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What are Alkaline Earth Metals – Definition
The name ‘earth’ was given by chemists to substances like alumina (Al2O3) which were quite stable to heat. The oxides of calcium, strontium, and barium were also quite stable to heat and did not decompose even on strong heating. These metal oxides and also the metals were called alkaline earth.
At present all the elements included in the group 2 family (except Be which has different characteristics) are known as alkaline earth metals. The oxides of these elements are similar to those of alkali metals and dissolve in water to form hydroxides or alkalies.
Alkaline Earth Metals Periodic Table
This group comprises of beryllium (4Be), magnesium (12Mg), calcium (20Ca), strontium (38Sr), barium (56Ba) and radium (88Ra).Just like alkali metals, alkaline earth metals properties show reactive nature and do not exist in the free state.
Beryllium does not occur in abundance and is also difficult to extract, therefore it is unfamiliar metal. Magnesium and calcium occur in considerable amounts in the earth’s crust. Strontium and barium do not occur as concentrated ores but are easy to extract. Radium is a radioactive rare element.
Characteristics of Alkaline Earth Metals
|They are malleable and ductile but less than the alkali metals. Atomic as well as ionic radii increases on moving down the group. The group 2 elements possess the smaller size and greater nuclear charge. The electrons are more tightly held. Hence the first ionization energy is greater than the alkali metals. Their compounds are diamagnetic and colorless, provided their anions are also colorless. On moving down the group, the electropositive character increases due to decreasing ionization energy. They have low electronegativity which decreases on moving down the group. The group 2 elements are the good conductor of heat and electricity.|
Alkaline Earth Metals Properties Comparison with Alkali metals
|When compared to alkali metals, alkaline earth metals are||When compared to alkali metals, alkaline earth metals have|
|(i) less reactive.||(i) smaller atoms and ions.|
|(ii) less electropositive.||(ii) stronger metallic bonds.|
|(iii) less metallic.||(iii) higher m.p. and b.p.|
|(iv) less reducing and less basic.||(iv) high density and are harder.|
Alkaline Earth Metals Properties
Because of their similarity in electronic configuration (ns2), they closely resemble their physical and chemical properties.
Like alkali metals, alkaline earth metal properties show highly reactive nature and hence do not occur free in nature. But unlike alkali metal compounds, alkaline earth metal compounds are insoluble due to high charge present on the ions. Hence group 2 elements occur as insoluble deposits in the earth’s crust in the form of silicates, carbonates, sulfates, and phosphates.
They are all silvery-white metals as they have a few valence electrons than the available orbitals. group 2 elements possess grayish luster when freshly cut. It becomes tarnish soon after their exposure to air. They are malleable and ductile but less than the alkali metals. Although they are sufficiently soft yet less than the alkali metals because metallic bonding (bonding between atoms in its crystal lattice) in these elements is stronger than in group IA elements.
Atomic volume, atomic and ionic radii
Because of the addition of an extra shell of electrons to each element from Be to Ra, the atomic volume increases from Be to Ra.
The atoms of these elements, although fairly large, are smaller than those of the corresponding elements of the group IA. This is due to the higher nuclear charge of these atoms which tends to draw the orbital electrons towards the nucleus. The ions are also large but smaller than those of the elements of group IA. This is again due to the fact that the removal of the two valence electrons to form M2+ ions increases the effective nuclear charge which pulls the electrons inwards and this reduces the size of the ion.
Atomic as well as ionic radii increases on moving down the group on account of the presence of an extra shell at each step.
Since the group 2 elements possess the smaller size and greater nuclear charge than the alkali metals, the electrons are more tightly held.Hence the first ionization energy is greater than the alkali metals.
The second ionization energy is much higher (nearly double) than the first ionization energy.
The general reactivity trend is Ba > Sr > Ca > Mg > Be.
Explanation with the following two ways.
(a) The first ionization energy represents the energy required to remove an electron from a neutral atom (M) while the second ionization energy represents the energy required to remove an electron from a positive ion (M) which, of course, is difficult than the former case.
(b) After the removal of one electron, the effective nuclear charge increases and hence the remaining outermost electron is held even more tightly leading to very high ionization energy. It is interesting to note that although the IE2 of the alkaline earth metals is much higher than the IE1, even then they form M2+ ions. The anomaly is explained due to their high heat of hydration in aqueous solution and high lattice energy in the solid state.
Due to the presence of two s-electrons in the outermost orbital, high heat of hydration of the dipositive ions and comparatively low values of IE2, the alkaline earth metals are bivalent. Moreover, since the bivalent ions have the inert gas configuration, it is very difficult to remove the third electron from the element and hence oxidation state higher than 2 are not encountered. Further, the divalent ion has no unpaired electron, hence their compounds are diamagnetic and colorless, provided their anions are also colorless. Thus, in short, it can be said that the chemistry of these elements is the chemistry of dipositive ions.
Due to their large size and comparatively low ionization energies, the group 2 elements are strongly electropositive elements. However, these are not as strongly electropositive as the alkali metals and hence, unlike alkali metals, these elements do not emit electron on exposure to light. On moving down the group, the electropositive character increases due to decreasing ionization energy.
Since these elements are electropositive, they have low electronegativity which decreases on moving down the group. The relatively high value of Be is due to its small size.
The ionic size increases from Be2+ to Ba2+, so the polarizing power decreases from Be2+ to Ba2+, i.e. Be2+ ion has maximum polarizing power and Ba2+, the least.Therefore, beryllium has a tendency to form covalent compounds.
Heat and electrical conductivity
Due to the presence of two loosely held valence electrons (per atom) which are free to move throughout the crystal structure, the group 2 elements are the good conductor of heat and electricity.
They are denser than the alkali metals because they can be packed more tightly due to their greater charge and smaller radii. The density decreases from Be to Ca and then increases again.
Melting and boiling points
They have high melting points than those of alkali metals. This is because these metals possess two valence electrons and are much strongly bonded in the solid state than the alkali metals.
Chlorides of alkaline earth metal, except that of Be and Mg, produce characteristic color to flame due to easy excitation of electrons to higher energy levels. Beryllium and magnesium atoms, due to their small size, bind their electrons more strongly, i.e. their ionization energies are high. Hence these possess high excitation energy and not excited by the energy of the flame to higher energy state with the result no color is produced in the flame.
Reducing properties (oxidation potential)
Alkaline earth metals have two electrons outside the noble gas configuration. Due to their large size and low ionization energy, they can easily lose outermost two electrons and hence undergo oxidation (loss of electrons) easily.
M → M2+ + 2e–
Thus they are strong reducing agents. Further since the ionisation energy decreases (i.e. electropositive character increases) from Be to Ba, the reducing character of the alkaline earth elements increases from Be to Ba which is evident from the values of their oxidation potential which increases from Be to Ba (recall that higher the value of oxidation potential or higher the negative value of reduction potential greater is the electropositive character and reducing property of the element).
Beryllium, on account of its relatively lower oxidation potential, liberates hydrogen from acids slowly; on the other hand, other elements having high values of oxidation potential react vigorously even with water.
Since the oxidation potential of group 2 elements is lower than that of alkali metals, they are less electropositive and weaker reducing agents than the alkali metals. The lower value of oxidation potentials of alkaline earth metals is due to their high sublimation and ionization energies.
The chemical reactions of the group 2 elements are quite comparable to that of alkali metals. But due to the smaller size and greater charge and hence high ionization energy.So, alkaline earth metal properties show much less reactive than the corresponding alkali metals properties. Further, since their ionization energies decrease with increase in atomic number, their reactivity increases from Be to Ba. Lastly, beryllium being extremely small in size has a unique chemical behavior.
Reaction with air
Their less reactivity than the alkali metals is evident by the fact that they are only slowly oxidized on exposure to air. However, when burnt in the air they form ionic oxides of the type MO, except Ba and Ra which give peroxides. Thus the tendency of the metal to form higher oxides like peroxide increases on moving down the group.
Reaction with water
These metals react slowly with water liberating hydrogen and forming metal hydroxides, e.g.
Ca+ 2H2O → Ca(OH)2 + H2
The reaction with water becomes increasingly vigorous on moving down the group. The reaction of beryllium with water is not certain, magnesium reacts only with hot water, while other metals react with cold water but, of course, not as energetically as the alkali metals.
Ba > Sr > Ca > Mg > Be (Reactivity with water)
The inertness of Be and Mg towards water is due to the formation of a protective thin layer of hydroxide on the surface of the metals.
Reaction with Hydrogen
All these elements, except beryllium, combine with hydrogen to form hydrides,(metal) MH2. Magnesium hydride (like BeH2) is covalent while other hydrides are ionic.
Reaction with halogens
All these elements combine with halogens at elevated temperatures forming halides, MX2. Beryllium halides are covalent, while the rest are ionic and thus dissolve in water and conduct electricity in aqueous solution and in the molten state. The solubility of halides (except fluoride) decreases on moving down the group; fluorides are almost insoluble in water.
Action with nitrogen
All these elements burn in nitrogen forming nitrides, (metal)M3N2 which react with water to liberate ammonia.
3Ca + N2 → Ca3N2
Ca3N2 + 6H2O→3Ca(OH)2 + 2NH3
The ease of formation of nitrides decreases on moving down the group.
Action with acids
On account of their high oxidation potentials, they readily liberate hydrogen from dilute acids. For example,
Mg + 2HCL—> MgCl2 + H2
The reactivity of group 2 elements increases on moving down the group. This is due to increase in electropositive character from Be to Ba. Thus beryllium reacts very slowly, Mg reacts very rapidly while Ca, Sr and Ba react explosively.
Formation of amalgam and alloys
They form an amalgam with mercury and alloys with other metals.
As described earlier, complex formation is favored by small size, highly charged ion and suitable empty orbitals; alkaline earth metal ions, not having these characteristics, do not have a significant tendency (although it is more than in the alkali metals by virtue of their double charge) to form complexes. However, beryllium, due to its small size, forms a number of stable complexes, e.g., [BeF2]–, [BeF4 ]2-,[Be(H2O)4]2+ etc.
Solubility in liquid ammonia
Like alkali metals, alkaline earth metal properties show a behavior to dissolve in liquid ammonia giving colored solutions. When the metal-ammonia solutions are evaporated, hexammoniates M(NH3)6 are formed. The ammoniates are good conductors of electricity and decompose at high temperatures. The tendency for the formation of ammoniates decreases with increases in the size of the metal atom, i.e., on moving down the group.
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