THE SOLID STATE OF MATTERChapter 10.5

CHAPTER 11Solutions & Colloids

THE DISSOLUTION PROCESSChapter 11.1

SECTION OBJECTIVES

• Describe the basic properties of solutions and how they form.

• Predict whether a given mixture will yield a solution based on the molecular properties of the components.

• Explain why some solutions either emit or absorb heat when they form.

SOLUTIONS• Solutions are almost always

homogenous.

• Typically the physical state of the solvent and solute are the same.

• The components of the solution are distributed on the molecular scale.

• The solute in the solution will not settle out or separate from the solvent.

• The concentration of the solute(s) can be varied extensively.

FORMATION OF SOLUTIONS

• The formation of solutions is an example of a spontaneous process - a process that occurs under the given conditions without adding external energy to the system.

• Two factors favor the spontaneous formation of solutions:

1. A decrease in the internal energy of the system.

2. An increase in the entropy (disorder) of the system.

SOLUTIONS• When a solute dissolves in a solvent,

the solute-solute interactions and solvent-solvent interactions are being replaced with solute-solvent interactions. The forces must be comparable in strength in order for a solution to occur.

• A system where these forces are equivalent is termed an ideal solution.

• Like dissolves like.

SOLUBILITY OF METHANOL IN WATER

Methanol is very soluble in water as they both form H-bonds as their principle intermolecular bonds.

SOLUBILITY OF ALCOHOLS IN WATER & HEXANE

Solubility measured in mol alcohol per kg solvent at 20°C

Hexane (C6H14)

Water (H2O)

HEATS OF SOLUTION & SOLUTION CYCLES

1. Solute particles separate from each other - endothermic

• Solute (aggregated) + heat → Solute (separated) ∆Hsolute> 0

2. Solvent particles separate from each other - endothermic

• Solvent (aggregated) + heat → Solvent (separated) ∆Hsolvent> 0

3. Solute and solvent particles mix - exothermic

• Solute (separated) + solvent (separated) → Solution + heat ∆Hmix < 0

∆Hsolution = ∆Hsolute + ∆Hsolvent + ∆Hmix

PROBLEM• Iridium has the highest density of any element. It

crystallizes in the face-centered cubic lattice with a unit cell edge length of 383.9 pm. • Calculate the atomic radius of iridium. • Calculate the density of an iridium crystal in kg/m3

HEAT OF SOLUTION

• The formation of solutions will be either an exothermic process or an endothermic process.

• The net heat of solution is the result of breaking and forming intermolecular bonds.

NON-LIQUID SOLUTIONS• Any mixture of substances, regardless of

their phases, can constitute a solution.

• Gas-Gas solutions are common, as all gasses are fully miscible with each other.

• Gas-Solid solutions occur when gas molecules dissolve into the spaces between atoms/molecules of the solid.

• H2 gas can be purified by passing it through Pd metal, only H2 fits through the openings.

Gas Hydrates - natural gas trapped in crystalline ice structures

SOLID-SOLID SOLUTIONS• Since particles in solids don’t move

significantly many solid-solid mixtures are heterogeneous.

• Alloys, mixtures of elements that have metallic character, can be of two different types.

• Substitution alloys - where one atom is replaced by another in the crystal structure.

• Interstitial alloys - where one atom fills spaces between the atoms of the crystalline solid.

QUESTIONIf a solute dissolves in an endothermic process...

... H bonds must exist between solvent and solute.

... strong ion-dipole forces must exist in the solutions.

... the solute must be a gas.

... the entropy of the solution is immaterial.

... the entropy of the solution must be greater than that of its pure components.

Answer:AB

CDE

ELECTROLYTESChapter 11.2

SECTION OBJECTIVES

• Define and provide examples of electrolytes.

• Distinguish between the physical and chemical changes that accompany the dissolution of ionic and covalent electrolytes.

• Relate electrolyte strength to solute-solvent interactions.

ELECTROLYTES

• Some compounds, when dissolved in water, result in solutions which can conduct electricity well. These are termed electrolytes.

• The electrolytes are ions in solution, which allow the electrical current to pass through the solution.

• Non-electrolytes are compounds which dissolve, but do not form ions, such as sucrose.

IONIC-ELECTROLYTES• Water and other polar solvents can

dissolve ionic compounds by forming ion-dipole intermolecular bonds.

• Ionic compounds are typically strong electrolytes, as 100% of the compound will dissociate (separate) into ions.

• Some ionic compounds however are insoluble, dissolution is not energetically favorable.

KCl(s) → K+(aq) + Cl–(aq)

COVALENT ELECTROLYTES• Some covalent compounds will react when dissolved in water to

form ions.

• The greater the extent of the reaction, the stronger the electrolyte.

• HCl dissociates completely - strong electrolyte (strong acid)

• Acetic acid only partially dissociates - weak electrolyte (weak acid)

Dissolution of acetic acid.

SOLUBILITYChapter 11.3

SECTION OBJECTIVES

• Describe the effects of temperature and pressure on solubility.

• State Henry’s law and use it in calculations involving the solubility of a gas in a liquid.

• Explain the degrees of solubility possible for liquid-liquid solutions.

SOLUBILITY EQUILIBRIUM• The solubility of a solute in a

solvent is a product of an equilibrium in the solution.

• For each solute there is a critical concentration that will saturatethe solvent - no more solute will dissolve.

• Unsaturated solutions can have more solute added, until the solution becomes saturated.

KNO3(aq) ⇌ KNO3(s)

SUPERSATURATED SOLUTIONS

• Solutions can be made with concentrations of solutes above the saturation limit - these are supersaturated.

• Typically this is done by heating the solution to increase the solubility, then with careful cooling the solute can remain in solution (temporarily).

SOLUBILITY OF GASSES• The solubility of gasses in liquids is

dependent on both the strength of the intermolecular forces, and external factors, such as temperature and pressure.

• Increasing temperature decreases the solubility of gasses - more kinetic energy to overcome intermolecular forces.

SOLUBILITY OF GAS• Changing the pressure of a gas will also alter its solubility.

• Higher pressures allow for more gas to be dissolved, reducing the pressure causes the gas to escape (opening a Coke bottle)

• Henry’s Law is used to express the impact of pressure on gas solubility.

Cgas = kH × Pgas

Cgas - solubility of gas in mol/LkH - const. in mol/L•atm (specific for

each gas/solvent)Pgas - pressure of the gas in atm

PROBLEMThe Henry’s Law constant (kH) for CO in water

at 25°C is 9.71×10-4 mol/L•atm. How many grams of CO will dissolve in 1.00 L of water if

the partial pressure of CO is 2.75 atm?

Cgas = kH × Pgas

LIQUID-LIQUID SOLUBILITY• Some liquids are infinitely soluble in each

other (e.g. ethanol & water). These liquids are termed miscible.

• Immiscible liquids are ones that do not dissolve in each other to any appreciable extent.

• The miscibility of liquids relates to the similarity of the intermolecular forces of the two liquids. “Like dissolves like”.

SOLID-LIQUID SOLUBILITY• Most soluble solids are more

soluble at higher temperatures.

• The relationship between temperature and solubility can be exploited to make supersaturated solutions.

Precipitation of sodium acetate in a

supersaturated solution (hand warmer).

COLLIGATIVE PROPERTIESChapter 11.4

SECTION OBJECTIVES

• Express concentration of solution components using mole fraction and molality.

• Describe the effect of solute concentration on various solution properties.

• Calculate the impact of colligative effects on vapor pressure, boiling point, melting point, and osmotic pressure.

CONCENTRATIONS

Term Symbol Ratio

Molarity M amount (mol) of solute/volume (L) solution

Molality m amount (mol) of solute/mass (kg) solvent

Parts by mass % (w/w) mass of solute/mass of solutionParts by volume % (v/v) volume of solute/volume of solution

Mole fraction X amount (mol) of solute/all moles (solute & solvent)

The amount of a solute in a solution, the concentration of the solute, can be expressed in a wide range of units.

Different units have different uses. Some concentrations may change with the conditions surrounding their measurements (e.g. temperature).

QUESTIONFor a given solution, which of the following

concentrations values will change as the temperature changes?

Mass percentMolality

Mole fractionMolarity

None of these will change.

CONVERSION BETWEEN CONCENTRATIONS

• Conversions of concentration units are very similar to other conversion calculations - just pay attention to the units.

• You will need to use the molar mass when converting between a mass based concentration and one based on moles.

• You will need to use the density when converting from a massbased concentration to one based on volume (or for the reverse).

• Keep in mind that the molality involves the quantity of solvent, whereas the others involve the quantity of solution.

PROBLEMBenzene and diethyl ether are both liquids at room temperature. A solution is prepared by dissolving

86.9 g of diethyl ether (C4H10O: 74.12 g/mol) in 425 g of benzene (C6H6: 78.108 g/mol)?

Determine:•The molarity of diethyl ether in this solution.•The molality of diethyl ether in this solution.•The mole fraction of diethyl ether in this solution.

The density of benzene is 0.877 g/mLThe density of diethyl ether is 0.713 g/mL

COLLIGATIVE PROPERTIES

• The presence of a solutein a solution gives rise to many properties different from those of the pure solvent.

• There are four properties that depend merely on the amount of solute and not the chemical nature of the solute to alter the solution.

• These colligative properties are:

• Vapor pressure lowering

• Boiling point elevation

• Freezing point depression

• Osmotic pressure

• The impact of solutes on a solution can be measured and predicted mathematically.

VAPOR PRESSURE DEPRESSION

• The vapor pressure of a nonvolatile, nonelectrolyte solution is always lowerthan that vapor pressure of the pure solvent.

•With some solute molecules at the surface, fewer molecules of solvent vaporize per unit time.

• There is also an entropic contribution, which favors maintaining the solution, as opposed to separating solvent into the gas phase.

QUANTIFYING VAPOR PRESSURE CHANGE

• Raoult’s Law - concerning the vapor pressure of a solvent above a solution (Psolvent)

• Psolvent = Xsolvent · P°solvent

• P°solvent is vapor pressure of pure solvent

• Xsolvent is the mole fraction of the solvent in solution

• P°solvent - Psolvent = ∆P = Xsolvent · P°solvent

CALCULATIONThe vapor pressure of water at 25.0°C is 23.7 torr. What will be the vapor pressure of water, if 6.34 g of urea (CO(NH2)2)

is dissolved in 135 g of water?

PHASE DIAGRAMS OF SOLUTIONS

• Boiling point elevation - boiling point occurs when the vapor pressure equals the external pressure. With lowered vapor pressure the solution needs to get to a higher temp. to reach the same pressure.

• Freezing point depression solids forming at lower temperatures is another colligative property. This arises from the decrease in entropy that occurs as solute is concentrated in less liquid, necessitating a greater decrease in enthalpybefore the solid can form.

BOILING & FREEZING POINT CHANGE CALCULATIONS

• Boiling point elevation & Freezing point depression

• ∆Tb = (Kb)(m) & ∆Tf = (Kf)(m)

• Kb and Kf are molal boiling and freezing point constants, specific to each substance in units of °C/m; these values are tabulated.

• m is the molal concentration of the solute in the solution.

PROBLEM• On a skiing trip three students put their bottles of alcohol into the

snow out the back door of the ski chalet to cool off. If the mixture in the bottle freezes, the bottle will shatter and be undrinkable.

• If it is -19.6°C in the snow, will any of the bottles below survive?

• Beer: ethanol content = 5.5% (v/v) = 0.997 m

• Sake: ethanol content = 21% (v/v) = 4.553 m

• Vodka: ethanol content = 40% (v/v) = 11.418 m

OSMOTIC PRESSURE•Osmotic pressure occurs when there are solutions with two

different concentrations of solutes on either side of a permeable membrane.

• The membrane only allows for the passage of solvent not solute.

• Solvent will move from lower solute concentration side to the higher solute concentration side in order to equilibrate the concentrations.

•Osmotic pressure is represented by π

• π = (nsolute/Vsolution)RT = MRT R = 0.08206 L•atm/mol•KT in Kelvinπ in atm

OSMOTIC PRESSURE

OSMOTIC PRESSURE IMPACT ON CELLS

Prue water (hypotonic)

Isotonic solution

Salt water (hypertonic

)

QUESTIONHuman blood has a molar concentration of

solutes of 0.30 M. What is the osmotic pressure of blood at 25°C?

0.012 atm0.62 atm6.8 atm7.3 atm

>10. atm

Answer:ABCDE

π=MRT

COLLIGATIVE PROPERTIES OF ELECTROLYTE SOLUTIONS

• For electrolyte solutions, the compound formulas tells us how many particles are in the solution.

• NaOH gives 2 particles (Na+,OH-)

• Ca3(PO4)2 gives 5 particles (3Ca2+, 2PO43-)

• However ionic particles don’t act perfectly, experimentally the magnitude of the colligative properties are lower than what is predicted.

• The van’t Hoff factor (i) tells us the “effective” number of ions that are in the solution.

THE VAN’T HOFF FACTOR• The van’t Hoff factor (i) is a correction value that modifies

the impact of ions on the various colligative properties.

• The van’t Hoff factor (i) is determined experimentally and tabulated.

i = measured value for electrolyte solution

expected value for nonelectrolyte solution

Vapor pressure lowering: ∆P = i(Xsolute × P°solvent)Boiling point elevation: ∆Tb = i(Kbm)

Freezing point depression: ∆Tf = i(Kfm)Osmotic pressure: π = i(MRT)

NONIDEAL BEHAVIOR OF ELECTROLYTE SOLUTIONS

• Ions in solution do not behave independently as we would expect.

• The larger the ion charge the greater the deviation from the ideal.

• This is the result of the ionic atmosphere that forms around ions in solution.

ION PAIRS• Oppositely charged ions in

solution are attracted to each other.

• At higher concentrations there is a greater likelihood of ions pairing in solution.

• Greater dilution of ions results in more “ideal” behavior.

COLLOIDSCHAPTER 11.5

(a) A solution is a homogeneous mixture that appears clear, such as the saltwater in this aquarium.(b) In a colloid, such as milk, the particles are much larger but remain dispersed and do not settle.

(c) A suspension, such as mud, is a heterogeneous mixture of suspended particles that appears cloudy and inwhich the particles can settle. (credit a photo: modification of work by Adam Wimsatt; credit b photo:

modification of work by Melissa Wiese; credit c photo: modification of work by Peter Burgess)

COLLOIDSColloids (or colloidal dispersions) exhibit properties intermediate between those of suspensions and solutions. The particles in a colloid are larger than most simple molecules; however, colloidal particles are small enough that they do not settle out uponstanding. The particles in a colloid are large enough to scatter light, a phenomenon called the Tyndall effect. This can make colloidal mixtures appear cloudy or opaque. Clouds are colloidal mixtures. They are composed of water droplets that are much larger than molecules, but that are small enough that they do not settle out.

EXAMPLES OF COLLOIDS

COLLOIDS - DISH SOAP

An emulsion of oil in water is a colloid and appears cloudy.