Solutions

1. Solutions

2. Definitions • Solution – Homogeneous mixture of two or more substances • Solute – Substance that is dissolved • Solvent – Substance that dissolves the solute • Aqueous Solution – water is the solvent Example: NaCl(aq) = NaCl + H2O (solute) (solvent)

3. Concentration • Concentration - Quantity of solute in a given measure (volume or mass) of solution • Concentration can be expressed as: • Molarity (M) • Percent Composition/Percent Mass • Parts per million (ppm)

4. Molarity • Number of moles of solute in one liter of solution M = moles of solute liters of solution Examples: • 1.0L of 6.0M HCl contains 6.0moles of HCl • 1.0L of 1.0M HCl contains 1.0mole of HCl

5. Molarity Examples • What is the molarity of a solution that contains 1.0 mole of solute in 5.0L of solution? • What is the molarity of a solution that contains 2.5 moles of NaCl in 3500mL of solution? • What is the molarity of a solution that contains 50.0g of AgNO3 in 2.0L of water? • How many grams of H2SO4 are present in 5.0L of 3.0M H2SO4?

6. Percent Composition/Percent Mass = mass of part x 100 mass of whole

7. Percent Composition Examples • If 5.0g of a 200.g sample is aluminium, what is the percent mass of aluminum? • 10.0g of NaCl is dissolved in 150.g of water. What is the percent mass of NaCl in the solution? • If 23% of a 58.5g sample is calcium, how many grams of calcium are in the sample?

8. Parts per million (ppm) = grams of solute x 1,000,000 grams of solution

9. ppm Examples • Calculate the concentration (in ppm) if 10.0g of NaCl is dissolved in 750g of solution. • Calculate the concentration (in ppm) if 1.25g of CuSO4 is dissolved in 550g of water. • How much Fe is in a 5000.g solution, if it has an iron concentration of 500.ppm?

10. Dilution M1V1 = M2V2 M = Molarity V = Volume

11. Dilution Examples • How many milliliters of 12M HCl must be added to water to make a 300.0mL solution of 6.0M HCl? • How many milliliters of 18M H2SO4 must be used to make a 1.0L sample of 1.0M H2SO4? • 15mL of 3.0M NaOH are diluted to a volume of 500.mL. What is the resulting concentration of the diluted solution?

12. Solubility Depends on: • The nature of the solute and solvent • “Like dissolves Like” • Polar solvents dissolve polar and ionic solutes • Nonpolar solvents dissolve nonpolar solutes

13. Solubility 2. Temperature • For solids, solubility increases as temperature increases • For gases, solubility increases as temperature decreases 3. Pressure • Gases are more soluble at higher pressures

14. Factors that affect the rate of dissolving • Size of the particles • More surface area (smaller pieces) = faster • Stirring • Amount of solute already dissolved • Temperature

15. Saturated Solutions • Contain the exact amount of solute that can be dissolved at a certain temperature and pressure • Solute is added under constant conditions until no more will dissolve • If additional solute is added it will not dissolve

16. Unsaturated Solutions • Contain less solute then can be dissolved at a certain temperature and pressure • If additional solute is added it will dissolve

17. Supersaturated Solutions • Contains more solute then should be dissolved (more than the saturated solution) • A saturated solution is made at a high temperature and allowed to gradually cool • Very rare, very unstable • If additional solute is added the “extra” solute will fall out of solution (re-crystallize)

18. Solubility Curves – Table G • Saturated Solution • on the line • Unsaturated Solution • below the line • Supersaturation Solution • above the line

19. Examples • 60g of NH4Cl is dissolved in 100g of water at 70oC. What type of solution is this? • How many grams of NaCl must be dissolved in 100g of water at 40oC to make a saturated solution?

20. Examples • How many grams of KNO3 must be dissolved in 200g of water at 55oC to make a saturated solution? • A solution contains 25g of KCl in 100g of water at 50oC. How much additional KCl must be added to make a saturated solution?

21. Table F – Solubility Guidelines • Explains if a combination of ions will be soluble or insoluble in water

22. Table F Examples • Indicate if the following are soluble or insoluble • LiNO3 • Li3PO4 • NaOH • Al(OH)3 • MgCO3

23. Colligative Properties • Properties that depend on the number of particles in solution Examples: Boiling Point, Freezing Point, Vapor Pressure

24. Boiling Point / Freezing Point • The addition of a solute to a solvent causes the • Boiling point to increase • Freezing point to decrease • The great the concentration of particles the greater the effect Example: The addition of antifreeze (ethylene glycol) to your radiator increases the boiling point and decreases the freezing point

25. Boiling Point • The addition of a solute raises the boiling point of the solvent • One mole of particles raises the boiling point of water by 0.52oC

26. Freezing Point • The addition of a solute lowers the freezing point of the solvent • One mole of particles lowers the freezing point of water by 1.86oC

27. Electrolytes / Nonelectrolytes • Electrolytes dissociate (break apart) in solution • Ionic Compounds, Acids, Bases Ex: NaCl in water breaks apart into Na+ and Cl- ion NaCl Na+ + Cl- • Nonelectrolytes do not dissociate in solution • Molecular substances Ex: sugar in water C6H12O6 C6H12O6 H2O H2O

28. Dissociation Factor • How many particles (ions) the molecule dissociates into Examples: • NaCl Na+ + Cl- (df = 2) • CaCl2 (df = ) • C6H12O6 (df = )

29. Examples • Equal amounts of which will lower the freezing point of water more? a. Sugar (C6H12O6) b. Salt (NaCl) • Explain why NaCl breaks into 2 particles (df = 2), sugar does not break apart (df = 1)

30. Examples • Which would have a lower freezing point, 3.0M CaCl2(aq) or 3.0M NaCl(aq)? Explain why. • Which will have a higher boiling point,6.0M HCl or 3.0M HCl? Explain why. • Which would have a lower freezing point, 6.0M HCl or 3.0M HCl? Expain why.