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MAIN GROUPS CHEMISTRY

Alkaline Earth Metals ( Group IIA)

GROUP 2 ELEMENTS : ALKALINE EARTH METALS

INTRODUCTION

The group 2 elements comprise

beryllium, magnesium, calcium,

strontium, barium and radium.

They follow alkali metals in the

periodic table.

These (except beryllium) are known

as alkaline earth metals.

The first element Be differs from

the rest of the members and shows

diagonal relationship to Al.

o Be, very toxic if its

compounds are inhaled

(destroys lungs)- minor

element in terms of

technical importance

o Mg, Ca, Sr, Ba are in

many common minerals and

in the ocean, e.g.

limestone (CaCO3),

dolomite (CaCO3∙ MgCO3)

o Ra all isotopes of this

element are radioactive


INTRODUCTION

These elements have two electrons in the s -orbital of the valence

shell.

Their general electronic configuration may be represented as

[noble gas] ns2.

Alkali earth metal are harder, denser and stronger than alkali

metals.

They are less reactive than alkali metals but are also too reactive to

be found in their free state in nature.


INTRODUCTION

Like alkali metals, the compounds of these elements are also ionic.

mostly form ionic compounds with(highly electronegative nonmetals

found in

Groups 6 (oxides,(sulfides) and 7 (halides), as well as P((phosphides) and

N (nitrides) from Group 5 and C (carbides) from Group 4,

however, Be forms covalent compounds and Mg sometimes forms

covalent compounds


Atomic and Ionic Radii

The atomic and ionic radii of the alkaline earth

metals are smaller than those of the

corresponding alkali metals in the same periods.

Within the group, the atomic and ionic radii

increase with increase in atomic number.


Ionization energies

The alkaine earth metals have low ionization enthalpies due to fairly large

size of the atoms.

Since the atomic size increases down the group, their ionization enthalpy

decreases.

The first ionisation enthalpies of the alkaline earth metals are higher than

those of the corresponding Group 1 metals.

This is due to their small size as compared to the corresponding alkali metals.

The second ionisation enthalpies of the alkaline earth metals are smaller

than those of the corresponding alkali metals.


Physical Properties

The melting and boiling points of these metals are higher than the

corresponding alkali metals due to smaller sizes.

The trend is, however, not systematic.

Because of the low ionisation enthalpies, they are strongly electropositive

in nature.

The electropositive character increases down the group from Be to Ba.

The alkaline earth metals like those of alkali metals have high electrical and

thermal conductivities which are typical characteristics of metals.


The atomic and physical properties of the alkaline earth metals


Property Group IA

(Alkali Metals )

Group IIA

(Alkaline Earth Metals)

Atomic radii larger smaller

Melting and boiling

point

lower

higher

Density lower higher

Ionization energy

lower higher

Hydration energy

lower higher

Lattice energy lower

higher


Flame test

In flame the electrons are excited to higher energy levels and when they

drop back to the ground state, energy is emitted in the form of visible light.

Calcium, strontium and barium impart characteristic brick red, crimson

and apple green colours respectively to the flame.

The electrons in beryllium and magnesium are too strongly bound to get

excited by flame. Hence, these elements do not impart any colour to the

flame.

The flame test for Ca, Sr and Ba is helpful in their detection in

qualitative analysis and estimation by flame photometry.


Hydration Enthalpies

Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions

decrease with increase in ionic size down the group.

Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+

The hydration enthalpies of alkaline earth metal

ions are larger than those of alkali metal ions. Thus,

compounds of alkaline earth metals are more

extensively hydrated than those of alkali metals.

Example; MgCl2 and CaCl2 exist as MgCl2.6H2O

and CaCl2·6H2O while NaCl and KCl do not form

such hydrates.


Alkaline Earth Metal and Water

The majority of alkali earth metals produce hydroxides when reacted with water.

The hydroxides of Ca, Sr, and Ba are only slightly soluble in water; however,

enough hydroxide ions are produced to make a basic environment.

The general reaction of Ca, Sr, and Ba with water is represented below, where

M represents the metal: M(s)+2H2O(l)⟶M(OH)2(aq) +H2 (g)

Mg reacts with water vapor to form magnesium hydroxide and hydrogen gas.

Be is the only alkaline earth metal that does not react with water. This is due to

its small size and high ionization energy in relation to the other elements in the

group.


Alkaline Earth Metal Hydrides and Water

The s-block metals form ionic hydrides when heated with hydrogen gas, for

example magnesium hydride: Mg(s) + H2(g) → MgH2(s)

which is a white crystalline solid which dissociates into Mg and H at 327oC.

However, BeH and MgH2 are intermediate/covalent.

With the exception of beryllium (Be), the alkaline metal hydrides react with

water to produce the metal hydroxide and hydrogen gas.

The reaction of these metal hydrides can be described below:

MH2 (s) +2H2O (l) ⟶M(OH) 2 (aq) +2H2 (g)


Alkaline Earth Metal and carbon

Group 1 metals do not react with carbon, but Group 2 metals react at high

temperatures to give the metal carbides:

Ca(s) + 2C(s) → CaC2(s)

which are ionic, containing the carbide ion: (C≡C)2–.

The carbides react with water to give acetylene gas:

CaC2(s) + 2H2O(l) → Ca(OH)2(s/aq) + C2H2(g)

Ignited Mg metal will also react with carbon dioxide in the absence of air:

Mg(s) + CO2(g) → 2MgO(s) + C(s)


Alkaline Earth Metal and Nitrogen

The Group 2 metals, all except Be react upon heating:

3Mg(s) + N2(g) → Mg3N2 (s)

The nitrides react with water to produce ammonia:

Mg3N2(s) + 6H2O(l) → 3Mg(OH)2(aq) + 2NH3(g)

Magnesium nitride is a yellow or yellow-green crystalline solid which

decomposes to Mg and N2 when heated.


Oxides of the alkaline earth metals

The reaction of Group 2 metals with oxygen is a redox reaction

2M(s) + O2(g) ==> 2MO(s) (M = Be, Mg, Ca, Sr, Ba)

The formation of the oxide expected when the element is heated or burned

in air.



Mg-Ba form 1:1 metal oxides, MO, which crystallize with NaCl-type

structures.

MgO is pretty insoluble in water, but the others react with water to form the

hydroxides, which tend to be insoluble.

MgO has a melting point of 2825°C; in its crystalline form is a great thermal

conductor and a lousy electrical conductor

When cut all the metals rapidly tarnish in air at RT, but combust when heated.

Be and Mg tarnish in air and this oxide layer prevents further reaction.

CaO is called “quicklime”; produced in huge quantities from calcium

carbonate (limestone)



Similarly to the alkali metal oxides, alkaline earth metal monoxides combine

with water to form metal hydroxide salts.

MO(s) + H2O(l) ==> M(OH)2(aq) (not a redox change, M = Be, Mg, Ca, Sr, Ba)

ionically: M2+O2–(s) + H2O(l) ==> M(OH)2(aq)

The exception to this general assumption is beryllium, whose oxide (BeO)

does not react with water.

Ionic oxides (I and II A element with O2) always react with water to give

basic solutions which increases in strength of basic character down the group.



The pH of the resulting solution ranges from ~pH 10 to ~pH 13 for

Mg(OH)2 to Ba(OH)2

One of the most familiar alkaline earth metal oxides is CaO or quicklime.

This substance is often used to treat water and to remove harmful SO2(g) from

industrial smokestacks.



All the oxides are basic and readily neutralised by acids (not a redox change).

MO(s) + 2HCl(aq) ==> MCl2(aq) + H2O(l) (M = Be, Mg, Ca, Sr, Ba)

to give the soluble chloride salt

ionically: M2+O2–(s) + 2H+

(aq) ==> M2+(aq) + H2O(l)

acid proton donation to the oxide ion base.

In each case the chloride Cl–, nitrate NO3– and sulphate SO4

2– are spectator

ions.

MO(s) + 2HNO3(aq) ==> M(NO3)2(aq) + H2O(l) (M = Be, Mg, Ca, Sr, Ba)

to give the soluble nitrate salt



MO(s) + H2SO4(aq) ==> MSO4(aq/s) + H2O(l) (M = Be, Mg, Ca, Sr, Ba)

to form the sulphate salt (soluble => insoluble)

but reaction increasingly slower for calcium oxide ==> barium oxide as

the sulphate becomes less insoluble.


Reaction with other Group 6 elements

In addition to forming oxides, the metals form sulphides on heating

• with sulphur (sulfur) e.g. MgS

• with selenides, e.g. MgSe

Magnesium sulphide, MgS, has the 6:6-coordinate structure of NaCl, but

much higher melting and boiling points, due to the greater ionic charges

producing stronger ionic bonding.

Mg(s) + S(s) → MgS(s)


Alkaline Earth Metal and halogen

All s-block metals react directly with chlorine, requiring moisture and/or

heat to initiate the reaction:

Ca(s) + Cl2(g) → CaCl2(s)

BeCl2 is more-or-less covalent, with a low melting point.

MgCl2 is intermediate between ionic and covalent.

Beryllium and Magnesium

Be is very small and has a high charge which makes for high charge density.

*A high charge density simply means that you have a lot of charge packed into a small volume.

• Be forms covalent bonds because of its high polarizing power.

• In the gas phase, Be is coordinatively unsaturated (electron-deficient).

• In solid phases, Be will act as a Lewis acid.

• Mg is found on the same diagonal as Li; therefore, Mg has similar properties to

Li

• Both Li and Mg are used as reagents in organic synthesis (Grignard’s reagent).

•All s-block metals except Be and Mg are stored under liquid paraffin to prevent

their reaction with oxygen. (Protective oxide layer)


The anomalous behaviour of beryllium

Be differs more from Mg than Li does from Na.

Be has a diagonal relationship to Al and both elements have some behaviours

in common.

1. Be has a higher melting/boiling point, higher density and much greater

hardness than Mg.

2. Like Li, Be has a small atomic radius and higher electronegativity and

electron affinity and a higher ionisation energy.

3. The small Be2+ ion has a higher polarising power and so all or most of its

compounds are largely covalent (and so often have lower melting/boiling points

due to weak intermolecular forces).



4. Be does not react with water and is resistant to acid.

due to a protective oxide film

Its compounds are more soluble in organic solvents.

5. Beryllium halides, BeX2, are hygroscopic solids (absorb moisture from the air)

that fume in air.

6. Beryllium is amphoteric (both basic and acidic).

7. Be is a poor reducing agent, due to its reluctance to lose its valence electrons.

8. Salts of large anions are very unstable and those that are stable are hydrated,

e.g. BeCO3·4H2O, BeSO4·4H2O,

both decompose on heating to give BeO.



9. No peroxide or superoxide is formed.

10. Be is poisonous since it has strong complexing power with O and N-ligands.

11. When treated with acetylene, Mg forms the acetylide, MgC2, whereas Be

forms the carbide, Be2C.



12. Be compounds tend to have coordination numbers of 4.

• The ion is too small to have higher coordination numbers.

• Example;

• [Be(H2O)4]2+ occurs in hydrated beryllium salts.

• BeO has 4:4 coordination.

• In the vapour phase, the chloride contains (BeCl2)2 molecules, with

chlorines bridging between the two Be atoms.



13. BeCl2 is a covalent chain polymer, existing as dimers, Be2Cl4, in the

vapour phase, which break-up into monomers, BeCl2, at higher temperatures

(AlCl3 is similar).

*Note the halide bridge structure with the halogens arranged tetrahedrally around

each Be.




AlCl3 has a similar structure, but existing as shorter Al2Cl6 molecules held

together by weak intermolecular van der Waal’s forces.

In the gas phase it forms Al2Cl6 dimers which break-up into AlCl3

monomers at higher temperatures.

AlBr3 and AlI3 are similar.


Beryllium Oxide

BeO has a very high melting point, making it useful for nuclear work and in

ceramics.

BeO is amphoteric,

for example, it acts as a base in the following (slow) reaction at very low pH:

BeO(s) + H2O(l) + 2H3O +

(aq) → Be(H2O)42+

in which the Be complexes with water, neutralising acid as it does so.

BeO is behaving as an acid in the following reaction:

BeO(s) + H2O(l) + 2OH–(aq) → Be(OH)4

2–(aq)

in which it neutralises hydroxide ions (alkali) forming the beryllate ion.


Beryllium Oxide

The hydroxides are not only basic, but they are alkali, an alkali being a

base which directly provides hydroxide, OH–(aq), ions on dissolving in water.

BeO has the wurtzite structure (hexagonal crystal structure).

A coordination number of 4

(4:4 so each Be2+ ion is surrounded by 4 of O2– ions and each O2– ion by 4

Be2+ ).

The other Group 2 oxides have the NaCl structure (see halides) with a

coordination number of 6

(6:6, so each metal ion is surrounded by 6 oxide ions and each oxide ion by 6

metal ions).

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