The dipositive oxidation state \(\left( {{{\rm{M}}^{{\rm{2 + }}}}} \right)\) is the dominant valence of Group \(2\) elements. Because of the greater nuclear charge and smaller size, alkaline earth metals create compounds that are primarily ionic but less ionic than alkali metal compounds. Read the article below to learn more about the general characteristics of compounds of alkaline earth metals.

Alkaline Earth Metals

Alkaline earth metals form compounds that are less ionic than alkali metal compounds because of their higher nuclear charge and smaller size. \({\rm{Be}}\) and \({\rm{Mg}}\) oxides and other salts are typically more covalent than those formed by heavier and larger members. The general characteristics of some of the compounds of alkaline earth metals are described below:

Oxides

The preparation and uses of oxides are explained below:

Oxides Preparation

The oxides of alkaline earth metals \({\rm{MO}}\) are obtained either by heating the metals in dioxygen or by thermal decomposition of their carbonates.

Oxides Properties and Uses

These have high melting points, very low vapour pressures, are excellent heat conductors, and are chemically inert. Because of these properties, these oxides are used for lining furnaces and hence are used as refractory materials.

Hydroxides

\({\rm{Ca,}}\) \({\rm{Sr,}}\) and \({\rm{Ba}}\) hydroxides are made by either treating the metal with cold water or reacting the corresponding oxides with water. The reaction of these oxides with \({{\rm{H}}_{\rm{2}}}{\rm{O}}\)is also known as slaking.

\({\rm{Be}}\)\({\left( {{\rm{OH}}} \right)_{\rm{2}}}\) and \({\rm{Mg}}{\left( {{\rm{OH}}} \right)_{\rm{2}}}\) are insoluble in water, and these are obtained from suitable metal ion solutions by precipitation with \({\rm{O}}{{\rm{H}}^{\rm{ – }}}\) ions.

Properties of Hydroxides

I. Basic Character:  \({\rm{Be}}{\left( {{\rm{OH}}} \right)_{\rm{2}}}\) is amphoteric due to its small size and high ionisation enthalpy. As a result, it dissolves in both acids and bases.

As we move down towards the group, their basic strength increases. Because of their low ionisation enthalpies, these hydroxides have a basic character.

II. Solubility in Water: In comparison to alkali metal hydroxides, alkaline earth metal hydroxides are less soluble in water. The solubility of alkaline earth metal hydroxides in water increases as the atomic number increases down the group. This is because, as the size of the cation increases, both lattice enthalpy and hydration enthalpy decrease down the group, but lattice enthalpy decreases faster than hydration enthalpy, and thus their solubility increases down the group.

Halides

The preparation and uses of halides are explained below:

Halides Preparation

At appropriate temperatures, alkaline earth metals combine directly with halogens to form halides, \({\rm{M}}{{\rm{X}}_{\rm{2}}}{\rm{.}}\)

These halides can also be prepared by the action of halogen acids \(\left( {{\rm{HX}}} \right)\) on metals, metal oxides, hydroxides and carbonates.

Properties of Halides

I. All beryllium halides are covalent and soluble in organic solvents. They are hygroscopic and emit fumes into the atmosphere as a result of hydrolysis. They produce an acidic solution when hydrolysed.

II. The halides of Be are covalent while halides of all other alkaline earth metals are ionic. Their ionic character, however, increases as the size of the metal ion increases. As the ionic character increases or the covalent character decreases, their tendency towards hydrolysis increases: \({\rm{BeC}}{{\rm{l}}_{\rm{2}}}{\rm{ < MgC}}{{\rm{l}}_{\rm{2}}}{\rm{ < CaC}}{{\rm{l}}_{\rm{2}}}{\rm{ < SrC}}{{\rm{l}}_{\rm{2}}}{\rm{ < BaC}}{{\rm{l}}_{\rm{2}}}\)

III. Except for \({\rm{BeC}}{{\rm{l}}_{\rm{2}}}{\rm{,}}\) all other chlorides in group \(2\) form hydrates, but their proclivity to form hydrates decreases down the group. \({\rm{MgC}}{{\rm{l}}_{\rm{2}}}{\rm{.6}}{{\rm{H}}_{\rm{2}}}{\rm{O,}}\) \({\rm{CaC}}{{\rm{l}}_{\rm{2}}}{\rm{.6}}{{\rm{H}}_{\rm{2}}}{\rm{O, SrC}}{{\rm{l}}_{\rm{2}}}{\rm{.2}}{{\rm{H}}_{\rm{2}}}{\rm{O,}}\) and \({\rm{BaC}}{{\rm{l}}_{\rm{2}}}{\rm{.2}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\) are a few examples.

IV. The hydrated chlorides, bromides and iodides of \({\rm{Ca,}}\) \({\rm{Sr}}\) and \({\rm{Ba}}\) can be dehydrated on heating but those of \({\rm{Be}}\) and \({\rm{Mg}}\) undergo hydrolysis. For example,

Uses of Halides

I. In the electrolytic extraction of magnesium, anhydrous \({\rm{MgC}}{{\rm{l}}_{\rm{2}}}\) is used.
II. Calcium fluoride, also known as fluorspar \(\left( {{\rm{Ca}}{{\rm{F}}_{\rm{2}}}} \right){\rm{,}}\) is by far the most important of all alkaline earth metal fluorides because it is the only large-scale source of fluorine.

Sulphates

The preparation and uses of sulphates are explained below:

Preparation of Sulphates

The sulphates of alkaline earth metals \(\left( {{\rm{MS}}{{\rm{O}}_{\rm{4}}}} \right)\) are prepared by the action of sulphuric acid on metals, metal oxides, hydroxides and carbonates.

Properties of Sulphates

I. The sulphates of alkaline earth metals are all white solids. Beryllium, magnesium, and calcium sulphates crystallise in hydrated form, i.e., \({\rm{BeS}}{{\rm{O}}_{\rm{4}}}{\rm{.4}}{{\rm{H}}_{\rm{2}}}{\rm{O, MgS}}{{\rm{O}}_{\rm{4}}}{\rm{.7}}{{\rm{H}}_{\rm{2}}}{\rm{O,}}\) and \({\rm{CaS}}{{\rm{O}}_{\rm{4}}}{\rm{.2}}{{\rm{H}}_{\rm{2}}}{\rm{O,}}\) whereas strontium and barium sulphates crystallise without crystallisation water.

II. Sulphate solubility in water decreases as one moves down the group, i.e., \({\rm{Be  >  Mg  >  >  Ca  >  Sr  >  Ba}}{\rm{.}}\) Thus, \({\rm{BeS}}{{\rm{O}}_{\rm{4}}}\) and \({\rm{MgS}}{{\rm{O}}_{\rm{4}}}\) are highly soluble, \({\rm{CaS}}{{\rm{O}}_{\rm{4}}}\) is sparingly soluble, but Sr, Ba, and Ra sulphates are practically insoluble.

III. When alkaline earth metal sulphates are heated, they decompose into their corresponding oxides and \({\rm{S}}{{\rm{O}}_{\rm{3}}}{\rm{.}}\)

Uses of Sulphates

I. The almost negligible solubility of \({\rm{BaS}}{{\rm{O}}_{\rm{4}}}\) in water is used in the detection and estimation of ions.
II. \({\rm{BaS}}{{\rm{O}}_{\rm{4}}}\)is both insoluble in \({{\rm{H}}_{\rm{2}}}{\rm{O}}\) and opaque to \({\rm{X}}\) -rays. Therefore, “barium meal’ is used to obtain a shadow of the stomach on an \({\rm{X}}\) -ray film which is useful in diagnosing stomach ulcers.

Carbonates

The preparation and uses of carbonates are explained below:

Preparation of Carbonates

Alkaline earth metal carbonates are obtained as white precipitates when
(i). A calculated amount of carbon dioxide is passed through the solution of the alkaline metal hydroxides

(ii). The solution of an alkaline earth metal salt, such as \({\rm{CaC}}{{\rm{l}}_{\rm{2}}}{\rm{,}}\) is enriched with sodium or ammonium carbonate.

Properties of Carbonates

I. All carbonates are ionic, but beryllium carbonate is prone to hydrolysis. It contains the hydrated ion, \({\left[ {{\rm{Be}}{{\left( {{{\rm{H}}_{\rm{2}}}{\rm{O}}} \right)}_{\rm{4}}}} \right]^{{\rm{2 + }}}}\) rather than \({\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}\) and hence is precipitated only in an atmosphere of \({\rm{C}}{{\rm{O}}_{\rm{2}}}{\rm{.}}\)

II. Solubility: Magnesium carbonates and other alkaline earth metal carbonates are sparingly soluble in water, and their solubility decreases down the group from \({\rm{Be}}\) to \({\rm{Ba}}{\rm{.}}\) \({\rm{MgC}}{{\rm{O}}_{\rm{3}}}{\rm{,}}\) for example, is slightly soluble in water, whereas \({\rm{BaC03}}\) is nearly insoluble.

III. Thermal stability: When heated, the carbonates of all alkaline earth metals decompose to form the corresponding metal oxide and \({\rm{C}}{{\rm{O}}_{\rm{2}}}{\rm{.}}\) However, as the electropositive character of the metal or the basicity of metal hydroxide increases from \({\rm{Be}}{\left( {{\rm{OH}}} \right)_{\rm{2}}}\) to \({\rm{Ba}}{\left( {{\rm{OH}}} \right)_{\rm{2}}}{\rm{,}}\) the temperature of decomposition, i.e., the thermal stability of these carbonates, increases down the group. For example,

Bicarbonates

The preparation and uses of Bicarbonates are explained below:

Preparation of Bicarbonates

Alkaline earth metal bicarbonates are made by passing \({\rm{C}}{{\rm{O}}_{\rm{2}}}\) through a suspension of metal carbonates in the water. All alkaline earth metal bicarbonates are only stable in solution and have never been isolated in their pure condition.

Uses of Bicarbonates

In qualitative analysis, the extremely low solubility of alkaline earth carbonates in water is used. For example,
I. The cations of group \({\rm{V }}\left( {{\rm{B}}{{\rm{a}}^{\rm{ + }}}{\rm{, S}}{{\rm{r}}^{{\rm{2 + }}}}{\rm{, C}}{{\rm{a}}^{{\rm{2 + }}}}} \right)\) in the qualitative analysis are precipitated as their insoluble carbonates from the solution of their soluble salts by adding \({\left( {{\rm{N}}{{\rm{H}}_{\rm{4}}}} \right)_{\rm{2}}}{\rm{C}}{{\rm{O}}_{\rm{3}}}\) in the presence of \({\rm{N}}{{\rm{H}}_{\rm{4}}}{\rm{Cl}}\) and excess of \({\rm{N}}{{\rm{H}}_{\rm{4}}}{\rm{OH}}{\rm{.}}\)
II. Soluble carbonates, i.e., alkali metal and \(\left( {{\rm{NH}}_4^ + } \right)\) ion carbonates, are detected by precipitating them as insoluble magnesium carbonate.
III. \({\rm{CaC}}{{\rm{O}}_{\rm{3}}}\) is used in the Solvay-ammonia process to produce \({\rm{N}}{{\rm{a}}_{\rm{2}}}{\rm{C}}{{\rm{O}}_{\rm{3}}}{\rm{,}}\) as well as in the production of glass and cement.

Nitrates

The preparation and uses of nitrates are explained below:

Preparation of Nitrates

The action of \({\rm{HN}}{{\rm{O}}_{\rm{3}}}\) on oxides, hydroxides, and carbonates results in the formation of alkaline earth metal nitrates in the solution, which can then be crystallised as hydrated salts.

Magnesium nitrate crystallises as \({\rm{Mg}}{\left( {{\rm{N}}{{\rm{O}}_{\rm{3}}}} \right)_{\rm{2}}}{\rm{.6}}{{\rm{H}}_{\rm{2}}}{\rm{O,}}\) while \({\rm{Ba}}{\left( {{\rm{N}}{{\rm{O}}_{\rm{3}}}} \right)_{\rm{2}}}\) crystallises as the anhydrous salt.

Beryllium nitrate is unusual because it forms basic nitrate in addition to the normal salt. This again shows that tendency to form hydrates decreases with increasing size and decreasing hydration enthalpy down the group.

Properties and Uses of Nitrates

I. All nitrates are water-soluble and decompose on heating to produce the corresponding oxides, with the evolution of a mixture of \({\rm{N}}{{\rm{O}}_{\rm{2}}}\) and \({{\rm{O}}_{\rm{2}}}{\rm{.}}\)

II. Strontium and barium nitrates are used in pyrotechnics for giving red and green flames, respectively.

Oxalates

\({\rm{Ca, Sr,}}\) and \({\rm{Ba}}\) oxalates are sparingly soluble in water, but their solubility increases from \({\rm{Ca}}\) to \({\rm{Ba}}{\rm{.}}\) Beryllium oxalate, on the other hand, is water-insoluble.

Summary

Because of their larger nuclear charge and smaller size, alkaline earth metals create compounds that are less ionic than alkali metal complexes. As oxides are utilised for lining furnaces and hence as refractory materials, these compounds have a variety of applications in industries and other settings. The ions are detected and estimated using \({\rm{BaS}}{{\rm{O}}_{\rm{4}}}\) extremely negligible solubility in water. Insoluble magnesium carbonate is used to identify soluble carbonates, such as alkali metal and \({\rm{N}}{{\rm{H}}_{\rm{4}}}\) ion carbonates.

FAQs on General Characteristics of Compounds of Alkaline Earth Metals

Q.1. What are the 2 general characteristics and gradation properties of alkaline earth metals?
Ans: General characteristics of alkaline earth metals are as follows:
I. They have a valence shell electronic configuration of \({\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{.}}\)
II. To form dipositive metal ions, they lose two valence electrons. As a result, they achieve a stable noble gas electronic configuration.
Gradation properties of alkaline earth metals are as follows:
I. \({\rm{Be}}\) and \({\rm{Mg}}\) react with air and water to form an inert layer of oxide on their surfaces. As a result, they are inert to air and water.
II. Alkaline earth metals form halides when they react with halogen at high temperatures.

Q.2. What kind of compounds do alkaline earth metals form?
Ans: Alkaline earth metals form ionic compounds because alkaline-earth metals easily lose electrons to become positive ions.

Q.3. What are the common physical and chemical features of alkaline earth metal?
Ans: Physical features of alkaline earth metals are as follows:
I. Alkaline earth metals are silvery-white, lustrous, and relatively soft but harder than alkali metals.
II. The boiling point and melting points of these metals are high due to their smaller size.
III. These are strongly electropositive in nature due to low ionisation enthalpies.
IV. The alkaline earth metals have high electrical and thermal conductivities.
Chemical features of alkaline earth metals are as follows:
I. Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.
II. Reactivity towards the halogens: All alkaline earth metals combine with halogens at high temperatures to form their halides.
III. Reactivity towards hydrogen: Except for beryllium, all elements combine with hydrogen to form hydrides when heated.
IV. Alkaline earth metals are strong reducing agents.
V. Alkaline earth metals readily react with acids by liberating dihydrogen.
VI. Alkaline earth metals dissolve in liquid ammonia to give a deep blue-black solution forming ammoniated ions.

Q.4. What are the main characteristics of alkali metals?
Ans: The main characteristics of alkali metals are as follows:
I. The alkali metals are shiny white and soft.
II. They can be easily cut with a knife.
III. Francium, the final metal in this group, is radioactive.
IV. Because alkali metals are extremely reactive, they can only be found in nature as compounds.
V. They are highly reactive metals that produce strong alkaline oxides and hydroxides.
VI. The oxidation state of all alkali metals is \({\rm{ + 1}}{\rm{.}}\)
VII. The alkali metals produce a distinct colour in the bunsen flame.

Q.5. What are the general characteristics of alkaline earth metals?
Ans: General characteristics of alkaline earth metals are as follows:
I. They have a valence shell electronic configuration of \({\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{.}}\)
II. To form dipositive metal ions, they lose two valence electrons. As a result, they achieve a stable noble gas electronic configuration.
III. Alkaline earth metal atomic and ionic radii are smaller than alkali metal atomic and ionic radii.
IV. Alkaline earth metals are highly electropositive in nature due to their lower ionisation enthalpy values. The electropositive nature increases from Be to Ba as one moves down the group.
V. The ionisation enthalpy value of alkaline earth metals is very low due to their large atomic size. The values of the first ionisation enthalpy are greater than those of alkali metals.

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