what change would increase the amount of solid solute able to be dissolved in liquid water?
13.2: Saturated Solutions and Solubility
- Page ID
- 21785
Learning Objectives
- To understand the relationship between solubility and molecular structure.
- To demonstrate how the strength of intramolecular bonding determines the solubility of a solute in a given solvent.
When a solute dissolves, its private atoms, molecules, or ions interact with the solvent, become solvated, and are able to lengthened independently throughout the solution (Figure \(\PageIndex{1a}\)). This is not, however, a unidirectional process. If the molecule or ion happens to collide with the surface of a particle of the undissolved solute, it may attach to the particle in a process called crystallization. Dissolution and crystallization continue as long as excess solid is present, resulting in a dynamic equilibrium analogous to the equilibrium that maintains the vapor pressure of a liquid. We can represent these opposing processes as follows:
\[ solute + solvent \underset{crystallization}{\stackrel{dissolution}{\longrightleftharpoons}} solution \]
Although the terms precipitation and crystallization are both used to describe the separation of solid solute from a solution, crystallization refers to the formation of a solid with a well-defined crystalline construction, whereas precipitation refers to the formation of whatever solid phase, ofttimes one with very modest particles.
Factors Affecting Solubility
The maximum amount of a solute that can dissolve in a solvent at a specified temperature and pressure is its solubility. Solubility is oft expressed as the mass of solute per volume (thou/50) or mass of solute per mass of solvent (g/g), or equally the moles of solute per book (mol/L). Even for very soluble substances, withal, there is commonly a limit to how much solute tin can dissolve in a given quantity of solvent. In general, the solubility of a substance depends on not only the energetic factors nosotros take discussed just besides the temperature and, for gases, the pressure. At 20°C, for example, 177 g of NaI, 91.ii g of NaBr, 35.ix one thousand of NaCl, and but iv.1 one thousand of NaF dissolve in 100 thou of water. At 70°C, however, the solubilities increase to 295 g of NaI, 119 grand of NaBr, 37.5 m of NaCl, and iv.8 grand of NaF. Equally you learned in Chapter 12, the lattice energies of the sodium halides increase from NaI to NaF. The fact that the solubilities decrease as the lattice energy increases suggests that the \(ΔH_2\) term in Figure 13.1 dominates for this series of compounds.
A solution with the maximum possible amount of solute is saturated. If a solution contains less than the maximum amount of solute, it is unsaturated. When a solution is saturated and backlog solute is present, the charge per unit of dissolution is exactly equal to the rate of crystallization (Effigy \(\PageIndex{1b}\)). Using the value simply stated, a saturated aqueous solution of NaCl, for example, contains 35.nine g of NaCl per 100 mL of water at 20°C. We tin gear up a homogeneous saturated solution past calculation excess solute (in this example, greater than 35.ix 1000 of NaCl) to the solvent (water), stirring until the maximum possible amount of solute has dissolved, and then removing undissolved solute past filtration.
The solubility of most solids increases with increasing temperature.
Because the solubility of almost solids increases with increasing temperature, a saturated solution that was prepared at a higher temperature unremarkably contains more dissolved solute than it would contain at a lower temperature. When the solution is cooled, it tin can therefore become supersaturated (Effigy \(\PageIndex{1c}\)). Similar a supercooled or superheated liquid, a supersaturated solution is unstable. Consequently, adding a small particle of the solute, a seed crystal, will commonly crusade the backlog solute to rapidly precipitate or crystallize, sometimes with spectacular results. The charge per unit of crystallization in Equation \(\ref{13.2.1}\) is greater than the rate of dissolution, so crystals or a precipitate form (Figure \(\PageIndex{1d}\)). In contrast, calculation a seed crystal to a saturated solution reestablishes the dynamic equilibrium, and the cyberspace quantity of dissolved solute no longer changes.
Video \(\PageIndex{1}\): hot water ice (sodium acetate) cute scientific discipline experiment. watered-downward sodium acetate trihydrate. Needle crystal is truly wonderful structures
Because crystallization is the reverse of dissolution, a substance that requires an input of estrus to form a solution (\(ΔH_{soln} > 0\)) releases that heat when it crystallizes from solution (\(ΔH_{crys} < 0\)). The amount of oestrus released is proportional to the amount of solute that exceeds its solubility. Two substances that take a positive enthalpy of solution are sodium thiosulfate (\(Na_2S_2O_3\)) and sodium acetate (\(CH_3CO_2Na\)), both of which are used in commercial hot packs, minor numberless of supersaturated solutions used to warm easily (see Figure thirteen.1.iii).
Interactions in Liquid Solutions
The interactions that determine the solubility of a substance in a liquid depend largely on the chemical nature of the solute (such every bit whether it is ionic or molecular) rather than on its physical country (solid, liquid, or gas). We volition first draw the full general case of forming a solution of a molecular species in a liquid solvent and then describe the formation of a solution of an ionic compound.
Solutions of Molecular Substances in Liquids
The London dispersion forces, dipole–dipole interactions, and hydrogen bonds that agree molecules to other molecules are generally weak. Even so, energy is required to disrupt these interactions. Every bit described in Section 13.i, unless some of that energy is recovered in the formation of new, favorable solute–solvent interactions, the increase in entropy on solution formation is non enough for a solution to form.
For solutions of gases in liquids, we tin can safely ignore the free energy required to dissever the solute molecules (\(ΔH_2 = 0\)) because the molecules are already separated. Thus nosotros demand to consider only the energy required to dissever the solvent molecules (\(ΔH_1\)) and the free energy released past new solute–solvent interactions (\(ΔH_3\)). Nonpolar gases such as \(N_2\), \(O_2\), and \(Ar\) take no dipole moment and cannot appoint in dipole–dipole interactions or hydrogen bonding. Consequently, the but way they can interact with a solvent is by means of London dispersion forces, which may be weaker than the solvent–solvent interactions in a polar solvent. Information technology is not surprising, then, that nonpolar gases are most soluble in nonpolar solvents. In this case, \(ΔH_1\) and \(ΔH_3\) are both small and of similar magnitude. In contrast, for a solution of a nonpolar gas in a polar solvent, \(ΔH_1\) is far greater than \(ΔH_3\). As a event, nonpolar gases are less soluble in polar solvents than in nonpolar solvents. For example, the concentration of \(N_2\) in a saturated solution of \(N_2\) in water, a polar solvent, is only \(vii.07 \times 10^{-4}\; G\) compared with \(4.5 \times x^{-3}\; Yard\) for a saturated solution of \(N_2\) in benzene, a nonpolar solvent.
The solubilities of nonpolar gases in water generally increase every bit the molecular mass of the gas increases, every bit shown in Table \(\PageIndex{ane}\). This is precisely the trend expected: as the gas molecules get larger, the strength of the solvent–solute interactions due to London dispersion forces increases, approaching the strength of the solvent–solvent interactions.
| Gas | Solubility (M) × ten−4 |
|---|---|
| He | 3.90 |
| Ne | 4.65 |
| Ar | 15.2 |
| Kr | 27.nine |
| Xe | 50.2 |
| H2 | 8.06 |
| Northward2 | seven.07 |
| CO | 10.6 |
| O2 | 13.9 |
| Due northtwoO | 281 |
| CH4 | 15.five |
Most all common organic liquids, whether polar or non, are miscible. The strengths of the intermolecular attractions are comparable; thus the enthalpy of solution is expected to be small-scale (\(ΔH_{soln} \approx 0\)), and the increase in entropy drives the formation of a solution. If the predominant intermolecular interactions in two liquids are very different from one another, however, they may be immiscible. For example, organic liquids such every bit benzene, hexane, \(CCl_4\), and \(CS_2\) (S=C=S) are nonpolar and have no ability to act as hydrogen bond donors or acceptors with hydrogen-bonding solvents such as \(H_2O\), \(HF\), and \(NH_3\); hence they are immiscible in these solvents. When shaken with water, they course separate phases or layers separated by an interface (Figure \(\PageIndex{two}\)), the region betwixt the two layers.
Just because two liquids are immiscible, all the same, does not mean that they are completely insoluble in each other. For example, 188 mg of benzene dissolves in 100 mL of water at 23.v°C. Adding more benzene results in the separation of an upper layer consisting of benzene with a modest amount of dissolved h2o (the solubility of water in benzene is only 178 mg/100 mL of benzene). The solubilities of simple alcohols in water are given in Tabular array \(\PageIndex{ii}\).
| Alcohol | Solubility (mol/100 g of \(H_2O\)) |
|---|---|
| methanol | completely miscible |
| ethanol | completely miscible |
| n-propanol | completely miscible |
| n-butanol | 0.11 |
| n-pentanol | 0.030 |
| n-hexanol | 0.0058 |
| due north-heptanol | 0.0008 |
Only the three lightest alcohols (methanol, ethanol, and n-propanol) are completely miscible with water. Every bit the molecular mass of the booze increases, then does the proportion of hydrocarbon in the molecule. Correspondingly, the importance of hydrogen bonding and dipole–dipole interactions in the pure alcohol decreases, while the importance of London dispersion forces increases, which leads to progressively fewer favorable electrostatic interactions with water. Organic liquids such as acetone, ethanol, and tetrahydrofuran are sufficiently polar to be completely miscible with water yet sufficiently nonpolar to exist completely miscible with all organic solvents.
The same principles govern the solubilities of molecular solids in liquids. For example, elemental sulfur is a solid consisting of cyclic \(S_8\) molecules that have no dipole moment. Because the \(S_8\) rings in solid sulfur are held to other rings by London dispersion forces, elemental sulfur is insoluble in water. It is, still, soluble in nonpolar solvents that have comparable London dispersion forces, such equally \(CS_2\) (23 g/100 mL). In contrast, glucose contains five –OH groups that can course hydrogen bonds. Consequently, glucose is very soluble in h2o (91 thousand/120 mL of water) but substantially insoluble in nonpolar solvents such as benzene. The structure of one isomer of glucose is shown here.
Low-molecular-mass hydrocarbons with highly electronegative and polarizable halogen atoms, such as chloroform (\(CHCl_3\)) and methylene chloride (\(CH_2Cl_2\)), have both pregnant dipole moments and relatively stiff London dispersion forces. These hydrocarbons are therefore powerful solvents for a wide range of polar and nonpolar compounds. Naphthalene, which is nonpolar, and phenol (\(C_6H_5OH\)), which is polar, are very soluble in chloroform. In contrast, the solubility of ionic compounds is largely determined not past the polarity of the solvent but rather by its dielectric constant, a measure out of its ability to separate ions in solution, as you will soon see.
Example \(\PageIndex{ane}\)
Identify the most important solute–solvent interactions in each solution.
- iodine in benzene solvent
- aniline (\(\ce{C_6H_5NH_2}\)) in dichloromethane (\(\ce{CH_2Cl_2}\)) solvent
- iodine in water solvent
Given: components of solutions
Asked for: predominant solute–solvent interactions
Strategy:
Place all possible intermolecular interactions for both the solute and the solvent: London dispersion forces, dipole–dipole interactions, or hydrogen bonding. Determine which is likely to exist the most important factor in solution formation.
Solution
- Benzene and \(\ce{I2}\) are both nonpolar molecules. The only possible attractive forces are London dispersion forces.
- Aniline is a polar molecule with a dipole moment of 1.half-dozen D and has an \(\ce{–NH_2}\) grouping that can act as a hydrogen bond donor. Dichloromethane is besides polar with a one.5 D dipole moment, but it has no obvious hydrogen bond acceptor. Therefore, the most important interactions betwixt aniline and \(CH_2Cl_2\) are likely to exist dipole-dipole interactions.
- H2o is a highly polar molecule that engages in extensive hydrogen bonding, whereas \(I_2\) is a nonpolar molecule that cannot deed as a hydrogen bond donor or acceptor. The slight solubility of \(\ce{I_2}\) in water (\(1.3 \times 10^{-3}\; mol/L\) at 25°C) is due to London dispersion forces.
Do \(\PageIndex{1}\)
Identify the most important interactions in each solution:
- ethylene glycol (\(HOCH_2CH_2OH\)) in acetone
- acetonitrile (\(\ce{CH_3C≡N}\)) in acetone
- n-hexane in benzene
- Answer a
-
hydrogen bonding
- Answer b
-
London interactions
- Reply c
-
London dispersion forces
Hydrophilic and Hydrophobic Solutes
A solute can exist classified as hydrophilic (literally, "water loving"), meaning that information technology has an electrostatic allure to water, or hydrophobic ("water fearing"), pregnant that information technology repels water. A hydrophilic substance is polar and ofttimes contains O–H or N–H groups that can form hydrogen bonds to water. For example, glucose with its five O–H groups is hydrophilic. In contrast, a hydrophobic substance may be polar but usually contains C–H bonds that practice not collaborate favorably with water, as is the case with naphthalene and n-octane. Hydrophilic substances tend to be very soluble in h2o and other strongly polar solvents, whereas hydrophobic substances are essentially insoluble in water and soluble in nonpolar solvents such as benzene and cyclohexane.
The difference betwixt hydrophilic and hydrophobic substances has substantial consequences in biological systems. For example, vitamins tin can be classified every bit either fat soluble or h2o soluble. Fatty-soluble vitamins, such equally vitamin A, are generally nonpolar, hydrophobic molecules. As a result, they tend to exist absorbed into fatty tissues and stored at that place. In contrast, water-soluble vitamins, such as vitamin C, are polar, hydrophilic molecules that circulate in the blood and intracellular fluids, which are primarily aqueous. H2o-soluble vitamins are therefore excreted much more rapidly from the body and must exist replenished in our daily diet. A comparing of the chemical structures of vitamin A and vitamin C chop-chop reveals why one is hydrophobic and the other hydrophilic.
Because h2o-soluble vitamins are rapidly excreted, the risk of consuming them in backlog is relatively small. Eating a dozen oranges a twenty-four hour period is likely to make you tired of oranges long before you suffer any ill effects due to their high vitamin C content. In dissimilarity, fat-soluble vitamins constitute a significant wellness hazard when consumed in large amounts. For example, the livers of polar bears and other large animals that live in cold climates contain large amounts of vitamin A, which have occasionally proven fatal to humans who accept eaten them.
Instance \(\PageIndex{2}\)
The following substances are essential components of the human being diet:
Using what you know of hydrophilic and hydrophobic solutes, classify each as water soluble or fat soluble and predict which are probable to be required in the nutrition on a daily basis.
- arginine
- pantothenic acrid
- oleic acrid
Given: chemical structures
Asked for: classification as water soluble or fatty soluble; dietary requirement
Strategy:
Based on the construction of each compound, decide whether information technology is hydrophilic or hydrophobic. If it is hydrophilic, information technology is likely to be required on a daily basis.
Solution:
- Arginine is a highly polar molecule with 2 positively charged groups and one negatively charged group, all of which can form hydrogen bonds with water. As a issue, information technology is hydrophilic and required in our daily diet.
- Although pantothenic acid contains a hydrophobic hydrocarbon portion, information technology also contains several polar functional groups (\(\ce{–OH}\) and \(\ce{–CO_2H}\)) that should interact strongly with h2o. It is therefore likely to be water soluble and required in the diet. (In fact, pantothenic acrid is virtually always a component of multiple-vitamin tablets.)
- Oleic acrid is a hydrophobic molecule with a single polar group at ane cease. It should be fatty soluble and not required daily.
Exercise \(\PageIndex{2}\)
These compounds are consumed by humans: caffeine, acetaminophen, and vitamin D. Identify each as primarily hydrophilic (water soluble) or hydrophobic (fat soluble), and predict whether each is probable to be excreted from the body quickly or slowly.
- Answer
-
Caffeine and acetaminophen are water soluble and rapidly excreted, whereas vitamin D is fat soluble and slowly excreted
Solid Solutions
Solutions are non express to gases and liquids; solid solutions as well be. For example, amalgams, which are usually solids, are solutions of metals in liquid mercury. Because most metals are soluble in mercury, amalgams are used in gold mining, dentistry, and many other applications. A major difficulty when mining gold is separating very minor particles of pure gold from tons of crushed rock. 1 manner to accomplish this is to agitate a break of the crushed stone with liquid mercury, which dissolves the gold (equally well as any metallic silver that might be present). The very dense liquid gold–mercury amalgam is then isolated and the mercury distilled abroad.
An alloy is a solid or liquid solution that consists of i or more than elements in a metallic matrix. A solid blend has a unmarried homogeneous phase in which the crystal structure of the solvent remains unchanged by the presence of the solute. Thus the microstructure of the blend is compatible throughout the sample. Examples are substitutional and interstitial alloys such as brass or solder. Liquid alloys include sodium/potassium and gold/mercury. In contrast, a fractional alloy solution has ii or more phases that tin be homogeneous in the distribution of the components, just the microstructures of the 2 phases are non the same. Every bit a liquid solution of lead and tin is cooled, for case, dissimilar crystalline phases form at dissimilar cooling temperatures. Alloys usually have properties that differ from those of the component elements.
Network solids such equally diamond, graphite, and \(\ce{SiO_2}\) are insoluble in all solvents with which they do not react chemically. The covalent bonds that concur the network or lattice together are simply too stiff to be broken under normal conditions. They are certainly much stronger than any conceivable combination of intermolecular interactions that might occur in solution. Most metals are insoluble in virtually all solvents for the same reason: the delocalized metallic bonding is much stronger than any favorable metal atom–solvent interactions. Many metals react with solutions such as aqueous acids or bases to produce a solution. However, as we saw in Section 13.1, in these instances the metallic undergoes a chemical transformation that cannot be reversed past but removing the solvent.
Solids with very strong intermolecular bonding tend to exist insoluble.
Solubilities of Ionic Substances in Liquids
Previously, you lot were introduced to guidelines for predicting the solubility of ionic compounds in water. Ionic substances are mostly most soluble in polar solvents; the higher the lattice energy, the more polar the solvent must be to overcome the lattice energy and dissolve the substance. Because of its high polarity, water is the most common solvent for ionic compounds. Many ionic compounds are soluble in other polar solvents, nevertheless, such as liquid ammonia, liquid hydrogen fluoride, and methanol. Because all these solvents consist of molecules that have relatively big dipole moments, they tin can interact favorably with the dissolved ions.
The ion–dipole interactions between \(\ce{Li^{+}}\) ions and acetone molecules in a solution of LiCl in acetone are shown in Figure \(\PageIndex{3}\). The energetically favorable \(\ce{Li^{+}}\)–acetone interactions brand \(ΔH_3\) sufficiently negative to overcome the positive \(ΔH_1\) and \(ΔH_2\). Considering the dipole moment of acetone (ii.88 D), and thus its polarity, is really larger than that of water (1.85 D), one might fifty-fifty expect that LiCl would be more soluble in acetone than in h2o. In fact, the reverse is truthful: 83 chiliad of LiCl dissolve in 100 mL of h2o at 20°C, but only about 4.1 grand of \(\ce{LiCl}\) dissolve in 100 mL of acetone. This apparent contradiction arises from the fact that the dipole moment is a property of a single molecule in the gas phase. A more useful mensurate of the ability of a solvent to dissolve ionic compounds is its dielectric constant (ε), which is the ability of a bulk substance to decrease the electrostatic forces between 2 charged particles. By definition, the dielectric constant of a vacuum is 1. In essence, a solvent with a high dielectric abiding causes the charged particles to behave every bit if they have been moved farther apart. At 25°C, the dielectric constant of water is fourscore.1, 1 of the highest known, and that of acetone is only 21.0. Hence h2o is better able to decrease the electrostatic attraction between \(\ce{Li^{+}}\) and \(\ce{Cl^{-}}\) ions, and so \(\ce{LiCl}\) is more soluble in water than in acetone. This behavior is in dissimilarity to that of molecular substances, for which polarity is the dominant cistron governing solubility.
A solvent's dielectric constant is the most useful measure of its ability to deliquesce ionic compounds. A solvent'south polarity is the dominant factor in dissolving molecular substances.
It is besides possible to deliquesce ionic compounds in organic solvents using crown ethers, cyclic compounds with the general formula \((OCH_2CH_2)_n\). Crown ethers are named using both the total number of atoms in the ring and the number of oxygen atoms. Thus 18-crown-six is an 18-membered ring with half dozen oxygen atoms (Figure \(\PageIndex{1a}\)). The cavity in the center of the crown ether molecule is lined with oxygen atoms and is large enough to be occupied by a cation, such as \(Yard^+\). The cation is stabilized by interacting with solitary pairs of electrons on the surrounding oxygen atoms. Thus crown ethers solvate cations inside a hydrophilic crenel, whereas the outer vanquish, consisting of C–H bonds, is hydrophobic. Crown ethers are useful for dissolving ionic substances such as \(KMnO_4\) in organic solvents such as isopropanol \([(CH_3)_2CHOH]\) (Figure \(\PageIndex{5}\)). The availability of crown ethers with cavities of different sizes allows specific cations to exist solvated with a high degree of selectivity.
Cryptands (from the Greek kryptós, pregnant "hidden") are compounds that tin can completely environment a cation with alone pairs of electrons on oxygen and nitrogen atoms (Figure \(\PageIndex{4b}\)). The number in the name of the cryptand is the number of oxygen atoms in each strand of the molecule. Similar crown ethers, cryptands can be used to prepare solutions of ionic compounds in solvents that are otherwise likewise nonpolar to dissolve them.
Summary
The solubility of a substance is the maximum amount of a solute that tin dissolve in a given quantity of solvent; it depends on the chemical nature of both the solute and the solvent and on the temperature and pressure. When a solution contains the maximum amount of solute that tin deliquesce under a given set of conditions, it is a saturated solution. Otherwise, it is unsaturated. Supersaturated solutions, which contain more than dissolved solute than allowed under particular atmospheric condition, are not stable; the add-on of a seed crystal, a small particle of solute, will usually cause the excess solute to crystallize. A organisation in which crystallization and dissolution occur at the same rate is in dynamic equilibrium. The solubility of a substance in a liquid is determined by intermolecular interactions, which also determine whether ii liquids are miscible. Solutes tin can exist classified as hydrophilic (h2o loving) or hydrophobic (water fearing). Vitamins with hydrophilic structures are water soluble, whereas those with hydrophobic structures are fat soluble. Many metals dissolve in liquid mercury to class amalgams. Covalent network solids and near metals are insoluble in nearly all solvents. The solubility of ionic compounds is largely determined by the dielectric abiding (ε) of the solvent, a measure of its ability to decrease the electrostatic forces betwixt charged particles. Solutions of many ionic compounds in organic solvents tin be dissolved using crown ethers, cyclic polyethers large plenty to accommodate a metal ion in the middle, or cryptands, compounds that completely surround a cation.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map:_Chemistry_-_The_Central_Science_%28Brown_et_al.%29/13:_Properties_of_Solutions/13.2:_Saturated_Solutions_and_Solubility
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