Chapter 13: Solutions

A solution is a system which contains two or more substances homogeneously (a single phase) dissolved in one another. When a solute dissolves in a solvent the resulting homogeneous mixture is called a solution.

Term

Definition

solvent

The solvent is the component whose phase is retained when the solution forms; if all components are the same phase, the one in the greatest amount is the solvent. If one of the components is water, water is always the solvent.

solute

The solute is the component(s) in the smallest amount.

dissolution

The process of a solute dissolving in a solvent is called dissolution (dissolving).

concentration

The concentration of the solution is the amount of solute in a given amount of solvent.

unsaturated

Is a solution which can dissolve more solute.

saturated

Is a solution with the dissolved solute in equilibrium with the undissolved solute.

supersaturated

Is a solution holding more dissolved solute than would be in equilibrium with the undissolved solute

The solvent is the component whose phase is retained when the solution forms; if all components are the same phase, the one in the greatest amount is the solvent. Other components are called solutes. The process of a solute dissolving in a solvent is called dissolution. We must specify the temperature of dissolution because temperature effects the extent of dissolution. Dissolution also depends on the nature of the solute and solvent.

The concentration of the solution is the amount of solute in a given amount of solvent.

Solutions can be unsaturated (a solution which can dissolve more solute), saturated (a solution with the dissolved solute in equilibrium with the undissolved solute) or supersaturated (a solution holding more dissolved solute than would be in equilibrium with the undissolved solute).

In general when a solution is formed there is no evidence of a chemical reaction, that is, no irreversible chemical reaction occurs between the components.

NaCl(s) --H2O --> Na+(aq) + Cl–(aq)

Na(s) + H2O(l) ---> Na+(aq) + -OH(aq) + H2(g)

When sodium chloride is added to water no dramatic chemical change is observed, while the opposite is true when sodium metal is added to water. If we evaporate off the water in the solution of sodium chloride we get the sodium chloride back again--unchanged. However, evaporating off the water after the addition of the sodium metal leaves a white solid which we know to be sodium hydroxide.

When a solution is formed it has particular properties;

1. It is a homogeneous mixture of two or more substances.

2. It can be colorless, or exhibit color, but is transparent.

3. The solute is uniformily distributed throughout the solvent and will not settle out in time.

4. The solute can be separated from the solvent by physical means.

Next let's see what happens when we add NaOH(s) to water and NH4NO3(s) to water. First we see that both are soluble. Then we note that when NaOH dissolves heat is released, but when NH4NO3 dissolves heat is absorbed. That is when NH4NO3 dissolves it is removing heat from the water which causes the temperature of the water to fall.

NaOH(s) --H2O --> Na+(aq) + OH-(aq) + heat

heat + NH4NO3(s) --H2O --> NH4+(aq) + NO3-(aq)

However, if a compound such as carbon tetrachloride is added to water the two liquids do not mix. Why do these things happen? How do we understand the solution process? How can we predict which components will mix and which will not. Whether heat is liberated or absorbed?

There are many solution which are possible. Some of these are more interesting to us then others in that their properties are reasonably well understood and can be described in terms which are understandable to introductory chemists, such as yourselves.

Solute

Solvent

gas

gas

gas

liquid

liquid

liquid

solid

liquid

gas

solid

liquid

solid

solid

solid

Nonreactive gases can be mixed in all proportions to yield a gaseous solution. Two components that are dissolve in one another in any proportion are said to be miscible. Solid solutions, alloys, are well known to us and are an important industry. Amalgams (solids in liquids) are also interesting as are the gases dissolving in solids (energy industry). However, these are not as well understood as liquid solutions. Solutions in which the solvent is liquid are the most interesting to us, particularily when the solvent is water.

When a solution is formed we need to be interested in several different things; first, what happens to the solute in the solvent? Do any changes occur and if so how do visualize those changes?

You must be able to write a chemical equation to represent the change. And you must be able to illustrate the change by drawing pictures of the interactions that are occurring between solute and solvent particles.

In general two types of changes will occur when a solute dissolves in a solvent. The substance will remain in its molecular form or it will ionize (dissociate). In our discussion we will cover each case.

Liquids dissolving in liquids

In liquid molecular solutions, when both the solute and solvent are covalent compounds, the intermolecular attractive forces are London dispersion, dipole-dipole and hydrogen bonding.

Lets try mixing some molecular substances and see what happens. When we mix carbon tetrachloride with benzene a solution is formed whereas when carbon tetrachloride is added to water the two do not mix. CCl4 and C6H6 have identical intermolecular attractive forces--London dispersion type. However water has hydrogen–bonding type of intermolecular attractive forces. The CCl4 molecules are unable to displace the hydrogen-bonded water molecules from one another because the interaction between two water molecules is stronger than the interaction between a water molecule and a carbon tetrachloride molecule.

When we mix ethanol with water the two liquids form a homogeneous solution because the intermolecular attractive forces are identical. In fact water and ethanol are miscible, that is, they will form a solution in any proportion. The solubility of molecular solutions depends on the similarity of intermolecular attractive forces--like dissolves like. When ethanol dissolves in water we can write a chemical equation which expresses the solution process. It is;

C2H5OH(l) --H2O --> C2H5OH(aq)

Because ethanol is a covalent compound and it does not dissociate in water we write the product as an aqueous species. The solution consists of molecules of ethanol and water in a mixture. It is the intermolecular attractive forces which are important in understanding the solution process. We must be able to separate the solution process into its component parts and evaluate what is happening.

The following interaction must be considered.

1) solvent-solvent intermolecular attractions

2) solute-solute intermolecular attractions

3) solute-solvent intermolecular attractions

When solute dissolve in the solvent, the particles of the solute must distribute themselves throughout the solvent. That is solute particles must occupy positions normally taken by solvent particles. Because molecules (particles) are packed close together in a liquid solvent the ease with which a solute particle displaces a solvent molecule depends on the relative forces of attraction of the solvent molcules for each other the solute particles for each other and the strength of the solute–solvent particles.

Recall that I recommended a simple experiment to demonstrate the energy changes associated with the solution process. We can better understand the energy changes by careful consideration of these three steps. If we begin with the separated solute and solvent, the first two steps, expanding the solute particles and solvent particles are both endothemic processes (See Figure above). If the energy released from the solute–solvent interactions is greater than the energy required to expand the solute ands solvent particles energy is released in the solution process, i.e. the solution warms up. If, on the other hand the energy released is less than the energy absorbed to expand the solute and solvent particles the solution process is endothermic, i.e. the solution cools off.

However, if the energy liberated from the solute-solvent interactions is too small compared to the energy required to separate the solute particles and the energy required to separate the solvent particles no solution results. If the calculated heat of solution is exothermic we can expect the homogeneous solution to be formed. However, if the calculated heat of solution is endothermic we can not know for sure whether a homogeneous solution will form. We must consider one other factor in the endothermic case. The other factor is related to the natural tendency towards disorder when mixing two pure substances. This natural tendency towards disorder must be considered when discussing the solution process. In every case this factor favors the formation of the solution. However, if the energy required for the solution process to occur is large, it is unlikely the solution will be formed.

Solids dissolving in liquids

When we try to dissolve a solid in a liquid the attractive forces are at a maximum in the solid. In order for the solid to dissolve in the liquid the solvent–solvent forces of attraction must be sufficient to overcome the attractive forces that hold the solid together. In molecular crystals the attractive forces are weak being of the London dispersion, dipole–dipole or hydrogen–bonding type. The solubility of molecular solids in molecular solvents is again governed by the like dissolves like principle.

Iodine dissolves in the carbon tetrachloride but not the water. The intermolecular attractive forces between I2 molecules are London dispersion type as are the intermolecular attractive forces between CCl4 molecules. However, when we add the I2 to H2O the nonpolar iodine molecules have a hard time separating the hydrogen-bonded water molecules. The water molecules do not interact as well with the I2 molecules as they do with themselves.

When we add glucose, C6H12O6, to water it dissolves because of the 'like' intermolecular attractive forces (hydrogen-bonding). We can write a chemical equation to describe the solution process for molecular solids;

C6H12O6(s) --H2O --> C6H12O6(aq)

The solution consists of glucose molecules distributed amoung the water molecules. It is important to note that the glucose does not ionize but remains as a molecular species.

When we try ionic solids we find that some ionic solids are soluble in water, and some are insoluble in water. Ionic solids are insoluble in nonpolar solvents. Ionic solids are held together by particularily strong electrostatic forces of attraction between the ions, so that only the most polar solvents are able to dissolve them. We already know how to write an equation to describe the dissolving of an ionic solid in water. We do so using the dissociation reaction. We also know how to use the solubility table to predict which ionic solids are soluble and which ae insoluble. Now we will gain some understanding of what is happening at the atomic level in the solution. When ionic solids dissolve in water the ions that are adjacent to each other in the solid become surrounded by the water molecules (hydrated). The attraction force that occurs between the ion and water is called an ion–dipole forces. The polar water molecules orient themselves so that the partially charged ends of the molecule are opposite the charge of the ions. So water molecules are oriented with their hydrogen atoms pointed at the anion and the oxygen atoms pointed at the cation. This process is called hydration. Hydration is more favored for small ions as compared to large ions. If this formation of the hydrated ions were the only factor than we would expect all ionic compounds to dissolve in water. However, that is not the case, and the problem is that we need to condsider the other factors in the solution process.

Recall that the solution process is governed by the solute-solute, solvent-solvent and solute-solvent intermolecular attractive forces. So far we have only considered hydration of the ions by the water molecules, that is the solute–solvent interactions. When we consider the solute interactions we begin to see some of the problems that can arise. The ions in a crystal are strongly attracted to each other and to dissolve it is necessary to overcome the electrostatic attraction between the oppositely charged ions. The lattice energy of a solid is a measure of the strength of those electrostatic attractions. The lattice energy works to keep the ions in the solid state and ionic compounds with large lattice energies are insoluble in water while compounds with small lattice energies are soluble. (1st Lecture) When potassium iodide dissolves in water the ions is the potassium iodide solid must be separated;

KI(s) -–H2O --> K+(g) + I-(g) ∆H1 = +632 kJ/mol

Now the gaseous ions are distributed in water according to the equation;

H2O(l) + K+(g) + I-(g) -–-> K+(aq) + I–(aq) ∆H2 + ∆H3 = -617 kJ

The sum of these two equations yields the overall solution process for KI dissolving in water. The value of ∆H2 + ∆H3 is call the hydration energy. In this example the hydration energy is not as large, in absolute terms as the energy required to separate the ions in the solute and the heat of solution is endothermic.

As you would expect the magnitude of ∆H1 depends on the magnitude of the charge on the ions. The greater the magnitude of the charge the greater the energy. We would than expect than singly charger ions (+1 and -1) would be more soluble than those ions with higher charges (+2, –2, +3 or -3). But this energy also depends on the distance between the ions. If the cation and anion are small they can get very close to one another and it is more difficult to separate the two charges. A large cation and a small anion are easier to separate than a small cation and a small anion.

So in the series of hydroxides Mg(OH)2, Ca(OH)2, Sr(OH)2 and Ba(OH)2 we find that Mg(OH)2 is less soluble than Ba(OH)2. It should be clear that to completely understand the solubility of ionic compounds requires a careful evaluation of all the interparticle interactions that are occuring. These include the energy ∆H1 and the hydration energy and the entropy.

Recall the two primary energy factors which effect how much of a solute will dissolve. (How soluble a substance is) These are;

1) Change in enthalpy (exothermic or endothermic) in the solution process

2) Change in entropy (measure of disorder)

When a solute is added to a solvent the formation of a solution is favored by decrease in the enthalpy (exothermic) and an increase in entropy (more dissorder). It turns out that for solutions the entropy is always favorable. Preparing a mixture always produces a more disordered system.

So it become the relative magnitude of the lattice energy (an endothermic process) and the hydration energy (an exothermic process) which govern the behavior of the solid/liquid solutions. So the overall solution process will be endothermic if the lattice energy is greater than the hydration energy and it will be exothermic if the hydration energy is greater than the lattice energy. The lattice energy of the ionic solid and the hydration energy of the ions are two terms which oppose one another in the solution process. For low charged species the lattice energy and the hydration energy are roughly the same, usually the lattice energy is less than the hydration energy and the solution process is slightly endothermic. Compensating the unfavorable enthalpy chnage is the favorable entropy change. So many ionic solids dissolve in water with an absorption of energy. We stated at the beginning that the solution process is favored when the enthalpy goes down (exothermic) and the entropy increases (more disorder). Also all solution processes that we will discuss result in the entropy increasing. So we can then conclude that the enthalpy is the determining factor) and if the solution process is exothermic (favors solubility) the solute will dissolve in the solvent and if the process is endothermic the solute will not dissolve in the solvent.

Gases in liquids

For gases the solution process is exothermic. We can understand this because there are no solute–solute interactions that must be broken in the solution process. Because dissolving gases is exothermic increasing the temperature will decrease the solubility of the gas. This is noticable as you heat water, bubbles of air (O2 and N2) form in the water. The solubility of O2 in fresh water is particularly important to aquatic life, thermal pollution which is an increase of the average temperature of an area of water decreases the the O2 dissolved in the water which makes it difficult for aquatic life to breath.

The solubility of gases also depend on the external pressure. Increasing the external pressure increases the solubility of a gas in a liquid. This principle is used in bottling carbonated beverages, such as beer, soft drinks, and some wines. These beverages are bottled under pressures of about 4 atm of CO2. When the bottle is opened the pressure above the liquid drops to 1 atm and carbon dioxide will bubble out of solution.

The principle is also used to understand the 'bends' a common problem in Sea Hunt. A lower depths more of the air breathed dissolves in the blood. As a diver rises to the surface this dissolved air begins to boil out of solution impairing blood circulation and nerve impluses. To prevent this problem divers must rise slowly or as many are doing use a helium–oxygen mixture, as helium is very insoluble in the blood. When an ionic substance or highly polar compound dissolves in water, water molecules orient themselves around the ions or polar compound, hydrating the species. (Show transparency of solution process.) In the solution process the hydration that occurs around the ion or polar compound involves ion-dipole or dipole-dipole bond fomation. While these interactions are weak, they are strong enough that the solution process is generally accomplished with a gain (exothermic) or loss (endothermic) of energy. The details of the parameters which govern the energy of the solution process involve a knowledge of the solute-solute attractive forces, the solvent-solvent attractive forces and the solute-solvent attractive forces. When a gas such as O2 dissolves in water the molecules of oxygen remain in the molecular form. However, when other gases, such as hydrogen chloride are added to water something else will happen. HCl completely dissociates into H+ and Cl- ion in water. The ions are solvated by the water molecules.