Chapter 2 – Solutions

Introduction

Solutions are all around us. Our atmosphere is a solution of gases in gases, carbonated beverages are solutions of gases in liquids, sweetened drinks are solutions of solids in liquids, and solder is a solution of solids in solids. Reactions are normally carried out in solutions of gases or liquids because the particles in these solutions are free to move and to collide. In this chapter, we define several terms relevant to the study of solutions and discuss some solution properties.

2.1. Concentration Units

Introduction

The properties of a solution depend upon the relative amounts of solvent and solute(s), which is given by the solute concentration in the solution. A concentrated solution is one in which the concentration of a solute is relatively large, while a dilute solution is one in which the solute concentrations are relatively low. However, the terms 'concentrated' and 'dilute' used in this manner are not quantitative. In this section, we present several different ways of reporting the concentrations of solutions in a quantitative manner.

Objectives

• •
  Determine the concentration of a solution from the amount of solute and the amount of solution or solvent.
• •
  Use the definition of a concentration unit to convert between the amount of solute and the amount of solution or solvent.
• •
  Prepare a solution of known molarity.
• •
  Use and understand the meaning of the prefixes m (milli), µ (micro) and n (nano).

2.1-1. Definitions

Before we start into the quantitative aspects of solutions, we review some of the terms introduced in Section 10.1 of Chemistry – A Molecular Science (Chapter 10).

• •
  A solution is a homogeneous mixture consisting of a solvent and one or more solutes. A homogeneous mixture is one whose composition is the same throughout. For example, sugar dissolves in water to make a homogeneous mixture in which the sugar concentration (sweetness) is the same throughout, so the mixture is a solution. Mixing water and oil forms a mixture whose composition varies from pure water to pure oil, so this mixture is heterogeneous and not a solution.
• •
  The solvent is that substance that dictates the phase of the solution. If a liquid is present, it is usually the solvent. This course focuses on aqueous solutions, which are solutions in which water serves as the solvent.
• •
  A solute is one of the substances that is dissolved in the solvent. Solutes of aqueous solutions can be solid (a sugar solution), liquid (alcoholic beverages contain liquid ethanol in water), and/or gas (carbonated beverages contain CO₂ in water).

When ionic substances dissolve in water, they break up into their component ions, i.e., they ionize or dissociate in water. A substance that ionizes or dissociates completely in water is said to be a strong electrolyte. Ionic compounds are strong electrolytes. NaCl dissociates completely into Na¹⁺ and Cl^1– ions in water, so it is a strong electrolyte. Molecular substances do not dissociate when in water and are classified as nonelectrolytes. Sugar dissolves in water as sugar molecules (C₁₂H₂₂O₁₁), so sugar is a nonelectrolyte. Acids and bases react with water to produce ions, so they too are electrolytes. Strong acids and bases react completely with water and are strong electrolytes, but weak acids and bases react only partially with water and are classified as weak electrolytes.

2.1-2. Molarity (M)

The molarity (M) of a solute is the number of moles of the solute in one liter of solution or the number of millimoles of solute in one milliliter of solution.

( 2.1 )

M =

moles of solute

liters of solution

=

n

V

=

mmoles of solute

mL of solution

Molarity (M)

The molarity of a solute allows us to convert between the number of moles (or millimoles) of that solute and the number of liters (or milliliters) of solution. To determine the number of moles of solute in a given volume of solution, we solve Equation 2.1

M =

moles of solute

liters of solution

=

n

V

=

mmoles of solute

mL of solution

for n, the number of moles: Alternatively, we can use the units of molarity and the factor-label method to determine the number of moles. If V is given in milliliters, then MV is millimoles of solute. We will frequently use millimoles because volumes in a chemistry laboratory are most often given in milliliters. We can also solve Equation 2.1

M =

moles of solute

liters of solution

=

n

V

=

mmoles of solute

mL of solution

for the volume of solution to determine what volume of solution contains a given number of moles of solute: Or, using the factor-label method,

2.1-3. Preparing a Solution of Known Molarity

The denominator of molarity is the volume of the solution, not the volume of solvent. Thus, a solution of known molarity must be made by adjusting the volume after the solute and solvent have been thoroughly mixed. This is done with a volumetric flask of the appropriate volume (Figure 2.3) as follows.

• 1
  Select a volumetric flask of the appropriate volume and determine the amount of solute that will be required. The flask in Figure 2.3 would be used to prepare 500 mL of solution.
• 2
  Add solvent to the volumetric flask until it is about half full.
• 3
  Add the required amount of solute and swirl the mixture until the solute has dissolved
• 4
  Add solvent to the flask up to the fill line of the flask.
• 5
  Invert the flask several times to assure homogeneity.

2.1-4. Preparing a Solution Exercise

2.1-5. Molarity to Solute Mass Exercise

2.1-6. Molarity of Ions Exercise

2.1-7. Common Prefixes

1 M solutions are considered fairly concentrated, and most of the solutions encountered in the chemistry laboratory are much less concentrated. Indeed, some processes require only very dilute concentrations. For example, the maximum allowed level of mercury in drinking water is 0.000002 M, and a testosterone level of only 0.00000001 M initiates puberty in human males. Prefixes are used to avoid the preceding zeros or exponents when expressing these very dilute solutions.

Unit	Symbol	Value
millimolar	mM	10–3 M	1 mmol·L–1
micromolar	µM	10–6 M	1 µmol·L–1
nanomolar	nM	10–9 M	1 nmol·L–1

Thus, the maximum allowed level of mercury in drinking water is 2 × 10^–6 M = 2 µM, and the testosterone level that initiates puberty is 10 × 10^–9 M = 10 nM.

2.1-8. Using Prefixes Exercise

2.1-9. Mole Fraction (X)

The mole fraction (X) of a substance in a solution is the number of the moles of that substance divided by the total number of moles of all substances (solutes and solvent) in solution. Because all of the substances comprise the whole mixture, the sum of all mole fractions (including the solvent) must equal one; i.e.,

	XA
A

= 1.

2.1-10. Mole Fraction Exercise

2.1-11. Mass Fraction (%, ppt, ppm)

The mass fraction relates the mass of a solute to the mass of the solution. It is commonly used for solutions where the solvent is not clearly defined.

( 2.3 )

mass fraction =

mass of solute

mass of solution

Mass Fraction

2.1-12. Mass Fraction Multipliers

Mass fractions represent a part of the whole, so they are always less than one. Indeed, mass fractions in dilute solutions can be so small that they are multiplied by powers of ten to eliminate the preceding zeros (see Table 2.2). Thus,

• •
  a mass fraction of 0.012 would be reported as (0.012)(100%) = 1.2%
• •
  a mass fraction of 1.2 × 10^–8 would be reported as (1.2 × 10^–8)(10⁹ ppb) = 12 ppb.

Name	Multiplier	Useful Units
mass fraction	1	g A g mixture
mass percent (%)	100	g A 100 g mixture
parts per thousand (ppT)	1000	g A 1000 g mixture
parts per million (ppm)	106	g A 106 g mixture
parts per billion (ppb)	109	g A 109 g mixture

2.1-13. Mass Fraction Exercise

2.1-14. Molality (m)

The molality (m) of a solute is the number of moles of that solute per kilogram of solvent. The volume of a solution varies with temperature, so the molarity (M) of the solution changes with temperature as well. Thus, molarity (M) is not a good unit of concentration to use in experiments involving temperature change. While volume is temperature dependent, mass is not, so molality (m), which involves no volumes, is commonly used for the concentration of any solution used in experiments that involve measurements resulting from temperature changes.

2.1-15. Molality Exercise

2.2 Changing Concentrations Units

Introduction

It is sometimes necessary to change the units of concentration. This routine procedure is most often done when the concentration of a stock solution is given in units that are not convenient for a particular experiment. The procedure presented here entails converting the units of the numerator and denominator of the given concentration to those of the desired concentration and then dividing.

Objectives

• •
  Convert between different concentration units.

2.2-1. Converting Concentration Units

All concentrations are ratios of two quantities. The numerator indicates the amount of solute and the denominator the amount of solution or solvent. Use the following steps to convert one concentration unit into another:

• 1
  Convert the given numerator into the numerator of the desired unit.
• 2
  Convert the given denominator into the denominator of the desired unit.
• 3
  Divide the result of step 1 by that of step 2 to obtain the desired concentration unit.

The numerators and denominators of the concentration units that we have dealt with are:

molarity	mass fraction	mole fraction	molality
moles of solute liters of solution	mass of solute mass of solution	moles of solute total moles (all solutes + solvent)	moles of solute kilograms of solvent

To convert between any of the above you need only some combination of the molar masses of solute and/or solvent, and/or the density of the solution.

Conversion	Requires
mol solute ↔ mass solute	molar mass of the solute
liters solution ↔ mass solution	density of solution
liters solution ↔ mol solute + mol solvent	density of solution, molar masses of solute and solvent
liters solution ↔ kg solvent	solution density, molar mass of solute
mass solution ↔ mol solute + mol solvent	molar masses of solute and solvent
mass solution ↔ kg solvent	molar mass of solute
mol solute + mol solvent kg solvent	molar masses of solute and solvent

2.2-2. Molarity to Mass Percent Exercise

2.2-3. Molarity to Molality Exercise

2.2-4. ppm to Molarity Exercise

2.2-5. Mass Percent to Molarity Exercise

2.3 Dilutions

Introduction

Stock solutions are often fairly concentrated and must be diluted before they are used. In this section, we show how to make a solution of known volume and molarity from a more concentrated stock solution.

Objectives

• •
  Calculate the concentration of a solution after dilution.
• •
  Prepare a dilute solution of known concentration from a more concentrated solution.

2.3-1. Dilution Lecture

2.3-2. Equations of Dilution

The amount of solute does not change with the addition of solvent, so we can write the following: where C is any concentration that indicates the amount of solute per unit volume of solution. Therefore, C can be molarity or g/mL, but not molality or mass percent. We solve Equation 2.5 for the final concentration, C_f to obtain the following where the ratio V_i/V_f is called the dilution factor. Thus, the final concentration equals the initial concentration times a dilution factor. If a solution is diluted successively several times, then the final concentration equals the initial concentration times the product of all of the successive dilution factors as shown in Equation 2.7.

2.3-3. Molarity after Dilution Exercise

2.3-4. Volume Required for Dilution Exercise

2.3-5. Concentration After Successive Dilutions Exercise

2.3-6. Concentration of Initial Solution Exercise

2.4 Determining Concentrations

Introduction

Analyzing solutions to determine their concentration is an important part of analytical chemistry. In this section, we discuss two important methods used to carry out such an analysis: spectrometry and titration.

Objectives

• •
  Use a color wheel to determine a complementary color.
• •
  Use Beer's Law to determine the molarity of a solution.
• •
  Determine the volume of one reactant that reacts with another reactant of known volume.
• •
  Determine the limiting reactant from a balanced chemical equation and the concentrations and volumes of the reactants.
• •
  Determine the concentrations of the excess reactants.

2.4-1. Complementary Colors

White light is the result of all visible colors. When white light shines on a colored substance, however, some of the colors are absorbed. Those colors that are not absorbed can be either reflected or transmitted to the eye. We perceive this reflected or transmitted portion as the color of the substance. Consequently, the characteristic color of a material is not the color of light that it absorbs; rather it is the mixture of the remaining, unabsorbed colors that are observed. In other words, it appears as its complementary color. The approximate relationship between observed and absorbed colors is summarized in a color wheel as shown in Figure 2.7. For example, a solution appears orange because it absorbs blue light. Thus, when white light shines on the solution, the blue portion of the spectrum is absorbed by the solution, which leaves only the orange portion to be detected by your eyes.

2.4-2. Spectrometry

Spectrometry, measuring the amount of light that a solute absorbs at some wavelength, is a convenient way to determine the concentration of a solute in a solution. The amount of light that is absorbed by the sample is called the absorbance (A) of the sample. Consider the experiment shown in Figure 2.8 below. Orange light of intensity I_o enters the cell, where some of it is absorbed by the blue solution, so the outgoing intensity I, is less than the initial I_o. The absorbance, which is defined as A = –log(I/I_o),depends upon three factors:

• •
  how strongly the solute absorbs at the wavelength of light used in the experiment, which is known as the molar absorptivity (ε, epsilon);
• •
  the concentration of the solute (c); and
• •
  the path length of the light through the solution (l).

The relationship between these factors is given by Beer's Law.

2.4-3. Determining an Unknown Concentration

The following procedure is used to determine an unknown concentration in a solution:

• 1
  Determine a good wavelength of light to use. The light should be absorbed by the solute to be determined, so it should be the complement of the color of the solution. The complement can be found as the color opposite the observed color on a color wheel. The light should not be absorbed by the rest of the solution, but the effect of the rest of the solution can be determined by measuring the absorbance of a blank. The blank should mimic the solution to be measured in all aspects except the presence of the solute to be measured. The absorbance of the blank is then subtracted from that of the sample to get a measurement of just the solute's absorbance.
• 2
  Measure the absorbance of a standard solution, which is simply a solution in which the concentration of the solute being measured is accurately known. The molar absorptivity of the solute can be determined from the measured absorbance, the known concentration, and the path length of the cell.
• 3
  Measure the absorbance of the sample. The unknown concentration can be determined from the measured absorbance, the molar absorptivity determined in step 2, and the known path length of the cell.

Many solutes absorb very strongly, and the prepared solutions frequently absorb so strongly that too little light passes through the cell. In these cases, the solutions must be diluted before they can be measured. Indeed, it is often the case that they must be diluted several times in order to get a good reading of the absorbance. The following exercise takes you through this process except that no blank is used. Follow the above steps and dilute the samples to obtain the good absorbance readings. Don't forget to take into account your dilutions when doing the calculations.

2.4-4. Determining Concentration Exercise

Thank you for your patience as this section undergoes re-construction.

2.4-5. Titrations

Spectrometry is an easy way to determine an unknown concentration only if the solute absorbs light at a convenient wavelength (usually one in the visible). The concentration of a solute can also be determined by measuring the stoichiometric amount of a solution of known concentration required to react with the unknown solute. This method is called titration. The known solution, whose volume is to be determined, is called the titrant, while the solute being analyzed is the analyte. The point at which the titrant and analyte are in stoichiometric amounts is called the equivalence point. The equivalence point is usually approximated by an end point, which is the point at which an indicator undergoes a color change.

2.4-6. Titration Method

Acid-base titrations are very common in the chemistry lab. The net chemical equation for the reaction of a strong acid and a strong base is: The above chemical equation shows that the reactants react in a 1:1 ratio, so the equivalence point is reached when the number of (milli)moles of base added equals the number of (milli)moles of acid present initially. However, the number of (milli)moles of a solute equals the volume of solution in (milli)liters times the molarity of the solute. Thus, the equivalence point in an acid base reaction is that point where If the acid is the unknown, then a known amount of acid, the analyte, would be added to the flask, so V_acid would be known. The concentration of the base (M_base), the titrant, would also be known and the volume of base required to reach the equivalence point (V_base) would be determined by the titration. M_acid would then be determined from Equation 2.9.

2.4-7. Titration Exercise

Thank you for your patience as this section undergoes re-construction.

2.4-8. Precipitation Reactions

The concentration of a solution can also be determined from the amount of precipitate that forms when an excess of another reactant is added. The procedure is the following:

• 1
  Use the solubility rules to decide upon a good reactant.
• 2
  Add an excess of the reactant to a known volume of the solution to be analyzed.
• 3
  Filter the solution to separate the precipitate from the solution. Add more reactant to the remaining solution to be certain no more precipitate forms.
• 4
  Dry and weigh the precipitate then determine the number of moles of precipitate that were produced.
• 5
  Use the results of step 4 to determine the number of moles of the solute that must have reacted.
• 6
  Determine the concentration of the original solution from the number of moles of solute present (step 5) and the original volume.

Rule 1	Compounds of NH41+ and group 1A metal ions are soluble.
Rule 2	Compounds of NO31–, ClO41–, ClO31– and C2H3O21– are soluble.
Rule 3	Compounds of Cl1–, Br1–, and I1– are soluble except those of Ag1+, Cu1+, Tl1+, Hg22+,and Pb2+.
Rule 4	Compounds of SO42– are soluble except those of Ca2+, Sr2+, Ba2+, and Pb2+
Rule 5	Most other ionic compounds are insoluble.

2.4-9. Precipitation Exercise

2.5 Colligative Properties

Introduction

The properties of a solution depend upon the concentrations of the solutes. Indeed, some solution properties depend only upon the concentration of particles in solution and not upon the identity of those particles. These properties are called colligative properties. In this section, we discuss the following colligative properties.

• •
  vapor pressure
• •
  boiling point
• •
  freezing point
• •
  osmotic pressure

Objectives

• •
  Determine the van't Hoff factor (i) for a compound.
• •
  Convert a molarity or molality into a colligative molarity or colligative molality.
• •
  Calculate the boiling point and melting point of a solution given its colligative molality and the boiling elevation and freezing point depression constants.
• •
  Describe the causes of vapor pressure lowering at the molecular level.
• •
  Calculate the vapor pressure of a solution given the mole fraction of the solute and the vapor pressure of the solvent.
• •
  Describe osmosis at the molecular level.
• •
  Determine the osmotic pressure of a solution of known concentration.
• •
  Explain why osmosis is important.

2.5-1. van't Hoff Factor and Colligative Concentrations

Colligative properties depend upon the concentration of particles in solution, which is different than the solute concentration when the solute is ionic and dissociates in water. The solution concentration is converted to a concentration of particles with the van't Hoff factor (i), which is the number of moles of particles produced when one mole of solute dissolves. For example, i = 3 for CaCl₂ because dissolving one mole of CaCl₂ produces three moles of ions: CaCl₂ → Ca²⁺ + 2 Cl^1–. The total particle concentration is called the colligative concentration. It is equal to the solute concentration times its van't Hoff factor. The following equations define the colligative molarity and colligative molality.

2.5-2. van't Hoff Exercise

2.5-3. Colligative Molality Exercise

2.5-4. Phase Diagram

Figure 2.10 shows the phase diagrams of a pure substance and of a solution in which the substance is the solvent. The two diagrams are identical except for the region in blue, which is either a solid or gas in the pure substance but is all liquid in the solvent. Thus, the phase diagram of the solvent differs from that of the pure substance in three ways:

• •
  The vapor pressure is lowered by ΔP at every temperature.
• •
  The boiling point is elevated (raised) by ΔT_b.
• •
  The freezing point is depressed (lowered) by ΔT_f.

The impact of these three effects is to increase the temperature-pressure range of the liquid state. The amount by which the liquid state is extended (shown in blue in the figure) depends only on the concentration of the solute particles, so freezing point depression, boiling point elevation, and vapor pressure lowering are all colligative properties, and we now examine the relationship between particle concentrations and their effect on each of these colligative properties. We then define a fourth colligative property, the osmotic pressure.

2.5-5. Vapor Pressure Lowering

The liquid ⇌ vapor equilibrium is a dynamic equilibrium that is established when the rate of evaporation equals the rate of condensation. The pressure of the vapor in equilibrium with the liquid at a given temperature is the vapor pressure of the liquid at that temperature. Evaporation occurs from the surface of the liquid, so the rate of evaporation depends upon the concentration of particles at the surface. Figure 2.11 compares a pure solvent with a solution in which the mole fraction of a nonvolatile solute is 0.2. A solute mole fraction of 0.2 means that 20% of the particles in solution are solute particles, so 20% of the sites on the surface are occupied by nonvolatile solute particles. Consequently, evaporation can occur from only 80% of the surface sites, which results in a 20% reduction in the rate of evaporation, which in turn, causes a 20% reduction in the vapor pressure. We conclude that ΔP, the amount by which the vapor pressure of the solvent at some temperature is lowered by the addition of a solute with a mole fraction of X_solute, is ΔP is the vapor pressure lowering; it is always positive because the vapor pressure of a solution is always lower than that of the pure solvent. P° is the vapor pressure of the pure solvent at the temperature under consideration. The vapor pressure of the solution is lower than that of the solvent by ΔP, so P = P° – ΔP = P° – X_soluteP° = (1 – X_solute)P°. However, 1 – X_solute = X_solvent, so we can rewrite Equation 2.11 as The above indicates that the vapor pressure of a solution is equal to the vapor pressure of the pure solvent times the fraction of the surface sites occupied by the solvent; i.e., the mole fraction of the solvent.

2.5-6. Vapor Pressure of Water

We will have made frequent use of the vapor pressure of water. The vapor pressure at 5 °C increments can be found in Table 2.5 and in the Resources.

T (°C)	P° (torr)		T (°C)	P° (torr)
0	4.6		50	92.5
5	6.5		55	118.0
10	9.2		60	149.9
15	12.8		65	187.5
20	17.5		70	233.7
25	23.8		75	289.1
30	31.8		80	355.1
35	41.2		85	433.6
40	55.3		90	525.8
45	71.9		95	633.9

2.5-7. Vapor Pressure Lowering Exercise

2.5-8. Boiling Point Elevation

The normal boiling point of a liquid is the temperature at which its vapor pressure is 1 atm. Although this is the definition of the normal boiling point, it is common to refer to it as simply the boiling point. The boiling point of water is 100 °C, so its vapor pressure at 100 °C is 1 atm, but the vapor pressure of an aqueous solution is less than 1 atm at 100 °C due to vapor pressure lowering. Consequently, an aqueous solution must be heated to a higher temperature to achieve a vapor pressure of 1 atm, so the boiling point of an aqueous solution is always higher than that of pure water. This reasoning can be applied to any solution, so we conclude that the boiling point of a solution is always higher than the boiling point of the pure solvent. The amount by which the boiling point of the solvent is increased by the addition of a nonvolatile solute is known as the boiling point elevation, ΔT_b, which can be determined from the following equation: k_b is the boiling point elevation constant, which has units of °C·m^–1 and depends on the solvent, and m_c is the colligative molality of the solute. ΔT_b is the amount by which the boiling point of the solvent (T°_b) is raised, so the boiling point of the solution (T_b) is given as

2.5-9. Freezing Point Depression

Spreading salt on icy streets and sidewalks melts the ice because solute particles reduce the freezing point of a solvent in much the same way as they reduce its vapor pressure. That is, solute particles block sites on the solid where solvent molecules might otherwise freeze, thereby reducing the rate of freezing. The amount by which the freezing point is lowered is called the freezing point depression, ΔT_f, and is shown in Equation 2.14: k_f is the freezing point depression constant, which has units of °C·m^–1 and depends on the solvent, and m_c is the colligative molality of the solute. ΔT_f is the amount by which the freezing point of the solvent (T°_f) is lowered, so the freezing point of the solution (T_f) is given as

2.5-10. Solvent Constants

The solvent constants relevant for colligative properties are shown in Table 2.6 and in the Resources.

Solvent	Freezing Point (°C)	kf (°C·m–1)	Boiling Point (°C)	kb (°C·m–1)
Acetic acid	16.6	3.90	117.9	3.07
Benzene	5.5	4.90	80.1	2.53
Chloroform	–63.5	4.70	61.7	3.63
Cyclohexane	6.6	20.0	80.7	2.79
Ethanol	–117.3	1.99	78.5	1.22
para-Xylene	11.3	4.30
Water	0.0	1.86	100.0	0.512

2.5-11. Antifreeze Exercise

2.5-12. Salt Exercise

2.5-13. Molar Mass from Freezing Point Exercise

2.5-14. Osmosis

A semipermeable membrane allows solvent molecules to pass through but denies passage of solute particles. The rate of passage of solvent molecules through the membrane is proportional to the concentration of the solvent molecules, so the rate is greater for a pure solvent than for a solution. When a solvent and a solution are separated by a semipermeable membrane, more solvent molecules move from the solvent to the solution than in the reverse direction. The net movement of the solvent molecules through the membrane from the solvent or a more dilute solution into a more concentrated one is called osmosis.

2.5-15. Osmotic Pressure

Osmosis of the solvent through the membrane from a solvent to a solution decreases the concentration of particles in the solution, but it also increases its volume while decreasing the volume of the solvent. These volume changes result in a height difference in the two columns as represented by h in Figure 2.13. The height difference results in a pressure differential at the membrane that increases the flow of solvent molecules from the solution side. Eventually, the combination of the increased pressure coupled with the decrease in concentration of the solute particles increases the flow of solvent particles from the solution back into the solvent to the point where the flow of solvent molecules through the membrane is the same in both directions. The pressure at which equilibrium is attained is called the osmotic pressure (Π) of the solution. The osmotic pressure that is developed is given in Equation 2.15 where R is the ideal gas law constant (R = 0.0821 L·atm/K·mol).

2.5-16. Osmotic Pressure of Sea Water Exercise

2.5-17. Molar Mass from Osmotic Pressure Exercise

2.5-18. Applications of Osmosis

Purification: Solvents can be purified by a process called reverse osmosis. If a pressure that exceeds the osmotic pressure is applied to a solution, solvent can be forced from the solution into the pure solvent. This process has been used to purify seawater in the Middle East. The challenge to purifying seawater lies in finding membranes that can withstand the large pressure (over 26 atm) required. Biology: Cell membranes are semipermeable. When the solution around a cell has the same colligative concentration as within the cell, the cell maintains its size. However, if the cell is placed in water, water passes into the cell and can rupture it. Placing the cell in a solution with a greater colligative molarity causes water to leave the cell. This is why drinking seawater does not quench your thirst. Water transport in plants: Water enters a tree through the membranes in the roots, but it evaporates from the leaves resulting in a substantial concentration difference between the roots and leaves. The large concentration difference can develop osmotic pressures of up to ~20 atm in tall trees.

2.6 Colloids

Introduction

Sometimes, rather than dissolving, one material will be suspended as small aggregates in another. Thus, sand will stay suspended in water as long as the water is stirred; but soon after the stirring has stopped, the sand settles to the bottom of the container. Because these mixtures do not meet the rigorous criterion of a solution, they are called dispersions or suspensions.

Objectives

• •
  Identify common colloidal suspensions.

2.6-1. Types of Colloids

Suspensions in which the particle size is very small (1 nm to 1 µm) are called colloidal suspensions or simply colloids. White paint is a colloidal suspension of SiO₂ and TiO₂ particles, which are used to make the paint opaque and white, respectively. Colloidal suspensions, which can be stable (will not settle) for years, are classified according to their composition. Whipped cream is a foam: a gas suspended in a liquid. Jellies and starch solutions are sols: solids suspended in a liquid. Milk is an emulsion: a liquid suspended in a liquid. Aerosols can be liquids suspended in a gas (hair sprays) or solids suspended in a gas (smoke). Fog is also an aerosol (water in air) and smog, the combination of smoke and fog, is also a colloidal suspension.

2.7 Exercises and Solutions

Select the links to view either the end-of-chapter exercises or the solutions to the odd exercises.

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