Alkaline Earth Metals are considered to have low electronegativity and electron affinity. The long form of the periodic table, beryllium, magnesium, calcium, strontium, barium, and radium are present in group \(2.\) All the elements of group \(2\) are called alkaline earth metals. This article covers the general characteristics of the second group of elements present in s-block, Alkaline Earth Metals.

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Alkaline Earth Metals: Occurence

Alkaline earth metals are highly reactive metals. Hence, it does not occur in state however, is widely distributed within the nature in combined state as silicates, carbonates, sulphates, and phosphates.

Electronic Configuration

All alkaline earth metal atoms have two electrons in their valence shell preceded by the noble gas configuration. Therefore, their general configuration may be written as (Noble gas) \({\rm{n}}{{\rm{s}}^2}\) where n represents the valence shell. The orbital electronic configuration of elements given as:

Elements	Atomic Number	Electronic Configuration	With Inert gas core
Beryllium \(({\rm{Be}})\)	\(4\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{He}}]2{{\rm{s}}^2}\)
Magnesium \(({\rm{Mg}})\)	\(12\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{Ne}}]3{{\rm{s}}^2}\)
Calcium \(({\rm{Ca}})\)	\(20\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{Ar}}]4{{\rm{s}}^2}\)
Strontium \(({\rm{Sr}})\)	\(38\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{Kr}}]5{{\rm{s}}^2}\)
Barium \(({\rm{Ba}})\)	\(56\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{6}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{Xe}}]6{{\rm{s}}^2}\)
Radium \(({\rm{Ra}})\)	\(88\)	\({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{6}}}{\rm{4}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{f}}^{{\rm{14}}}}{\rm{5}}{{\rm{s}}^{\rm{2}}}{\rm{5}}{{\rm{p}}^{\rm{6}}}{\rm{5}}{{\rm{d}}^{{\rm{10}}}}{\rm{6}}{{\rm{s}}^{\rm{2}}}{\rm{6}}{{\rm{p}}^{\rm{6}}}{\rm{7}}{{\rm{s}}^{\rm{2}}}\)	\([{\rm{Rn}}]7{{\rm{s}}^2}\)

General Characteristics of Alkaline Earth Metals

Now let’s discuss alkaline earth metals characteristics in detail.

Physical Characteristics

1. Atomic and Ionic Radii – Due to the high nuclear charge of alkaline earth metals, the electrons are attracted strongly towards the nucleus. As a result, the atomic and ionic radii are smaller than those of the corresponding alkali metals; on moving down the group, atomic and ionic radii increase due to adding an extra shell of electrons in each succeeding element, increasing the screening effect.

2. Melting and Boiling Point- They have higher melting and boiling points than those of alkali metals. However, down the group, there is no regular trend in their melting and boiling points.

3. Ionization Energy- Alkaline earth metal has low ionization energies due to the relatively large size of the atoms; since the atomic size increases down the group, the ionization energy decreases. A comparison of the ionization of energies of the members of group \(1\) and \(2\) shows that the members present in the second group have higher values than those of group \(1\) because they have smaller sizes, and electrons are more attracted towards the nucleus of the atoms. It noted that although \({\rm{I}}{{\rm{E}}_1}\) and values of alkaline earth metals are higher than those of alkali metals, the \({\rm{I}}{{\rm{E}}_2}\) values of alkaline earth metals are much smaller than those of alkali metals. Therefore, the ionization energy values of sodium and magnesium are given below.

Element	\({\rm{I}}{{\rm{E}}_1}\)	\({\rm{I}}{{\rm{E}}_2}\)
\({\rm{Na}}\)	\(496\,{\rm{KJ}}\,{\rm{mo}}{{\rm{l}}^{ – 1}}\)	\(4562\,{\rm{KJ}}\,{\rm{mo}}{{\rm{l}}^{ – 1}}\)
\({\rm{Mg}}\)	\(737\,{\rm{KJ}}\,{\rm{mo}}{{\rm{l}}^{ – 1}}\)	\(1540\,{\rm{KJ}}\,{\rm{mo}}{{\rm{l}}^{ – 1}}\)

4. Electro Positive and Metallic Character- Because of the lower ionization energies of alkaline earth metals, they are strongly electropositive in nature. However, they are not as electropositive solid as the alkali metals of group \(1\) because of comparatively higher ionization energies. Thus, the electropositive character increases down the group that is from beryllium to barium.

5. Characteristic Flame Colourization- Except for \({\rm{Be}}\) and \({\rm{Mg,}}\) the alkaline earth metals salts impart distinctive colour to the flame.

6. Hydration Enthalpy- Like Alkaline metals, the divalent cations of the alkaline earth metals also tend to get hydrated. The negative hydration enthalpies are more due to the smaller size of cations than the cations of the alkali metals present in the same period. For example, \({\rm{L}}{{\rm{i}}^{\rm{ + }}}\left( {\Delta {\rm{hydH = }}\,{\rm{ – 506}}\;{\rm{kJ}}\;{\rm{mo}}{{\rm{l}}^{{\rm{ – 1}}}}} \right);{\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}\left( {\Delta {\rm{hydH = }}\,{\rm{ – 2494}}\;{\rm{kJ}}\;{\rm{mo}}{{\rm{l}}^{{\rm{ – 1}}}}} \right).\) The hydration enthalpies decrease down the group since the ionic size increases.

\({\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}{\rm{ > M}}{{\rm{g}}^{{\rm{2 + }}}}{\rm{ > C}}{{\rm{a}}^{{\rm{2 + }}}}{\rm{ > S}}{{\rm{r}}^{{\rm{2 + }}}}{\rm{ > B}}{{\rm{a}}^{{\rm{2 + }}}}\)

Alkaline earth metals have more significant hydration enthalpies. So these are extensively hydrated than alkali metals. Therefore, for example, \({\rm{MgC}}{{\rm{l}}_2}\) exists as \({\rm{MgC}}{{\rm{l}}_2} \cdot 6{{\rm{H}}_2}{\rm{O}}\) while \({\rm{CaC}}{{\rm{l}}_2}\) is \({\rm{CaC}}{{\rm{l}}_2} \cdot 6{{\rm{H}}_2}{\rm{O}}.\) On the other hand, the corresponding salts of alkali metals and \({\rm{NaCl}}\) and the \({\rm{KCl}}\) are anhydrous since their cations are not hydrated so extensively.

7. Density- Alkaline earth metals are denser than the alkali metals present in the same period because these are more closely packed due to their smaller size and stronger metallic bonds. However, the trend in the density is not uniform. It initially increases from beryllium to calcium and then decreases from calcium to barium. The irregular trends in densities are due to the differences in the crystal structures of these metals.

8. Nature of Bond Formed- Like alkali metals, alkaline earth metals predominantly form ionic compounds that are less ionic than the corresponding alkali metal compounds. The tendency to form ionic compounds increases down the group. The first member, beryllium, however, forms covalent compounds magnesium also shows some trend for covalency. All other elements include ionic compounds.

Chemical Properties

Alkaline earth metal is smaller in size when compared to the corresponding alkali metals. However, they have higher cohesive energy, higher ionization potential and higher enthalpy of sublimation as well. The reactivity of alkaline earth metals is high however, it is relatively less when compared to the reactivity of alkali metals. It will be quite apparent from the following discussion concerning these elements.

1. Reaction with oxygen (Formation of oxides)- The alkaline earth metals are less electropositive than alkali metal reactive with air or oxygen slowly upon heating to form oxides. However, radium forms peroxides.

\({\rm{2M + }}{{\rm{O}}_2} \to {\rm{2MO}}{\mkern 1mu} ({\rm{M}}{\mkern 1mu} {\rm{ = }}{\mkern 1mu} {\rm{Be}},{\rm{Mg}}{\mkern 1mu} {\rm{or}}{\mkern 1mu} {\rm{Ca}})\)

\({\rm{M + }}{{\rm{O}}_2} \to {\rm{M}}{{\rm{O}}_2}{\mkern 1mu} ({\rm{M}}{\mkern 1mu} {\rm{ = }}{\mkern 1mu} {\rm{Be}},{\rm{Mg}}{\mkern 1mu} {\rm{or}}{\mkern 1mu} {\rm{Ca}})\)

The reactivity with oxygen increases as we move down the group due to the increasing electropositive character of the elements.
The decomposition of their carbonates can prepare the monoxide.

\(\mathop {{\rm{MC}}{{\rm{O}}_{\rm{3}}}}\limits_{{\rm{Carbonate}}} \to {\rm{MO + C}}{{\rm{O}}_{\rm{2}}}({\rm{M}}\,{\rm{ = }}\,{\rm{Mg}},{\rm{Ca}},{\rm{Sr}}\,{\rm{or}}\,{\rm{Ba}})\)

Among the oxides, beryllium oxide is amphoteric, while oxides of other elements are basic. The amphoteric character of beryllium oxide can be supported by reacting with acids and alkalis.

\({\rm{BeO}} + {\rm{HCl}} \to {\rm{BeC}}{{\rm{l}}_2} + {{\rm{H}}_2}{\rm{O}}\) (basic)

\({\rm{BeO}} + {\rm{NaOH}} \to {\rm{N}}{{\rm{a}}_2}{\rm{Be}}{{\rm{O}}_2} + {{\rm{H}}_2}{\rm{O}}\) (Acidic)

The oxides of beryllium and magnesium are almost insoluble in water, while oxides of the rest of the metals dissolve in water to form hydroxides.

\({\rm{CaO}} + {{\rm{H}}_2}{\rm{O}} \to {\rm{Ca}}{({\rm{OH}})_2} + {\rm{Heat}}\)

The insolubility of beryllium oxide and magnesium oxide in water is due to the considerable lattice energies.

2. Combination with hydrogen – Formation of hydrides- All the elements except beryllium combined with hydrogen upon heating form ionic hydrides \({\rm{M}}{{\rm{H}}_{\rm{2}}}\).

\({\rm{Ca + }}{{\rm{H}}_2} \to {\rm{Ca}}{{\rm{H}}_2}\)

3. Action with water – Formation of hydroxides- The alkaline earth metals have a lesser tendency to react with water than alkali metals. It is because they combine with water to form hydroxides.

\({\rm{Ca}} + 2{{\rm{H}}_2}{\rm{O}} \to {\rm{Ca}}{({\rm{OH}})_2} + {{\rm{H}}_2}\)

Like alkali metals, the hydroxides are the basis in nature and basic strength increases down the family.

The basic character of hydroxides is due to the low ionization energies of these metals.  Because of low and ionization energies, the \({\rm{M}} – {\rm{O}}\) bond in \({\rm{M}}{({\rm{OH}})_2}\) is weak, and it can cleave to give \({\rm{O}}{{\rm{H}}^ – }\) ions in solution. Since the ionization energy decreases down the family, the basic strength of the hydroxide increases. For example, barium hydroxide is amphoteric, and magnesium hydroxide is mildly basic, while others are strong bases.

4. Formation of Carbonates- The alkaline earth metals can be prepared by passing carbon dioxide in limited supply through the solution of their hydroxides.

\({\rm{Ca}}{({\rm{OH}})_2} + {\rm{C}}{{\rm{O}}_2} \to {\rm{CaC}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{O}}\)

5. Formation of Sulphates- It can form by diluting sulphuric acid on metals, metal oxides, metal hydroxides and metal carbonates.

\({\rm{Ca}} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + {{\rm{H}}_2}\)

\({\rm{CaO}} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + {{\rm{H}}_2}{\rm{O}}\)

\({\rm{Ca}}{({\rm{OH}})_2} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + 2{{\rm{H}}_2}{\rm{O}}\)

\({\rm{CaC}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + {{\rm{H}}_2}{\rm{O}} + {\rm{C}}{{\rm{O}}_2}\)

The sulphates of these metals are all white solids and stable to heat. Beryllium sulphate and magnesium sulphate are readily soluble in water. Therefore, their solubilities decrease on going down the group. It is due to the higher lattice energies of alkaline earth metal sulphates than alkaline earth metal sulphates. On moving down the group, the hydration energy decreases with an increase in the size of the metal ion, and consequently, their solubilities decrease.

6. Complex Formation- Generally, the members of the family do not form complexes. However, smaller lions like \({\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}\) and \({\rm{M}}{{\rm{g}}^{2 + }}\) form complexes with the electron donor species. For example, the stable complexes of beryllium are \({\left( {{\rm{Be}}{{\rm{F}}_{\rm{3}}}} \right)^{\rm{ – }}}{\rm{,}}{\left( {{\rm{Be}}{{\rm{F}}_{\rm{4}}}} \right)^{{\rm{2 – }}}}\) and   \({\left( {{\rm{Be}}{{\left( {{{\rm{H}}_{\rm{2}}}{\rm{O}}} \right)}_{\rm{4}}}} \right)^{{\rm{2 + }}}}\).

Chlorophyll, the complex compound of magnesium, plays an essential role in photosynthesis calcium, Strontium, and barium form complexes only with potent complexing agents like acetylacetone and EDTA.

Alkaline Earth Metal Uses

Uses of Beryllium

1. Alloys of beryllium are quite helpful. For example, \({\rm{Cu – Be}}\) alloy uses in making high-strength springs.
2. The metal is used for making windows of \({\rm{X – }}\)ray tubes.
3. It is a source of neutrons when bombarded with alpha- particles.
4. Beryllium and its compounds find applications in making computer parts.

 Uses of Magnesium

1. It is used as a reducing agent in the preparation of silicon and boron from their oxides.
2. Magnesium is used in the preparation of Grignard reagents, which are alkyl magnesium halides. These are organic compounds and used in the synthesis of a large number of compounds.
3. It was employed as a deoxidizer in metallurgy and removed the last traces of air from radio values.
4. Magnesium ribbon, along with barium peroxide, is used as an ignition mixture in the aluminothermic process.
5. Magnesium is a constituent of the green pigment in plants known as chlorophyll. It is responsible for the synthesis of carbohydrates in plants in the presence of sunlight due to the reaction.

Uses of Calcium

1. Calcium is a powerful reducing agent used to extract several metals from their oxides.
2. Calcium has a strong affinity for both oxygen and nitrogen present in the air. Therefore, it is commonly used to remove traces of air from vacuum tubes.
3. It is preferred over sodium for removing traces of water from alcohol since it does not react with alcohol at all.
4. It is used for removing sulphur from petroleum.

Barium is used as a deoxidizer in the metallurgy of copper, lubricant for anode rotors in \({\rm{X – }}\)ray tubes, and spark-plug alloys-radium salts used in radiotherapy to treat cancer.

Summary

From what we discussed so far, it is clear that group II elements and their compounds show gradation in their physical and chemical properties. Hence, their inclusion in group II of the periodic table is justified. In this article, we studied the second group of metals- their occurrence, properties, and trends. We also learned the uses of various alkaline earth metals.

Frequently Asked Questions (FAQs) – Alkaline Earth Metals

Frequently asked questions related to alkaline earth metals are listed as follows:

Q.1. What are the characteristics of alkaline earth metals? Ans: Alkaline earth metals exhibit gradation in their physical and chemical properties. Compared to alkali metals, these metals are harder with higher melting and boiling points. They show a +2 Oxidation state. Some appear white, but magnesium and beryllium appear greyish. Alkaline earth metals give different colours with a flame test such as strontium gives crimson colour barium gives apple green colour calcium gives Brick red colour.

Q.2. Why are they called alkaline earth metals? Ans: The oxides of these elements, for example, lime (CaO), strontian (SrO), baryta (BaO), etc., are thermally stable and exhibit basic character. Therefore, these elements are known as alkaline earth metals or alkaline earth.

Q.3. What are the six alkaline earth metals? Ans: The six elements are beryllium, magnesium, calcium, strontium, barium, and radium.

Q.4. What are alkaline earth metals used for? Ans: Alloys of beryllium are quite helpful—for example, Cu–Be alloy used in making high strength springs. Magnesium is used in the preparation of Grignard reagents, which are alkyl magnesium halides. Barium is used as a deoxidizer in the metallurgy of copper, lubricant for anode rotors in X–ray tubes, and spark-plug alloys. Calcium has a strong affinity for both oxygen and nitrogen present in the air. Therefore, it is used to remove traces of air from vacuum tubes.

Q.5. What is special about alkaline earth metals? Ans: The elements form a well-graded series of highly reactive metals but are less reactive than group 1. These are typically divalent and generally include colourless ionic compounds. The oxides and hydroxides are less basic than those of group 1; hence their oxo salts are less stable to heat. Magnesium is an important structural metal and used in large amounts. Several compounds are used in vast quantities: limestone is used to make quicklime and cement, and large quantities of Chalk are also used. Other compounds used on a larger scale include gypsum CaSO4, fluorite CaF2 and barytes BaSO4.