[Explained] Physical and Chemical Properties of Alkali Metals

Cell Phone Batteries are made up of Lithium
Cell Phone Batteries are made up of Lithium


 The group 1 or Group 1 A of the periodic table consists if six elements including hydrogen. These elements are 

  • Hydrogen
  • Lithium 
  • Sodium 
  • Potassium 
  • Rubidium 
  • Cesium 
  • Francium 
They are collectively known as alkali metals. In this article we discuss the physical and chemical properties of the Alkali (Group 1) metal elements.

Element Symbol Atomic Number Electron Configuration Brief Representation of Electron Configuration
Lithium Li 3 1s2 2s1 [He] 2s1
Sodium Na 11 1s2 2s2 2p6 3s1 [Ne] 3s1
Potassium K 19 1s2 2s2 2p6 3s2 3p6 3d10 4s1 [Ar] 4s1
Rubidium Rb 37 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1 [Kr] 5s1
Cesium Cs 55 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s1 [Xe] 6s1
Francium Fr 87 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 6s2 6p6 7s1 [Rn] 7s1

Physical Properties of Alkali Metals

The important physical properties of alkali metals are given below:

Physical State

Alkali Metals show typical silvery white metallic luster when freshly cut. The metallic luster fades rapidly due to oxidation by atmospheric air. They are soft, malleable and ductile. They are so soft that they can be cut with a knife. Lithium is the hardest of all the elements present in the alkali group.

Atomic and Ionic Radii

The atomic and ionic radii of alkali metals show the following characteristics:

The ionic radii of alkali metals ions are smaller than the atomic radii of the corresponding atoms.  

For example the ionic radius of Na+ ion is 102 pm whereas the atomic radius of Na atom is 186 pm.

Alkali metals possess only one electron in their valence shell. During the formation of cation, the valence s electron is lost. The cation thus formed has one electrons shell less than the parent atom. The removal of an electron shell decreases the size. 

      Na           ⟶     Na+ + e-

     1s22s22p63s1              1s22s22p6         

Moreover, the removal of an electron from the valence shell increases the effective nuclear charge experienced by the remaining electrons. Thus, the remaining electrons are pulled closer to the nucleus resulting in a further decrease in the size of the ion. 

The combined effect of the decrease in the number of the electron shells and an increase in the effective nuclear charge is responsible for the smaller size of alkali metal cations as compared to those of the corresponding alkali metal atoms. 

The atomic and ionic radii of alkali metals are the largest in their respective periods.

Each alkali metal atom is the first element of its period. As one moves from the left to right in a period, the differentiating electrons are added in the same electron shell and the nuclear charge increases with increase in the atomic number. Thus, in going from left to right in a period, the number of shells remains the  same but nuclear charge increases with each succeeding element. 

Thus, the electrons in the valence shell experience a greater pull towards the nucleus. This results in the successive decrease of the atomic and ionic radii with increase in the atomic number. This is why the atomic and ionic radii of alkali metals are the largest in their respective periods. 

The atomic and ionic radii of alkali metals increase on moving down the group i.e. they increase in going from Li to Cs.

As one moves from Li to Cs in group 1, a new electron shell is added at each element and the nuclear increases in the atomic number. The addition of an electron shell at each element tends to increase the size of the atom but the increase in the nuclear charge has a tendency to decrease the size of the atom or ion. Thus, the two factors oppose each other. 

The increase in the number of shells increases the screening effect of the inner electrons on the valence s-electron. This results in the expansion of the electron cloud. As the screening effect is quite  large, it over weighs the contractive effect of the nuclear charges with increase in the atomic number. The net result is an increase in the atomic and ionic radii of the alkali metals in going from Li to Cs.

Table Salt contains Sodium as a constituent
Table Salt contains Sodium as a constituent


Density

Alkali metals possess quite low densities as compared to other metals. Li, Na, and K are lighter than water. The density of the alkali metals increases from Li to Cs.

Although alkali metals possess close packed structures. You might be wondering why alkali metals have low density? It is due to the large size of atoms that their densities are low. On moving down the group, the atomic mass as well as the size of the atoms increase. 

The increase in atomic mass overweighs the effect of the increase in the size of the atoms. Therefore, densities of alkali metals increase in going from Li to Cs. However, potassium is an exception and is lighter than sodium.

Melting and Boiling Points

The alkali metals possess low melting and boiling points. The melting and boiling points decrease on moving down the group. The answer to why the melting and boiling points of alkali metals are low is given below.

The alkali metals possess only one electron in their valence shells. Therefore, the inter atomic forces responsible for the binding of atoms in the lattice which come into existence due to formation of metallic bonding are weak. 

Consequently, alkali metals possess low melting and boiling points. On moving down the group, the size of the atoms increases without any increase in the number of valence electrons. This further decreases the inter atomic forces. Therefore, the melting and boiling points further decreases in moving from Li to Cs.

Ionization enthalpy or ionization energy

The ionization energy values follows the following trends:

Alkali metals possess very low values of ionization energy. The ionization energy of an alkali metal atom is lowest in the period.

The alkali metal atoms possess electronic configuration of the type [Noble Gas] ns1.

The noble gas core shields the valence s-electron from the nucleus. Therefore in alkali metals the valence electron is loosely held by the nucleus and can be removed easily by supplying a small amount of energy. This is why alkali metals possess quite low ionization energies.

The ionization enthalpy of alkali metals decrease progressively in going from Li to Cs.

In going from Li to Cs. the distance of the valence electron from the nucleus increases progressively due to the addition of a new shell at each succeeding element. The increase in the number of shells causes an increase in the screening effect which consequently decreases the effective nuclear charge experienced by the valence electron. 

This facilitates an easier removal of the valence electron. This is why the ionization energies of alkali metals decrease on moving down the group. 

The second ionization energies of the alkali metals are very high

When and electron is removed from an alkali metal atom, the cation formed has a stable noble gas configuration. For example. 

Li+ = 1s2

Na+= 1s2 2s2 2p6

The noble gas configuration is a very stable configuration. The removal of an electron from such as configuration is very difficult and requires a large amount of energy. This why the second ionization energies of alkali metals are very high.

Alkali Metal ion Electron Configuration
Li+ 1s2
Na+ 1s2 2s2 2p6
K+ 1s2 2s2 2p6 3s2 3p6

Bananas contain a lot of Potassium
Bananas contain a lot of Potassium

 

Electronegativity

Alkali metals possess low values of electronegativity. In general, the electronegativity of alkali metals decrease in going down the group. 

Due to large size and low nuclear charge, the alkali metal atoms are unable to attract electrons towards them. This is why they possess low values of electronegativity. In going from Li to Cs, the size of atom increases further. Consequently, the electronegativity of alkali metals decreases on going down the group. 

Electropositive Character (also known as Metallic character)

 The alkali metals are strongly electropositve. Each alkali metal atom is the most electropositive atom in its period. Due to strong electropositive character alkali metals exhibit strong metallic character. The electropositive character of alkali metal metals increases in going from Li and Cs.

The tendency of an atom to form positive ions by losing its valence electron s determine its electropositive character. Since alkali metals possess very low values of ionization energies, they have strong tendency to lose their valence electrons. This is why they show strong electropositive or metallic character. 

As the ionization energies of alkali metals decrease progressively in going from Li to Cs, their tendency of losing valence electrons also increases progressively. Consequently, the electropositive character increases on moving down the group.

Oxidation State

An alkali metals exhibit only +1 oxidation state in their compounds. They do not show variable oxidation states as shown by several other elements of their periods.

The alkali metal atoms possess only one electron ns1 in their valence shells and can lose it readily due to low ionization energies. On losing the valence electron they form a monopositive cation and thus exhibit +1 oxidation state.

    M ⟶  M+ + e-

The monopositive cation formed has the configuration of a nearest noble gas. As the noble gas configuration is a very stable configuration, the cation formed does not allow the further removal of electrons easily. This is why alkali metal atoms do not exist in higher oxidation states and exhibit only =+1 state.

The alkali metal cations (M+) have no unpaired electrons. Therefore, they are colorless and diamagnetic in nature.

Hydration Energy

The alkali metal cations have a strong tendency to get hydrated.

    M+ (g) + aq (excess water) ⟶  M+ (aq)

The process of hydration is exothermic and the energy involved is called hydration energy. The hydration energy of alkali metal cations decreases in going from Li+  to Cs+.

The hydration energy depends upon the charge-radius ratio(q/r). Since the radius of alkali metal cations increases in going from Li+ to Cs+, the hydration energy decreases in going down the group. 

Hydration energy is a measure of the tendency of an ion to undergo hydration. This is why the tendency of alkali metal cations to undergo hydration in going from Li+  to Cs+. Li+ ions are most heavily hydrated in aqueous solutions.

Flame Coloration

Alkali metals and their salts show characteristic colors when heated in a non-luminous flame. The color imparted to the flame darkens on moving down the group.


Element Color
Li Crimson
Na Golden Yellow
K Pale Violet
Rb Violet
Cs Violet

When an alkali metal or its salt is heated in a flame, its electrons get excited to higher energy levels due to absorption of energy. The excited states are short lived. When the excited electrons return back to their normal states, the energy is emitted. The emitted energy corresponds to the visible region and therefore a characteristic color is imparted to the flame.

Photoelectric Effect

Alkali metals emit electrons when irradiated with with light. The phenomenon of the emission of electrons from the from the surface of a metal on irradiating the metal surface with electromagnetic radiation is called photoelectric effect and the electrons thus ejected are called photo electrons.

As mentioned above, alkali metals possess low ionization energies. If the energy of the light falling on the surface of the metal is greater than or equal to the work function (the minimum energy required to overcome the attractive forces responsible for binding electrons to the metal), the electrons present on the metal surface get ejected in the form of photoelectrons. 

Due to low ionization energies, alkali metals particularly potassium and cesium have low values of work function and emit photoelectrons easily when exposed to light suitable frequency.

Nature of the Compounds

Alkali metals form ionic compounds. The formation of ionic compounds by alkali metals may be attributed to their electropositive nature. Due to low ionization energies, alkali metals readily form cations by losing their valence electrons. Consequently, they form ionic bonds with non metals.

Cesium in a vial
Cesium in a vial



Lattice Energy of Compounds

The salts of alkali metals possess high values of lattice energy. The lattice energy values decrease in going down the group.


Salt Lattice Energy(kJ mol-1)
LiCl -840
NaCl -776
KCl -701
RbCl -682
CsCl -630

The change in enthalpy involved in condensing required number of gaseous positive and negative ions to form one mole of lattice of an ionic compound is termed as the lattice energy of the compound.

The compounds of alkali metals (salts) consists of cations and anions and thus are ionic in nature. They are held together by strong electrostatic forces of attraction. This is why a large amount of energy is releases during the condensation of ions to form the lattice. Consequently, the lattice energies of alkali metal salts are quite high. 

Since the lattice energy of a salt is inversely proportional to the sum of ionic radii, the values of lattice energy decreases on moving from Li to Cs.

 Chemical Properties of Alkali Metals

Alkali metals are highly reactive metals. The high chemical reactivity of these elements may be attributed to their low ionization energy and low heats of atomization. The reactivity if alkali metals increases from going from Li to Cs.

Let us look at some important chemical properties of alkali metals which have been described below:

Action of Air

Alkali metals are so reactive that they get tarnished rapidly when exposed to air due to the formation of oxides,hydroxides and finally carbonates at their surface. For example, 

  4Na (s) + O2 ⟶ 2Na2O

  Na2(s) + H2O (l)   2NaOH

   2NaOH (s) + CO2 (g)  Na2CO3 + H2O (l)

 

Due to their reactivity towards air and moisture, alkali metals can not be stored in open air. They are always stored in an hydrocarbon solvent such as kerosene oil which prevents them from coming in contact with air and moisture.

Reaction with Oxygen

Alkali metals burn vigorously when heated in oxygen and form different types of oxides depending upon the nature of the metal. Lithium forms monoxide (LI2O), sodium forms a peroxide, while the other alkali metals (Potassium, Rubidium, Cesium) form superoxides having the general formula MO2. Thus,

                                                        4Li + O2    2Li2O

                                                        2Na + O2   Na2O2

                                                         K + O2    KO2

The formation of different types of oxides by different alkali metals can be explained on the basis of their ionic sizes. Lithium ion is the smallest of all alkali metal ions. Due to small size, it has a strong positive field around it. The strong positive field around lithium ion attracts the negative charge so strong that it does not permit the monoxide ion, O2- to combine with another oxygen atom to form peroxide ion, O22-

On the other hand, the weaker positive field around larger sodium ion permits the O2- ion to combine with another oxygen to form O22- ion. This is why lithium forms only monoxide, whereas sodium forms mainly peroxide. The larger K+,Rb+ and Cs+ possess still weaker positive fields around them and allow peroxide ions to further combine with oxygen to form superoxide ions, O2-. This is why potassium, rubidium and cesium form mainly superoxides.

Thus it is generally said that a small cation can stabilize a small anion, whereas a large cation can stabilize a large anion.




Avocado Contains Potassium
Avocado Contains Potassium

Nature of Oxides

The normal oxides of alkali metals (monoxides) are basic oxides. They are highly soluble in water and form highly alkaline solutions due to formation of hydroxides.

   M2O + H2⟶ 2MOH

The peroxides and superoxides are oxidizing agents. They react with water to form a hydrogen peroxide or oxygen or both.

   Na2O2 + H2 2NaOH + H2O2

  2KO2 + 2H2O   2KOH +  H2O2 + O2

Reaction with Di-hydrogen

Alkali metals react with H2 at about 673K (Li at 1073 K) to form saline hydrides having high melting points.

    2M  +  H2     2M+ H-

Reaction with Water

 Alkali metals react readily and violently with water to form hydroxides with the liberation of hydrogen gas.

    2M  +  H2⟶ 2MOH + H2O

Lithium reacts with water somewhat slowly but the reaction of alkali metals with water is so vigorous that the hydrogen liberated catches fire immediately. This is why alkali metals can not be kept under water.

The reactivity of alkali metals with water increases on going down the group in going from Li to Cs.

Nature of Hydroxides

The alkali metal hydroxides are white crystalline solids, highly soluble in water. They are highly basic in nature and form the strongest bases known.

Basic Strength

The basic strength of hydroxides increases on moving down the group form LiOH to CsOH.

       LiOH< NaOH< KOH< RbOH< CsOH

This may be attributed to an increase in the electropositive character of alkali metals on moving down the group.

      2LiOH + Heat  Li2O + H2

 Stability

Alkali metal hydroxides except LiOH are thermally stable. LiOH decomposes on heating.

Reaction with Halogens

Alkali metals readily combine with halogens to form halides of the type MX.

     2M + X2 ⟶ 2MX

The reactivity if alkali metals towards a particular halogen increases in going from Li to Cs, while the reactivity of halogens towards a particular alkali metal follows the order                                                                                                                                                                               F2>Cl2>Br2>I2.

Except lithium halides, alkali metal halides are ionic in nature. Lithium halides (particularly LiI) possess partial covalent character due to polarization of anions by small Li+ cation. Alkali metals halides except LiF are free soluble in water. LiF is insoluble in water due to its high lattice energy. They possess high melting and boiling points and conduct electric current in the fused state.

Partial Covalent Character in Lithium Halides

The partial covalent character in lithium halides can be explained as follows.

Lithium ion is the smallest of all alkali metal cations. Due to small size, it has a high charge density. Due to high density of a positive charge, lithium ion has a tendency to pull electrons of the neighboring anion towards itself. This distorts the electron cloud of the neighboring anion. the distortion of the electron cloud of an anion by a cation is called polarization. Polarization leads to the lowering of charge on both the ions due to partial neutralization of charge and introduces a partial covalent character in the molecule.

The polarization in an ionic compound is governed by a set of rules called the Fajan's Rules.

They can be stated as follows: 

Smaller the cation and larger the anion, greater is the extent of polarization in an ionic compound.

For example, the extent of polarization and therefore the extent of covalent character in lithium halides follows the order LiI > LiBr > LiCl

Higher the charge on cation and anion, greater is the extent of polarization.

For example the polarizing power of Na+,Mg2+,Al3+ ions follows the order
Al3+>Mg2+>Na+

 Lesser the dielectric constant of the medium, greater is the extent of polarization.

 This is because the solvents of high dielectric constant such as water tend to weaken the electrostatic forces of attraction existing between ions and this decrease the tendency of polarization in the molecule.

The polarizing power of a cation having outermost shell configuration of the type s2p6d+10 is more than that of the cation having configuration of the type s2p6.

For example, the polarizing power of Cu+ ion having electron configuration 3s23p53d10 is more than that of Na+ ion which has configuration 2s22p6 This is why CuCl is more covalent than NaCl.

Reducing Nature

Alkali metals act as strong reducing agents. This is due to their values of ionization energies. The low values of ionization energies help them to lose electrons easily. This is why all alkali metals act as strong reducing agent. The reducing character of alkali metals decreases in going from Li to Cs.

Lithium is the first element of group 1. It has the highest value of ionization energy. Therefore,  it should have been the weakest reducing agent. However, lithium is the strongest reducing agent as indicated by its high oxidation potential value. The anomalous behavior can be explained as follows.

The oxidation potential of a metal is the sum of the energies involved in the following processes.

 $$M(s) \xrightarrow {Sublimation} M (g)$$

  $$M(s) \xrightarrow {Sublimation} M (g)^2 + {e^-}$$

 $$M^2 (g) \xrightarrow {Hydration} M^+ (aq)$$

Among the alkali metal cations, Li+ undergoes hydration to the maximum extent due to its smallest state. The hydration energy of Li+ is maximum and therefore a large amount of energy is released during its hydration. The large amount of energy released during hydration compensates the energy needed to remove the electron from Li atom (I.E.) This facilitates the release of electron which makes lithium strongest reducing agents that hydrogen. Therefore, they react with compounds containing acidic hydrogen atoms such as alcohols, acetylene etc. to liberate hydrogen.

 2Li + 2C2H5OH  ⟶ 2C2H5OLi + H2

2Na + 2C2H5OH   2C2H5ONa + H2

2Na + H ー C=C ー H  NaC = CNa + H2

 

Solubility in Liquid Ammonia

All alkali metals dissolve in liquid ammonia to form deep colored solutions which are highly conducting.

When an alkali metal is dissolved in liquid ammonia, ammoniated cations and ammoniated electrons are formed as follows:

 M + (x+y) NH3  M++ (NH3)x + e- (NH3)y

The blue color of the solution is due to the presence of ammoniated electrons which get excited on absorption of energy from visible light. Further, due to the presence of ammoniated cations and ammoniated electrons, the solutions of alkali metals in liquid ammonia are highly conducting.

On standing the blue colored solutions of alkali metals in liquid ammonia slowly decompose to give hydrogen gas.

 2M+ (NH3 2MNH2 + H2 + (2x-2) NH3

In concentrated solution, the blue color changes to bronze color and the solution becomes diamagnetic.

Nature of Salts of Oxoacids

Alkali metals being the most electropositive metals form salts with all the oxoacids. The salts are soluble in water and are quite stable even at high temperatures. The carbonates and most of the bicarbonates are highly stable towards heat. The stability of carbonates and bicarbonates increases on moving down the group. The carbonate of lithium is less stable than its bicarbonate does not exist as a solid.


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