I did a project for lab chemistry where my aim was to see which substance, from a selection of 4 chlorides, recrystallises from a solution to form the most crystals. The 4 chlorides were KCl, NaCl, MgCl2 and CaCl2.
KCl and NaCl didn't form any crystals, they formed a saturated solution whereas MgCl2 and CaCl2 formed crystals. It seems as if the charges on the positve ion of the chloride are important here; but why? How do I explain my results? Conclusion??
If you had removed all the moisture, potassium and sodium chloride would have formed crystals. The solubility of those is much higher than magnesium and calcium chloride, so as moisture is removed, the magnesium and calcium start to form crystals first. Now as to why the solubility is different for alkali and alkaline earth chlorides, books have been written on that. It was the forte of inorganic chem texts in the 20's and 30's, as well as some in depth study of the rare earth series.
I remember my first chem text as a youngster that I read had hundreds of pages on characteristics of the elements, and relating that to periodicity. I wish you would have had aluminum chloride in your experiment.
Here is my suggestion: look at a table or chart of solubilities, look at trends on the chlorides by period of the metal ion.
To explain your results, you need to consider the concept of solubility and the solubility rules for salts. The solubility of a substance refers to its ability to dissolve in a solvent, such as water. This is influenced by several factors, including the nature of the solute and solvent, temperature, and pressure.
In your experiment, you observed that KCl and NaCl did not form any crystals but instead formed saturated solutions. This means that the solute (KCl and NaCl) dissolved completely in the solvent (water) up to a point where no more solute could be dissolved at that particular temperature. This indicates that KCl and NaCl have high solubility in water.
On the other hand, MgCl2 and CaCl2 formed crystals, indicating that they have limited solubility in water. The difference in behavior can be explained by the charges on the positive ions (cations) in each chloride compound.
KCl and NaCl are both composed of monovalent cations (K+ and Na+) paired with monovalent anions (Cl-). These cations have a charge of +1, meaning they have a weaker attraction to the surrounding water molecules due to their smaller charge. As a result, the cations can be easily hydrated by water molecules, leading to strong ion-dipole interactions and high solubility.
MgCl2 and CaCl2, on the other hand, are composed of divalent cations (Mg2+ and Ca2+) paired with monovalent anions (Cl-). These cations have a charge of +2, meaning they have a stronger attraction to the surrounding water molecules due to their larger charge. The stronger ion-dipole interactions between the cations and water molecules make it harder for the cations to be hydrated, leading to lower solubility.
Based on these explanations, you can conclude that the charges on the positive ions of the chlorides are indeed important in determining their solubility in water. Monovalent cations have higher solubility due to their weaker attraction to water molecules, while divalent cations have lower solubility due to their stronger attraction.