INTRODUCTION:

 The members of Group IA or 1, except hydrogen are called the alkali metals,
because they react with water to form alkaline solutions.
 The group IA of the periodic table consists of six elements-lithium, sodium,
potassium, rubidium, cesium and francium besides hydrogen. These elements
are collectively called as alkali metals and group IA is known as alkali group as
the hydroxides of these metals are soluble in water and these solutions are highly
alkaline.
 The word alkali has been derived from Arabic word 'Alquili' meaning the ashes of
plants from which compounds of sodium and potassium were first isolated.
 They are soft enough to cut with a knife.

 The alkali metals are the most reactive of all metals, reacting readily with water
and air. The alkali metals, being reactive in nature, are never found free in nature
but always found in combined state. The alkali metals are so reactive that, they
must be stored under oil. This prevents the reaction of alkali metals with oxygen
and water vapor in the air.
 Of all the alkali metals, only sodium and potassium are found in abundance in
nature, i.e., they are seventh and eighth most abundant elements by weight in
earth's crust. Their compounds are very common and have been known and
used from very early times.
 The last member, francium, occurs only in traces as a radioactive decay product
because its half life period is very small, i.e., 21 minutes.

ELECTRONIC CONFIGURATION:
 The outermost shells of these elements have one electron and the penultimate
(next to outermost) shells contain 8 electrons except in the first member, lithium,
which contains 2 electrons. which is the atomic number of helium. Since, the last
electron enters at s-orbital, these are called s-block elements.
 Lithium shows somewhat abnormal properties as its electronic configuration is
slightly different than the rest of the members. Because of their similarity in
electronic configuration, (noble gas) ns2 , they are placed in the same sub
 group, i.e., IA of the periodic table and closely resemble in their physical and
chemical properties.

PHYSICAL PROPERTIES:

Property Li Na K Rb Cs Fr
Ionisation energy (KJ/mol) 520 496 419 403 376 375
Hydration enthalpy -506 -406 -330 -310 -276 ------
Metallic radius (pm) 152 186 227 248 265 ------
Ionic radius (pm) 76 102 138 152 167 180
Melting point (K) 454 371 336 312 302 ------
Boiling point (K) 1615 1156 1032 961 944 ------
Density (g/cm3) 0.53 0.97 0.86 1.53 1.90 ------
Standard potential E0 -3.04 -2.714 -2.925 -2.93 -2.927 ------

(a) All the alkali elements are silvery white solids. These are soft in nature and can be cut
with the help of knife except the lithium. When freshly cut, they have a bright luster
which quickly fades due to surface oxidation.
The silvery luster of alkali metals is due to the presence of highly mobile electrons of
the metallic lattice. There being only a single electron per atom, the metallic bonding is
not so strong. As a result of this, these metals are soft in nature. However, the softness
increases with increase of atomic number because there is continuous decrease of
metallic bond strength on account of increase in atomic size. Bigger is the size of metal
kernel weaker is the metallic bonding.
(b) Atomic and ionic radii: Atomic radius decreases when we move across the period in
the periodic table. This is because of increase in effective nuclear charge due to the
increase in atomic number (no of protons increases) and constancy in the no of shells.
Group IA atoms are the largest in their horizontal periods in the periodic table. When the
outermost electron is removed , it becomes a positive ion, because of this the size
decreases considerably. There are two reasons for this
(i) The outermost shell of electron has been completely removed.
(ii) The positive charge on the nucleus is now acting on lesser number of electrons,
attraction increases which brings contraction in size.
Atomic as well as ionic size increases from Li to Fr due to the presence of one extra shell
of electrons. This is because of increase in shell no.

(c) Density: All are light metals. The densities are low. Lithium, sodium and potassium are
lighter than water, for this very reason they float on water. Density gradually increases in
moving down from Li to Cs. Potassium is, however, lighter than sodium.
The reason for the low values is that these metals have high atomic volumes. The
abnormal value of potassium is due to unusual increase in atomic size, i.e., atomic
volume.

(d) Melting and boiling points: 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.
(e) Ionization energies and electropositive character: Due to their large size, the
outermost electron is far from the nucleus and can easily be removed. Their ionization
energies or ionization enthalpies are relatively low. Thus, the metals have a great
tendency to lose the ns electron to change into M+ ions. These metals are highly
electropositive in nature. As the ionization enthalpy decreases from Li to Cs, the
electropositive character increases, i.e., metallic character increases. The reactivity of
these metals increases from Li to Cs.
Photoelectric effect:
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 as photoelectric effect. K and
Cs are used in photoelectric cells which are sensitive to blue light.
(f) Oxidation states: The alkali metals can lose their ns¹ electron quite easily to form
univalent positive ion, M+ . The ion has a stable configuration of an inert gas.
The energy required to pull out another electron from M+ ion is very high, i.e., the
second ionization enthalpy values are high.
Consequently, it is not possible for alkali metals to form M2+ ions under ordinary
conditions. These metals thus show only one oxidation state, i.e., +1 oxidation state.
These metals are univalent in nature and show electro valency, i.e., form electrovalent
compounds.
Since the electronic configurations of M+ ions are similar to those of inert gases, these
ions have no unpaired electrons and consequently are colorless and diamagnetic in
nature.
(g) 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 fact that cations get hydrated by
water molecules.
The smaller the cation, the greater is the degree of its hydration. Thus, the degree of
hydration of M+ ions decreases from Li+ to Cs+ . Consequently the radii of the hydrated
ion decreases from Li+ to Cs+.
Li+ has a smaller size and high charge density meaning many water molecules will
surround the cation, resulting with bigger hydrated radii than Cs.
(h) Electronegativity: The tendency to attract electrons is low as the alkali metals are
electropositive. The electro negativity, thus, decreases from Li to Cs as the
electropositive character increases.
(i) 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.
(j) Specific heats: The specific heat values decrease from Li to Cs. This is due to the
decrease in the metallic bond strength from Li to Cs.
(I) Flame coloration: The alkali metals and their salts impart a characteristic color to
flame.
Li Na K Rb Cs
Crimson red Golden yellow Pale violet Red violet Violet

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 minimum in the case of Li+ and
increases in order Li+, Na+, K+, Rb+, Cs+. Thus, the frequency of the light emitted increases
in accordance with the formula E = hν. The frequency of light in lithium is minimum
which corresponds to red region of the spectra.
(m) Reducing nature: An element, which acts as a reducing agent, must have low
ionisation energy. Alkali metals act as strong reducing agents as their ionization energy
values are low. Since ionisation energy decreases on moving down from Li to Cs, the
reducing property increases in the same order. Thus, Li is weakest reducing agent while
Cs is the strongest reducing agent amongst alkali metals in free state. The tendency of
an element to lose an electron in solution is measured by its standard oxidation
potential value (E). Since alkali metals have high Eo values, these are strong reducing
agents. Because their standard oxidation potentials are so strongly positive, the alkali
metals can even reduce water.
However, it is observed that Li is the strongest reducing agent amongst alkali metals in
solution as E0 OX value of Li is maximum.
It looks surprising at first sight that lithium having high value of ionisation energy amongst
alkali metals acts as strongest reducing agent in solution. This can be explained if we
understand the fact that ionisation energy is the property of an isolated atom in
gaseous state while oxidation potential is concerned when the metal atom goes into
the solution.
The ionisation energy involves the change of gaseous atom to gaseous ion,
M(g) → M+(g) + e-
while oxidation potential involves the following change:
M(s) → M+(aq.) + e-
The above change occurs in three steps:
(i) M(s) → M(g) - sublimation energy
(ii) M(g) → M+(g) + e- ionisation energy
(iii) M+(g) + H₂O→ M+(aq.) + hydration energy
Sublimation energy is nearly same for all the alkali metals.

Since ionisation energy and hydration energy both are highest in case of lithium, with
greater ease the following overall change occurs in lithium and it acts as a strongest
reducing agent :
M(S)→ M+(aq.) + e
It is, therefore, concluded that highest reducing power of lithium in solution is due to its
large heat of hydration.