5.2: Solubility
- Page ID
- 372208
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)Solubility and its related terminologies
The majority of solutes do not dissolve in water or other solvents in all proportions.
What is solubility
The maximum proportion of the solute that can dissolve in a given amount of the solvent, usually expressed in grams of solute in 100 grams of solvent, is called the solubility of the solute in the solvent.
- Substances that make a solution when mixed in any proportion are called miscible. For example, ethanol and water are miscible.
- Some substances make a solution when mixed in some proportion but not in all proportions; these are called partially miscible. For example, n-butanol can mix in water up to 7.3 g n-butanol/100 ml water.
- Substances that almost do not dissolve in each other are called immiscible. For example, n-Hexane is immiscible in water.
- The solution that has not yet reached its solubility limits and can dissolve more solute added is called an unsaturated solution.
- The solution that has reached its solubility limits and can not dissolve if more solute is added to it is called a saturated solution.
When a solute is added to a saturated solution, it does dissolve, but, at the same time, the dissociated components recombine to form the crystals of the solute, i.e., recrystallize, at the same rate, so that there is no net dissolution, as illustrated in Fig. 5.2.1.
Effect of temperature on solubility
The solubility of most of the solids and liquids increases as the temperature increases.
For example, when a saturated solution of sugar in water is heated, it can dissolve more sugar. Fig 5.2.2 show the solubility vs temperature curves for some compounds.
When a hot saturated solution is cooled, the concentration of solute in the solution becomes above the solubility limits, making a supersaturated solution.
The supersaturated solution is unstable and ultimately crystallizes out the excess solute leaving any impurities in the solution, as illustrated in Fig. 5.2.3. This process is called re-crystallization, which is used to purify the solutes.
The crystallization of excess solute from a supersaturated solution is responsible for some medical problems like gout and kidney stone. Gout is the crystallization of uric acid in the cartilage, tendons, and soft tissues when the concentration of the uric acid in blood plasms exceeds its solubility limit of ~7 mg/100 mL at 37 oC. It causes redness, swelling, and pain in the affected area, as illustrated in Fig. 5.2.4. Kidney stones are solid materials formed in the urinary tubes, as illustrated in Fig. 5.2.5. Kidney stones are the result of the crystallization of excess calcium phosphate, calcium oxalate, or uric acid in the urine.
Opposite to the solids and liquids, the solubility of gases generally decreases with an increase in temperature.
That is why the carbonated water releases dissolved gas when heated, causing pressure increase, which, in turn, causes the bursting of the soda can
Effect of pressure on solubility
The pressure has almost no effect on the solubility of solids and liquids but has a strong effect on the solubility of gases.
The solubility of gases in liquids is directly proportional to the pressure of the gas above the liquid.
An increase in pressure causes a decrease in the gas volume that increases the gas concentration. More frequent collision of the gas molecules with the gas-liquid boundary in a concentrated solution causes an increase in the rate of dissolution of the gas in the liquid, as illustrated in Fig 5.2.6. The opposite happens when the gas pressure decreases. For example, carbon dioxide starts bubbling out when a soda can is open because the gas escapes resulting in a decrease in the gas pressure above the liquid and a decrease in the solubility of the gas in water.
Solubility guidelines for dissolution of ionic compounds in water
If the solubility of a compound is less than 0.01 mol/L, it is considered insoluble.
The solubility of ionic compounds in water depends on the nature of the compound. For example, lead(II)iodide (PbI2) and silver chloride (AgCl) are insoluble in water because the solubility of PbI2 is 0.0016 mol/L of the solution and the solubility of AgCl is about 1.3 x 10-5 mol/L of solution. Potassium iodide (KI) and Pb(NO3)2 are soluble in water. When aqueous solutions of KI and Pb(NO3)2 are mixed, the concentration of PbI2 in the mixture goes above its solubility limits, and it precipitates out, as illustrated in Fig. 5.2.7.
There are no fail-proof guidelines for predicting the solubility of ionic compounds in water. However, the following guideline can predict the solubility of most ionic compounds.
- Salts of alkali metals (Li1+, Na1+, K1+, Rb1+, Cs1+ ) and ammonia (NH41+) are soluble. No exceptions.
- Salts of nitrate (NO31-), acetate (CH3COO1-), and perchlorate (ClO41-) are soluble. No exceptions.
- Salts of chloride (Cl1-), bromide (Br1-), and Iodide (l1-) are soluble, except when the cation is Pb2+, Hg22-, or Ag1+. (Remember the acronym “LMS” based on the first letter of the element name, or phrase ‘Let Me See” to recall Lead, Mercury, and Silver.)
- Sulfates (SO42-) are soluble except when the cation Pb2+, Hg22-, Ag1+, or a heavy alkaline earth metal ions: calcium (Ca2+), barium (Ba2+), or strontium (Sr2+). ((recall “Let Me See” for Lead, Mercury, and Silver. Remember the acronym “CBS” based on the first letter of the element name, or the phrase “Come By Soon” to recall calcium, barium, and strontium.)
- Hydroxide (OH1-) and sulfides (S2-) are insoluble except when the cation is an alkali metal, ammonia, or a heavy alkaline earth metal ions: Ca2+, Ba2+, and Sr2+. (Recall the phrase “Come By Soon” to recall calcium, barium, and strontium.)
- Carbonates (CO32-), phosphates (PO43-), and oxide (O2-) are insoluble except when the cation is an alkali metal, ammonia
- If there is a conflict between the two guidelines, then the guideline listed first has priority. For example, the salts of insoluble ions become soluble when the cation is an alkali metal, ammonia (rule#1).
Fig. 5.2.8 shows precipitates of some insoluble ionic compounds formed by mixing aqueous solutions of appropriate soluble ionic compounds. The precipitation can be predicted, as illustrated in Fig. 5.2.9. List the ions of the soluble ionic compounds and then cross-combine the cations of one with the anion of the other to make the potential products. If any of the potential products is an insoluble ionic compound, it precipitates out.
