lithium being small in size forms certain complexes but this tendency decreases as the size increases.

13. Diagonal Relationship: Similarities with Magnesium

Lithium shows a resemblance with magnesium, an element of group IIA. This resemblance is termed a diagonal relationship.

The reasons for the diagonal relationship are the following:

• Electronegativities of Li and Mg are quite comparable (Li = 1.00 Mg = 1.20)
• Atomic radii and ionic radii of Li and Mg are not very much different.
  Atomic radii (Å)               Li^-1.23,              Mg^-1.36
  Ionic radii (Å)                  Li⁺-00.60,          Mg²⁺-0.65
• Both have high polarizing power (ionic potential) Polarizing power
• Cations with large ionic potentials tend to polarize the anions and give partial covalent character to compounds.

Lithium resembles magnesium in the following respects.

• Both Li and Mg are harder and have higher melting points than the other metals in their respective groups.
• Li like Mg decomposes water slowly to liberate hydrogen.
• 2Li + 2H₂O ——→ 2LiOH + H₂; Mg + 2H₂O ——→ Mg(OH)₂ + H₂ 
• Both elements combine with nitrogen on heating.
• 6Li + N₂ ——→ 2Li₃N; 3Mg + N₂ ——→ Mg₃ + N₂ 
• Both the nitrides are decomposed by water with the evolution of ammonia (NH₃)
• Li₃N + 3H₂O ——→ 3LiOH + NH₃; Mg₃N₂ + 6H₂O ——→ 3Mg(OH)₂ + 2NH₃
• Both Li and Mg combine with carbon on heating
• 2Li + 3H₂O ——→ 3Li₂C₂; Mg + 2C ——→ MgC₂ 
• Both the carbides yield C₂H₂ with water.
• Lithium forms the only monoxide when heated in oxygen.
  Mg also forms the monoxide
• 4Li + O₂ ——→ 2Li₂O; 2Mg + O₂  ——→ 2MgO
• Both are less soluble in water.
• Hydroxides of Li and Mg are weak bases and are slightly soluble in water. Both decompose on heating
• 2LiOH ——→ Li₂O + H₂O; Mg(OH)₂ ——→ MgO + H₂O
• Lithium fluoride, phosphate, oxalate, and carbonate like the corresponding salts of Mg, are sparingly soluble in water.
• Carbonates of Li and Mg decompose on heating.
• Li₂CO₃ ——→ Li₂O + CO₂; MgCO₃  ——→ MgO + CO₂
• Nitrates of Li and Mg decompose on heating giving a mixture of nitrogen dioxide and oxygen.
• 4Li₂ NO₃ ——→ 2Li₂O + 4NO₂ + O₂ ; 2Mg(NO₃)₂ ——→ 2MgO + 4NO₂ + O₂

Compounds of Sodium

Na₂O 2NaNO₃ + 10 Na ——→ 6Na ₂O + N₂ ; 2NaNO₂ + 6 Na ——→ 4Na ₂O + N₂

Properties 

NO₂O + H₂O ——→ 2NaOH

2. Sodium Peroxide, Na₂O₂ 

It is formed by heating sodium more than air free from moisture and carbon dioxide or over pure oxygen.
Properties2Na 2O₂ + 2H₂O ——→ 4NaOH + O₂; Na 2O₂ + H₂SO₄ ——→ Na ₂SO₄ + H₂O₂

3. Sodium Hydroxide (Caustic Soda), NaOH 

It is one of the important chemicals and is manufactured on a very large scale forming an important chemical industry. It is most conveniently manufactured by one of the following processes:

(a) Methods involving sodium carbonate as a starting material: Two methods are used. These are 

(i) Causticisation process (Gossage process)
Na₂Co₃ + Ca(Oh)₂ ⇆ CaCo₃ + 2NaOH

(ii) Lowig’s process 

Na₂CO₃ + Fe₂O₃ ——→ 2NaFe₂ + CO₂; 2Na₂FeO₂ + H2O ——→ 2NaOH + Fe₂O₃

(b) Methods involving sodium chloride as starting material (electrolysis of brine)

At Cathod  2H₂O + 2e ⇆ H₂ + 2OH^- ; Na⁺ + OH^-  ⇆ NaOH

At anode  

Properties: 

• Strong alkali: NaOH ⇆ Na⁺ + OH^-
  (i) NaOH + HCl ——→ NaCl + H₂O;
  (ii) 2NaOH + CO₂ ——→ Na₂CO₃ + H₂O
  (iii) Al₂O₃ + 2NaOH ——→ 2NaAlO₂ + H2O
• Action on non-metals Cl₂ + 2NaOH ——→ NaCl + NaClO + H₂O
  3Cl₂ + 6NaOH ——→ 5NaCl + NaClO₃ + 3H₂O
  P₄ + 3NaOH + 3H₂ O ——→ 3NaH₂PO₂ + PH₃
• Action on Metals Zn + 2NaOH ——→ Na₂ZnO₂ + H₂;
• Action on Salts
  (i) Ni(NO₃)₂ + 2NaOH ——→ Ni(OH)₂ + 2NaNO₃
  (ii) Insoluble hydroxides which dissolve more than NaOH
  ZnSO₄ + 2NaOH ——→ Zn(OH)₂ + Na₂SO₄
  (iii) Unstable hydroxides 2AgNO₃ + 2NaOH ——→ 2AgOH + 2NaNO₃
  2AgOH ——→ Ag₂O + H₂O

4. Sodium Carbonate or Washing Soda (Na₂CO₃.10H₂O)

(a) Le-Blanc Process

(i) Conversion of NaCl into 

Na₂CO₄ NaCl + H₂SO₄ ——→ NaHSO₄ + HCl

NaHSO₄ + NaCl ——→ Na₂SO₄ + HCl

(b) Solvay ammonia soda process

NH₃ + H₂O + CO₂ ——→ NH₄HCO₃

NaCl + NH4HCO₃ ——→ NaHCO₃ + NH4Cl

Chemical Properties of Sodium Carbonate

• Action of acids
  Na₂CO₃ + HCl → NaHCO₃ + NaCl; NaHCO₃ + HCl → NaCl + H₂O + CO₂ 
• The action of water & carbon dioxide
  Na₂CO3 + H₂O + CO₂ → 2NaHCO₃ 
• Action of silica
  When the mixture of sodium carbonate and silica is fused, sodium is formed. Sodium silicate is called soluble glass or water glass as it is soluble in water.

• Action on salt of non-alkali metals
  The solution of carbonates reacts with metal salts (except alkali metal Salt) to form insoluble normal or basic carbonates.
  2MgCl₂ + 2Na₂CO₃ + H₂O   → MgCO₃ · Mg(OH)₂ + 4NaCl + CO₂

5. Sodium Bicarbonates (NaHCO₃) 

Na₂CO₃ + CO₂ + H₂O → 2NaHCO₃ (Sparingly soluble)
Properties
NaHCO₃ + H₂O → NaOH + H₂CO₃; 2NaHCO₃ → Na₂CO₃ + H₂O + CO₂

6. Sodium thiosulphate (Na₂S₂O₃.5H₂O)

(i) Spring’s reaction Na₂S + I₂ + Na₂SO₃ → Na₂S₂O₃ + 2NaI
Properties 

(i) Oxidation
2Na₂S₂O₃ + I₂ → 2NaI + Na₂S₄O₆ 

7. Sodium ammonium hydrogen phosphate

(NaNH₄HPO₄)
 Preparation
NH₄Cl + Na₂HPO₄ + 4H₂O → NaNH₄HPO₄.4H₂O + NaCl

Properties (i) Na(NH₄)HPO₄ → NaPO₃ + NH₃ + H₂O

8. Sodium Cyanide (NaCN)

From calcium cyanide 

Ca(CN)₂ + C + Na₂CO₃ → 2NaCN + CaCO₃

Properties:

It forms complex cyanides with the salts of copper cadmium, zinc, iron, cobalt, nickel etc.
Some examples are given below.

(i) AgNO₃ + NaCN → AgCN + NaNO₃; AgCN + NaCN → Na[Ag(CN)₂]

Compounds of Potassium

Carnalite               KClMgCl₂ 6H₂O
Kainite                  KCl.MgSO₄.MgCl₂.3H₂O
Indian Saltpetre    KNO₃
Felspar                 K₂O.Al₂O₃.6SiO₂ (clay)
Polyhalite             K₂SO₄MgSO₄.CaSO₄.6H₂O
Sylvine                 KCl

Potassium Iodide (KI)

(i) KOH + HI → KI + H₂O

(ii) 3I₂ + 6KOH → 5KI + KIO₃ + 3H₂O

Complex Formations

• An unusual compound [Na(Cryptanad-222)]⁺ Na^– has an interesting feature in that it contains Na^–, the iodide ion (Negative charge on metal).