Properties down group 2d.recall the trends in solubility of the hydroxides and sulfates of group 2 elementsSULPHATEsolubility DECREASES down the groupHYDROXIDEsolubility INCREASES down the group

Properties down group 2d.recall the trends in solubility of the hydroxides and sulfates of group 2 elementsSolubility is determined by two factors:Lattice dissociation enthalpy (energy needed to break up a crystal lattice)Hydration enthalpy (energy released when ions are hydrated)

Ionic size has an effect on these factors.Solubility of HydroxidesThere is a decrease in lattice dissociation enthalpy down the group.This outweighs the change in enthalpy of hydration.As a result, there is an INCREASE in SOLUBILITY down the group.insolubleslightly solublesolublesoluble

Properties down group 2d.recall the trends in solubility of the hydroxides and sulfates of group 2 elementsSolubility is determined by two factors:Lattice dissociation enthalpy (energy needed to break up a crystal lattice)Hydration enthalpy (energy released when ions are hydrated)

Ionic size has an effect on these factors.Solubility of SulphatesMagnesium and calcium sulphates are soluble.

Anion >> Cation: so lattice enthalpy does not vary much down group as cation size changes.

However, hydration enthalpy decreases down the group.

Therefore, solubility decreases down the group.solublesolubleinsolubleinsoluble

Properties down group 2e.recall the trends in thermal stability of the nitrates and the carbonates of the elements in groups 1 and 2 and explain these in terms of size and charge of the cations involved

Properties down group 2e.recall the trends in thermal stability of the nitrates and the carbonates of the elements in groups 1 and 2 and explain these in terms of size and charge of the cations involvedThermal Stability of CarbonatesCarbonates of Group 1 are thermally stable:Exception is lithium carbonate which decomposes to give the oxide:

All group 2 carbonates decompose to form stable oxides. E.g.

Properties down group 2e.recall the trends in thermal stability of the nitrates and the carbonates of the elements in groups 1 and 2 and explain these in terms of size and charge of the cations involvedExplanation of stability of carbonatesPolarising the carbonate ionThe positive ion attracts the delocalised electrons in the carbonate ion towards itself. The carbonate ion becomes polarised.If this is heated, the carbon dioxide breaks free to leave the metal oxide.

The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. As the positive ions get bigger as you go down the Group, they have less effect on the carbonate ions near them. 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.In other words, as you go down the Group, the carbonates become more thermally stable.(A similar explanation can be used for the stability of nitrates.)

The effect of heat on the Group 2 carbonatesAll the carbonates in this Group undergo thermal decomposition to give the metal oxide and carbon dioxide gas.Thermal decomposition is the term given to splitting up a compound by heating it.All of these carbonates are white solids, and the oxides that are produced are also white solids.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 carbonates become more stable to heat as you go down the Group.

The effect of heat on the Group 2 nitratesAll the nitrates in this Group undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen.The nitrates are white solids, and the oxides produced are also white solids. Brown nitrogen dioxide gas is given off together with 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.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 nitrates also become more stable to heat as you go down the Group.

SummaryBoth carbonates and nitrates become more thermally stable as you go down the Group. The ones lower down have to be heated more strongly than those at the top before they will decompose.Explaining the trend in terms of the polarising ability of the positive ionA small 2+ ion has a lot of charge packed into a small volume of space. It has a highcharge densityand will have a marked distorting effect on any negative ions which happen to be near it.A bigger 2+ ion has the same charge spread over a larger volume of space. Its charge density will be lower, and it will cause less distortion to nearby negative ions.The structure of the carbonate ionIf 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. 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. We say that the charges aredelocalised.

The next diagram shows the delocalised electrons. The shading is intended to show that there is a greater chance of finding them around the oxygen atoms than near the carbon.

Polarising the carbonate ionNow imagine what happens when this ion is placed next to a positive ion. The positive ion attracts the delocalised electrons in the carbonate ion towards itself. The carbonate ion becomes polarised.

If this is heated, the carbon dioxide breaks free to leave the metal oxide.How much you need to heat the carbonate before that happens depends on how polarised the ion was. If it is highly polarised, you need less heat than if it is only slightly polarised.The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. As the positive ions get bigger as you go down the Group, they have less effect on the carbonate ions near them. 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.In other words, as you go down the Group, the carbonates become more thermally stable.What about the nitrates?The argument is exactly the same here. The small positive ions at the top of the Group polarise the nitrate ions more than the larger positive ions at the bottom. Drawing diagrams to show this happening is much more difficult because the process has interactions involving more than one nitrate ion. You wouldn't be expected to attempt to draw this in an exam.

Using an enthalpy cycleYou 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.Confusingly, there are two ways of defining lattice enthalpy. 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 enthalpyis the heat needed to split one mole of crystal in its standard state into its separate gaseous ions. For example, for magnesium oxide, it is the heat needed to carry out 1 mole of this change:

Solubility of the hydroxidesThe hydroxides becomemore solubleas you go down the Group.This is a trend which holds for the whole Group, and applies whichever set of data you choose.Some examples may help you to remember the trend:Magnesium hydroxideappears to be insoluble in water. However, if you shake it with water, filter it and test the pH of the solution, you find that it is slightly alkaline. This shows that there are more hydroxide ions in the solution than there were in the original water. Some magnesium hydroxide must have dissolved.Calcium hydroxidesolution is used as "lime water". 1 litre of pure water will dissolve about 1 gram of calcium hydroxide at room temperature.Barium hydroxideis soluble enough to be able to produce a solution with a concentration of around 0.1 mol dm-3at room temperature.

olubility of the sulphatesThe sulphates becomeless solubleas you go down the Group.The simple trend is true provided you include hydrated beryllium sulphate in it, but not if the beryllium sulphate is anhydrous.The Nuffield Data Book quotes anyhydrous beryllium sulphate, BeSO4, asinsoluble(I haven't been able to confirm this from any other source), whereas the hydrated form, BeSO4.4H2O is soluble. (The Data Books agree on this - giving a figure of about 39 g dissolving in 100 g of water at room temperature.)Figures for magnesium sulphate and calcium sulphate also vary depending on whether the salt is hydrated or not, but nothing like so dramatically.Two common examples may help you to remember the trend:You are probably familiar with the reaction between magnesium and dilute sulphuric acid to give lots of hydrogen and a colourless solution ofmagnesium sulphate. Notice that you get a solution, not a precipitate. The magnesium sulphate is obviously soluble.You may also remember thatbarium sulphateis formed as a white precipitate during the test for sulphate ions in solution. The ready formation of a precipitate shows that the barium sulphate must be pretty insoluble. In fact, 1 litre of water will only dissolve about 2 mg of barium sulphate at room temperature.

Solubility of the carbonatesThe carbonates tend to becomeless solubleas you go down the Group.None of the carbonates is anything more than very sparingly soluble. Magnesium carbonate (the most soluble one I have data for) is soluble to the extent of about 0.02 g per 100 g of water at room temperature.I can't find any data for beryllium carbonate, but it tends to react with water and so that might confuse the trend.The trend to lower solubility is, however, broken at the bottom of the Group. Barium carbonate is slightly more soluble than strontium sulphate.There are no simple examples which might help you to remember the carbonate trend.

The E valuesYou will remember that the oxidising ability of the halogens decreases as you go down the Group in the Periodic Table.The E values of the four halogens from fluorine to iodine are: Remember that the electrode potentials give a measure of the positions of the equilibria. The more positive, the further the equilibrium lies to the right.In the fluorine case, the E value is almost as positive as they get. That means that fluorine will very readily pick up electrons to make fluoride ions.Fluorine will therefore remove electrons from other things extremely well. Taking electrons away from something is oxidising it. So fluorine is a very powerful oxidising agent indeed.As you go down the rest of the group, the E values become less positive, and so the oxidising ability decreases.

Why does chlorine oxidise iodide ions to iodine?The two E values are: When you couple two of these equilibria together in a test tube, the more positive one will tend to move to the right, and the more negative one (or less positive one) to the left.That is exactly what you want to happen to turn iodide ions into iodine. The chlorine E is more positive, and so chlorine molecules take electrons from the iodide ions to turn them into iodine.

Why won't bromine oxidise chloride ions to chlorine?The two E values are: The chlorine equilibrium lies further to the right because it is the more positive. That means that if you couple the two equilibria together, you would expect the chlorine one to move to the right and the bromine one to the left.But if you start with bromine and chloride ions, the two equilibria are already as far in those directions as possible. To get a reaction, they would have to move in a direction opposite to that predicted by the E values. That can't happen.Which halogens could you use to oxidise Fe2+to Fe3+?The E value for the iron(II) / iron(III) system is . . . and what you want to do is to drive it to the left to turn Fe2+into Fe3+. To do that you would need to couple it with something with a more positive E value.If you look at the halogen list above, you will see that fluorine, chlorine and bromine are all capable of oxidising iron(II) to iron(III), but iodine isn't.

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