What periodic trends are observed for the relative solubilities of the compounds

Question

What periodic trends are observed for the relative solubilities of the compounds formed from the Group 2 metal cations? Choose one anion and compare the solubilities of the compounds formed when each of the Group 2 metal cations is combined with the anion you chose. Suggest a reason for the observed trends using your knowledge of the periodic table. (How do the cations differ from one another?) Choose a second anion and again compare the relative solubilities of the compounds it formed with the Group 2 metal cations. Do the compounds containing this anion show the same trend in solubility as the compounds containing the first anion you discussed?

Explanation / Answer

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Unlike alkaline salts, which used to be soluble, here they are just the opposite, many are insoluble (or rather poorly soluble). Salts with monocharged anions, such as chlorides and nitrates, are usually soluble; those with more than one charge are usually insoluble, but nonetheless sulfates which have a double negative charge, it gets soluble to insoluble as it is lowered into the group.

As with the alkali metals, entropic and enthalpic terms are quite similar, and this gives problems; however, the values are much higher compared to those of the alkali. For example, if anhydrous calcium chloride in water is tossed, the dissolution process is exothermic and heats the water quite appreciably. Sometimes, the roles are reversed, for example, the entropic factors are that favor sodium chloride is soluble, but instead on magnesium chloride do not favor. The fact is that we talk about relatively high energy factors, which numerically differ very little from each other.

the alkaline ionic radii are smaller than those of the alkali in the same period ... even so alkaline charge are twice. The charge / ion radius ratio is a measure of "polarizing power", ie, the ability to alter the ion orbitals of an atom is about.

In other words, the ions of this group are very small (less the higher the group) and also have a relatively large ccharge, so that this is highly concentrated. it explains whether or not they have some salts of these elements, hydration salts and the similarity between some elements of different groups such as lithium and magnesium, for side, and beryllium and aluminum, on the other.

The basic strength of hydroxides in water increases to lower in the group with decreasing charge density ion, so that the Be (OH) 2 is amphoteric, Mg (OH) 2 is a weak base, Ca (OH) 2 and Sr (OH) 2 are moderately strong bases and Ba (OH) 2 has a basicity approaches the alkali hydroxide.

Be2 + (aq) + 2 OH- (aq) Be (OH) 2 (s)

Be (OH) 2 (s) + 2 OH- (aq) [Be (OH) 4] 2- (aq)

The solubilities of the hydroxides in water increases with the size of the metal, as expected since the anion OH- is small. This difference in solubilities is profited in the Dow process for obtaining magnesium. Due to the high insolubility of magnesium hydroxide, in solid, finely divided and mixed with water forms a suspension which is used in medicine as an antacid. In the stomach, the solid acid neutralizes excess without there being virtually free OH- that would cause very severe burns.

Mg (OH) 2 (s) + 2 H + (aq) Mg2 + (aq) + 2 H2O (l)

Of these, calcium hydroxide is the most commercially important because it is the cheapest strong base in the market.

Sulfates .

The solubility of sulfate decreases with decreasing in the group:

Be> Mg >> Ca > Sr > Ba

As expected, the Be and Mg are soluble due to the high hydration entapías its ions , while slightly soluble Ca and Sr , Ba and Ra are virtually insoluble . Of these the most commercially important calcium is in form of hemihydrate , CaSO4.1 / 2H2O , is used as building material, plaster. Sulfates decompose by the heat and causing the corresponding oxides , as occurred with carbonates , the decomposition temperatures increase with the size of the alkaline earth ion

MgSO 4 ( s ) + Q MgO ( s ) + SO 3 ( g )

Sulphates are reduced to sulphides with coal at elevated temperature. Most of the barium compounds prepared from their sulfide

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