Properties of Solutions

7 Properties of Solutions

Vocabulary

Boiling Point:
Ion-Dipole Forces: Attractive forces between ions and polar molecules.
Molarity (M): The number of moles of solute per liter of solution; expressed as M = \frac{\text{moles of solute}}{\text{liters of solution}}.
Parts Per Million (ppm): A ratio expressing the mass of a solute to the total mass of the solution, multiplied by 1,000,000; expressed as ppm = \frac{\text{grams of solute}}{\text{grams of solution}} \times 1,000,000.
Percent by Volume: The ratio of the volume of solute to the total volume of the solution, expressed as a percentage; expressed as \text{percent by volume} = \frac{\text{volume of solute}}{\text{volume of solution}} \times 100\%.
Percent Mass: The mass of a solute divided by the total mass of the solution, expressed as a percentage; expressed as \text{percent mass} = \frac{\text{mass of part}}{\text{mass of whole}} \times 100\%.
Saturated: A solution that contains the maximum amount of solute that can dissolve at a specific temperature.
Solute: The substance that is dissolved in a solution and is present in a smaller amount.
Solution: A homogeneous mixture of substances in the same physical state.
Solvent: The substance that dissolves the solute and is present in the greater amount.
Supersaturated: A solution that holds more solute than is present in a saturated solution at the same temperature; unstable and will precipitate solute if disturbed.
Unsaturated: A solution that contains less solute than the maximum amount it can hold at a specific temperature.
Vapor: A substance in the gas phase that is normally a solid or liquid at room temperature.
Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid or solid phase.

Topic Overview

Most materials encountered daily are mixtures, not pure substances. This topic explores solutions, a specific and important type of mixture in chemistry. Solutions are crucial because most chemical reactions occur within them. The discussion covers the nature of solutions, their properties, and methods to express concentration.

Solutions

A solution is defined as a homogeneous mixture where one substance is uniformly spread throughout another in the same physical state. This involves atoms, ions, or molecules of one substance distributed evenly within another.

Types of Solutions

Solid in Solid: E.g., Brass (zinc and copper). When metals are mixed, the solution is called an alloy.
Gas in Gas: E.g., Air.
Liquid Solutions: Most commonly, a solid or liquid dissolved in a liquid. The discussion focuses on liquid solutions.

Solute and Solvent

Solute: The substance being dissolved, present in a smaller amount (e.g., sodium nitrate in water).
Solvent: The substance that dissolves the solute, present in a greater amount (e.g., water). Water solutions are called aqueous solutions, denoted as (aq) in equations.
NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)

Characteristics of Liquid Solutions

Homogeneous mixtures.
Clear and do not disperse light.
May or may not have color.
Do not settle on standing.
Pass through a filter.

Solubility Factors

Solubility refers to how much of a solute will dissolve in a certain amount of solvent at a specific temperature.

Nature of Solute and Solvent

Ionic compounds like NaCl dissolve in water due to the attraction between positively and negatively charged ions and the polar water molecules. This attraction is called ion-dipole forces, stronger than the forces between the ions themselves.

Ionic and polar substances dissolve in polar solvents.
Nonpolar substances dissolve in nonpolar solvents.

"Like dissolves like" describes this phenomenon.

Temperature

The solubility of most solids in water increases with temperature.
The solubility of gases in liquids decreases with increasing temperature.

Pressure

Pressure has little effect on the solubility of solid or liquid solutes.
The solubility of gases in liquids increases with increasing pressure (e.g., carbon dioxide in soda).

Solubility Curves

Solubility Curves table shows the number of grams of a substance that can be dissolved in 100 g of water at temperatures between 0°C and 100°C.

Each line represents the maximum amount of that substance that can be dissolved at a given temperature.
All of the lines that show an increase in solubility as temperatures increase represent solids being dissolved in water.
Although these lines on the graph show an increase in solubility as temperature increases, a few solids, such as cesium sulfate, become less soluble as temperature increases.
Three lines show decreasing solubility with increasing temperature. These three lines represent the gases NH3, HCl, and SO2.
The solubility of all gases decreases with increasing temperature.

Unsaturated, Saturated, and Supersaturated Solutions

Unsaturated: Contains less solute than the maximum it can hold at a given temperature. More solute can be dissolved without changing the temperature.
Saturated: Contains the maximum amount of solute that will dissolve at a specific temperature. Additional solute will not dissolve and will settle at the bottom.
Supersaturated: Holds more solute than is present in a saturated solution at that temperature. These solutions are quite unstable and the addition of a single solid crystal of the substance will cause additional solid to form, and the solution will return to a saturated condition.
The only way to make a supersaturated solution is to cool a saturated solution in which there are no crystals or impurities, such as dust, present.

Solubility Tables

The Solubility Guidelines of Reference Tables for Physical Setting/Chemistry contains some guidelines for the solubility of common ionic compounds. The table shows that all compounds of the ammonium and the nitrate ion are soluble. All of the halide ions, such as Cl-, form compounds that are soluble, but three exceptions are listed. Silver chloride is not soluble, nor are Pb^{2+} nor Hg_2^{2+} chlorides, and they are precipitates if they form in a double-replacement reaction. This table is useful in predicting whether or not a precipitate will form when two ionic solutions are mixed. A reaction will take place if one or both of the products is listed as insoluble.

Concentration of Solutions

Solutions are homogeneous mixtures with variable compositions. Concentration is the specific amount of solute in a solution.

Molarity (M)

Molarity (M) is the number of moles of solute per liter of solution; expressed as M = \frac{\text{moles of solute}}{\text{liters of solution}}.

To find the molarity of a solution given the mass of the solute, convert the mass to moles and use the molarity formula.

Percent by Mass

Percent mass expresses the concentration of a solute as a percentage of the total mass of the solution; expressed as \text{percent mass} = \frac{\text{mass of part}}{\text{mass of whole}} \times 100\%.

Percent by Volume

Percent by volume expresses the concentration of a solute in a liquid solution as a percentage of the total volume; expressed as \text{percent by volume} = \frac{\text{volume of solute}}{\text{volume of solution}} \times 100\%.

Parts per Million (ppm)

Parts per million (ppm) is a ratio between the mass of a solute and the total mass of the solution, multiplied by 1,000,000; expressed as ppm = \frac{\text{grams of solute}}{\text{grams of solution}} \times 1,000,000. Useful for extremely dilute solutions.

Preparation of a Solution of Known Concentration

Add the desired amount of solute to a volumetric flask.
Add some distilled water and mix until the solute is dissolved and the solution is homogeneous.
Fill the volumetric flask to the mark on the neck of the flask, stopper, and again mix to ensure that the solution is homogeneous.

Colligative Properties

The freezing and boiling points of water change upon the addition of nonvolatile solutes. The amount of the lowering of the freezing point is not dependent on the nature of the added particle but only on the total number of dissolved particles. One mole of any particles will have the same effect on the freezing point: One mole of particles lowers the freezing point of 1000 g of water by 1.86°C.

Molecular vs Ionic

Molecular: One mole of sugar dissolved in water yields one mole of particles in solution.
C{12}H{22}O{11}(s) \rightarrow C{12}H{22}O{11}(aq)
Ionic: One mole of sodium chloride produces two moles of particles (ions) and will depress the freezing point of water twice as much as a mole of sugar.
NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)
The greater the number of ions, the greater the effect on the freezing point.

One mole of particles will elevate the boiling point of 1000 g of water by 0.52°C.

Vapor Pressure

In any sample of a liquid, some of the particles at the surface have sufficient energy to escape from their neighboring molecules and enter the gas phase. When a substance that is normally a solid or a liquid at room temperature enters the gas phase it is called a vapor.

As the temperature of a liquid increases, the particles have more energy, and more particles escape from the surface. These vapor particles are gaseous particles and exert pressure in the gaseous phase. The pressure that a vapor exerts is called vapor pressure.

Boiling Point

As the temperature of a liquid rises, vapor pressure increases. Finally the vapor pressure becomes equal to atmospheric pressure. At this point the gas may vaporize, not only on the surface but at any point in the container. A bubble of vapor below the surface has enough pressure that it does not collapse from the atmospheric pressure pushing against it. When a bubble can occur at any point in the liquid, the process is called boiling. The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid is 101.3 kPa, standard atmospheric pressure.

Equivalent pressures are 1 atm, 760 mm Hg, and 760 torr. The heat required to change 1 mol of a substance from a liquid at its boiling point to 1 mol of a vapor is termed the heat of vaporization.