Observing phenomena affecting solubility: Common Ion

_{(Licence: Public Domain, Author: PublicDomainPictures) Source: Pixabay}

Observing phenomena affecting solubility: Determination of solubility in the presence of a common ion.

Do you know what the common ion effect is?

The chemistry contemplates very interesting topics, daily you can observe the solubility of different compounds, such as common salt (sodium chloride), sugar (sucrose) or any other similar compound. However, not all compounds have high solubility, so some are called insoluble compounds. Solubility is a very relevant concept for chemistry, with very diverse applications; from the use in the industry of many factories, to marketing campaigns for the advertising of a new product, for example, a detergent. In this post, the influence of a common ion in solution, calcium hydroxide (Ca(OH)2) and calcium chloride (CaCl2), will be presented in two experiments; they are compounds containing calcium (Ca), therefore, both being in one solution, they change the solubility balance of the little soluble compound, in this case calcium hydroxide. This phenomenon not only occurs with these compounds, but also in nature quite frequently.

Theoretical Basis of the Common Ion Effect

Some compounds such as NaCl have a relatively high solubility, other compounds have a low solubility, the part that dissociates in water may even be negligible, so these compounds are referred to as insoluble or slightly soluble.

These compounds have an associated constant of ionic products, or "Ksp", which is defined as the product of the molar concentrations (in equilibrium) of the constituent ions, each of which is elevated to the power of the stoichiometric coefficient in the balanced equilibrium equation. Given a general dissolution, it follows that the Ksp will be:

_{General expression of the Kps - Own authorship}

In this way, the solubilized part of a slightly soluble compound is calculated, taking into account its constant Ksp (experimentally determined for a given solvent, temperature and pressure). An example might be calcium hydroxide:

_{Example of Ksp, for calcium hydroxide - Own authorship}

Where the solubilised part is represented by the letter 'S' together with the corresponding stoichiometric number, its Ksp:

_{Mathematical expression to calculate solubility from the Ksp of a compound - Own authorship}

One effect that causes a decrease in the solubility of a compound is the common ion. When a compound is added to provide the solution with an ion in common with another compound that is not very soluble, the equilibrium is shifted to the left by the principle of Le Châtelier, increasing the part of the compound that remains undissociated and thus the solubility decreases. The Ksp value of the species remains unchanged.

_{Chemical equation that exposes the effect of the common ion - Own authorship}

In this case the common ion is Ca+2, which decreases the solubility of the insoluble compound.

_{Calcium Hydroxide - (Licence: Public Domain, Author:Walkerma ) Source: Wikimedia Commons}

In order to verify the above, an experimental procedure will be presented below, which will compare qualitatively the presence and absence of a common ion and its effect on the solubility of a slightly soluble compound:

Solubility of calcium hydroxide in water in the presence of CaCl2 (common ion)

In a 250ml beaker a solution containing approximately 1.5 grams of calcium chloride (CaCl2) and one gram of calcium hydroxide (Ca(OH)2) was prepared, along with 150 ml of distilled water. It was gently mixed while each compound was added slowly and in the order indicated above to allow the maximum amount of solute to dissolve, and allowed to stand so that the non-differentiated solid could settle.
The temperature was taken and the solution was filtered by the simple bending method, always covering the solution to avoid possible carbonation.
Then with a standardized solution of HCl at 0.0488 M the system was prepared to titrate the previously filtered solution, using bromocresol green as an indicator (which has a yellow to blue shift in the range 4.04-5.6 on the pH scale).

Following the above procedure, the following experimental results were tabulated:

_{Experimental results obtained - Own authorship}

Knowing that the solubility of calcium hydroxide at 25° C is 0.167 g in 100 ml of water, as the aliquots are 25ml, then the solubilized part corresponds to 0.0418 g. (value taken from CRC HandBook of Chemistry and Physics, 86th Edition, CRC Press 2005)

According to the titration, the titrated moles of oxydryl ions correspond to 1,030(6)x10^(-3) mole on average.

According to the dissociation reaction of the compound, the solubilized part of the oxydrile ion is double stoichiometrically to calcium hydroxide, so the moles corresponding to calcium hydroxide are:

_{Calculation of the solubilized moles - Own authorship}

Being the grams of calcium hydroxide solubilized:

_{Calculation of the solubilized grams - Own authorship}

This quantity corresponds to the solubility in 25ml of water, similarly for 100ml the solubility will be: 0.153(4) g/100ml. This quantity represents 91.39% of the theoretical quantity for 25°C. It is important to note that the temperature at the time of the practice was approximately 24°C, therefore, these temperature variations (since it could later increase) could systematically affect the measurements performed.

Analogous to the above, it can be calculated, according to the data of table 2, that the solubility value in this experience is 0.122g/100ml.

This quantity represents 73.21% of the theoretical solubility in 100ml of water, at 25°C and with the presence of a common ion. The decrease in the solubility of the compound after the addition of Ca+2 as a common ion is evident.

Conclusion

Based on the results obtained, it was possible to quantitatively calculate the effect of the common ion on the solubility of Ca(OH)2 which decreased by approximately 20%.

As the solubility is relative to the temperature, this could be a source of systematic errors since it was not constant, and although it was not 25°C, it could vary in the course of practice, thus influencing the solubility of the compound is a question.

In synthesis, the titration technique made it possible to calculate the concentrations of the oxidryl ions, and then to calculate the solubilized grams of the compound, using stoichiometry and its molecular weight. Even with some associated systematic errors it was possible and satisfactory to observe how much a common ion affects the solubility of a compound.

I hope that this brief theoretical explanation allows us to understand that the solubility of a compound depends on many factors, including the solvent, but also on the other components found in it. This is one of my favorite effects, and I think it manifests itself many times in experimental chemistry as well as in nature. Thank you for reading.