Nitrates of alkaline-earth metals and LiNO3 decompose on heating to form oxides, nitrogen to form oxides, nitrogen dioxide and oxygen. Magnesium and calcium nitrates normally have water of crystallisation, and the solid may dissolve in its own water of crystallisation to make a colourless solution before it starts to decompose. 3. Note: If you are interested, you could follow these links to benzene or to organic acids. The ones lower down have to be heated more strongly than those at the top before they will decompose. This is a rather more complicated version of the bonding you might have come across in benzene or in ions like ethanoate. It describes and explains how the thermal stability of the compounds changes as you go down the Group. It describes and explains how the thermal stability of the compounds changes as you go down the Group. Now imagine what happens when this ion is placed next to a positive ion. The larger compounds further down require more heat than the lighter compounds in order to decompose. 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} (substitute Na, K etc where Li is). 2LiNO3 +Heat -> Li 2 O +2NO 2 +O 2 2Ca (NO 3) 2 +Heat -> 2CaO +4NO 2 +O 2 Thermal stabilities of nitrates of group-1 and group-2 metals increase on moving down the group from top to bottom. The shading is intended to show that there is a greater chance of finding them around the oxygen atoms than near the carbon. The effect of heat on the Group 2 carbonates. The thermal stability of the nitrates follows the same trend as that of the carbonates, with thermal stability increasing with proton number. For the purposes of this topic, you don't need to understand how this bonding has come about. This is a rather more complicated version of the bonding you might have come across in benzene or in ions like ethanoate. The carbonates and nitrates of group 2 elements carbonates become more thermally stable as you go down the Group. The argument is exactly the same here. That implies that the reactions are likely to have to be heated constantly to make them happen. Here's where things start to get difficult! Now imagine what happens when this ion is placed next to a positive ion. Although the inter-ionic distance will increase by the same amount as you go from magnesium carbonate to calcium carbonate, as a percentage of the total distance the increase will be much less. Forces of attraction are greatest if the distances between the ions are small. \end{gathered}. A small 2+ ion has a lot of charge packed into a small volume of space. (2) 2 X (N O 3) 2 (s) → 2 X O (s) + 4 N O 2 (g) + O 2 (g) Down the group, the nitrates must also be heated more strongly before they will decompose. All other group 1 carbonates are stable in Bunsen flame. You would observe brown gas evolving (NO2) and the White nitrate solid is seen to melt to a colourless solution and then resolidify 2Mg(NO3)2→ 2MgO + … Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals. All the carbonates in this group undergo thermal decomposition to the metal oxide and carbon dioxide gas. Thermal Stability Group 2 In this Group 2 tutorial we look at the thermal stability of metal nitrates and carbonates and the trends down groups 1 and 2. The carbonates become more stable to heat as you go down the Group. Thermal Stability of Group 1/2 Nitrates (4:38) Flame tests (9:14) Uses of Group 2 Compounds AS: GROUP 7 (4B) GROUP 7 OVERVIEW Group 7 Properties & Trends (6:55) Testing for Halide Ions Reactions of Group … Exactly the same arguments apply to the nitrates. For the sake of argument, suppose that the carbonate ion radius was 0.3 nm. The decomposition temperature of - and -substituted derivatives is found to be linearly related to the Hammett substituent constant σ. Since both 2-methyl-2-butanol nitrate and 2-methyl-2-propanol nitrate exhibited low thermal stability, they were not distilled from the reaction solvent diethyl ether. The nitrates are white solids, and the oxides produced are also white solids. Two factors are involved in dissolving: 1. The cycle we are interested in looks like this: You can apply Hess's Law to this, and find two routes which will have an equal enthalpy change because they start and end in the same places. The rest of group 1 follow the same pattern. Brown nitrogen dioxide gas is given off together with oxygen. The nitrates also become more stable to heat as you go down the Group. Explaining the trend in terms of the energetics of the process. Its charge density will be lower, and it will cause less distortion to nearby negative ions. It has a high charge density and will have a marked distorting effect on any negative ions which happen to be near it. Both carbonates and nitrates of Group 2 elements become more thermally stable down the group. If you think carefully about what happens to the value of the overall enthalpy change of the decomposition reaction, you will see that it gradually becomes more positive as you go down the Group. Since the ionic radius of the metal ion increases, this will reduce the distortion to the NO3^ - electron cloud. In order to make the argument mathematically simpler, during the rest of this page I am going to use the less common version (as far as UK A-level syllabuses are concerned): Lattice enthalpy is the heat needed to split one mole of crystal in its standard state into its separate gaseous ions. Even for hydroxides we have the same observations. If this is the first set of questions you have done, please read the introductory page before you start. Only lithium carbonate and group 2 carbonates decompose (in Bunsen flame, 1300K). In the oxides, when you go from magnesium oxide to calcium oxide, for example, the inter-ionic distance increases from 0.205 nm (0.140 + 0.065) to 0.239 nm (0.140 + 0.099) - an increase of about 17%. For the purposes of this topic, you don't need to understand how this bonding has come about. The peroxy nitrates shown in Table II are observed to fall into two classes of thermal stability. Similar to lithium nitrate, alkaline earth metal nitrates also decompose to give oxides. Group 2 nitrates become more thermally stable down the group. All the nitrates in this Group undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen. The rest of Group 2 follow the same pattern. In the carbonates, the inter-ionic distance is dominated by the much larger carbonate ion. The next diagram shows the delocalised electrons. (You wouldn't see the oxygen also produced). Although the inter-ionic distance will increase by the same amount as you go from magnesium carbonate to calcium carbonate, as a percentage of the total distance the increase will be much less. Magnesium and calcium nitrates normally have water of crystallisation, and the solid may dissolve in its own water of crystallisation to make a colourless solution before it starts to decompose. Thermal decomposition is the term given to splitting up a compound by heating it. You have to supply increasing amounts of heat energy to make them decompose. Again, if "X" represents any one of the elements: As you go down the Group, the nitrates also have to be heated more strongly before they will decompose. The carbonate ion becomes polarised. The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. Which of these statements is correct? The stability appears to depend on whether or not the peroxy nitrate group (—OONO2) is attached to a carbonyl group (C=O). I can't find a value for the radius of a carbonate ion, and so can't use real figures. Going down group II, the ionic radii of cations increases. Lattice enthalpy: the heat evolved when 1 mole of crystal is formed from its gaseous ions. b) lower c) A white solid producing a brown gas and leaving a white solid. On that basis, the oxide lattice enthalpies are bound to fall faster than those of the carbonates. The thermal stability/reducibility of metal nitrates in an hydrogen atmosphere has also been studied by temperature-programmed reduction (TPR). The oxide ion is relatively small for a negative ion (0.140 nm), whereas the carbonate ion is large (no figure available). All the carbonates in this Group undergo thermal decomposition to give the metal oxide and carbon dioxide gas. You should look at your syllabus, and past exam papers - together with their mark schemes. Unfortunately, in real carbonate ions all the bonds are identical, and the charges are spread out over the whole ion – although concentrated on the oxygen atoms. If you aren't familiar with Hess's Law cycles (or with Born-Haber cycles) and with lattice enthalpies (lattice energies), you aren't going to understand the next bit. The nitrates are white solids, and the oxides produced are also white solids. Confusingly, there are two ways of defining lattice enthalpy. We say that the charges are delocalised. Detailed explanations are given for the carbonates because the diagrams are easier to draw, and their equations are also easier. For reasons we will look at shortly, the lattice enthalpies of both the oxides and carbonates fall as you go down the Group. All of these carbonates are white solids, and the oxides that are produced are also white solids. If you worked out the structure of a carbonate ion using "dots-and-crosses" or some similar method, you would probably come up with: This shows two single carbon-oxygen bonds and one double one, with two of the oxygens each carrying a negative charge. The ones lower down have to be heated more strongly than those at the top before they will decompose. This means that the enthalpy change from the carbonate to the oxide becomes more negative so more heat is needed to decompose it. 2. The ones lower down have to be heated more strongly than those at the top before they will decompose. In other words, as you go down the Group, the carbonates become more thermally stable. THERMAL STABILITY OF THE GROUP 2 CARBONATES AND NITRATES Go to the main page. Drawing diagrams to show this happening is much more difficult because the process has interactions involving more than one nitrate ion. The solubility of these sulphates decreases as we descend the group, with barium sulphate being insoluble in water. Here's where things start to get difficult! A bigger 2+ ion has the same charge spread over a larger volume of space. THERMAL STABILITY OF THE GROUP 2 CARBONATES AND NITRATES This page looks at the effect of heat on the carbonates and nitrates of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium. Its charge density will be lower, and it will cause less distortion to nearby negative ions. All the carbonates in this Group undergo thermal decomposition to give the metal oxide and carbon dioxide gas. The oxide ion is relatively small for a negative ion (0.140 nm), whereas the carbonate ion is large (no figure available). Remember that the reaction we are talking about is: You can see that the reactions become more endothermic as you go down the Group. Enthalpy of hydration. Strontium Nitrate Strontium has a greater ionic radius than beryllium since it is affected by more electrostatic forces of attraction due to more protons in its nucleus and more electron shells. How much you need to heat the carbonate before that happens depends on how polarised the ion was. 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. The thermal stability of ring-substituted arylammonium nitrates has been investigated using thermal methods of analysis. if you constructed a cycle like that further up the page, the same arguments would apply. To compensate for that, you have to heat the compound more in order to persuade the carbon dioxide to break free and leave the metal oxide. For nitrates we notice the same trend. That implies that the reactions are likely to have to be heated constantly to make them happen. \text{Mg}O_{s} \longrightarrow \text{Mg}^{2+}_{(g)} + O^{2-}_{(g)} \\{\Delta}H_{\text{lattice}} = +3889~kJ~mol^{-1} The effect of heat on the Group 2 nitrates. The nitrate ion is less polarised and the compound is more stable. The nitrate ion is bigger than an oxide ion, and so its radius tends to dominate the inter-ionic distance. Products: barium oxide, nitrogen dioxide (nitrogen(IV) oxide) and oxygen d) lower 2. 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. Explaining the relative falls in lattice enthalpy. The term we are using here should more accurately be called the "lattice dissociation enthalpy". 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