Alkali Metals: We all know that the Long Form Periodic Table is based on the Modern Periodic Law, which states that elements’ physical and chemical properties are a periodic function of their atomic numbers. Therefore, based upon the sub-shell in which the last electron enters, the table is divided into four blocks – s, p, d, and f. This article will cover the general characteristics of the first group of elements of the s-block that is Alkali Metals.

What are Alkali Metals?

Group I of the modern Periodic table comprises six elements- Lithium \({\rm{(Li)}}\) Sodium \(\left( {{\rm{Na}}} \right)\) Potassium \(\left( {\rm{K}} \right)\) Rubidium \(\left( {{\rm{Rb}}} \right)\) Cesium \(\left( {{\rm{Cs}}} \right)\) Francium \(\left( {{\rm{Fr}}} \right)\) besides hydrogen. These six elements are alkali metals because they readily dissolve in water to form hydroxides which are strongly alkaline. The word alkali has derived from the Arabic word ‘Alquili,’ which means the ashes of plants from which certain elements were initially isolated.

Though hydrogen is placed in the first group of the periodic table, it is not an alkali metal. Instead, it is included because of the similarity of the electronic configuration with these elements.

Occurrence of Alkali Metals

Alkali metals are highly reactive metals and hence, do not occur in the state in nature. Nevertheless, all of the discovered alkali metals occur in nature in the form of compounds- in order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and francium, which is very rare due to its high radioactivity. Francium occurs only in minute traces in nature as an intermediate step in some obscure side branches of the natural decay chains.

Electronic Configuration of Alkali Metals

The alkali metals are included in the s-block since the elements have one electron each in the s-subshell of their atoms. They contain only one s-electron outside the noble gas core. They thus have \(\mathrm{ns}^{1}\) as general electronic configuration. The orbital electronic configuration of elements is given as:

Elements	Atomic Number	Electronic Configuration	With Inert Gas Core
Lithium (Li)	\(3\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{1}}}\)	\([\mathrm{He}] 2 \mathrm{~s}^{1}\)
Sodium (Na)	\(11\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{1}}}\)	\({\rm{[Ne]3}}{{\rm{s}}^{\rm{1}}}\)
Potassium (K)	\(19\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{s}}^{\rm{1}}}\)	\({\rm{[Ar]4}}{{\rm{s}}^{\rm{1}}}\)
Rubidium (Rb)	\(37\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{5}}{{\rm{s}}^{\rm{1}}}\)	\({\rm{[Kr]5}}{{\rm{s}}^{\rm{1}}}\)
Caesium (Cs)	\(55\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{6}}}{\rm{6}}{{\rm{s}}^{\rm{1}}}\)	\({\rm{[Xe]6}}{{\rm{s}}^{\rm{1}}}\)
Francium (Fr)	\(87\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{f}}^{{\rm{14}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{6}}}{\rm{5}}{{\rm{d}}^{{\rm{10}}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}{\rm{6}}{{\rm{p}}^{\rm{6}}}{\rm{7}}{{\rm{s}}^{\rm{1}}}\)	\({\rm{[Rn]7}}{{\rm{s}}^{\rm{1}}}\)

Since all these elements have similar valence shells or outer electronic configurations, they have similar physical and chemical properties.

Trends in Atomic and Physical Properties

The following table shows some atomic and physical properties of Alkali metals.

1. Atomic radii – The atomic radii of alkali metals are the largest in their respective period and increase down the group as the number of energy shells increases.
2. Ionic radii – Alkali metal atoms form monovalent cations by losing one electron from the outermost shell. The cationic radius is smaller as compared to the atom. Therefore, ionic radii increase on moving down the group.
3. Ionization enthalpies – Ionization enthalpy is a measure of the tendency of an atom to lose electrons in the gaseous state. They have the lowest ionization enthalpies in their respective periods, and these decrease progressively on moving down the group.
4. Electronegativity – Electronegativity is the property of an element that tends to attract a relative electron for its atom for a shared pair in a bond. They have a low electron attracting tendency. As a result, electronegativity tends to decrease down the group.
5. Oxidation state – All the members of alkali metals are strongly electropositive and show a +1 oxidation state. The electropositive character increases down the group.
6. Metallic character – All the members are metals that are soft, malleable, and ductile solids. These can easily be cut by a knife.
7. Melting and boiling point – The melting and boiling points of alkali metals are very low and decrease with the increase in atomic number, increasing the group.
8. Nature of bonds formed – All alkali metals form ionic (electrovalent) bonds. The ionic character increases down the group.
9. Density – Alkali metals are very light. The first three elements are even lighter than water. Therefore, the density tends to increase down the group.

Chemical Properties

Alkali metals are highly reactive metals. Some of the important chemical properties are as follows:

1. Reaction in Air – Alkali metals tarnishes in the air due to the formation of oxide, hydroxide, or carbonates on the surface.
\({\rm{4M + }}{{\rm{O}}_{\rm{2}}} \to {\rm{2}}{{\rm{M}}_{\rm{2}}}{\rm{O}}\)
\({{\rm{M}}_2}{\rm{O}} + {{\rm{H}}_2}{\rm{O}} \to 2{\rm{MOH}}\)
\(\mathrm{MOH}+\mathrm{CO}_{2} \rightarrow \mathrm{M}_{2} \mathrm{CO}_{3}+\mathrm{H}_{2} \mathrm{O}\)
Therefore, it is always advised not to keep alkali metals in the air.

2. Reaction with Oxygen – Alkali metals combine with oxygen upon heating to form different oxides depending on the nature of the alkali metal. \(4 \mathrm{Li}+\mathrm{O}_{2} \rightarrow 2 \mathrm{LiO}_{2}\)
\(2 \mathrm{Na}+\mathrm{O}_{2} \rightarrow \mathrm{Na}_{} \mathrm{O}_{2}\)

3. Reaction with Water – Alkali metals reacts with water, liberating hydrogen gas and forming their hydroxides. The reaction is exothermic.
\(2 \mathrm{Li}+\mathrm{H}_{2} \mathrm{O} \rightarrow 2 \mathrm{LiOH}+\mathrm{H}_{2}\)
\(2{\rm{Na}} + {{\rm{H}}_2}{\rm{O}} \to 2{\rm{NaOH}} + {{\rm{H}}_2}\)

4. Reaction with Hydrogen – All alkali metals combine with hydrogen on heating to form colourless crystalline hydrides of ionic nature.
\({\rm{2M + }}{{\rm{H}}_{\rm{2}}} \to {\rm{2}}{{\rm{M}}^{\rm{ – }}}{{\rm{H}}^{\rm{ + }}}\) where M stands for alkali metals.

5. Reaction with Halogen – Alkali metals combines with halogens directly to form their respective halides. \(2{\rm{M}} + {{\rm{X}}_2} \to 2{{\rm{M}}^ + }{{\rm{X}}^ – }\) where M is an alkali metal, and X is a halogen.

6. Reducing Character – The reducing nature of an element is expressed in terms of its electrons releasing tendency. The members of alkali metals are powerful reducing agents since they have low ionization energy. These metals, therefore, liberate hydrogen from water, acids, and acetylene.

Extraction of Alkali Metals

Alkali metals generally are extracted by the electrolysis of their ores. Only the electrolytic method is suitable because of the highly reactive nature of alkali metals.

Uses of Alkali Metals

1. Lithium
   (i) An alloy of lithium with lead is used in electrical cables
   (ii) Lithium hydroxide is used in the manufacture of some lubricants.(iii) It is used as a deoxidizer in the purification of copper and nickel.
2. Sodium
   (i) An alloy of sodium amalgam is used in the preparation of several compounds.
   (ii) Sodium vapour lamps are used for lighting.
   (iii) In the molten state, it is used in nuclear reactors as a heat transfer medium.
3. Potassium
   (i) It is used as a reducing agent in chemical reactions.(ii) Potassium hydroxide is used in the manufacture of toilet soaps. It is also an excellent absorber of carbon dioxide.
   (iii) Some salts of potassium are used as fertilizers.
4. Caesium
   (i) It is used in photoelectric cells.
   (ii) It finds its utility as a rocket propellant or fuel.
   (iii) It is used as a heat transfer fluid in power generators.

FAQs on Alkali Metals

Q.1. What are alkali metals?
Ans: The alkali metals are a group of chemical elements from the s-block of the periodic table with similar properties: they appear silvery and can also be cut with a plastic knife. Alkali metals are highly reactive at standard temperature and pressure and readily lose their outermost electron to form cations with charge +1.

Q.2. Is Hydrogen an alkali metal?
Ans: No, hydrogen is not an alkali metal. It has been included because of the similarity of the electronic configuration with the elements of alkali metals. 

Q.3. What are the three properties of alkali metals?
Ans: The three properties of alkali metals are as follows:
They have one valence electron in the outermost shell, which they seek to lose to have a fully-filled outer shell. This property is what makes them so reactive. They are soft enough to be cut with a knife. When exposed to air, they tarnish due to oxidation.

Q.4. Which alkali metal can easily be cut?
Ans: Sodium, an alkali metal, is very soft and can easily be cut by a knife.

Q.5. What are the three facts about alkali metals?
Ans: The three facts about alkali metals are:
1. Because they are so reactive with air and water, they generally are stored in oil.
2. Caesium and rubidium are used to make atomic clocks. Caesium clocks are considered the most accurate of all clocks.
3. Sodium and potassium both play an important role in biological life on Earth. We cannot live without them.

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