Remember that the reaction we are talking about is: You can see that the reactions become more endothermic as you go down the Group. In the carbonates, the inter-ionic distance is dominated by the much larger carbonate ion. For example, for magnesium oxide, it is the heat needed to carry out 1 mole of this change: Note: In that case, the lattice enthalpy for magnesium oxide would be -3889 kJ mol-1. If it is highly polarised, you need less heat than if it is only slightly polarised. Brown nitrogen dioxide gas is given off together with oxygen. Remember that the reaction we are talking about is: You can see that the reactions become more endothermic as you go down the Group. The effect of heat on the Group 2 carbonates. All the carbonates in this group undergo thermal decomposition to the metal oxide and carbon dioxide gas. The Effect of Heat on the Group 2 Nitrates All the nitrates in this Group undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen. On that basis, the oxide lattice enthalpies are bound to fall faster than those of the carbonates. The enthalpy changes (in kJ mol-1) which I calculated from enthalpy changes of formation are given in the table. On that basis, the oxide lattice enthalpies are bound to fall faster than those of the carbonates. In a Unit 2 question it asks: Calcium nitrate decomposes in a similar way to magnesium nitrate, but at ahigher temperature. I can't find a value for the radius of a carbonate ion, and so can't use real figures. What factors affect this trend? The shading is intended to show that there is a greater chance of finding them around the oxygen atoms than near the carbon. if you constructed a cycle like that further up the page, the same arguments would apply. Figures to calculate the beryllium carbonate value weren't available. But they don't fall at the same rate. If "X" represents any one of the elements: As you go down the Group, the carbonates have to be heated more strongly before they will decompose. The reason, once more, is that the polarising power of the M2+decreases as ionic radius increases. Hydrides liberate hydrogen at anode on electrolysis. You have to supply increasing amounts of heat energy to make them decompose. If this is heated, the carbon dioxide breaks free to leave the metal oxide. The effect of heat on the Group 2 nitrates. Confusingly, there are two ways of defining lattice enthalpy. (substitute Na, K etc where Li is). The ones lower down have to be heated more strongly than those at the top before they will decompose. Both carbonates and nitrates become more thermally stable as you go down the Group. Explaining the trend in terms of the polarising ability of the positive ion. If this is heated, the carbon dioxide breaks free to leave the metal oxide. The inter-ionic distances in the two cases we are talking about would increase from 0.365 nm to 0.399 nm – an increase of only about 9%. XCO_{3(s)} \longrightarrow XO_{(s)} + CO_{2(g)}, 2X(NO_3)_{2(s)} \longrightarrow 2XO_{(s)} + 4NO_{2(g)} + O_{2(g)}, \begin{gathered} THERMAL STABILITY OF THE GROUP 2 CARBONATES AND NITRATES. The small positive ions at the top of the Group polarise the nitrate ions more than the larger positive ions at the bottom. Only lithium carbonate and group 2 carbonates decompose (in Bunsen flame, 1300K). You can dig around to find the underlying causes of the increasingly endothermic changes as you go down the Group by drawing an enthalpy cycle involving the lattice enthalpies of the metal carbonates and the metal oxides. This means you polarize the electron cloud less, producing stronger ionic bonds. The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. If you calculate the enthalpy changes for the decomposition of the various carbonates, you find that all the changes are quite strongly endothermic. A small 2+ ion has a lot of charge packed into a small volume of space. 3. The lattice enthalpy of the oxide will again fall faster than the nitrate. if you constructed a cycle like that further up the page, the same arguments would apply. A smaller 2+ ion has more charge packed into a smaller volume than a larger 2+ ion (greater charge density).. Start studying Thermal stability of Group II nitrates, carbonates and hydroxides. The larger compounds further down require more heat than the lighter compounds in order to decompose. A bigger 2+ ion has the same charge spread over a larger volume of space. The nitrates are white solids, and the oxides produced are also white solids. The rest of group 1 follow the same pattern. How much you need to heat the carbonate before that happens depends on how polarised the ion was. 2. 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