Just because two liquids are immiscible, however, does not mean that they are completely insoluble in each other. For example, mg of benzene dissolves in mL of water at Because water is less dense than the perfluoroheptane, the water layer floats on top.
The goldfish is swimming in the water layer. Figure from the Wikipedia.. The solubilities of simple alcohols in water are given in Table 9. Only the three lightest alcohols methanol, ethanol, and n -propanol are completely miscible with water.
As the molecular mass of the alcohol increases, so 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 be 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.
In contrast, glucose contains five —OH groups that can form hydrogen bonds. 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 2 Cl 2 , have both significant dipole moments and relatively strong 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 6 H 5 OH , which is polar, are very soluble in chloroform. In contrast, the solubility of ionic compounds is largely determined not by the polarity of the solvent but rather by its dielectric constant , a measure of its ability to separate ions in solution, as you will soon see.
Identify the most important solute—solvent interactions in each solution. Given: components of solutions. Asked for: predominant solute—solvent interactions. Identify 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 be the most important factor in solution formation. Identify the most important interactions in each solution:. A solute can be classified as hydrophilic A substance attracted to water.
Hydrophilic substances are polar and can form hydrogen bond s to water. Hydrophobic substances do not interact favorably with water. A hydrophilic substance is polar and often 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 do not interact favorably with water, as is the case with naphthalene and n -octane.
Hydrophilic substances tend to be very soluble in water 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 between hydrophilic and hydrophobic substances has substantial consequences in biological systems.
For example, vitamins can be classified as either fat soluble or water soluble. Fat-soluble vitamins, such as vitamin A, are mostly nonpolar, hydrophobic molecules.
As a result, they tend to be absorbed into fatty tissues and stored there. 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.
Water-soluble vitamins are therefore excreted much more rapidly from the body and must be replenished in our daily diet. A comparison of the chemical structures of vitamin A and vitamin C quickly reveals why one is hydrophobic and the other hydrophilic. Because water-soluble vitamins are rapidly excreted, the risk of consuming them in excess is relatively small. Eating a dozen oranges a day is likely to make you tired of oranges long before you suffer any ill effects due to their high vitamin C content.
In contrast, fat-soluble vitamins constitute a significant health 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 have eaten them.
The following substances are essential components of the human diet:. Using what you know of hydrophilic and hydrophobic solutes, classify each as water soluble or fat soluble and predict which are likely to be required in the diet on a daily basis.
Given: chemical structures. Asked for: classification as water soluble or fat soluble; dietary requirement. Based on the structure of each compound, decide whether it is hydrophilic or hydrophobic. If it is hydrophilic, it is likely to be required on a daily basis. 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 likely to be excreted from the body rapidly or slowly.
Answer: Caffeine and acetaminophen are water soluble and rapidly excreted, whereas vitamin D is fat soluble and slowly excreted. Solutions are not limited to gases and liquids; solid solutions also exist. For example, amalgams 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 small particles of pure gold from tons of crushed rock.
One way to accomplish this is to agitate a suspension of the crushed rock with liquid mercury, which dissolves the gold as well as any metallic silver that might be present. When we do place solutes and solvents together, there is what we call the solution process.
You can think of it as being similar to what you would experience if you tried to squeeze into an already packed elevator. Everyone has to adjust to "find their space" again. Now just like in the elevator, molecules will adjust differently dependent on the type of molecule making an entrance. And also like in an elevator there will come a point when no more people can be added. For a solution, this point is called the saturation point and the solution itself is called a saturated solution.
At the point of saturation, no more solute will dissolve in the solvent. Rather the process of dissolving and precipitation are both occurring simultaneously and at the same rate. Generally speaking only certain molecules will dissolve in water to begin with. The old phrase "like dissolves like" or "birds of a feather flock together" is very true with respect to what degree solutes are soluble or miscible in different solvents. At very low concentrations, almost all molecules are somewhat soluble in all solvents.
But by trend, ionic and polar solutes are more soluble in polar solvents and non-polar molecules are soluble in non-polar mostly organic solvents. The units of concentration we just discussed are used to describe the degree to which a solute is soluble in a solvent. When you place a non-polar molecule in a polar solvent like oil in water the molecules try to minimize surface contact between them. This is actually the basis for the cells in our bodies.
The lipids oily fatty acids form our cell membranes so that their non-polar tails face inward away from the polar cytoplasm and the polar heads face towards the polar cytoplasm. Although much of the explanation for why certain substances mix and form solutions and why others do not is beyond the scope of this class, we can get a glimpse at why solutions form by taking a look at the process by which ethanol, C 2 H 5 OH, dissolves in water.
Ethanol is actually miscible in water, which means that the two liquids can be mixed in any proportion without any limit to their solubility. Much of what we now know about the tendency of particles to become more dispersed can be used to understand this kind of change as well. Picture a layer of ethanol being carefully added to the top of some water Figure below. Because the particles of a liquid are moving constantly, some of the ethanol particles at the boundary between the two liquids will immediately move into the water, and some of the water molecules will move into the ethanol.
In this process, water-water and ethanol-ethanol attractions are broken and ethanol-water attractions are formed. The attractions that form between the ethanol and water molecules are also hydrogen bonds Figure below.
Because the attractions between the particles are so similar, the freedom of movement of the ethanol molecules in the water solution is about the same as their freedom of movement in the pure ethanol. The polar vitamins, as well as the polar water molecules, have strong intermolecular forces that must be overcome in order for a solution to be formed, requiring energy.
When these polar molecules interact with each other i. Hence, the overall enthalpy change energetics is small. The small enthalpy change, coupled with a significant increase in randomness entropy change when the solution is formed, allow this solution to form spontaneously. Nonpolar vitamins and nonpolar solvents both have weak intermolecular interactions, so the overall enthalpy change energetics is again small. Hence, in the case of nonpolar vitamins dissolving in nonpolar lipid solvents, the small enthalpy change, coupled with a significant increase in randomness entropy change when the solution is formed, allow this solution to form spontaneously as well.
For a nonpolar vitamin to dissolve in water, or for a polar vitamin to dissolve in fat, the energy required to overcome the initial intermolecular forces i. Hence, in these cases, the enthalpy change energetics is unfavorable to dissolution, and the magnitude of this unfavorable enthalpy change is too large to be offset by the increase in randomness of the solution.
Therefore, these solutions will not form spontaneously. There are exceptions to the principle "like dissolves like," e. In general, it is possible to predict whether a vitamin is fat-soluble or water-soluble by examining its structure to determine whether polar groups or nonpolar groups predominate. In the structure of calciferol Vitamin D 2 , shown in Figure 3 below, we find an —OH group attached to a bulky arrangement of hydrocarbon rings and chains.
As a result, we might expect carbon tetrachloride to be very soluble in water. However, water molecules form strong hydrogen bonds with one another, causing them to stick tightly to one another. Since the water molecules have very strong intermolecular forces with each other and interact only weakly with carbon tetrachloride via London dispersion forces? Let's imagine what happens when a polar solute such as sodium chloride is placed in a nonpolar solvent such as carbon tetrachloride.
Because CCl 4 doesn't have a partial charge, it won't attach itself to the sodium or chloride ions. As we've mentioned before, the sodium and chloride ions in NaCl are strongly attracted to one another because of their opposite charges. This very weak solvent-solute interaction, as well as the very strong attraction between neighboring solute particles, causes sodium chloride to be insoluble in carbon tetrachloride.
If we place a nonpolar solid into a nonpolar liquid, "like dissolves like" implies that the solid will dissolve. However, the only forces that will cause the liquid to be attracted to the solid are weak London dispersion forces. Why should the solid dissolve? Let's imagine that we have placed a chunk of carbon tetrabromide in a beaker containing carbon tetrachloride. The carbon tetrabromide molecules in the solid are held together by very weak London dispersion forces, as are the carbon tetrachloride molecules in the solvent.
One might expect, then, that there is no particular reason for the solute to dissolve. As it turns out, there's another force involved.
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