How To Determine Solubility 14 Steps With Pictures

Ions with negative charges are called anions, while ions with positive charges are cations. Normally, the number of electrons in an atom is equal to the number of protons, canceling out the electrical charges.

Some ionic compounds aren’t stuck together very well; these are soluble since the water will pull them apart and dissolve them. Other compounds are bonded more strongly, and are insoluble since they can stick together despite the water molecules. Some compounds have internal bonds that are similar in strength to the pull of the water. These are called slightly soluble, since a significant amount of compounds will be pulled apart, but the rest will stay together.

For example, to check Strontium Chloride (SrCl2), look for Sr or Cl in the bold steps below. Cl is “usually soluble,” so check underneath it for exceptions. Sr is not listed as an exception, so SrCl2 must be soluble. The most common exceptions to each rule are written beneath it. There are other exceptions, but you are unlikely to encounter them in a typical chemistry class or laboratory.

Exception: Li3PO4 is insoluble.

Exceptions: Ag(OAc) (silver acetate) and Hg(OAc)2 (mercury acetate) are insoluble. AgNO2- and KClO4- are only “slightly soluble. "

Exception: If any of these pair with the ions silver Ag+, mercury Hg22+, or lead Pb2+, the result is insoluble. The same is true of less common compounds made from pairing with copper Cu+ and thallium Tl+.

Exceptions: The sulfate ion forms insoluble compounds with the following ions: strontium Sr2+, barium Ba2+, lead Pb2+, silver Ag+, calcium Ca2+, radium Ra2+, and diatomic silver Ag22+. Note that silver sulfate and calcium sulfate dissolve just enough that some people call them slightly soluble.

Exceptions: Remember the alkali metals (Group I-A) and how they love forming soluble compounds? Li+, Na+, K+, Rb+, and Cs+ all form soluble compounds with the hydroxide or sulfide ions. In addition, hydroxide forms soluble salts with the alkali earth (Group II-A) ions: calcium Ca2+, strontium Sr2+, and barium Ba2+. Note that the compounds resulting from hydroxide and an alkali earth do have just enough molecules that stay bonded to sometimes be considered “slightly soluble. "

Exceptions: These ions form soluble compounds with the usual suspects, the alkali metals Li+, Na+, K+, Rb+, and Cs+, as well as with ammonium NH4+.

For example, if you’re dissolving lead iodide, or PbI2, write down its product solubility constant.

For example, a molecule of PbI2 splits into the ions Pb2+, I-, and a second I-. (You only need to know or look up the charge on 1 ion, since you know the total compound will always have a neutral charge. ) Write the equation 7. 1×10–9 = [Pb2+][I-]2 The equation is the product solubility constant, which can be found for the 2 ions in a solubility chart. Since there are 2 I- ions, I- is to the second power.

In our example, we need to rewrite 7. 1×10–9 = [Pb2+][I-]2 Since there is 1 lead ion (Pb2+) in the compound, the number of dissolved compound molecules will be equal to the number of free lead ions. So we can set [Pb2+] to x. Since there are 2 iodine ions (I-) for each lead ion, we can set the number of iodine atoms equal to 2x squared. The equation is now 7. 1×10–9 = (x)(2x)2

For example, if our lead iodide compound was being dissolved in a solution with 0. 2 M lead chloride (PbCl2), we would rewrite our equation as 7. 1×10–9 = (0. 2M+x)(2x)2. Then, since 0. 2M is such a higher concentration than x, we can safely rewrite it as 7. 1×10–9 = (0. 2M)(2x)2.

The following is for solubility in pure water, not with any common ions. 7. 1×10–9 = (x)(2x)2 7. 1×10–9 = (x)(4x2) 7. 1×10–9 = 4x3 (7. 1×10–9) ÷ 4 = x3 x = ∛((7. 1×10–9) ÷ 4) x = 1. 2 x 10-3 moles per liter will dissolve. This is a very small amount, so you know this compound is essentially insoluble.