Properties of Solutions

Learning objectives Define terms solute, solvent and solution Distinguish between solutions and heterogeneous mixtures Distinguish among non-, weak and strong electrolytes Describe factors that affect solubility Describe Henry’s law and its application to explain common phenomena involving gases Perform calculations of solution concentration using various definitions Use molarity in stoichiometry calculations Describe basis of Raoult’s law and colligative properties Calculate solute concentrations in colligative property context Explain basis of osmotic pressure

Definitions of a solution A homogeneous mixture of two or more substances Solute is the component that is dispersed in the solvent – usually the minority component Solvent is the dispersing component – usually the majority component Sometimes definitions can become blurred: water (solvent) dissolves much greater than its own mass of sugar (solute)

Formation of a solution Crystals are held together by strong ionic bonds Polar water molecules exert attractive forces on ions Hydration of the ions by water molecules overcomes lattice energy Crystal lattice disperses

Like dissolves like All gases mix with each because there are no intermolecular forces Solids and liquids mix if intermolecular forces between unlike substances (adhesive forces) are similar to forces between like substances (cohesive forces)

Electrolytes Electrolytes are substances that dissociate into ions in solution – ionic compounds (sodium chloride) Strong electrolytes are completely ionized Weak electrolytes are partly ionized Non-electrolytes are those substances that produce no ions (sugar)

Colloids and solutions Both appear clear and uniform Solution is homogeneous; colloid is heterogeneous Colloid contains particles suspended in the liquid 1 – 200 nm diameter Colloid particles scatter light, solute particles do not

Factors affecting solubility Difference in polarity between solute and solvent – like dissolves like Temperature Solid solutes: depends on balance of several factors – can increase, decrease or stay the same Gases: solubility always decreases with temperature Pressure Solids: little influence Gases: solubility always increases with pressure (Henry’s law)

Saturation A saturated solution is one which is in equilibrium with undissolved solute – it has reached limit of solubility Supersaturation arises when amount of substance in solution is greater than that predicted on basis of saturation. An essential condition for the growth of crystals

Henry’s Law The number of moles of gas dissolved in a liquid is proportional to the partial pressure of the gas Exchange of CO2 and O2 in respiration depends on Henry’s Law. In the lungs, the O2 partial pressure is higher than that of CO2 In the blood, the CO2 pressure is higher after respiration

Real world applications 1: Henry’s Law and sodie pop The quantity of gas dissolved in a liquid depends directly on the pressure of that gas above the liquid Under pressure the CO2 in the liquid is kept in solution Open the cap and the CO2 rapidly escapes

Real world applications 2: The science of breathing The gas laws explain the mechanics of breathing: the transport of oxygen from the lungs and exchange with carbon dioxide produced in the body.

Measuring concentration Concentration = amount of solute/amount of solution Weight/volume percent Mass solute in g/volume of soln in mL x 100% Weight/weight percent Mass solute in g/mass solution in g x 100%

Molarity Concentration is usually expressed in terms of molarity: Moles of solute/liters of solution (M) Moles of solute = molarity x volume of solution Moles = M x V

Molarity and concentration M = moles solute/liter of solution Dilution M1V1 = M2V2 Dilution factor = V2/V1 (V2>V1) M2<M1

Example What is molarity of 50 ml solution containing 2.355 g H2SO4? Molar mass H2SO4 = 98.1 g/mol Moles H2SO4 = .0240 mol (2.355 g/98.1 g/mol) Volume of solution = 50 mL/1000 mL/L = .050 L Concentration = moles/volume = .0240 mol/.050 L = 0.480 M

Solution stoichiometry How much volume of one solution to react with another solution Given volume of A with molarity MA Determine moles A Determine moles B Find target volume of B with molarity MB

Titration Use a solution of known concentration to determine concentration of an unknown Must be able to identify endpoint of titration to know stoichiometry Most common applications with acids and bases

Colligative properties Properties that depend upon the concentration of solute particles but not their identity Vapor pressure lowering Freezing point depression Boiling point elevation Osmotic pressure

Raoult’s law When nonvolatile solute is added to solvent, vapor pressure of solvent decreases in proportion to concentration of solute Freezing point goes down Boiling point goes up

Freezing and melting are dynamic processes At equilibrium, rate of freezing = rate of melting

Adding salts upsets the equilibrium Fewer water molecules at surface: rate of freezing drops Ice turns into liquid Lower temperature to regain balance Depression of freezing point

The same model explains elevated boiling point Condensation and evaporation are dynamic processes Replacing some of the liquid water with salt reduces rate of evaporation – leads to condensation Raise temperature to recover balance

Mathematical base Freezing point depression Boiling point elevation ΔTf = kf x solute concentration Boiling point elevation ΔTb = kb x solute concentration

Units of concentration Effect depends upon number of particles not mass of particles, so concentration must be in moles. Molality (m) is used in these situations Moles solute/kg solvent Temperature independent measure of concentration

Type of solute important Covalent solute produces one particle per molecule: C6H12O6 (s) → C6H12O6 (aq) Ionic solutes produce >1 particle per formula unit: NaCl (s) → Na+(aq) + Cl-(aq) (2 particles) CaCl2(s) → Ca2+(aq) + 2Cl-(aq) (3 particles)

Osmotic pressure Transport across semipermeable membranes: Solvent particles admitted but solute particles rejected Osmosis involves passage of water molecules across a membrane

Osmosis Transport of water molecules from dilute solution to more concentrated one Imbalance of concentration provides driving force Osmotic pressure is the pressure required to oppose this flow

Osmotic pressure Osmotic pressure is written as: πV = nRT Temperature Volume No moles Gas constant

Calculating osmotic pressure But n/V = concentration in moles per liter =M π = MRT (T in Kelvin) But what molarity? Need to know moles of particles C6H12O6 = 1 mole particles NaCl = 2 moles of particles CaCl2 = 3 moles of particles Osmolarity refers to concentration of particles for osmotic pressure determination

Osmotic pressure and cells Concentration in cells depends on osmosis Concentration outside cell > inside (hypertonic) – crenation Concentration outside cell < inside (hypotonic) - hemolysis Crenation Hemolysis