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

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

7.9. The properties and chemistry of the carbonates (CO₃^2–) & 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) + CO_2(g) ==> M₂CO_3(aq) + H₂O_(l)    (M = Li, Na, K, Rb, Cs)

  • ionic equation: 2OH^–_(aq) + CO_2(g) ==> CO₃^2–_(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:

    • M₂CO_3(aq) + H₂O_(l) + CO_2(g) ==>  2MHCO_3(aq)    (M = Li, Na, K, Rb, Cs) 

    • ionic equation: CO₃^2–_(aq) + H₂O_(l) + CO_2(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 270^oC: 2NaHCO_3(s) ==> Na₂CO_3(s) + H₂O_(l) + CO_2(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 CO₃^2– 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.

  • M₂CO_3(aq) + 2HCl_(aq) ==> 2MCl_(aq) + H₂O_(l) + CO_2(g) to give the chloride salt*

    • ionically for any soluble carbonates: CO₃^2–_(aq) + 2H⁺_(aq) ==> H₂O_(l) + CO_2(g)  

  • M₂CO_3(aq) + 2HNO_3(aq) ==> 2MNO_3(aq) + H₂O_(l) + CO_2(g) to give the soluble nitrate salt

  • M₂CO_3(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + H₂O_(l) + CO_2(g) to give the soluble sulphate salt

  • M₂CO_3(aq) + 2CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l) + CO_2(g) to give the soluble ethanoate salt

  • MHCO_3(aq) + HCl_(aq) ==> MCl_(aq) + H₂O_(l) + CO_2(g) to give the soluble chloride salt

    • ionically for any soluble hydrogencarbonates: HCO₃^–_(aq) + H⁺_(aq) ==> H₂O_(l) + CO_2(g)

  • MHCO_3(aq) + HNO_3(aq) ==> MNO_3(aq) + H₂O_(l) + CO_2(g) to give the nitrate salt

  • 2MHCO_3(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + 2H₂O_(l) + CO_2(g) to give the sulphate salt

  • MHCO_3(aq) + CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l) + CO_2(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) + CO_2(g) ==> MCO_3(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:

  • MCO_3(s) + H₂O_(l) + CO_2(g)    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 CO₃^2–)–acid (H⁺ from HCl etc.) reactions giving a salt, water and carbon dioxide e.g. for M = Mg, Ca, Sr, Ba

  • MCO_3(s) + 2HCl_(aq) ==> MCl_2(aq) + H₂O_(l) + CO_2(g) to give the chloride salt    (M = Mg, Ca, Sr, Ba)

    • ionically for all four examples: M²⁺CO₃^2–_(s) + 2H⁺_(aq) ==> M²⁺_(aq) + H₂O_(l) + CO_2(g)  

  • MCO_3(s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + H₂O_(l) + CO_2(g) to give the nitrate salt

  • MCO_3(s) + H₂SO_4(aq) ==> MSO_4(aq) + H₂O_(l) + CO_2(g) to give the sulphate salt

  • MCO_3(s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + H₂O_(l) + CO_2(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

    • MgCl_2(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

  • BaCl_2(aq) + Na₂SO_4(aq) ==> 2NaCl_(aq) + BaSO_4(s) 

  • or

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

  • Ba(NO₃)_2(aq) + MgSO_4(aq) ==> Mg(NO₃)_2(aq) + BaSO_4(s) 

    • or ionically in each case Ba²⁺_(aq) + SO₄^2–_(aq) ==> BaSO_4(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

  • MgCl_2(aq) + Na₂CO_3(aq) ==> 2NaCl_(aq) + MgCO_3(s) 

    • or ionically: Mg²⁺_(aq) + CO₃^2–_(aq) ==> MgCO_3(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)

    • CaSO_4(aq) + Na₂CO_3(aq) ==> Na₂SO_4(aq) + CaCO_3(s)

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

      • so, ionically: Ca²⁺_(aq) + CO₃^2–_(aq) ==> CaCO_3(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., X^n– could be halide, oxide, hydroxide, sulphate, carbonate anion etc., and n is the charge on ion:

    • (1) Mⁿ⁺_aXn^–_b(s) ==> aMⁿ⁺_(g) + bX^n–_(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  X^n–_(g) + aq ==> X^n–_(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 XO₃^n– 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').

  • 

  • 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: Li₂CO_3(s) ==> Li₂O_(s) + CO_2(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) + CO_2(g)    (M = Be, Mg, Ca, Sr, Ba)

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

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

• Group 1 Nitrates:

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

  • 4LiNO_3(s) ==> 2Li₂O(s) + 4NO_2(g) + O_2(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.

  • 2MNO_3(s) ==> 2MNO_2(s) + O_2(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) + 4NO_2(g) + O_2(g)    (M = Be?, Mg, Ca, Sr, Ba)

  • Thermal T_decomp 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  T_decomp sequence is CsOH > RbOH > KOH > NaOH > LiOH

      • and for group 2 the  T_decomp 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

other radii: oxide ion O^2– = 0.140 nm, carbonate ion CO₃^2– = 0.176 nm

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

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

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

  • Note ΔS^Ø = ΔS^Ø_system and is NOT ΔS^Ø_surroundings ...

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

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

• LE^Ø_MCO3 = lattice enthalpy/lattice energy of the group 2 carbonate MCO₃ (at 298K, 1 atm. pressure)

• LE^Ø_MO = 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

    • ΔG^Ø = ΔH^Ø – TΔS^Ø    (terms defined above)

    • The criteria for equilibrium is when ΔG = 0

    • therefore at equilibrium: ΔH^Ø – TΔS^Ø = 0,   and rearranging terms and signs gives

    • TΔS^Ø = ΔH^Ø,   therefore   T = ΔH^Ø / ΔS^Ø

    • = 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.

    • ΔS^Ø_total = ΔS^Ø_system + ΔS^Ø_surroundings

    • The criteria for equilibrium is that ΔS^Ø_total = 0

    • therefore at equilibrium: 0 = ΔS^Ø_system + ΔS^Ø_surroundings,  rearranging so ...

    • at equilibrium: –ΔS^Ø_surroundings = ΔS^Ø_system

    • since ΔS^Ø_surroundings =  –ΔH^Ø / T i.e. minus the enthalpy change divided by the absolute temperature

    • then: –(–ΔH^Ø / T) = ΔS^Ø_system

    • ΔH^Ø / T = ΔS^Ø_system,  rearranging ...

    • gives: T_{decomposition}  = ΔH^Ø / ΔS^Ø_system

    • 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 ...

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

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

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

    • T_decomp(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 S_gas >> S_liquid > S_solid

      • 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⁺CO₃^2– lattice is endothermic, but the formation of the M²⁺O^2– 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 ~100^oC 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 T_decomp 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 T_decomp 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 T_decomp 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)] ...

      • 2MNO_3(s) ==> 2MNO_2(s) + O_2(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) + 4NO_2(g) + O_2(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.

Website content © Dr Phil Brown 2000+. All copyrights reserved on revision notes, images, quizzes, worksheets etc. Copying of website material is NOT permitted. Exam revision summaries & references to science course specifications are unofficial.

 Doc Brown's Chemistry 

*