Revision notes Periodic Table: chemical reactions of compounds of s-block Groups 1 & 2 Advanced Inorganic Chemistry

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Doc Brown's Chemistry  Advanced Level Inorganic Chemistry – Periodic Table Revision Notes on group I and group II s-block metals

Part 7. s–block Groups 1 Alkali Metals and Group 2 Alkaline Earth Metals – Sections 7.9 to 7.12

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 Advanced level multiple choice quiz on the basics of  s–block chemistry

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Sub-index for this page on Group 1 and 2 s-block metals

7.9 Describes the chemistry of the Group 1 Alkali Metal and Group 2 Alkaline Metal carbonates & hydrogencarbonates e.g. their formation and reaction with acids.

7.9 carbonates & hydrogen carbonates

7.10 Gives the solubility trends of hydroxides, nitrates, sulfates and carbonates.

 7.10 Solubility trends of hydroxides, nitrates, sulfates, carbonates

7.11 Describes the thermal decomposition of the carbonates and offers two explanations of stability trend of the Groups 1/2 carbonates & nitrates

7.11 Thermal decomposition, stability trend of carbonates & nitrates

7.12 Describes the uses of group 1–2 metallic elements and their compounds.

7.12 Uses of group 1–2 metals and compounds

Pd s block elements d blocks and f blocks of metallic elements  p block elements
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Gp7/17 Gp0/18


2 3Li




ZSymbol, z = atomic or proton number

highlighting position of Group 1 and Group 2 elements

outer electrons: 1 ns1 and ns2

5B 6C 7N 8O 9F 10Ne
3 11Na




13Al 14Si 15P 16S 17Cl 18Ar
4 19K




21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb




39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs




57-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At 86Rn
7 87Fr




89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Nh 114Fl 115Mc 116Lv 117Ts 118Og

outer electrons: Group  1 ns1 and Group 2 ns2

7.9. The properties and chemistry of the carbonates (CO32–) & hydrogencarbonates (HCO3)

The carbonates and hydrogencarbonates are white ionic solids

  • Group 1 carbonates M2CO3: Formed on bubbling carbon dioxide into excess hydroxide solution

    • 2MOH(aq) + CO2(g) ==> M2CO3(aq) + H2O(l)    (M = Li, Na, K, Rb, Cs)

    • ionic equation: 2OH(aq) + CO2(g) ==> CO32–(aq) + H2O(l) 

    • The carbonates are white solids, quite soluble in water, and, apart from lithium, thermally stable to red–heat. (see section 7.11)

    • Hydrated sodium carbonate, Na2CO3.10H2O, is known as washing soda and is used to soften water by precipitating magnesium and calcium salts as their less soluble carbonates (see section 7.10).

  • Group 1 hydrogencarbonates MHCO3: Formed on bubbling excess carbon dioxide into the hydroxide solution

    • The reaction above happens first and then:

      • M2CO3(aq) + H2O(l) + CO2(g) ==>  2MHCO3(aq)    (M = Li, Na, K, Rb, Cs) 

      • ionic equation: CO32–(aq) + H2O(l) + CO2(g) ==>  2HCO3(aq) 

      • They are white solids, slightly soluble in water, weakly alkaline and readily decompose on heating to form the carbonate and carbon dioxide gas.

      • e.g. at 270oC: 2NaHCO3(s) ==> Na2CO3(s) + H2O(l) + CO2(g) 

      • An alternative to yeast in baking is to use sodium hydrogencarbonate ('sodium bicarbonate' or 'baking soda'). The rising action is also due to carbon dioxide gas formation from reacting with an acid (e.g. an organic acid like tartaric acid) and nothing to do with enzymes:

        • self–raising baking powder = carbonate base + a solid organic acid, giving

        • sodium hydrogencarbonate + acid ==> sodium salt of acid + water + carbon dioxide

  • Group 1 carbonates and hydrogencarbonates are readily neutralised by acids:

    • these are base (proton accepting CO32– or HCO3)–acid (H+ from HCl etc.) reactions giving a salt, water and carbon dioxide

    • In all cases M = Li, Na, K, Rb, Cs e.g.

    • M2CO3(aq) + 2HCl(aq) ==> 2MCl(aq) + H2O(l) + CO2(g) to give the chloride salt*

      • ionically for any soluble carbonates: CO32–(aq) + 2H+(aq) ==> H2O(l) + CO2(g)  

    • M2CO3(aq) + 2HNO3(aq) ==> 2MNO3(aq) + H2O(l) + CO2(g) to give the soluble nitrate salt

    • M2CO3(aq) + H2SO4(aq) ==> M2SO4(aq) + H2O(l) + CO2(g) to give the soluble sulphate salt

    • M2CO3(aq) + 2CH3COOH(aq) ==> CH3COOM(aq) + H2O(l) + CO2(g) to give the soluble ethanoate salt

    • MHCO3(aq) + HCl(aq) ==> MCl(aq) + H2O(l) + CO2(g) to give the soluble chloride salt

      • ionically for any soluble hydrogencarbonates: HCO3(aq) + H+(aq) ==> H2O(l) + CO2(g)

    • MHCO3(aq) + HNO3(aq) ==> MNO3(aq) + H2O(l) + CO2(g) to give the nitrate salt

    • 2MHCO3(aq) + H2SO4(aq) ==> M2SO4(aq) + 2H2O(l) + CO2(g) to give the sulphate salt

    • MHCO3(aq) + CH3COOH(aq) ==> CH3COOM(aq) + H2O(l) + CO2(g) to give the ethanoate salt

    • * The group 1 carbonates e.g. sodium carbonate can be titrated with standardised hydrochloric acid using methyl orange indicator (red in acid, yellow in alkali, the end point is a sort of 'pinky orange').

  • Group 2 carbonates MCO3: formed on bubbling carbon dioxide into the hydroxide solution or 'slurry', but beryllium carbonate is not stable (another anomaly). Non of them are very soluble.

    • M(OH)2(aq) + CO2(g) ==> MCO3(s) + H2O(l)     (M = Mg, Ca, Sr, Ba)

    • When M = Ca, this the reaction of limewater when positively testing for carbon dioxide gas.

    • They can also be prepared by a double decomposition precipitation reaction (see section 7.10).

    • The carbonates decompose on heating to give the oxide and carbon dioxide and exhibit a clear thermal stability trend (see section 7.11).

  • Group 2 hydrogencarbonates M(HCO3)2: formed when excess carbon dioxide is bubbled through a slurry of the carbonate and they are only stable in solution and their reaction with acids is not important:

    • MCO3(s) + H2O(l) + CO2(g)  (c) doc b  M(HCO3)2(aq)     (M = Mg, Ca, Sr, Ba)

      • This the reaction of 'temporary hard water' formation from calcium and magnesium carbonate minerals like limestone and dolomite. Boiling the solution reverses the reaction, so removing the metal cations from solution, and referred to as removing 'temporary hardness'.

  • Group 2 carbonates MCO3 readily neutralised by acids to form salt, water and carbon dioxide:

    • These are Bronsted–Lowry base (proton accepting CO32–)–acid (H+ from HCl etc.) reactions giving a salt, water and carbon dioxide e.g. for M = Mg, Ca, Sr, Ba

    • MCO3(s) + 2HCl(aq) ==> MCl2(aq) + H2O(l) + CO2(g) to give the chloride salt    (M = Mg, Ca, Sr, Ba)

      • ionically for all four examples: M2+CO32–(s) + 2H+(aq) ==> M2+(aq) + H2O(l) + CO2(g)  

    • MCO3(s) + 2HNO3(aq) ==> M(NO3)2(aq) + H2O(l) + CO2(g) to give the nitrate salt

    • MCO3(s) + H2SO4(aq) ==> MSO4(aq) + H2O(l) + CO2(g) to give the sulphate salt

    • MCO3(s) + 2CH3COOH(aq) ==> (CH3COO)2M(aq) + H2O(l) + CO2(g) to give the ethanoate salt

  • The thermal decomposition of carbonates and nitrates is covered in detail in section 7.11

TOP OF PAGE and sub-index

7.10. Solubility Trends of Group 2 compounds – linked to preparations

  • All the nitrates, M(NO3)2, are soluble in water.    (M = Be, Mg, Ca, Sr, Ba)

  • The hydroxides M(OH)2, get more soluble down the group:    (M = Be, Mg, Ca, Sr, Ba)

    • If more or less insoluble, they can be made by adding sodium hydroxide solution to a solution of a soluble salt of M e.g.

      • magnesium chloride + sodium hydroxide ==> sodium chloride + magnesium hydroxide

      • MgCl2(aq) + 2NaOH(aq) ==> 2NaCl(aq) + Mg(OH)2(s)

        • or ionically: Mg2+(aq) + 2OH(aq) ==> Mg(OH)2(s)

      • Magnesium hydroxide is almost insoluble in water i.e. sparingly soluble.

      • Calcium hydroxide is slightly soluble in water, so–called 'limewater' used in the simple test for carbon dioxide gas.

      • Barium hydroxide is moderately soluble in water.

  • The sulphates, MSO4, get less soluble down the group.     (M = Be?, Mg, Ca, Sr, Ba)

    • Magnesium sulphate is very soluble in water, in fact it was first crystallised from spring water e.g. the chalk downs of Southern England, hence known as Epsom Salts.

      • the heptahydrate salt  MgSO4.7H2O

    • Calcium sulphate is slightly soluble in water.

    • If more or less insoluble, they can be made by adding dilute sulphuric acid or sodium sulphate solution to a solution of a soluble salt of M.

    • This reaction is used as a test for a sulphate by adding an acidified  barium chloride/dil. hydrochloric acid or barium nitrate/dil. nitric acid solution to a solution of the suspected sulphate. A dense white precipitate of barium sulphate forms in a positive result and also illustrates the preparation too e.g.

    • barium chloride + sodium sulphate ==> sodium chloride + barium sulphate

    • BaCl2(aq) + Na2SO4(aq) ==> 2NaCl(aq) + BaSO4(s) 

    • or

    • barium nitrate + magnesium sulphate ==> magnesium nitrate + barium sulphate

    • Ba(NO3)2(aq) + MgSO4(aq) ==> Mg(NO3)2(aq) + BaSO4(s) 

      • or ionically in each case Ba2+(aq) + SO42–(aq) ==> BaSO4(s)

      • Why is the acidification necessary? The addition of dilute hydrochloric acid is to prevent the precipitation of other insoluble salts like barium sulphite which would be confusing and make the test less specific.

  • The carbonates, MCO3, get less soluble down the group.     (M = Mg, Ca, Sr, Ba)

    • If insoluble, they can be made by adding sodium carbonate solution to a solution of a soluble salt of M e.g. the 'double decomposition' ...

    • magnesium chloride + sodium carbonate ==> sodium chloride + magnesium carbonate

    • MgCl2(aq) + Na2CO3(aq) ==> 2NaCl(aq) + MgCO3(s) 

      • or ionically: Mg2+(aq) + CO32–(aq) ==> MgCO3(s) 

      • You can also use the nitrate and in the case of magnesium, its sulphate too.

      • The spectator ions are Na+ and the chloride or sulphate etc. anion from the original group II salt.

      • Hydrated sodium carbonate, Na2CO3.10H2O, is known as washing soda and is used to soften water by using the above reaction.

      • e.g. calcium sulphate (gypsum) + sodium carbonate ==> sodium sulphate (soluble) + calcium carbonate (insoluble)

      • CaSO4(aq) + Na2CO3(aq) ==> Na2SO4(aq) + CaCO3(s)

        • (ionic equation similar to above, sodium ions and sulfate ions are 'spectators')

        • so, ionically: Ca2+(aq) + CO32–(aq) ==> CaCO3(s) 

  • Explanation of solubility trends (usually dealt with later in course e.g. in UK A2 advanced level)

    • The simplest approach is to consider the two enthalpy change trends.

      • The process of dissolving can be analysed in terms of two theoretical stages e.g. for simple cation–anion ionic compound.

        • In the arguments outlined below Mn+ could be Gp1 or Gp2 metal cation etc., Xn– could be halide, oxide, hydroxide, sulphate, carbonate anion etc., and n is the charge on ion:

      • (1) Mn+aXnb(s) ==> aMn+(g) + bXn–(g) (breaking the lattice apart into its constituent ions)

      • This process is always endothermic, and is called the lattice enthalpy. Its usually defined in the opposite direction by saying it is 'the energy released when 1 mole of an ionic lattice is formed from its constituent gaseous ions' (at 298K, 1 atmos./101kPa pressure).

      • * The lattice enthalpy decreases down the group as the cation radius increases (anion radius constant for a particular series e.g. sulphates). Therefore, energetically, the solvation in terms of lattice energy is increasingly favoured down the group. 

      • (2) Mn+(g) + aq ==> Mn+(aq)  and  Xn–(g) + aq ==> Xn–(aq) 

        • Representing the solvation–hydration of ions.

      • The equations above represent to the two 'hydration enthalpies', the heat released when an isolated gaseous ion becomes solvated by water to form an aqueous solution (1M, 298K, 1 atmos./101kPa pressure)

      • * The hydration enthalpy for the cation decreases down the group as the radius gets larger. Therefore, energetically, the solvation is less favoured down the group as the cation radius increases.

    • * In both cases the numerical enthalpy value increases the smaller the radii as charges closer, and the greater the ionic charge (constant for a series), both factors increase the electrical attraction of either cation–anion in the crystal or ion–water in aqueous solution.

  • We therefore have two competing trends!

    • So, one approach is to say which 'energy change' trend outweighs the other to explain the solubility trend ...

    • e.g. for Group 2 hydroxides, energetically, the decrease in lattice enthalpy more than compensates for the decrease in the hydration enthalpy of the M2+ cation as it gets larger down the group so leading to greater solubility.

  • Unfortunately the above is hardly an explanation of a correct prediction! and neither is entropy taken into consideration.

    • The explanations offered are argued after the fact and unsatisfactory!

    • There is no simple explanation possible and ultimately the solubility is dependent on the entropy changes, a notoriously difficult concept area.

    • If there was an appropriate AS–A2 answer, it would be in the textbooks by now!

    • See below on Jim Clarks website for an intelligent discussion on the matter. Jim's Group 2 pages solubility descriptions and trends and discussions and theory of solubility

TOP OF PAGE and sub-index

7.11. Thermal decomposition & stability trends of Group 1 and Group 2 compounds

  • One theory of the thermal instability trend

    • The lower down the metal in the group the more thermally stable is its hydroxide, nitrate, carbonate or sulphate etc.

    • This is because the polarising power of the cation increases up the group with the smaller ionic radius,

      • AND, in most cases discussed here, the smaller the cation the greater the lattice enthalpy of the oxide formed on decomposition (meaning the oxide is more thermodynamically stable up the group).

      • Particularly for the tiny Li+ and Be2+ ions, the polarising effect considerably reduces the stability of their compounds (e.g. BeCO3 is quite unstable and Li2CO3 decomposes on gentle heating).

    • The 'polarising power' of a cation is a measure of its electric field effect to attract and distort electron charge on a neighbouring anion: 

      • The cation polarising power increases with increase in charge on the ion or decreasing the radius of the ion, both of which increase the intensity of the electric field - a high charge density.

      • This high charge density effect polarises the O-H in hydroxides, the N-O bond in nitrates and C-O bond in the carbonate ion.

      • Within a group of cations e.g. groups 1 or 2, the effect of the charge density  increases up the group as the atomic/ionic radius gets smaller (for the same positive charge). So the hydroxide, carbonate or nitrate tend to get less thermally stable up the group.

    • The 'polarisability of an anion' is how easily the electron charge clouds are 'distorted' by a neighbouring cation.

      • The anion is more easily distorted the larger the anion radius and the higher its charge.

  • The general 'polarising effect' is shown in the diagram below.

    • Think of the Mn+ cation as the Gp1 or Gp2 cation and the XO3n– anion as the nitrate ion or the carbonate ion (and the think the same way for a hydroxide ion or a sulphate – in fact any 'oxyanion').

    • (c) doc b

    • The electrical field of the cation distorts or polarises the anion, and at the decomposition temperature, a 'residual' oxide ion is attracted to the cation and the rest of the original larger anion is released as a gas or gases.

    • Notes:

      • (i) The residual oxide ion is smaller and less polarisable.

      • (ii) These reactions eventually become favourable at higher temperature because of the large increase in the 'systems' entropy when gases formed.

      • (iii) the smaller oxide ion means the resulting oxide has a higher lattice enthalpy than the carbonate or nitrate etc. and increases up the group making the decomposition more favourable.

      • Point (iii) forms the basis of a purely thermodynamic argument to explain the thermal stability trend and even predict decomposition temperatures (see the end of this section).

    • The trend down Groups I & II of increasing thermal stability of the MCO3 carbonate is also paralled by an increase in ionic character of the carbonate (also nitrates and hydroxides).

      • As you go up Groups I & II the radius of the Group I & II ions, (M+ & M2+), decreases, therefore the ion's electric field effect is much stronger, i.e. increase in polarising power of the metal ion.

      • Therefore the polarising effect of the increasingly smaller Mn+ ion on the outer electron clouds of the carbonate ion (or any other ion) increases, distorting the electron clouds of the carbonate ion towards the metal ion.

      • This increases the potential for orbital overlap i.e. increase in covalent character of the metal ion - carbonate ion bond up the group.

      • Therefore down the group, there is a decrease in covalent character as the polarising metal ion gets larger.

      • Hence the increase in ionic character of the carbonates (and nitrates/hydroxides) down the Groups 1 & 2, the same trend for increasing thermal stability of the carbonate and the metal ion radius argument applies to both trends.

        • BUT, we haven't finished yet! - a further 'intellectual' complication!

        • The increase in thermal stability down the group goes against the trend of lattice enthalpy!

        • As the Group I/II cation gets larger, the lattice enthalpy decreases, and so the metal ion - carbonate ion bond should be weaker and so the ionic compound thermally less stable.

          • Lattice enthalpy increases with the greater the charges or the smaller the radii of the ions, again this is an electrical field effect, both factors increasing the attraction between the cations and anions in the crystal lattice.

        • BUT, the lattice enthalpy trend doesn't correctly predict the thermal stability trend, so you can argue that the decrease in polarising power of the Gp I/II metal ion down the group must have a greater effect than the decrease in lattice enthalpy.

        • The above are reasonable conceptual arguments, BUT, these arguments ignore the lattice energy of the oxide formed on decomposition of the carbonate, nitrate, hydroxide.

        • Which is why thermodynamics usually (always?) has the final say!

        • Thermodynamic argument for carbonates - there is quite a bit to get your head round in this section!

  • Trends in thermal stability:

    • In all cases, for a particular set of e.g. Gp1 or Gp2 compounds, the thermal stability increases down the group as the ionic radius of the cation increases, and its polarising power decreases.

    • Group 1 compounds tend to be more thermally stable than group 2 compounds because the cation has a smaller charge and a larger ionic radius, and so a lower polarising power, particularly when adjacent metals on the same period are compared.

  • Group 1 Carbonates:

    • lithium carbonate readily decomposes: Li2CO3(s) ==> Li2O(s) + CO2(g) 

    • but the others are quite stable to red heat, so again lithium is anomalous by its comparative carbonate instability.

  • Group 2 Carbonates:

    • The carbonates thermally decompose into the metal oxide and carbon dioxide gas.

    • MCO3 ==> MO(s) + CO2(g)    (M = Be, Mg, Ca, Sr, Ba)

    • Thermal Tdecomp order BaCO3 > SrCO3 > CaCO3 > MgCO3 > BeCO3

    • This is the reaction that converts calcium carbonate (limestone) into calcium oxide (quicklime) in a limekiln at about 900oC. beryllium BeCO3 is unstable at room temperature, magnesium carbonate MgCO3 decomposes at about 400oC, strontium carbonate SrCO3 at 1280oC and barium carbonate BaCO3 at 1360oC.

  • Group 1 Nitrates:

    • lithium nitrate is the least stable and decomposes readily on heating to form lithium oxide, nitrogen dioxide and oxygen.

    • 4LiNO3(s) ==> 2Li2O(s) + 4NO2(g) + O2(g) 

    • The other group 1 Alkali Metal nitrates [NO3, nitrate(V)] decompose to form the nitrite [NO2, nitrate(III)] salt and oxygen gas. Lithium is anomalous due to the particularly high polarising power of the Li+ ion.

    • 2MNO3(s) ==> 2MNO2(s) + O2(g)   (M = Na, K, Rb, Cs)

    • The nitrites, or nitrate(III) ions, are very thermally stable white solids, soluble in water giving neutral solutions in water.

  • Group 2 Nitrates:

    • For M = Mg, Ca, Sr, Ba the nitrate decompose to form the metal oxide, nasty brown nitrogen dioxide [nitrogen(IV) oxide] gas and oxygen gas when strongly heated.

    • 2M(NO3)2(s) ==> 2MO(s) + 4NO2(g) + O2(g)    (M = Be?, Mg, Ca, Sr, Ba)

    • Thermal Tdecomp order Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2 > BeNO3)2

  • Group 1 and Group 2 hydroxides

    • The general thermal decomposition equations are ...

      • Group 1:  2MOH(s) ==> M2O(s) + H2O(g)   (M = Li, Na, K, Rb, Cs)

      • Group 2: M(OH)2(s) ==> MO(s) + H2O(g)    (M = Be?, Mg, Ca, Sr, Ba)

      • The thermal stability trend is just the same as for carbonates, nitrates (and even sulphates) i.e. they become more stable down the group with increasing atomic number of the metal M.

        • So for group 1 the  Tdecomp sequence is CsOH > RbOH > KOH > NaOH > LiOH

        • and for group 2 the  Tdecomp sequence is Ba(OH)2 > Sr(OH)2 > Ca(OH)2 > Mg(OH)2 > Be(OH)2

  • Comparing the stabilities of Group 1 and Group 2 compounds:

    • Group 1 compounds are more stable than group 2 compounds, particularly when group adjacent elements on the same period.

    • The reason being the polarising effect of the group I cation M+, is much less than the polarising power of the smaller and more highly charged group II M2+ ion, particularly when comparing adjacent metals on the same period.

      • The group 2 cation is both smaller and more highly charged than the corresponding group 1 cation.

A thermodynamic discussion on the thermal stability of s–block compounds

I'm referring to oxyanion compounds like carbonates, sulfates and hydroxides

  • What ensues is a completely alternative explanation of the thermal stability trends of the oxyanion compounds of alkali metals and alkaline earth metals without any reference to the relative polarising effect of the cation.

  • It may seem curious at first sight, but large cations can stabilise large anions in a crystal lattice (and vice versa).

    • The decomposition temperatures of thermally unstable compounds containing large anions e.g. carbonates, increases with cation radius.

    • The stabilizing influence of a cation can be explained in terms of trends in lattice enthalpies (lattice energies).

    • The arguments given below are purely in terms of the thermodynamics and explain the Group 2 carbonate stability trend without reference to the polarising power of the cation (which is the argument required by most UK GCE A level syllabuses).

  • An initial discussion of the Group II carbonate stability trend illustrates the points made above.

    • A study of the thermal decomposition temperatures of the Group 2 Alkaline Earth carbonates proves most instructive.

    • The general decomposition equation for group II carbonates to give the group II oxide is

      • MCO3(s) ==> MO(s) + CO2(g)

      • M = Be, Mg, Ca, Sr and Ba

      • Beryllium carbonate is not very stable (can be stabilised in an atmosphere of CO2) and radium carbonate would be rather too radioactive to study!

      • MCO3 and MO are sometimes referred to as the alkaline–earth carbonates and alkaline–earth oxides.

  • Most of the data tabulated below was obtained from 'Inorganic Chemistry' 2nd edition, by Shriver, Atkins and Langford and the Nuffield Science Book of Data (revised edition 1988) plus internet research for research papers quoting as up to date values as I could find.

Thermodynamic and decomposition temperature data for Group II carbonates of the periodic table.

the basis for the purely thermodynamic argument for the Group II carbonate thermal stability trend

Decomposition data Be Mg Ca Sr Ba
ΔGØ (kJ mol–1) ? +48.3 +130.4 +183.8 +218.1
ΔHØ (kJ mol–1) ? +100.6 +178.3 +234.6 +269.3
ΔSØ (J K–1 mol–1) ? +175.0 +160.6 +171.0 +172.1
Tdecomp (oC, K) theoretical ? 302, 575 837, 1110 1099, 1372 1292, 1565
Typical quoted decomposition temperatures ~100oC 400 oC 900 oC 1280 oC 1360oC
LEØMCO3 (kJ mol–1) ? 3180 2987 2720 2615
LEØMO (kJ mol–1) 4443 3960 3489 3248 3011
ΔLE(MO–MCO3) ? 780 502 528 396
M2+ radius in nm Be2+ = 0.034 Mg2+ = 0.078 Ca2+ = 0.100 Sr2+ = 0.127 Ba2+ = 0.143

other radii: oxide ion O2– = 0.140 nm, carbonate ion CO32– = 0.176 nm

  • ΔGØ = standard Gibbs free energy change for the thermal decomposition of the carbonate MCO3 (at 298K, 1 atm)

  • ΔHØ  = standard enthalpy change for the thermal decomposition of the carbonate MCO3 (at 298K, 1 atm)

  • ΔSØ = standard entropy change for the thermal decomposition of MCO3 (at 298K, 1 atm)

    • Note ΔSØ = ΔSØsystem and is NOT ΔSØsurroundings ...

    • ... see theory of the method of calculating the decomposition temperature below.

  • Tdecomp = decomposition temperature in Celsius and Kelvin when the equilibrium pressure pCO2 = 1 atm (101kPa)

  • LEØMCO3 = lattice enthalpy/lattice energy of the group 2 carbonate MCO3 (at 298K, 1 atm. pressure)

  • LEØMO = lattice enthalpy/lattice energy of the group 2 oxide MO (at 298K, 1 atm. pressure)

  • ΔLE(MO–MCO3) is the difference between the lattice enthalpies of the group 2 oxide and the corresponding group 2 carbonate

  • The decomposition temperature was calculated as follows ...

    • (i) From the Gibbs free energy equation

      • ΔGØ = ΔHØ – TΔSØ    (terms defined above)

      • The criteria for equilibrium is when ΔG = 0

      • therefore at equilibrium: ΔHØ – TΔSØ = 0,   and rearranging terms and signs gives

      • TΔSØ = ΔHØ,   therefore   T = ΔHØ / ΔSØ

      • = decomposition temperature (K) to give an equilibrium pressure of 1 atm of carbon dioxide gas

    • (ii) From the total entropy change equation

      • This is if your course doesn't involve free energy, or just an alternative method depending on what data is given or available.

      • ΔSØtotal = ΔSØsystem + ΔSØsurroundings

      • The criteria for equilibrium is that ΔSØtotal = 0

      • therefore at equilibrium: 0 = ΔSØsystem + ΔSØsurroundings,  rearranging so ...

      • at equilibrium: –ΔSØsurroundings = ΔSØsystem

      • since ΔSØsurroundings =  –ΔHØ / T i.e. minus the enthalpy change divided by the absolute temperature

      • then: –(–ΔHØ / T) = ΔSØsystem

      • ΔHØ / T = ΔSØsystem,  rearranging ...

      • gives: Tdecomposition  = ΔHØ / ΔSØsystem

      • i.e. the identical expression derived from the Gibbs free energy expression.

    • (iii) All you have to do now is substitute in the numerical values from the data table ...

      • ... and note that ΔG and ΔH are usually in kJ BUT S or ΔS values are usually in J so don't forget to multiply the ΔG and ΔH values by 1000, therefore ...

      • Tdecomp(MgCO3) = 100600 / 175.0 = 575 K

      • Tdecomp(CaCO3) = 178300 / 160.6 = 1110 K

      • Tdecomp(SrCO3) = 234600 / 171.0 = 1372 K

      • Tdecomp(BaCO3) = 269300 / 172.1 = 1565 K

    • Assumptions and comments

      • The calculations have been based on the enthalpy, free energy and entropy value changes at 298K.

      • Enthalpy values (H) do vary with temperature and entropy values (S) increase with temperatures.

      • These factors have been ignored in the calculation BUT the values seem to be roughly born out by experiment as far as I can gather from the values quoted in the literature.

      • Note that the entropy change is almost constant because the increase in entropy is primarily due to the formation of 1 mole of carbon dioxide gas in each case.

        • Remember Sgas >> Sliquid > Ssolid

        • Giving a large increase in entropy, the formation of a gas is a powerful driving force to facilitate the decomposition of these essentially stable compounds BUT this factor applies almost equally to all the group 2 carbonates, therefore entropy cannot be used to explain the stability trend.

        • To explain the thermal stability trend thermodynamically, we must look at the enthalpy changes and the lattice enthalpies of the carbonate and the oxide residue.

  • The thermodynamic argument to explain the thermal stability trend of Group 2 carbonates.

    • To follow the argument you need to x–reference with the numerical values in the data table above.

    • Irrespective of the validity of the theoretical values calculated for the group 2 carbonate decomposition temperatures, what is clearly predicted is that they become more thermally stable down the group i.e. with increase in atomic number of the metal.

    • This increase in stability trend matches the experimental values which in turn are of the same order as those calculated theoretically despite the decrease in lattice enthalpy of the carbonate down the group!

    • The first point to be made is that the endothermic enthalpy of reaction increases down the group.

    • This is itself a clear indication that the decomposition is becoming much less energetically favourable down the group and remember the entropy change is almost constant down the group.

    • The pivotal point in the argument–explanation rests on the differences between the lattice enthalpies (LEs) of the decomposing carbonate and the oxide product as you descend the group.

    • The difference in their LEs is primarily the reason for the rise in the endothermic enthalpy of reaction leading to the rising decomposition temperature.

    • Down the group the lattice enthalpy of both the carbonate and the oxide decrease because the cation radius increases.

      • Lattice enthalpy is a function of two factors (other than the spatial positions of the ions)

        • (i) the charge on the positive and negative ions attracting each other, both constant in this case and ...

        • (ii) the radius of the ions. Here the carbonate ion and the oxide ion radii are constant, but the cation radius is increasing with atomic number of the group 2 metal.

          • This increase the nuclear (+) ... (–) ion distance, reducing the force of attraction and hence reducing the lattice enthalpy of the oxide down the group.

    • However, generally speaking (3/4 values!), as you go down the group the lattice enthalpy of the oxide decreases more rapidly than the lattice enthalpy of the carbonate.

    • This means the difference between the two enthalpies becomes less and less down the group making the enthalpy more and more positive/endothermic and resulting in an increasingly higher temperature to effect the thermal decomposition (to the extent of producing an equilibrium partial pressure of 1 atmosphere of carbon dioxide gas).

      • Breaking up the M2+CO32– lattice is endothermic, but the formation of the M2+O2– lattice is exothermic and numerically greater than for MCO3 and particularly the lower the atomic number of the metal (i.e.) higher up the group).

    • Note that, although I do not have the comparable data for beryllium, the quoted lattice energy for beryllium oxide (BeO) is very high due to the very small beryllium cation Be2+, and therefore extrapolating up the group, you would expect beryllium carbonate to have a much lower decomposition temperature.

      • This is usually quoted as ~100oC and completely fits in with both the theoretical thermal stability trend and the experimental thermal decomposition values quoted in the literature–textbooks.

  • Further extension of the ideas – looking at other thermal stability trends

    • The anhydrous Group 2 sulphates show a similar thermal stability trend to the carbonates...

      • i.e. for the reaction: MSO4(s) ==> MO(s) + SO3(g)

      • the Tdecomp is in the order BaSO4 > SrSO4 > CaSO4 > MgSO4

    • The effect of the cation radius also shows up when comparing the thermal stability of Group 1 carbonates and Group 2 carbonates.

      • Comparing the two thermal decomposition reactions ...

      • Because of the greater charge on the Group 2 cation (M2+) compared to the Group 1 cation (M+) the lattice enthalpy of the Group 2 oxide is much greater than for the Group 1 oxide.

      • So, for the s–block metals on the same period, for Tdecomp the trend is M2CO3 > MCO3

      • The lattice enthalpies are ...

        • 2478 kJ mol–1 for Na2O (1239 kJ per mol Na)

        • and 3960 kJ mol–1 for MgO (3960 kJ per mol Mg)

      • The very high lattice enthalpy of MgO compared to that the LE for Na2O contributes to a much less endothermic enthalpy of decomposition for the Group 2 carbonate compared to the Group 1 carbonate and hence a lower decomposition temperature for the MCO3.

    • The stability trend for Group 1 alkali metal carbonates is similar to that of the Group 2 carbonates ...

      • i.e. for Tdecomp the trend is K2CO3 > Na2CO3 > Li2CO3 etc. ...

      • ... for exactly the same reasons argued above for MCO3 stability.

    • Comparing the thermal stability of Group 1 nitrates [nitrate(V)] and Group 2 nitrates [nitrate(V)].

      • In group 1, only lithium nitrate readily decomposes to the oxide ...

        • 4LiNO3(s) ==> 2Li2O(s) + 2NO2(g) + O2(g)

      • whereas all the other nitrates initially give the thermally stable nitrite [nitrate(III)] ...

        • 2MNO3(s) ==> 2MNO2(s) + O2(g)   (M = Na, K, Rb, Cs)

      • The relatively much smaller size of the lithium cation (Li+) produces a particularly high lattice enthalpy for lithium oxide compared to the other group 1 oxides, hence the direct formation of the oxide.

      • However, due to the much higher MO lattice enthalpies, the oxide is formed directly in each case for the group 2 nitrates ...

      • 2M(NO3)2(s) ==> 2MO(s) + 4NO2(g) + O2(g)  (M = Mg, Ca, Sr and Ba)

      • and the thermal stability trend will be Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2

      • as in the case of the group 2 carbonates and sulfates etc.

TOP OF PAGE and sub-index

7.12. some examples of the uses of Group 1 and 2 Metals and their Compounds.

  • MCl & MCl2 The Group 1 and Group 2 chlorides are used as sources of metal extraction by electrolysis.

  • Na & Mg Sodium and magnesium are then used to extract titanium from its chloride by displacement.

  • Na Sodium vapour is used in the yellow–orange street lamps.

  • NaCl Sodium chloride 'common salt' is used as a food flavouring and preservative, source of chlorine, hydrogen, sodium metal and sodium hydroxide via electrolytic processes.

  • NaHCO3 is used in baking powders – heat or a weak organic acid (e.g. citric acid) is used in baking powders to form carbon dioxide gas to produce the 'rising' action in baking.

  • Na2CO3 Sodium carbonate is used in the manufacture of glass and the treatment of hard water.

  • NaOH Sodium hydroxide, an important strong alkali, is used in the manufacture of sodium salts, soaps, detergents, bleaches, rayon.

  • KNO3 Potassium nitrate is used in NPK fertilisers. 

  • Mg Magnesium metal is used in the manufacture of alloys, particularly those of aluminium.

  • Mg(OH)2 Magnesium hydroxide is used in antacid indigestion powders to neutralise excess stomach (hydrochloric) acid. When magnesium hydroxide mixed with water it is known commercially as 'milk of magnesia' as an antacid remedy avoiding the use indigestion tablets.

  • CaCO3 Calcium carbonate (limestone) and calcium oxide (quicklime, from thermal decomposition of limestone in kiln) are both used in agriculture to reduce the acidity of soil to improve its fertility.

  • CaCO3 Limestone is used directly as building and road foundation material.

  • CaCO3 Limestone is heated with clay (aluminium silicates) to make cement.

  • BaSO4 Barium sulphate is used in medicine for X–ray colonoscopy of the bowel ('barium meal'), the dense white solid shows up clearly as a white or dark shadow and hence the physical topography of the intestines.

  • and there are lots of other examples if you dig around.


GCSE Level periodic table notes (for the basics) and GCSE Level alkali metal notes (for the basics)

INORGANIC Part 7 s–block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals  sub–index: 7.1 Introduction to s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals  * 7.2 Group 1 data and graphs * 7.3 Group 2 data and graphs * 7.4 General trends down groups I & II and formulae *7.5 Oxygen reaction & oxides of s–block metals * 7.6 Water reaction & hydroxides of group 1/2 metals * 7.7 Acid reaction & salts of group1/2 metals * 7.8 chlorine reaction & halides of group I/II metals * 7.9 carbonates & hydrogen carbonates of s–block metals * 7.10 Solubility trends of groups 1/2 OH, NO3,SO4,CO3 compounds * 7.11 Thermal decomposition and stability of group 1 and group 2 carbonates & nitrates * 7.12 Uses of s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals and their compounds Each page has a matching sub-index

Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr AND important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots All 11 Parts have their own sub-indexes near the top of the pages

Group numbering and the modern periodic table

The original group numbers of the periodic table ran from group 1 alkali metals to group 0 noble gases (= group 8). To account for the d block elements and their 'vertical' similarities, in the modern periodic table, groups 3 to group 0/8 are numbered 13 to 18. So, the p block elements are referred to as groups 13 to group 18 at a higher academic level, though the group 3 to 0/8 notation is still used, but usually at a lower academic level. The 3d block elements (Sc to Zn) are now considered the head (top) elements of groups 3 to 12.

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