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The diagram shows what happens with an ion from Group 2, carrying two positive charges: If this system is heated, the carbon dioxide breaks free, leaving a metal oxide. Solubility of the carbonates increases as you go down Group 1. Now imagine what happens when this ion is placed next to a positive ion. 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. These are made by passing hydrogen gas over the heated metal. In other words, the charges are delocalized. The carbonates get more soluble as you go down Group 1, but tend to get less soluble down Group 2. There is no clear solubility trend observed down this group. The carbonates of alkali metals are stable towards heat. Solubility of the hydroxides increases down Group 1. It is also difficult to get reliable results if you heat these carbonates in the lab. You will often find that the lithium compounds behave similarly to Group 2 compounds, but the rest of Group 1 are in some way different. For UK A level purposes all you would need to do is talk about how the polarising ability of the positive ion increases as it gets smaller or more charged. Its charge density will be lower, and it will cause less distortion to nearby negative ions. Lv 6. For example, a typical Group 2 carbonate like calcium carbonate decomposes like this: $CaCO_3 (s) \rightarrow CaO(s) + CO_2$. Legal. A saturated solution has a concentration of about 1.3 g per 100 g of water at 20°C. The decomposition temperatures again increase as you go down the Group. For the purposes of this topic, you don't need to understand how this bonding has come about. Most carbonates decompose on heating to the metal oxide and carbon dioxide. Unit 1: THE LANGUAGE OF CHEMISTRY, STRUCTURE OF MATTER AND SIMPLE REACTIONS. Just learn that Group 1 compounds tend to be more soluble than their Group 2 equivalents. Imagine that this ion is placed next to a positive ion. are carbonates soluble in water? Any attempt to extract them from solution causes them to decompose to the carbonate, carbon dioxide and water as shown: $Ca(HCO_3)_2 (aq) \rightarrow CaCO_3 (s) + CO_2 (g) + H_2O (l)$. The hydroxides. None of the carbonates is anything more than very sparingly soluble. Solubility of the carbonates. For example, a typical Group 2 nitrate like magnesium nitrate decomposes like this: In Group 1, lithium nitrate behaves in the same way - producing lithium oxide, nitrogen dioxide and oxygen. The hard way is in terms of the energetics of the process; the simple way is in terms of the polarizing ability of the positive ions. The metal is deposited at the cathode as expected. The least soluble hydroxide in Group 1 is lithium hydroxide - but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C. Detailed explanations are given for the carbonates because the diagrams are easier to draw. As the positive ions get bigger down the group, they have less effect on the carbonate ions near them. This page discusses a few compounds of the Group 1 elements (lithium, sodium, potassium, rubidium and cesium), including some information about the nitrates, carbonates, hydrogen carbonates and hydrides of the metals. So perhaps there is nothing special about the hydrogen mentioned above. And, again, the Group 1 compounds will need to be heated more strongly than those in Group 2 because the Group 1 ions are less polarising. For example, for lithium hydride: Two of the most common reactions include electrolysis and reactions with water. I'm not sure what the purpose of the hydrogen is. The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. Solubility in water: * Except Li 2 CO 3, The group-1 carbonates are fairly soluble in water. A bigger positive ion has the same charge spread over a larger volume of space. The hydrides of Group 1 metals are white crystalline solids which contain the metal ions and hydride ions, H-. Group 1 compounds are more stable to heat than the corresponding compounds in Group 2. This is too difficult to talk about at this level - and I'm not going to do it! The diagram shows what happens with an ion from Group 2, carrying two positive charges. Solubility of the carbonates increases as you go down Group 1. Please provide a detailed answer. For UK A level purposes, the important thing to remember is that Group 1 compounds tend to be more soluble than the corresponding ones in Group 2. The rest of the Group 1 carbonates don't decompose at Bunsen temperatures, although at higher temperatures they will. They have exactly the same crystal structure as sodium chloride - that's why they are called saline or salt-like hydrides. Magnesium carbonate (the most soluble Group 2 carbonate) has a solubility of about 0.02 g per 100 g of water at room temperature. The least soluble hydroxide in Group 1 is lithium hydroxide - but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C. They produce the metal nitrite and oxygen, but no nitrogen dioxide: $2XNO_3 (s) \rightarrow 2XNO_2(s) + O_2 (g)$. The hard way is in terms of the energetics of the process; the simple way is to look at the polarising ability of the positive ions. Solubility of the carbonates. Unless otherwise noted, LibreTexts content is licensed by CC BY-NC-SA 3.0. 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