MAIN GROUPS CHEMISTRY Alkaline Earth Metals ( Group IIA)
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 (CaCO 3 ), dolomite (CaCO 3 MgCO 3 ) 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] ns 2. 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. Be 2+ > Mg 2+ > Ca 2+ > Sr 2+ > Ba 2+ 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; MgCl 2 and CaCl 2 exist as MgCl 2.6H 2 O and CaCl 2 6H 2 O 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) +2H 2 O (l) M(OH) 2(aq) +H 2 (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) + H 2(g) MgH 2(s) which is a white crystalline solid which dissociates into Mg and H at 327 o C. However, BeH and MgH 2 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: MH 2 (s) +2H 2 O (l) M(OH) 2 (aq) +2H 2 (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) CaC 2(s) which are ionic, containing the carbide ion: (C C) 2. The carbides react with water to give acetylene gas: CaC 2(s) + 2H 2 O (l) Ca(OH) 2(s/aq) + C 2 H 2(g) Ignited Mg metal will also react with carbon dioxide in the absence of air: Mg (s) + CO 2(g) 2MgO (s) + C (s)
Alkaline Earth Metal and Nitrogen The Group 2 metals, all except Be react upon heating: 3Mg (s) + N 2(g) Mg 3 N 2 (s) The nitrides react with water to produce ammonia: Mg 3 N 2(s) + 6H 2 O (l) 3Mg(OH) 2(aq) + 2NH 3(g) Magnesium nitride is a yellow or yellow-green crystalline solid which decomposes to Mg and N 2 when heated.
Oxides of the alkaline earth metals The reaction of Group 2 metals with oxygen is a redox reaction 2M (s) + O 2(g) ==> 2MO (s) (M = Be, Mg, Ca, Sr, Ba) The formation of the oxide expected when the element is heated or burned in air.
Oxides of the alkaline earth metals 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)
Oxides of the alkaline earth metals Similarly to the alkali metal oxides, alkaline earth metal monoxides combine with water to form metal hydroxide salts. MO (s) + H 2 O (l) ==> M(OH) 2(aq) (not a redox change, M = Be, Mg, Ca, Sr, Ba) ionically: M 2+ O 2 (s) + H 2 O (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 O 2 ) always react with water to give basic solutions which increases in strength of basic character down the group.
Oxides of the alkaline earth metals 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 SO 2(g) from industrial smokestacks.
Oxides of the alkaline earth metals All the oxides are basic and readily neutralised by acids (not a redox change). MO (s) + 2HCl (aq) ==> MCl 2(aq) + H 2 O (l) (M = Be, Mg, Ca, Sr, Ba) to give the soluble chloride salt ionically: M 2+ O 2 (s) + 2H + (aq) ==> M 2+ (aq) + H 2 O (l) acid proton donation to the oxide ion base. In each case the chloride Cl, nitrate NO 3 and sulphate SO 2 4 are spectator ions. MO (s) + 2HNO 3(aq) ==> M(NO 3 ) 2(aq) + H 2 O (l) (M = Be, Mg, Ca, Sr, Ba) to give the soluble nitrate salt
Oxides of the alkaline earth metals MO (s) + H 2 SO 4(aq) ==> MSO 4(aq/s) + H 2 O (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) + Cl 2(g) CaCl 2(s) BeCl 2 is more-or-less covalent, with a low melting point. MgCl 2 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 Be 2+ 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).
The anomalous behaviour of beryllium 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, BeX 2, 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. BeCO 3 4H 2 O, BeSO 4 4H 2 O, both decompose on heating to give BeO.
The anomalous behaviour of beryllium 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, MgC 2, whereas Be forms the carbide, Be 2 C.
The anomalous behaviour of beryllium 12. Be compounds tend to have coordination numbers of 4. The ion is too small to have higher coordination numbers. Example; [Be(H 2 O) 4 ] 2+ occurs in hydrated beryllium salts. BeO has 4:4 coordination. In the vapour phase, the chloride contains (BeCl 2 ) 2 molecules, with chlorines bridging between the two Be atoms.
The anomalous behaviour of beryllium 13. BeCl 2 is a covalent chain polymer, existing as dimers, Be 2 Cl 4, in the vapour phase, which break-up into monomers, BeCl 2, at higher temperatures (AlCl 3 is similar). *Note the halide bridge structure with the halogens arranged tetrahedrally around each Be.
The anomalous behaviour of beryllium AlCl 3 has a similar structure, but existing as shorter Al 2 Cl 6 molecules held together by weak intermolecular van der Waal s forces. In the gas phase it forms Al 2 Cl 6 dimers which break-up into AlCl 3 monomers at higher temperatures. AlBr 3 and AlI 3 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) + H 2 O (l) + 2H 3 O + (aq) Be(H 2 O) 2+ 4 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) + H 2 O (l) + 2OH (aq) Be(OH) 2 4 (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 Be 2+ ion is surrounded by 4 of O 2 ions and each O 2 ion by 4 Be 2+ ). 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).