Chapter 4 Solutions _ Their Properties

CHAPTER 4CHAPTER 4SOLUTIONS SOLUTIONS

& & THEIR PROPERTIESTHEIR PROPERTIES

SOLUTIONSSOLUTIONS

Learning objectives:Learning objectives: Discuss different kinds of solutionsDiscuss different kinds of solutions

Solutions• most important class of homogeneous mixtures• contain particles with diameters in range 0.1-2 nm• transparent, do not separate on standing (salt water, sugar water)

Colloids• contain particles with diameters in the range 2-500nm• murky, do not separate on standing (milk, fog,)

Suspensions• having larger particles than colloids• Not truly homogeneous• Particles separate on standing (blood, paint)

Types of Homogeneous Mixtures

Solutes and Solvents• A solution consists of a solute and a solvent: solute the substance which is being dissolved. solvent the substance (usually a liquid) that

dissolves the solute (usually, the solvent is the most abundant component in the mixture).

• Aqueous solution are solutions in which the solvent is water.

Solute Solvent Solution

Kinds of Solutions

Metal alloys such as sterling silver(Ag and Cu), brass (Cu and Zn) andbronze (Cu and Sn); waxes

SolidSolid

Dental amalgam (mercury in silver)SolidLiquidH2 in palladium metalSolidGasSolid

solutions

Seawater (NaCl and other salts inwater)

LiquidSolid

Gasoline (Mixture of hydrocarbons),vodka (ethanol and water)

LiquidLiquid

Carbonated water (CO2 in water)LiquidGasLiquid

solutions

Air (O2, N2, Ar, CO2, H2O, and othergases)

GasGasGaseous solution

ExamplesSolvent SoluteSolutionPhase

ENERGY CHANGES & THE ENERGY CHANGES & THE SOLUTION PROCESSSOLUTION PROCESS

Learning objectives:Learning objectives: Identify intermolecular forces in solution Identify intermolecular forces in solution Explain types of solution interactionsExplain types of solution interactions Identify intermolecular forces in solutionsIdentify intermolecular forces in solutions Describe dissolution of Describe dissolution of NaClNaCl in waterin water Explain the rule of thumb "like dissolves like.Explain the rule of thumb "like dissolves like.““ Define the entropy and enthalpy of solutionDefine the entropy and enthalpy of solution

Intermolecular Forces in SolutionsRelative strengths of intermolecular forces must be considered between solute and solvent particles that promote or prevent the formation of a solution.

Intermolecular Forces (in order of decreasing strength):

ion-dipole forces — solvent molecules cluster around ions in hydration shells, disrupting the bonding in the crystal lattice.

Hydrogen bonds— substances with O—H and N—H bonds are often soluble in water because of H-bonding (unless the molecules are large).

dipole-dipole forces —polar solutes interact well with polar solvents through attraction of partial charges.

ion - induced dipole forces — responsible for the attraction between Fe2+ and O2 molecules in the bloodstream.

dipole - induced dipole forces —responsible for the solvation of gases (nonpolar) in water (polar).

London (dispersion) forces —the principal attractive force in solutions of nonpolar substances (e.g., petroleum).

Intermolecular Forces in Solutions

Solution Interactions

Solutions form when solvent-solvent, solute-solute, andsolute-solvent forces are similar.

Solution may or may not form depending on relative disparity

Solvent-solute interaction < Solvent-solvent and solute-solute interaction

Solution formsSolvent-solute interaction = Solvent-solvent and solute-solute interaction

Solution formsSolvent-solute interaction > Solvent-solvent and solute-solute interaction

Relative Interaction and Solution Formation

Solution Interactions

Dissolution of NaCl in water

The General Solubility RuleThe general rule in solubility is that “like dissolves like”Water, a polar molecule, dissolves ethanol, which is alsopolar, but does not dissolve hexane and dichloromethane which are both nonpolar.

Ethanol and water are miscible—completely soluble in each other in all proportions.

Hexane and water are immiscible—they do not mix with each other at all.

Hexane and dichloromethane are miscible with each other.

Entropies of solutionsEntropy is a measure of thedisorder or energyrandomization in a system.

Entropies of solution areusually positive becausemolecular randomnessusually increases when (a) a solid dissolves in a

liquid or (b) one liquid dissolves in

another.

Enthalpy of solution• Enthalpy of solution measures how much energy is either absorbed

or released when a solution is prepared.

• The value of ΔHsoln is the sum of three terms:

Solvent-solvent interactions:Energy is required (+ ΔH) to overcome intermolecular forces between solvent molecules because the molecules must be separated and pushed apart to make room for solute particles

Solute-solute interactions:Energy is required (+ ΔH) to overcome intermolecular forces holding solute particles together in a crystal.

Solvent-solute:Energy is released (- ΔH) when solvent molecules cluster around solute particles and solvate them

Energy Changes and the Solution Process

The solute-solvent interactions are greater than the sum of the solute-solute and solvent-solvent interactions.

Energy Changes and the Solution Process

The solute-solventinteractions are less than the sum of the solute-soluteand solvent-solventinteractions.

Enthalpy of solution

UNITS OF CONCENTRATIONUNITS OF CONCENTRATION

Learning objectives:Learning objectives: Interconvert units of concentrationInterconvert units of concentration Perform calculations using solution density, Perform calculations using solution density,

molarity,molemolarity,mole fraction, weight percent, parts per fraction, weight percent, parts per million, parts per billion, and million, parts per billion, and molalitymolality..

Units of Concentration• Concentration: The amount of solute present in a given

amount of solution.

• Molarity (M):

• Mole Fraction (X):

SOLUTION of Literssolute of MolesMolarity

moles of number Total Aof MolesAX

• Mass percent: The ratio of the mass of a solute to the mass of a solution, multiplied by 100%.

% bymassof solute =mass of solute

mass of solution 100%

mass of solution =mass of solute +mass of solvent

Units of Concentration

• Parts per Million (ppm):

• One ppm gives 1 gram of solute per 1,000,000 g or one mg per kg of solution.

For solid samples: ppm = µg/g = mg/kg

• For dilute aqueous solutions this is about 1 mg per liter of solution.

For liquid samples: ppm = µg/mL = mg/L

610xolutionsofmassTotal

soluteofMassUnits of Concentration

• Parts per Billion (ppb):

For solid samples: ppm = µg/g = mg/kg

For liquid samples: ppm = µg/mL = mg/L

• Molality (m):

SOLVENT of Kilogramssolute of Moles=Molality

Units of Concentration910x

olutionsofmassTotalsoluteofMass

Units of Concentration

FACTORS AFFECTING FACTORS AFFECTING SOLUBILITYSOLUBILITY

Learning objectives:Learning objectives: Define saturated, unsaturated and supersaturated Define saturated, unsaturated and supersaturated

solutionssolutions Describe crystallization processDescribe crystallization process Discuss the effect of temperature and pressure on Discuss the effect of temperature and pressure on

solubilitysolubility State HenryState Henry’’s Law and its exampless Law and its examples

Saturated and Unsaturated Solutions• Saturated: Contains the maximum amount of solute that

will dissolve in a given solvent.

• Unsaturated: Contains less solute than a solvent has the capacity to dissolve.

• Supersaturated: Contains more solute than would be present in a saturated solution; these solutions are unstable, and a slight disturbance causes the “extra”solute to precipitate out.

• Crystallization: The process in which dissolved solute comes out of the solution and forms crystals.

a) A supersaturated solution of sodium acetate in water

b) When a tiny seed crystal is added,larger crystals begin to grow and precipitate from the solution until equilibrium is reached

Precipitation from a supersaturated solution

Solubilities are temperature-dependent.The solubility of most molecular and ionic solids increases with temperature, although some are almost unchanged, and some decreaseFor a solute with ΔHsoln > 0:

solute + solvent + heat saturated solution solubility increases with temperature.

For a solute with ΔHsoln < 0:solute + solvent saturated solution + heat solubility decreases with temperature.

Effect of Temperature on Solubility

Solids:

Solubilities of some common solids in water as a function of temperature. Most substances become more soluble as temperature rises, although the exact relationship is often complex and nonlinear.

Effect of Temperature on Solubility

Effect of Temperature on SolubilityGases:

• Solubilities of some gases in water as a function of temperature. • Most gases become less soluble in water as the temperature rises. Soft drinks become “flat” as they warm up and lose carbon dioxide. Aquatic life is affected by decreasing amounts of dissolved oxygen

as a result of thermal pollution.

Effect of Pressure on the SolubilityPressure has little effect on the solubility of solids and liquids, but has a large effect on gases.At a given pressure, there is an equilibrium between the

gas which is dissolved in the solution and the gas in the vapor phase.If the pressure increases, more gas dissolves to reduce the “extra” pressure; the new equilibrium is established with more gas dissolved.

• Henry’s Law:The solubility(Sgas, in mol/L) of a gas is proportional to the pressure of the gas (Pgas, in atm) over the solution

(The Henry’s law constant, k is a proportionality constant, unique to each gas, at a given temperature, with units of mol L-1 atm-1.)

Effect of Pressure on the Solubility

Sgas = k . Pgas

Examples of Henry’s – law behavior:• When a can of soda is opened, bubbles of gas fizzing out

of solution because the pressure of CO2 in the can drops and CO2 suddenly becomes less soluble.

• If a deep sea diver comes up to the surface too quickly, N2which has dissolved in his bloodstream at higher pressures comes back out of solution.

The N2 forms bubbles which block capillaries and inhibit blood flow, resulting in a painful, and potentially lethal, condition called the “bends.”

Less soluble gases, such as He, are often used in the breathing mixtures to reduce this problem.

Effect of Pressure on the Solubility