CRITERI
5. CONTEXT
5.1.3. Relació entre el Decret 119/2015 i l’EpJG
o These are Bronsted–Lowry base (proton accepting CO3
2–)–acid (H+ from HCl etc.) reactions giving a salt,
The thermal decomposition of carbonates and nitrates is covered in detail in section 7.11
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)
o 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)
o 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 o Calcium sulphate is slightly soluble in water.
o 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.
o 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.
o barium chloride + sodium sulphate ==> sodium chloride + barium sulphate o BaCl2(aq) + Na2SO4(aq) ==> 2NaCl(aq) + BaSO4(s)
o or
o barium nitrate + magnesium sulphate ==> magnesium nitrate + barium sulphate o Ba(NO3)2(aq) + MgSO4(aq) ==> Mg(NO3)2(aq) + BaSO4(s)
or ionically in each case Ba2+(aq) + SO4 2–
(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)
o 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' ...
o magnesium chloride + sodium carbonate ==> sodium chloride + magnesium carbonate o 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)
Explanation of solubility trends (usually dealt with later in course e.g. in UK A2 advanced level) o 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 – the n's may be different or the same):
(1) Mn+aXn–b(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.
o * 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!
o So, one approach is to say which 'energy change' trend outweighs the other to explain the solubility trend ...
o 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.
o The explanations offered are argued after the fact and unsatisfactory!
o There is no simple explanation possible and ultimately the solubility is dependent on the entropy changes, a notoriously difficult concept area.
o If there was an appropriate AS–A2 answer, it would be in the textbooks by now!
o 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
7.11. Thermal decomposition & stability trends of Group 1 and Group 2 compounds
One theory of the thermal instability trend
o The lower down the metal in the group the more thermally stable is its hydroxide, nitrate, carbonate or sulphate etc.
o 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) o 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 effect.
o 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.
o 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').
o o 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.
o 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).
Trends in thermal stability:
o 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.
o 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:
o lithium carbonate readily decomposes: Li2CO3(s) ==> Li2O(s) + CO2(g)
o but the others are quite stable to red heat, so again lithium is anomalous by its comparative carbonate instability.
Group 2 Carbonates:
o The carbonates thermally decompose into the metal oxide and carbon dioxide gas.
o MCO3 ==> MO(s) + CO2(g) (M = Be, Mg, Ca, Sr, Ba)
o Thermal Tdecomp order BaCO3 > SrCO3 > CaCO3 > MgCO3 > BeCO3
o 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:
o lithium nitrate is the least stable and decomposes readily on heating to form lithium oxide, nitrogen dioxide and oxygen.
o 4LiNO3(s) ==> 2Li2O(s) + 4NO2(g) + O2(g)
o 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.
o 2MNO3(s) ==> 2MNO2(s) + O2(g) (M = Na, K, Rb, Cs)
o The nitrites, or nitrate(III)'s, are very thermally stable white solids, soluble in water giving neutral solutions in water.
Group 2 Nitrates:
o 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.
o 2M(NO3)2(s) ==> 2MO(s) + 4NO2(g) + O2(g) (M = Be?, Mg, Ca, Sr, Ba)
o Thermal Tdecomp order Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2 > BeNO3)2
Group 1 and Group 2 hydroxides
o 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