solubility of group 2 chlorides

I'm going to "suggest" that BeCl2 is the least soluble because the bonds are less polar, with a greater covalent character. This is where the explanation usually stops, but to stop at this point is very misleading because it won't explain all the facts! Compound - solubility --- molar solubility--- percent ionic character, BeCl2 == 15.1 g/100 mL .... 1.89 mol / L .... 47%, MgCl2 == 54.3 g/100 ml ..... 5.70 mol / L......58%, CaCl2 == 74.5 g/100m........ 6.71 mol / L......69%, SrCl2 == 53.8 g/100 mL ..... 3.39 mol / L......71%, BaCl2 == 35.8 g/100 mL .... 1.72 mol / L......72%. And indeed, BeCl2 has the lowest solubility in terms of grams per 100 mL, but then comparing masses isn't the best way to make the comparison since beryllium is also the lightest element in group 2. (b) sulphates = The sulphates become less soluble as you go down the Group. The way those changes happen will vary from one type of compound to another. In this case, the enthalpy of solution will become more positive (or less negative). Mg 2+ (aq) reacts with NaOH to form a white precipitate because Mg(OH) 2 is insoluble (only sparingly soluble). You might have expected exactly the opposite to happen. BaSO4 is the least soluble. I do a similar activity with my students and we find, just as we have found here, that the polarity of the bonds as measured by the percent ionic character isn't a good indicator of solubility. The enthalpy of solution figures for the Group 2 carbonates are: (source: Chemistry Data Book by Stark and Wallace; values in kJ mol-1). Cl, Br, I: All chlorides, bromides, and iodides are soluble except those of silver, mercury, and lead (e.g., AgCl, Hg 2 Cl 2, and PbCl 2). The outer The entropy change is becoming less negative (or perhaps even at this stage, positive). does the entropy increase when sodium chloride dissolve in water? Generally, Group 2 elements that form compounds with single charged negative ions (e.g. The most obvious thing that's wrong is that it won't explain why some compounds (like magnesium carbonate, and most of the other Group 2 carbonates) don't dissolve in water even though their enthalpies of solution are mainly negative. For large negative ions like sulphate or carbonate, the hydration enthalpy of the positive ions falls faster than the lattice enthalpy. The substances are listed in alphabetical order. Instead, compare the molar solubilities. . With sulphates, for example, the percentage increase in the inter-ionic distance as you go from magnesium to calcium sulphate isn't as great as it would be with a smaller negative ion like hydroxide. The solubilities of the Group 1 chlorides (in moles of solute saturating 100 g of water at 298 K) compared with their enthalpies of solution are: There is no obvious relationship connecting the relative movements of these solubility values with the enthalpy of solution figures. The bigger the ions, the more distance there is between them, and the weaker the forces holding them together. And indeed, BeCl2 has the lowest solubility in terms of grams per 100 mL, but then comparing masses isn't the best way to make the comparison since beryllium is also the lightest element in group 2. This page looks at the usual explanations for the solubility patterns in the hydroxides, sulphates and carbonates of Group 2. Net Ionic Equation Definition. 1.3.2 (b) Reactivity of Group 2 Elements. Silver acetate is sparingly soluble. All the Group 2 carbonates are very sparingly soluble. with metals Cl-, Br-, I-, etc. The reasons for the discrepancies lie in the way the numbers are calculated. Entropy is given the symbol S. If a system becomes more disordered, then its entropy increases. The Nuffield Data Book quotes anyhydrous beryllium sulphate, BeSO 4 , as insoluble (I haven't been able to confirm this from any other source), whereas the hydrated form, BeSO 4 .4H 2 O is soluble. You could, however, make a reasonable suggestion as to why the solubility trend in the carbonates is broken at barium. The nitrates, chlorates, and acetates of all metals are soluble in water. The Nuffield Data Book doesn't have any hydration enthalpy values. So . These can be combined mathematically to give an important term known as free energy change. That means that you have two entropy effects to consider. Unfortunately, the enthalpy of solution values for the Group 1 chlorides as calculated above don't agree with the values given in the same Data Book: The discrepancies are enough to disrupt any pattern (such as there is!). In the sodium chloride case, you don't have to have very much increase in entropy to outweigh the small enthalpy change of +3.9 kJ mol-1. Instead of milling around pretty much at random, they become attracted to the ions present and arranged around them. If you are unfortunate enough that your examiners expect you to explain this, use past papers, mark schemes and examiner's reports if they are available, and find out exactly what your examiners expect you to say. That would seem to support the decrease in solubility as you go down the Group quite nicely. which makes you more jittery coffee or tea? In these cases, the entropy of the system must fall when the compounds dissolve in water - in other words, the solution in water is more ordered than the original crystal and water! The chlorides, bromides, and iodides of all metals except lead, silver, and mercury(I) are soluble … Taking the sign of enthalpy of solution at face value, you get some bizarre results. You can see that the enthalpy of solution changes from NaCl to KCl because the lattice enthalpy and hydration enthalpy of the positive ion fall by different amounts. For example, although it might be possible to account for the lack of pattern in the solubilities of the Group 1 chlorides (and also the bromides) by a mathematical application of these effects, trying to do it in general terms defeats me completely! I see that as quite dangerous. DOI: 10.1016/j.jct.2016.09.031. Remember that where you have a big negative ion, its size dominates the inter-ionic distance and so doesn't allow the lattice enthalpy to change much. This gives the enthalpy of solution values we've already looked at (values in kJ mol-1): But the entropy change will also be varying as you go down the Group. Energy has to be supplied to break up the lattice of ions, and energy is released when these ions form bonds of one sort or another with water molecules. The overall effect is a complex balance between the way the enthalpy of solution varies and the way the entropy change of solution alters. It would be quite untrue to say that the more endothermic the change, the less soluble the compound! Clearly, trying to correlate solubility simply with the enthalpy change of solution doesn't work. Don't expect this page to be easy - it is probably best avoided unless your syllabus specifically asks for these explanations! In this case, we are defining lattice enthalpy as the heat needed to convert 1 mole of crystal in its standard state into separate gaseous ions - an endothermic change. . Solubility of chlorides of alkali metal decrease down the group. (Don't expect the explanation to be instantly understandable though!). And indeed, BeCl2 has the lowest solubility in terms of grams per 100 mL, but then comparing masses isn't the best way to make the comparison since beryllium is also the lightest element in group 2. What happens if the enthalpy change is positive - as for example when sodium chloride dissolves in water (+3.9 kJ mol-1, using the values in one of the tables above)? 3. The relationship between enthalpy of solution and solubility. Get your answers by asking now. That means that the enthalpy of solution will become less positive (or more negative). Where you have a big negative ion, this inter-ionic distance is largely controlled by the size of that negative ion. Yes, it does! For example, if each of the numbers in the calculations we did earlier on this page was out by just 5 kJ, each answer could vary by +/- 15 kJ - completely disrupting the patterns! Chem help, chemical bonds and wavelength. Alkali metals (Group I) Na +, K +, etc. Should I call the police on then? As an approximation, for a reaction to happen, the free energy change must be negative. At barium carbonate, the effect of increasing entropy must be enough to make it more soluble than strontium carbonate. In order to see whether a change is possible or not, you have to think about a combination of the enthalpy change and the entropy change. 2. Sodium chloride and the other Group 1 chlorides dissolve despite the fact that their enthalpies of solution are positive, and yet magnesium carbonate (and most of the other Group 2 carbonates) are very sparingly soluble, but have exothermic enthalpies of solution. Problems in relating the sign of the enthalpy change to solubility. You can't therefore reliably use the data available to calculate the trends you want with sufficient accuracy to make sense.

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