Properties of Solutions

1. Properties of Solutions

2. 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

3. 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)

5. 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

6. 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)

7. 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)

9. 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

10. 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)

11. 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

12. 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

13. 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

14. 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.

15. 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%

16. 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

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

18. 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

19. 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

20. 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

21. 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

22. 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

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

24. 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

25. 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

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

27. 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

28. 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)

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

30. 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

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

32. 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

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