General Characteristics of the Compounds of the Alkali Metals

General Characteristics of the Compounds of the Alkali Metals: You must be well aware that elements in the periodic table are arranged in the increasing order of their atomic number. This arrangement also aligns the elements according to the similarity in their properties. The very first group of the periodic table is IA which is also known as Alkali metals, and all the elements in this group show similar properties. In this article, we will explore a little more about the general characteristics of the compounds formed by alkali metals. We will get an idea about their oxides, hydroxides, carbonates, and nitrates.

What are Alkali Metals?

The periodic table’s Group \(1\) elements are known as Alkali Metals. Lithium, sodium, potassium, rubidium, caesium, and francium are all part of it. They are also referred to as active elements. They have the least nuclear charge in their respective periods. They have a tendency to lose their one valence electron in the final shell and establish strong ionic connections with anions.

General Characteristics of the Compound of the Alkali Metals

Oxides and Hydroxides

Alkali metal oxides, peroxides, and superoxides dissolve readily in water due to the property of alkali metals. When such oxides are dissolved in water, they produce corresponding hydroxides, which are essentially very strong alkalis.

Thus, peroxides and superoxides act as oxidizing agents because they easily react with water to form hydrogen peroxide and oxygen. All alkali metal hydroxides are white crystalline solids. They are the strongest of all bases and easily dissolve in water, evolving a lot of heat.

Basic Strength

As we move down the group \(\text {Li}\) to \(\mathrm{Cs}\), the basic strength of these hydroxides increases. Because of the low ionisation energies that decrease down the group, alkali metal hydroxides behave as strong bases. The decrease in ionisation energies weakens the bond between the metal and the hydroxide ion, and the \(\text {M}-\text {O}\) bond in \(\text {M}-\text {O}-\text {H}\) can easily break, yielding \(\text {M}^{+}\) and \(\text {OH}^{-}\). This causes an increase in the concentration of hydroxyl ions in the solution, resulting in more basic properties.

Solubility and Stability

Except for lithium hydroxide, all of these hydroxides are highly soluble in water and thermally stable. Because alkali metals and their hydroxides are strongly basic, they tend to react with all acids, resulting in the formation of salts.

For example-

\(\mathrm{NaOH}+\mathrm{HCl} \rightarrow \mathrm{NaCl}+\mathrm{H}_{2} \mathrm{O}\)

\(2 \mathrm{NaOH}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}+2 \mathrm{H}_{2} \mathrm{O}\)

Halides of Alkali Metals

The alkali metals combine directly with halogens under the right conditions to form halides with the general formula \(\text {MX}\). These halides can also be synthesised by reacting to aqueous halogen acids \((\text {HX})\) with metal oxides, hydroxides, or carbonate.

1. \(\mathrm{M}_{2} \mathrm{O}+2 \mathrm{HX} \rightarrow 2 \mathrm{MX}+\mathrm{H}_{2} \mathrm{O}\)
2. \(\mathrm{MOH}+\mathrm{HX} \rightarrow \mathrm{MX}+\mathrm{H}_{2} \mathrm{O}\)
3. \(\mathrm{M}_{2} \mathrm{CO}_{3}+2 \mathrm{HX} \rightarrow 2 \mathrm{MX}+\mathrm{CO}_{2}+\mathrm{H}_{2} \mathrm{O}\)

All of the halides are colourless, high melting crystalline solids with high negative formation enthalpies. The value of halides decreases in the following order:

Fluoride > Chloride > Bromide > Iodide

Fluorides are thus the most stable, whereas iodides are the least stable.

The trends in melting points, boiling points, and solubility of alkali metals halides can be understood in terms of polarization effects, lattice energy, and hydration of ions.

Polarization Effects

The ionic and covalent properties of alkali metal halides are compared. When a cation approaches an anion, the anion’s electron cloud is attracted towards the cation and thus distorted. This is known as polarization. The ability of the cation to polarize the anion is referred to as its polarizing power, and the anion’s tendency to become polarized is referred to as its polarizability. The greater the polarization produced, the greater the concentration of electrons between the two atoms, resulting in a decrease in ionic character or an increase in covalent character.

In general, the covalent character of any compound is determined by the following factors-

1. Size of the cations
   
   The smaller the cation, the greater is the polarizing power, and therefore, the greater is its covalent character. As the size of the cation increases, the covalent character decreases.
   \(\text {LiCl} > \text {NaCl} > \text {KCl} > \text {RbCl} > \text {CsCl}\)
   
2. Size of the anions
   
   The larger the anion, the more polarizable it is. The covalent character of lithium halides in order is-
   \(\text {Lil} > \text {LiBr} > \text {LiCl} > \text {LiF}\)
   This explains why lithium iodide has a higher covalent bond than lithium fluoride.
   
3. Charge of the ions
   
   If the cation has a higher charge, its polarizing power increases , and thus, the covalent character increase. Similarly, when an anion’s charge increases, so do its polarizing capacity, making it more covalent.
   
4. Electronic configuration of the cation
   
   If two cations have the same charge and size, the one with the pseudo noble gas configuration, which has \(18\) electrons in the outermost shell, has more polarizing power than the cation with the noble gas configuration, which has \(8\) electrons.

Lattice Energy

The amount of energy required to separate one mole of a solid ionic compound into its gaseous ions is defined as lattice energy. Clearly, the higher the lattice energy, the higher the melting point of the alkali metals halide and the lower its solubility in water.

Hydration Energy

When one mole of gaseous ions combines with water to form hydrated ions, it releases a certain amount of energy that is known as hydration energy. The greater the hydration energy of the ions, the greater is the compound’s solubility in water. Furthermore, the extent of hydration is determined by the size of the ions. The smaller the ion, the more hydrated it is, and thus greater is its hydrated ionic radius and the lower is its ionic mobility (Conductance).

Thus, the melting point and solubility in water or organic solvents of halides of alkali metals can be explained.

A delicate balance of lattice enthalpy and hydration enthalpy determine the ultimate solubility of a compound in water. For example, the low solubility of \(\text {LiF} \left(0.27 \mathrm{~g} / 100 \mathrm{~g} \mathrm{H}_{2} \mathrm{O}\right)\) is due to its high lattice energy \(\left(-1005 \mathrm{KJmol}^{-1}\right.)\), whereas the low solubility of \(\operatorname{CsI}\left(44 \mathrm{~g} / 100 \mathrm{~g} \mathrm{H}_{2} \mathrm{O}\right)\) is due to both ions lower hydration energy \(\left(-600 \mathrm{KJmol}^{-1}\right.)\).

Except for fluorides, the solubility of most alkali metal halides decreases as one descends the group because the decrease in hydration energy is greater than the corresponding decrease in lattice energy.

Because of their small size and high electronegativity, lithium halides, with the exception of \(\text {LiF}\), are predominantly covalent and thus soluble in covalent solvents such as alcohol, acetone, ethyl acetate, and pyridine. In contrast, because it is ionic, \(\mathrm{NaCl}\) is insoluble in organic solvents.

Lithium halides are soluble in water due to the high hydration energy of the \(\text {Li}^{+}\) ion, except for \(\text {LiF}\), which is only sparingly soluble due to its high lattice energy. Because the lattice energies of the same alkali metal ion decrease as the size of the halide ion increases, the melting point of the same alkali metal decreases in the order- \(\text {Fluoride} > \text {Chloride} > \text {Bromide} > \text {Iodide}\).

The melting points of lithium halides are lower than those of the corresponding sodium halides for the same halide ion, and they continue to decrease as we move down the group from \(\text {Na}\) to \(\text {Cs}\).

The low melting point of \(\operatorname{LiCl}(887 \mathrm{~K})\) in comparison to \(\mathrm{NaCl}\) is most likely because \(\text {LiCl}\) is covalent in nature and \(\text {NaCl}\) is ionic.

Salts of Oxoacids

Since alkali metals are highly electropositive, their hydroxides are very strong bases and form salts with all oxoacids \(\left(\mathrm{H}_{2} \mathrm{SO}_{4}, \mathrm{H}_{2} \mathrm{CO}_{3}, \mathrm{HNO}_{3}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{HNO}_{2}\right.\) and so on). They are generally water-soluble and heat stable. Alkali metal carbonates \(\left(\text {M}_{2} \text {CO}_{3}\right)\) are remarkably stable up to \(1273 \mathrm{~K}\), after which they melt and eventually decompose to form oxides. However, \(\mathrm{Li}_{2} \mathrm{CO}_{3}\) is much less stable and decomposes quickly.

\(\mathrm{Li}_{2} \mathrm{CO}_{3} \stackrel{\Delta}{\rightarrow} \mathrm{Li}_{2} \mathrm{O}+\mathrm{CO}_{2}\)

This is most likely due to the large size difference between lithium and carbonate ions, which causes the crystal lattice to become unstable.

Alkali metals are also strongly basic, forming solid bicarbonates. While the solution of \(\mathrm{NH}_{4} \mathrm{HCO}_{3}\) bicarbonate exists. No other metals form solid bicarbonates. Lithium bicarbonate, on the other hand, is not solid.

All the carbonates and bicarbonates are water-soluble, and their solubility increases rapidly as the group size decreases. This is because as one moves down the group, the lattice energies decrease faster than the hydration energies.

Other Compounds of the Alkali Metals

Sodium Bicarbonate \(\left(\mathrm{NaHCO}_{3}\right)\)

Carbon dioxide can be absorbed by a concentrated solution of sodium carbonate, yielding sparingly soluble sodium bicarbonate. We can see this by using the following chemical equation:

\(2 \mathrm{Na}_{2} \mathrm{CO}_{3}+\mathrm{CO}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow 2 \mathrm{NaHCO}_{3}\)

Quite sparingly, this happens to be soluble in water. When heated between \(250^{\circ} \mathrm{C}\) and \(300^{\circ} \mathrm{C}\) It becomes pure anhydrous sodium carbonate, which can then be used to standardize acids.

Sodium Chloride (NaCl)

We commonly refer to it as ‘common salt’ because it occurs in abundance in nature as rock salt or halite. The most abundant source of sodium chloride is seawater, which contains \(3\%\) sodium chloride. In large shallow pits, the seawater is exposed to the sun and air to obtain salt.

The crystallisation of salt occurs as a result of the gradual evaporation of water. Later, the solution is saturated with a current of dry hydrogen chloride, causing crystals of pure sodium chloride to separate. Sodium chloride is a colourless crystalline salt that is insoluble in alcohol but highly soluble in water.

Sodium Sulphate \(\left(\mathrm{Na}_{2} \mathrm{SO}_{4}\right)\)

On an industrial scale, anhydrous salt, also known as salt cake, is produced by heating strongly sodium chloride with conc. sulphuric acid.

\(\mathrm{NaCl}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{NaHSO}_{4}+\mathrm{HCl}\)

\(\mathrm{NaCl}+\mathrm{NaHSO}_{4} \rightarrow \mathrm{Na} \mathrm{SO}_{4}+\mathrm{HCl}\)

Glauber’s salt is also known as hydrated sodium sulphate, \(\mathrm{Na}_{2} \mathrm{SO}_{4} \cdot 10 \mathrm{H}_{2} \mathrm{O}\), is prepared from the salt cake by crystallisation from water at temperatures below \(32^{\circ} \mathrm{C}\). This temperature represents the transition temperature for \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{SO}_{4} .10 \mathrm{H}_{2} \mathrm{O}\).

It’s a colourless salt that crystallises in large monoclinic prisms. It is highly soluble in water.

Potassium Chloride \((\text {KCl})\)

It is made from nearly pure potassium chloride that is fused carnallite, separated from the melt, leaving fused \(\mathrm{MgCl}_{2}\) behind. The salt is a colourless cubic crystal-like solid that dissolves in water. Its solubility rises in direct proportion to the temperature.

Potassium Sulphate \(\left(\text {K}_{2} \text {SO}_{4}\right)\)

We can obtain it by strongly heating potassium chloride with concentrated sulphuric acid.

\(\mathrm{KCl}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{KHSO}_{4}+\mathrm{HCl}\)

\(\mathrm{KCl}+\mathrm{KHSO}_{4} \rightarrow \mathrm{K}_{2} \mathrm{SO}_{4}+\mathrm{HCl}\)

It is a colourless crystalline salt that is less water-soluble than sodium sulphate.

Summary

We can conclude that-

1. Oxides are formed when alkali metals react with air.
2. In the presence of oxygen, they produce superoxides and peroxides. In nature, all monoxides are basic. When they come into contact with water, they produce hydroxides.
3. All alkali metal hydroxides, including sodium hydroxide, are strong bases.
4. When alkali metals react with halogens, alkali halides are formed.
5. Alkali metal hydration enthalpies decrease from top to bottom in a group.
6. All alkali metals react with oxo-acids to form salts. Oxoacid is a compound that contains hydrogen, oxygen, and at least one other element that can dissociate to form hydrogen ions.

We also studied some of the compounds of alkali metals like sodium chloride, sodium bicarbonate, potassium chloride, etc.

FAQs on General Characteristics of the Compounds of the Alkali Metals

Q.1. Name the alkali metal which forms superoxides when heated in excess of air and why?
Ans: Potassium forms superoxides when heated in excess of air. This is due to the stabilization of large size cation by large size anion.

Q.2. Why does lithium form only lithium oxide and not peroxide or superoxide?
Ans: Due to the small size of the lithium, it has a strong positive field around it. In combination with the oxide anion, the positive field of lithium restricts the spread of positive charge towards another oxygen atom and thus prevents the formation of higher oxides.

Q.3. Why are alkali metal halides soluble in water?
Ans: Alkali metal halides are soluble in water due to their high ionic character and low lattice energy.

Q.4. What are the characteristics of alkali metals?
Ans: Some of the general characteristics of alkali metals are as follows-
i. They are found in column \(1\text {A}\) of the periodic table.
ii. They have one electron in their outermost layer of electrons.
iii. They can be easily ionized.
iv. Silvery, soft, and less dense.
v. Have low melting points.
vi. They are incredibly reactive.

Q.5. What type of compounds are formed from alkali metals?
Ans: The alkali metals tend to form ionic solids in which the alkali metal has an oxidation number of \(+1\). This is because they have the tendency to lose their one valence electron present in the last shell and form strong ionic bonds with their anions. Therefore, the compounds that are formed from alkali metals are oxides, hydroxides, halides, etc.

Q.6. What are the characteristics of oxides and hydroxides of alkali metals?
Ans: The characteristics of oxides and hydroxides of alkali metals are that they are strong bases and the basicity of alkali metal oxides and hydroxides increase on moving down the group.

Q.7. What are the uses of alkali metals?
Ans: Some of the uses of alkali metals are as follows-
i. Lithium is used in the purification of copper and nickel, where lithium is utilised as a deoxidizer.
ii. Both primary and secondary batteries are made of lithium.
iii. Sodium is used in a sodium vapour lamp.
iv. Sodium is employed in nuclear reactors in molten conditions.
v. Potassium salts are utilised in fertilisers.
vi. Potassium is used as a reducing agent.
vii. Caesium is used in rocket propellant and photographic cells.

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