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The lattice enthalpies of both carbonates and oxides fall as you go down the Group because the positive ions are getting bigger. The oxide ion is relatively small for a negative ion (0.140 nm), whereas the carbonate ion is large (no figure available). A small 2+ ion has a lot of charge packed into a small volume of space. If it is highly polarised, you need less heat than if it is only slightly polarised. Today we're covering: Properties of Group 2 compounds Reactions Oxides with water Carbonates with acid Thermal decomposition Carbonates Nitrates Solubility Hydroxides Sulfates Let's go! Solubility Rules . The argument is exactly the same here. The size of the lattice enthalpy is governed by several factors, one of which is the distance between the centres of the positive and negative ions in the lattice. 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. How much you need to heat the carbonate before that happens depends on how polarised the ion was. Explaining the relative falls in lattice enthalpy. For the sake of argument, suppose that the carbonate ion radius was 0.3 nm. The nitrate ion is bigger than an oxide ion, and so its radius tends to dominate the inter-ionic distance. The small cations at the top of the group polarize the nitrate ions more than the larger cations at the bottom do. 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. Almost all inorganic nitrates are soluble in water.An example of an insoluble nitrate is Bismuth oxynitrate.Removal of one electron yields the nitrate radical, also called nitrogen trioxide NO This page examines at the effect of heat on the carbonates and nitrates of the Group 2 elements (beryllium, magnesium, calcium, strontium and barium). I had explained all of the trends except one, group 2 nitrates. All group 2 nitrates and chlorides are soluble, but the solubility of the group 2 sulphates decreases down the group -Magnesium sulphate is classed as soluble -Calcium sulphate is classed as slightly soluble -Strontium and barium sulphate are insoluble 3.19 Recall the general rules which describe the solubility of common types of substances in water: all common sodium, potassium and ammonium salts are soluble; all nitrates are soluble; common chlorides are soluble except those of silver and lead… A/AS level. More polarization requires less heat. ... As you descend group II hydroxide solubility increases. They are in Group 2 (Acids, Inorganic Oxidizing). No headers. You need to find out which of these your examiners are likely to expect from you so that you don't get involved in more difficult things than you actually need. Magnesium carbonate (the most soluble Group 2 carbonate) has a solubility of about 0.02 g per 100 g of water at room temperature. 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%. The small positive ions at the top of the Group polarise the nitrate ions more than the larger positive ions at the bottom. For example, for magnesium oxide, it is the heat needed to carry out 1 mole of this change: $MgO_{(s)} \rightarrow Mg^{2+}_{(g)} + O^{2-}_{(g)}$. For the sake of argument, suppose that the carbonate ion radius was 0.3 nm. It reacts with cold water to produce an alkaline solution of calcium hydroxide and hydrogen gas is released. This page offers two different ways of looking at the problem. The lattice enthalpy of the oxide will again fall faster than the nitrate. SOLUBILITY OF COMPOUNDS (GROUP 1) Solubility of a compound mainly depends on two factors . Magnesium carbonate, for example, has a solubility of about 0.02 g per 100 g of water at room temperature. I can't find a value for the radius of a carbonate ion, and so can't use real figures. The rates at which the two lattice energies fall as you go down the Group depends on the percentage change as you go from one compound to the next. However, in a reaction with steam it forms magnesium oxide and hydrogen. 2Mg(NO 3) 2 → 2MgO + 4NO 2 + O 2 Explaining the trend in terms of the polarising ability of the positive ion. Contents This is a rather more complicated version of the bonding you might have come across in benzene or in ions like ethanoate. You should look at your syllabus, and past exam papers - together with their mark schemes. All the Group 2 carbonates are very sparingly soluble. Most of the precipitation reactions that we will deal with involve aqueous salt solutions. The nitrates are white solids, and the oxides produced are also white solids. The increasing thermal stability of Group 2 metal salts is consistently seen. The carbonates become more stable to heat as you go down the Group. Remember that the reaction in question is the following: $XCO_{3(s)} \rightarrow XO_{(s)} + CO_{2(g)}$. We say that the charges are delocalised. Here's where things start to get difficult! All of these carbonates are white solids, and the oxides that are produced are also white solids. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. Detailed explanations are given for the carbonates because the diagrams are easier to draw, and their equations are also easier. Watch the recordings here on Youtube! You will need to use the BACK BUTTON on your browser to come back here afterwards. Figures to calculate the beryllium carbonate value weren't available. So what causes this trend? Down the group, the nitrates must also be heated more strongly before they will decompose. The larger compounds further down require more heat than the lighter compounds in order to decompose. The amount of heating required depends on the degree to which the ion is polarized. If "X" represents any one of the elements, the following describes this decomposition: Down the group, the carbonates require more heating to decompose. Covers the elements beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr) and barium (Ba). These compounds are white solids and brown nitrogen dioxide and oxygen gases are also given off when heated. The oxide lattice enthalpy falls faster than the carbonate one. 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