Solubility of alkali earth metals/solvation energy,periodic trends, IMFs

[Kraniak]

Homework Statement

Look at the solubility guidelines (Table 4.1). Using your knowledge of periodic trends and the energetics of formation of solutions, explain why cations derived directly from alkali metals are usually soluble, but this is less the case for alkaline Earth metals. Also, discuss how these would be affected in the case of solvation with sulfide ions. You should be clear about the identity of the relevant intermolecular forces as well as periodic trends leading to such differences.

Homework Equations

No equations, the Solubility guidelines referred to is a fairly standard solubility guide telling what ions in water will typically form precipitates. It looks a lot like this:

http://wps.prenhall.com/wps/media/objects/3080/3154577/blb0405/blb04t02.gif

The Attempt at a Solution

I was thinking that the solvation energy might be stronger than the lattice energy of the ionic compounds that might form from the +1 charged alkali metals than in the +2 charged alkali metals, since ionic bond strength is proportionate to charge. But many alkali Earth metals are soluble in water with the sulfide ion, so my idea doesn't seem to make sense there. Thanks in advance.

 

[nate808]

Thank you for your question regarding the solubility of cations derived from alkali metals and alkaline Earth metals. I can offer some insights into this topic.

Firstly, let's consider the periodic trends in the formation of solutions. As you mentioned, the solvation energy, or the energy released when ions are surrounded by water molecules, is an important factor in determining solubility. Alkali metals, such as lithium, sodium, and potassium, have low ionization energies and relatively small ionic radii. This makes it easier for them to lose electrons and form ions, and their small size allows for more efficient hydration by water molecules. Therefore, the solvation energy for these cations is relatively high, making them more soluble in water.

On the other hand, alkaline Earth metals, such as magnesium, calcium, and strontium, have higher ionization energies and larger ionic radii. This means that it takes more energy for them to lose electrons and form ions, and their larger size makes it more difficult for water molecules to surround and stabilize them. Thus, the solvation energy for these cations is lower, resulting in lower solubility in water.

Now, let's consider the effect of solvation with sulfide ions. Sulfide ions are larger and more polarizable than water molecules, which means they can form stronger interactions with cations. This can lead to the formation of complexes between the cations and sulfide ions, which can disrupt the lattice structure and increase solubility. However, this effect is more pronounced for alkaline Earth metals, as their larger size and lower solvation energy make them more susceptible to complex formation with sulfide ions.

In conclusion, the solubility of cations derived from alkali metals and alkaline Earth metals can be attributed to a combination of factors, including ionization energy, ionic radius, and solvation energy. The presence of sulfide ions can also impact solubility, especially for alkaline Earth metals. I hope this explanation has helped to clarify the differences in solubility between these two groups of elements.