Chapter 12 Solutions

1. Chapter 12 Solutions • Solute (smaller amount) • + Solvent (larger amount) • Solution • Solute Solvent • S L (salt and water) • L L (alcohol and water) • G L (carbonated drink –CO2) • S S (any alloy – sterling silver- copper and silver ) • G G ( air) • Water is universal solvent • (aq) dissolved in water

2. Energy • Separate solute and solvent particles –endothermic (+ ΔE) • Solvation -attraction between solvent and solute particles exothermic (- ΔE) • Hydration (when water is solvent) • Hydration energy (supplies energy to break bonds)

3. Will a solid dissolve? • E to separate solute + Δ E • E to separate solvent + ΔE • Energy of solvation (hydration) - ΔE • Usually if overall Δenergy is negative the solute will dissolve • Usually if overall Δenergy is positive the solute will not dissolve

4. Minimum energy (exothermic) • Maximum entropy (disorder) dissolving increases entropy • Solids can dissolve if the Δ energy is + but small in value because of increasing in entropy • Solids do not dissolve If the Δ energy is + and large

5. Like dissolves Like • Polar sub dissolve Polar sub • Non polar sub dissolve Non polar sub • Polar and non polar usually do not dissolve. Why?

6. Soaps and detergents • Soaps have a long nonpolar tail (hydrophopbic –does not like water) • and a polar end (hydrophilic –does like water) • Nonpolar substances like grease dissolve in the hydrophobic tail • The hydrophilic end dissolves in water • See page 450 and 451

7. Solubility- the maximum amount of solute that can be dissolved in agiven amount of solvent Most solids will increase in solubility as the temperature increases When the maximum amount of solute is dissolved the solution is SATURATED

8. Solubility of a gas in liquid (temp) • Gases will decrease in solubility as the temperature increases

9. Solubility of a gas in liquid (pressure) • Gases will increase in solubility as the pressure increases

10. Concentrations of Solutions • Molarity (M) = moles of solute/liter of solution • % by mass = (grams of solute/grams of solution)x100 % • % by volume =( volume of solute/volume of solution) x100 % • molality (m) = moles solute/kg of solvent

11. Explain how to prepare 400. ml of a .150 M solution of NaCl. (from solid NaCl) • Moles? • Grams? • To make 400 mL of a .150 M solution of NaCl add ___________ grams of NaCl to a flask and add enough water until the total volume of solution is __________ mL

12. Explain how to prepare 400. ml of a .150 M solution of NaCl. (from a stock solution of. 500M) dilution NaCl) • Moles? • Volume? • To make 400 mL of a .150 M solution of NaCl from a .500 M solution of NaCl measure _____________mL of the stock solution and add enough water until the total volume of solution is __________ mL

13. Dilution M stx V st = M dilx V of dil • Moles( in stock) = Moles (in dilute) • Explain how to prepare 400. ml of a .150 M solution of NaCl. (from a stock solution of. 500M) dilution NaCl)

14. Molarity of ions in solution • What is the concentration of each ion in a .25 M sodium phosphate solution?

15. Stoichiometry and Molarity • What volume of .80M Al(NO3)3 is needed to completely react 2.0 grams of calcium? • Balanced equation • Find moles of Ca • Find moles of Al(NO3)3 • Convert to volume

16. Heating/cooling curve of pure substance

17. Boiling/Freezing points of solutions • Pure substances –definite B.P. and F.P solutions do not (B.P is higher, F.P is lower) • What happens when a solute is added to a solvent? (salt to water)

18. What does the solute do to the V.P? • Adding a solute lowers the vapor pressure of the liquid which raises the B.P

19. Heating/cooling curve of pure substance

20. Colligative property • Property that depends on the number of particles • Ex. Boiling point elevation and freezing point depression • ΔTb = Kbx moles of solute particles/kilograms of solvent • ΔTf = Kfx moles of solute particles/kilograms of solvent

21. Molecular vs ionic solutes • C6H12O6(s) -------> C6H12O6 (aq) • NaCl(s) ------->Na+(aq) + Cl- (aq)

22. If 5.85 gNaCl are added to 100 g of water what temperature will the water begin boiling? • How many moles of NaCl? • How many moles of particles? • ΔTb = Kbx moles of solute particles/kilograms of solvent

23. Heating/cooling curve of pure substance

24. Line A-B heating (or cooling) of solid- temperature change q = cm ΔT K.E changing q= heat energy c= specific heat of solid m= mass of solid ΔT= change in temperature Line B-C melting ( or freezing) - heat of fusion- amount of energy added to melt (or removed to freeze) at the substance’s M.P. (F.P) energy may be given per gram or mole- ex. kJ/mole or kJ/g P.E change No K.E change -no temperature change Line C-D heating (or cooling) of liquid- temperature change q = cm ΔT K.E changing q= heat energy c= specific heat of liquid m= mass of liquid ΔT= change in temperature

25. Line D-E boiling( or condensing) - heat of vaporization –amount of energy added to boil (or removed to condense) at the substance’s B.P. energy may be given per gram or mole- ex. kJ/mole or kJ/g P.E change No K.E change-no temperature change Line E-F heating ( or cooling) of gas- temperature change q = cm ΔT K.E increasing q= heat energy c= specific heat of gas m=mass of liquid ΔT= change in temperature

26. For H2O specific heat (c) Solid = 2.13 J/goC Liquid=4.18 J/goC Gas=2.01 J/goC heat of fusion = 334 J/gram or 6.01 kJ/mol heat of vaporization =2228 J/gram 40.1 kJ/mole M.P. (F.P.) = 0 oC B.P. = 100 oC Ex. If you have 100.0 grams of water at -5.0 C, how much energy will it take to change it to 25o C? Sketch a heating /cooling graph and determine how many energy changes the substance will undergo