Periodic discussion and general characteristics of alkali metals

Electronic configurations

All the alkali metals have one electron in their outermost 's' orbitals preceded by the noble gas configuration. Thus, the general configuration of alkali metals may be written as [Noble gas] ns1 where 'n' represents the valence shell. The electronic configurations of alkali metals are:

The electronic configurations of alkali metals

ElementSymbolAtomic No.Electronic configuration
LithiumLi3[He]2s1
SodiumNa11[Ne]3s1
PotassiumK19[Ar]4s1
RubidiumRb37[Kr]5s1
CesiumCs55[Xe]6s1
FranciumFr87[Rn]7s1

GENERAL CHARACTERISTICS OF ALKALI METALS

Atomic and ionic radii

Being the first elements of each period, alkali metals have the largest atomic and ionic radii in their respective periods. As we move within a period, the atomic radius and ionic radius tend to decrease due to increase in the effective nuclear charge. On moving down the group, there is increase in the number of shells and, therefore, atomic and ionic radii increase.

Physical propertyLiNaKRbCs
Atomic radius (pm)152186227248264
Ionic radius (pm)76102138152167

Ionization energies

  • Alkali metals have the lowest ionization energy in each period. Within the group, as we go down, the ionization energies of alkali metals decrease due to their atomic size being the largest in their respective periods. In large atoms the valence electrons are loosely held by the nucleus and are easily lost, leading them to have low ionization energies and acquiring stable noble gas configurations. On moving down the group, the atomic size increases and the number of inner shells also increases, increasing the magnitude of screening effect and consequently, the ionization energy decreases down the group.
  • The second ionizations energies of alkali metals are very high. The removal of an electron from alkali metals causes the formation of monovalent cations having very stable electronic configurations (same as that of noble gases). Therefore, it becomes very difficult to remove the second electron from the stable noble gas configurations, giving very high second ionization energy values (IE2).
Physical propertyLiNaKRbCs
Ionization Energy I520496419403376
KJ mol -1II72984562305126332230

Melting and boiling points

All alkali metals are soft and have low melting and boiling points. As alkali metals have only one valence electron per metal atom, the energy binding the atoms in the crystal lattice of the metal is low. Consequently, the metallic bonds in these metals are not very strong and their melting and boiling points decrease on moving down the group.

Physical propertyLiNaKRbCs
Melting point (K)453.5370.8336.2312.0301.5
Boiling point (K)16201154.41038.5961.0978.0

Density

The densities of alkali metals are low as compared to other metals, with Li, Na and K being even lighter than water (K is lighter than Na). While alkali metals do have close packing of metal atoms in their lattice the large size of their atoms cause them to have low densities. As we move down the group from Li to Cs, even though there is an increase in atomic size, the simultaneous increase in atomic mass compensates more than the increase in atomic size. The densities (mass/volume) of alkali metals thus gradually increase from Li to Cs. However, potassium is lighter than sodium probably due to increase in atomic size of potassium.

Physical propertyLiNaKRbCs
Density (g cm -1)0.530.970.861.531.90

Electropositive or metallic character

The electropositive character of an element is expressed in terms of the tendency of its atom to release electrons:

All the alkali metals are strongly electropositive or metallic in character, since they have low ionization energies and their atoms readily lose the valence electron. As the ionization energies decrease down the family, the electron releasing tendency or electropositive character is expected to increase down the family.

Physical propertyLiNaKRbCs
Eo value (V)- 3.03- 2.71- 2.93- 2.93- 2.92

Oxidation states

All alkali metals have only one electron in their valence shell. They exhibit an oxidation state of +1 in their compounds and can lose the single valence electron readily to acquire the stable configuration of a noble gas. Thus, they form monovalent ions, M+(e.g., Li+, Na+, K+, Rb+, Cs+). Since the second ionization energies are very high, they cannot form divalent ions. Thus, alkali metals are univalent and form ionic compounds.

Characteristic flame coloration

As the alkali metals have very low ionization energies, the energy from the flame of a bunsen burner is sufficient to excite the electrons of alkali metals to higher energy levels. The excited state being unstable, these electrons return to their original energy levels, emitting extra energy, which gives characteristic flame colorations. The different colors of the alkali metals can be explained on the basis of amount of energy absorbed for excitation of the valence electron.

Physical propertyLiNaKRbCs
Flame colourcrimson redyellowpale violetvioletbluish

Photoelectric effect

When electromagnetic radiation strikes on the surface of alkali metals, they emit electrons. This is called the photoelectric effect. This occurs as alkali metals have low ionization energies, which allows the electrons to be easily ejected when exposed to light. Among alkali metals, cesium has lowest ionization energy and hence it can show photoelectric effect to the maximum extent.

Nature of the compounds

The compounds of the alkali metals are ionic in nature. Alkali metals form cations readily by losing the valence electrons (due to the low ionization energies and large atomic sizes). They go on to form ionic bonds with the non-metals of the 'p' block.

Lattice energies

The lattice energies of alkali metal salts are very high because strong electrostatic forces of attraction hold up the cations and anions formed.

Lattice energy gives a measure of the forces of attraction between the ions. It is defined as the amount of energy required to break one mole of a crystal into its free ions.

definition of lattice energy

The larger the forces of attraction, the greater will be lattice energy. The lattice energy also depends upon the size of the ion and its charge. For the cation of same valency, ionic solids having the same anion will display a decrease in the lattice energy with increase in size of the cation. This is due to decrease in forces of attraction between them.

Chemical properties of alkali metals
Alkali metals exhibit a high chemical reactivity because of their
  • low ionization energies
  • low heat of atomization.

Since the value of ionization energy decreases down the group (from Li to Cs) the reactivity of alkali metals increases from Li to Cs. All alkali metals are highly reactive towards the more electronegative elements such as oxygen and halogens. Some characteristic chemical properties of alkali metals are described.

ACTION WITH AIR

Formation of oxides

All the alkali metals on exposure to air or oxygen burn vigorously, forming oxides on the surface of the metals. Lithium forms monoxide (Li2O), sodium forms the peroxide (Na2O2) and the other elements form superoxides.

(MO2: M = K, Rb, Cs).

Lithium forms monoxide

sodium forms the peroxide

alkali metals on exposure to air

The fact that a small cation can stabilize a small anion and a large cation can stabilize a large anion explains the formation and stability of these oxides. The size of Li+ ion is very small and it has a strong positive field around it. It can combine with only small anion, O2- ion, resulting in the formation of monoxide (Li2O). Conversely, the Na+ ion is a larger cation and has a weak positive field around it and thus can stabilize a bigger peroxide ion, O22- or [-O-O-]2- which is also surrounded by a weak negative field. Similarly, the other ions K+, Rb+, Cs+ are still larger, having very weak positive field. Thus these ions stabilize a bigger superoxide O2- anion and form superoxides.

The alkali metal oxides are basic in nature because they dissolve in water to form alkali metal hydroxides.

alkali metal oxides form alkali metal hydroxides

Peroxides give hydrogen peroxide also

Peroxides give hydrogen peroxide

ACTION WITH HYDROGEN

Formation of hydrides

All alkali metals react with hydrogen to form hydrides that are ionic in nature (M+H-).

alkali metals react with hydrogen to form hydrides
  • Reactivity of alkali metals with hydrogen increases from Li to Cs.
  • Ionic character of the hydrides increases from Li to Cs. The decrease in ionization energy down the group permits easy availability of electrons to hydrogen, forming H- ion.
  • Stability of hydrides decreases from Li to Cs because the M-H bond becomes weak as the size of the alkali metal increases from Li to Cs. This causes the stability of hydrides to decrease.
  • The hydrides behave as strong reducing agents and their reducing nature increases down the group.

ACTION OF HALOGENS

Formation of halides

Alkali metals combine readily with halogens to form ionic halides M+X-. For example,

Formation of halides

formation of sodium chloride

Reactivity of alkali metals with halogen increases down the group because of corresponding increase in electropositive character (decrease in ionization energy).

All metal halides are ionic crystals. Lithium iodide is slightly covalent as it has the smallest cation, which exerts maximum polarizing power and iodide ion being the largest anion can be polarized to the largest extent.

All alkali metal halides are soluble in water except LiF, as it has a high lattice energy combined with a small cation and anion, making it insoluble in water.

ACTION TOWARDS WATER

Formation of hydroxides

Hydroxides and hydrogen gas result from the reaction of alkali metals with water.

action of sodium with water

action of potassium with water

Sodium and the other members of the family react so rapidly with water that the hydrogen gas evolved immediately catches fire. Only lithium reacts slowly. Thus, alkali metals are normally kept in kerosene oil and cannot be kept either in air or in water.

The hydroxides of alkali metals are strongly basic as the M-OH bond in the hydroxides of alkali metals is very weak and can easily ionize to form M+ and OH- ions. Further, the strength of the base increases as the ionization energy decreases down the group because the bond between metal and oxygen becomes weaker. This increases the basic strength of the hydroxides accordingly. Thus, NaOH is a stronger base than LiOH and so on.

Solutions in liquid ammonia

Alkali metals dissolve in liquid ammonia to give deep blue solutions that are conducting in nature. This happens because the alkali metal atom readily lose the valence electron in ammonia solution. Both the cation and the electron combine with ammonia to form ammoniated cation and ammoniated electron.

electron combine with ammonia to form ammoniated cation and ammoniated electron

The ammoniated electron is responsible for the blue color of the solution. The solution is made conducting in nature by both ammoniated cation and ammoniated electron.

On standing, the solution slowly liberates hydrogen as:

formation of metal amine

Hydration of ions

The alkali metal ions are highly hydrated. The smaller the size of the ion, the greater is the degree of hydration. Thus, Li+ ion gets much more hydrated than Na+ ion which in turn is more hydrated than K+ion and so on. The extent of hydration decreases down the group.

As a result of larger hydration of Li+ ion than Na+ ion, the effective size of Li+ ion is more than that of Na+ion. Further the ionic radii in water (called hydrated ionic radii) decreases in the order:

Li+ > Na+ > K+ > Rb+ > Cs+

hydrated ionic radii

As a result, the hydrated Li+ ion being largest ionic size, has the lowest mobility in water. On the other hand, the hydrated Cs+ ion being smallest in size has the highest mobility in water.

Reducing nature

Alkali metals are strong reducing agents as indicated by the large negative values of their reduction potentials. This is so because of the ease with which they lose electrons. All of them are better reducing agents than hydrogen (E = zero). Therefore, these metals react with compounds containing acidic hydrogen atoms such as alcohol and acetylene, liberating hydrogen.

action of lithium with ethanol

action of sodium with actylene