1. Chapter 13 Solutions Chemistry 100

2. Definition When one substances is dispersed uniformly throughout another at the molecular level, we have a solution. Each substance in a solution is a component. The solvent is generally the component present in the greatest amount. The other components are called solutes

3. Types of Solutions All gases mix. Generally, we do not think of a mixture of two gases as a solution. Gases dissolve in liquids - oxygen in water Liquids dissolve in liquids - alcohol in water Solids dissolve in liquids - sugar in tea Gases dissolve in solids - hydrogen in palladium Liquids dissolve in solids (rare) - Hg in silver Solids dissolve in solids - alloys (brass, solder, etc)

4. Energy considerations Begin with solid S and liquid L End with S dissolved in L The particles in the solid must separate to dissolve.  Overcome intermolecular forces or electrostatic forces in ionic compounds! Let  H 1 be the enthalpy for this process

5. Solvent-SolventInteractions The solvent molecules must separate to allow room for the solute molecules. Intermolecular forces must be overcome in L Let  H 2 be the enthalpy for this process.   H 1 and  H 2 will be positive - energy is added to the system

6. Solute-Solvent Interactions The solute and solvent particles interact. Example: Ions dissolved in water give rise to ion-dipole interactions as the water molecules surround the ions.  Hydration Let  H 3 be the enthalpy for this process   H 3 will be negative - energy is released when two objects that attract get closer together

7. Enthalpy of Solution For the overall process of S + L  Solution  H soln =  H 1 +  H 2 +  H 3 The first two will be endothermic processes (positive) values since they need to take in energy. The last process is exothermic (negative value) since heat is released when two attracting molecules are brought together.

8. Like Dissolves Like The enthalpy of solution will be  positive if  H 3 has a small (negative) value.  negative when  H 3 has a large (negative) value.  H 3 will have a  large value with an ionic solid and a polar liquid  small value with a non-polar solute is and a polar solvent  If  H soln very endothermic the solution will not form

9. Spontaneity A process is said to be spontaneous when it occurs without external help Endothermic process can be spontaneous provides  H is not too large (more later) Example: NaOH plus water:  H = -44 kJ/mole NH 4 NO 3 plus water:  H = +26 kJ/mol

10. Entropy of Solution State A consists of solid and liquid as separate phases State B consists of the solid dissolved in the liquid. State B has a higher entropy! Which has least energy? That depends on the the  H values discussed earlier - the net value is the enthalpy of solution If enthalpy of solution is negative (exothermic), state B (solution) is preferred If it is only slightly endothermic, state B can still be preferred

11. Concentration terms Dilute - not a lot of solute. Concentrated - a large amount of solute. Chemists try not to talk about a strong solution - as in: a strong cup of tea Concentration can be expressed quantitatively is many ways: molarity, molality, percentage, mole fraction, etc

12. Molarity The molarity is the number of moles of solute in 1 litre of solution.  M = moles of solute / volume (litres) solution A 1 M solution is fairly concentrated (often 100 s of grams of solute in a litre) Millimolar:  A 0.001M solution is 1  10 -3 moles/L  or 1 millimole/L or 1mM

13. Molality The molality is the number of moles of solute dissolved in 1 kg of solvent.  m = moles of solute / volume (litres) solution Molalities are related to mass fractions (% mass, see later). Molality is independent of the temperature.  A 1.0 m (for molal) solution is 1.0 moles of solute dissolved in 1.00 kg of solvent.

14. Normality Chemists has stopped using normality as a concentration unit but the health profession has not. Simple rule to convert molarity to normality  Multiply molarity by ion charge to get normality  A 0.12 M Na + solution is a 0.12 N Na + solution  A 0.25 M Ca 2+ solution is a 0.50 N Ca 2+ soln. This will not be on the exam but nurses need to know it.

15. Mass percentage A solution that is 36% HCl by mass, contains 36 g HCl in 100 g of solution. Milk is not a true solution, but 2% milk refers to 2 g of milk fat in 100 g of milk. You may see figures on packing such as 3% w/w where w is short for weight (mass)

16. Parts per million Percentage refers to parts in 100. Thus 3% means 3 parts solute in 100 parts of solution. Parts per million - how many parts of solute in a million parts of solution

17. ppm Let our solution have that has a Ag + ion concentration of 3 ppm 3 g Ag + in 1,000,000 g of solution 3 mg Ag + in 1,000 g of solution (1/1000th of each) But 1,000 g of water solution has a volume of approximately 1 litre. So our solution has 3mg Ag + / litre ppm is approximately the same as mg/L

18. Ppb - parts per billion This unit is used for highly toxic materials We saw that 1 ppm = 1 mg / L In the same way, 1 ppb = 1  g /L where 1  g is 1 microgram or 1  10 -6 g

19. Percentage by volume When a solute is a liquid, solution concentrations s are sometimes expressed as percentages by volume Thus a 20%v/v solution of hydrogen peroxide contains 20 ml of H 2 O 2 in every 100 ml of solution.

20. Mole fraction If a solution consists of m A moles of A, m B moles of B, m C moles of C, etc. Mole fraction of A X A = m A / (m A + m B + m C ) Clearly, X A + X B + X C = 1

21. Terminology Solute Solvent Solubility Saturated solution Unsaturated solution

22. Factors That Affect Solubility Solute-solvent interaction Pressure effect when solute is a gas Henry’s Law: C g = kP g k is a constant for the gas at a given temp. Temperature effect most solids are more soluble in warm water gases are less soluble in warm water

23. Colligative Properties Properties that depend on solute concentration (collection) not the kind of solute.  Depression of the vapour pressure  Depression of the freezing point  Elevation of the boiling point  Osmosis

24. Explanation of lowering of VP There are serious problems with this so-called explanation

25. Depression of VP Raoult’s Law: How VP of solution relates to VP of pure solvent P A = X A P 0 A Note that when X A = 1 (pure solvent), P A = P 0 A Solutions that obey Raoult’s law are called ideal solutions

26. Colligative Properties

27. Freezing & boiling point changes For dilute solutions of non-volatile solutes  T f = mK f K f varies with solvent  T b = mK b K b varies with solvent For electrolytes, m is the molality of particles: 1m NaCl is 2m in particles (n=2) 1m CaCl 2 is 3m in particles(n=3) Use  T = nmK for solutions of electrolytes

28. Osmosis

29. The movement of water through a semi- permeable membrane from dilute side to concentrated side  the movement is such that the two sides might end up with the same concentration Osmotic pressure: the pressure required to prevent this movement

30. Osmosis Ideal gas law: PV = nRT Osmosis Law:  V = nRT  = (n/V)RT  = MRT Remember to use osmotic pressure in ATMS so use can use R = 0.0821 Lt.-atm/mole-K

31. Terminology Isotonic: having the same osmotic pressure Hypertonic: having a higher osmotic pressure Hypotonic: having a lower osmotic pressure Hemolysis: the process that ruptures a cell placed in a solution that is hypotonic to the cell’s fluid Crenation: the opposite effect

32. Colloids In a solution the particles are small: single molecules or ions In colloids the particles are large: collection of molecules or a giant molecule Particle size 10 to 2000Å Colloids scatter light because of large particles

33. Types of Colloids Gas as the dispersion agent Dispersed phase Gasnone (all are solutions) Liquidaerosolfog Solidaerosolsmoke

34. Types of Colloids Liquid as the dispersion agent Dispersed phase Gasfoamwhipped cream Liquidemulsionmilk Solidsolpaint

35. Types of Colloids Solid as the dispersion agent Dispersed phase Gassolid foam marshmellow Liquidsolid emulsionbutter Solidsolid solruby glass

36. Stability of colloids What prevents the dispersed particles from coalescing? Hydrophobic colloids: the dispersed phase reacts with the water with forces similar to those present in solutions Hydrophobic: The particles have ions on the surface; the charges keep the particles from coming together

39. Detergents Long chain molecules One end is hydrophobic The other is hydrophilic The hydrophobic ends attach to dirt particles (grease) The hydrophilic ends keep the particles in solution