1. + Ions in Aqueous Solutions and Colligative Properties Chemistry 1 (Chapter 13)

2. + Exam Analysis Next Exam… REMINDER-- Tests #2 and #3 will NOT replace Test #1 if… You have a 3 or less on your packet the day of Test #2!!

3. + Unit Objectives: 1. I can use a solubility table to predict products in chemical reactions. 2. I can explain and correctly construct molecular, complete ionic, and net ionic equations (spectator ions). 3. I understand and can apply and calculate the four colligative properties.

4. + Questions From Readings Explain what happens when an ionic substance is dissolved in water. (examples to strengthen response)

5. + Dissociation reactions … AgNO 3  KBr  BaSO 4 

6. + Questions from Reading What is a precipitation reaction? How can you determine if the reaction occurs?

7. + Precipitate Reactions (ppt) Precipitate: An insoluble solid compound is formed during a reaction. Anions are exchanged between two cations. (This is a double displacement reaction) To be a ppt. rxn, both must occur: 1. Both reactants must be aqueous (aq) 2. At least one product must be a solid (s) You MUST include phase labels with your equation. (See solubility table)

8. + Predicting Solubility of Compounds Use a solubility table to determine if a substance is going to be soluble (aqueous) or insoluble (solid) in water a) Hg 2 Cl 2 b) KI c) lead (II) nitrate

9. + Practice Solubility 8 Minutes!

10. + Rules for ppt. reactions 1.Write a balanced chemical equation. 2.Use the solubility table to place phase labels to each formula. 3.If one of the products is a solid and the reactants are aqueous the reaction is classified as a precipitate reaction. 4.If all of the products are (aq) then the reaction is NOT a ppt rxn and is classified as double displacement.

11. + Q: For each of the following decide if a ppt. will occur. A) Aqueous solutions of sodium chloride and iron (II) nitrate are mixed. B) Aqueous solutions of aluminum sulfate and sodium hydroxide are mixed.

12. + Check for Understanding For the following reactions, predict the identity of the precipitate formed. Write the correct formula of the precipitate on the space. If no precipitate is likely, write No Reaction. BaCl 2 and K 2 SO 4 ______________________ CuCl 2 and AgNO 3 ______________________ (NH 4 ) 3 PO 4 and CaS______________________ KCl and Ca(NO 3 ) 2 ______________________

13. + Ionic Equations Molecular Equation: Chemical equation in which the reactants and products are written as if they were molecular substances, even though they may exist in solutions as ions. Must include phase labels (s, l, g, aq) This provides you with the big picture Example: Al 2 (SO 4 ) 3(aq) + 6NaOH (aq) ----- 2Al(OH) 3(s) + 3Na 2 SO 4(aq)

14. + The next 2 types of equations are only completed for precipitate reactions!!! Complete ionic equation: Shows all the particles in the solution as they realistically exist. Break apart the aqueous substances into their ions. Do NOT break apart s, l, or g!! When writing a complete ionic equation include the Amount, Symbol and CHARGE!

15. + Example of a Complete Ionic Equation Molecular Equation: Al 2 (SO 4 ) 3(aq) + 6NaOH (aq) --- 2Al(OH) 3(s) +3Na 2 SO 4(aq) Complete Ionic Equation: Amount, Symbol, Charge

16. + Last step for precipitate reaction: Net ionic equation: Ionic equations that include only the particles that participate in the reaction. This tells us what substances actually formed something new in the reaction. Cross out the spectator ions

17. + Writing Net Ionic Equation Molecular Equation: Al 2 (SO 4 ) 3(aq) + 6NaOH (aq) --- 2Al(OH) 3(s) +3Na 2 SO 4(aq) Complete Ionic Equation: Amount, Symbol, Charge 2Al +3 + 3SO 4 -2 + 6Na +1 + 6OH -1 ---- 2Al(OH) 3 + 6Na +1 + 3SO 4 -2 Net Ionic Equation : What formed during the reaction?

18. + One more example of a precipitate reaction Calcium hydroxide reacts with sodium carbonate to produce calcium carbonate and sodium hydroxide Molecular Equation: Complete Ionic Equation: Net Ionic Equation:

19. + Check for Understanding Why are complete ionic equations more informative than molecular equations for reactions of ions in aqueous solutions? What is the difference between a complete/total ionic equation and a net ionic equation? Why are spectator ions left out of the net ionic equation? What substance is designated with an (s) in the net ionic equation? What state designation do the other substances have? Why is it necessary to balance the molecular equation before writing the total and net ionic equation?

20. + Practice, Practice, Practice Unit Packet

21. + Check for Understanding Explain the difference between ionization and dissociation. What determines how much a solute ionizes in solution? Explain how to tell the difference between a strong electrolyte and a weak electrolyte.

22. + Check for Understanding Explain the difference between ionization and dissociation. Ionization occurs when ions are formed from the solute particles due to the action of the solvent. Dissociation occurs when an ionic compound dissolves into ion. Difference is ionization’s ions are formed from molecular compounds, not ionic compounds. What determines how much a solute ionizes in solution? The strength of the solute molecules The strength of attraction between solute and solvent Explain how to tell the difference between a strong electrolyte and a weak electrolyte. The degree of ionization or dissociation is what determines strength of electrolyte, not the amount of solute dissolved.

23. + Colligative Properties 1. Vapor Pressure Lowering 2. Freezing-Point Depression 3. Boiling-Point Elevation 4. Osmotic Pressure

24. + Colligative Properties Properties that depend upon the concentration of solute particles, but not the properties of the solute. Vapor Pressure Lowering Boiling Point Elevation Freezing Point Depression Osmotic Pressure

25. + Vapor Pressure Lowering Vapor Pressure: the pressure exerted in a closed container by liquid particles that have escaped the liquid’s surface and entered the gaseous state. Adding a solute to a solvent lowers the solvent’s vapor pressure.

26. + Vapor Pressure Lowering Fewer liquid molecules are available to escape from the liquid to become gaseous in any solution vs. the pure solvent. Does not depend upon the properties of the solute, only the concentration=colligative. This lowers the vapor pressure in solutions as compared to pure solvents When the vapor pressure is lower, the solution remains liquid over a larger temperature range (solutions have lower freezing points and higher boiling points than pure solvents). True for nonvolatile substances-substances with little tendency to become a gas under existing conditions.

27. + Freezing Point Depression

28. + Freezing-Point Depression Molal (m) freezing-point constant (K f ) is the freezing-point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solution For water, this value is -1.86°C/m This means the freezing point of a 1m solution of any nonelectrolyte solute in water is 1.86°C lower than the freezing point of water (0°C). Using this number, freezing-point depression of any solution can be determined. Freezing-point depression/ Δ t f is the difference between the freezing point of the pure solvent and a non-electrolyte solution

29. + Freezing-Point Depression Δ t f = K f mi Formula can be used to solve for the freezing- point of the molal concentration of a solution. Δ t f = the change in the freezing point temperature K f = the molal freezing point constant (pg. 448 in text) m = the molality of the solution i = the number of ions present (covalent = 1, ionic = total number of ions)

30. + Solute Impact on Colligative Properties SoluteIons# Particles NaCl HCl MgCl 2 Ca 3 (PO 4 ) 2 K 2 SO 4 C 12 H 22 O 11

31. + Example Problem 1 Calculate the freezing point depression of adding 150 g of NaCl into 250 g of water. Calculate molality Apply freezing-point depression equation

32. + Example Problem 2 Calculate the freezing point depression of adding 100. g of CH 3 OH into 500. g of the non-electrolyte, camphor. Calculate molality Use the freezing-point equation

33. + Boiling Point Elevation

34. + Boiling-Point Elevation The boiling point is the temperature at which the vapor pressure of a liquid is equal to the prevailing atmospheric pressure. The boiling point of a solution is higher than the boiling point of the pure solvent (due to lower vapor pressure, less particles available to become gaseous). The molal boiling point constant, K b, is the boiling-point elevation of the solvent in 1-molal solution of a nonvolatile, nonelectrolyte solution In water, K b = 0.51°C/m The boiling point elevation, Δ t b, is the difference between the boiling point of the solution and the boiling point of the pure solvent

35. + Boiling-Point Elevation Δ t b = K b mi Use this equation to solve for boiling-point elevation or molal concentration of solutions Δ t b = the change in the boiling point temperature K b = the molal boiling point constant (pg 448 text) m = molality of the solution i = the number of ions present (covalent =1, ionic = total number of ions)

36. + Example 1 Calculate the boiling point elevation of adding 150.g of NaCl into 250. g of water. Calculate molality Use boiling-point elevation equation New boiling point?

37. + Example 2 Calculate the boiling point elevation of adding 100. g of CH 3 OH into 500.g of the non-electrolyte, camphor. Calculate molality Use the boiling-point elevation equation

38. + Osmotic Pressure The movement of a solvent through a semipermeable membrane from the side of lower solute concentration to the side of higher solute concentration is osmosis. In a U-tube, the side with higher concentration would move up as water moves to that side Only occurs in solution if each side contains a different concentration Osmotic pressure is the external pressure that must be applied to stop osmosis The greater the concentrations of solution (regardless of solute), the greater the osmotic pressure of the solution.

39. + Electrolytes and Colligative Properties Electrolytes behave differently than non-electrolytes and result in changes in colligative properties. The freezing point is lower and the boiling point is higher. Due to the separation of the ions when electrolytes are dissolved in a solution. 1 mole of NaCl, for example, has 2 particles (Na+ and Cl-), thus results in 2 moles of solute Nonelectrolyte molecules stay together in solution. Some electrolytes may form more than 2 particles, thus change the colligative properties more.

40. + Electrolytes and Colligative Properties Actual values for electrolyte solutions The actual values are not quite those that are calculated. The freezing point/boiling point are not quite twice the expected or three times the expected. Due to attractive forces between the ions– interferes with movement and affects freezing/boiling Compounds with ions that have higher charges show less change than those with lower charges (+2, +3, -2 ….vs -1, +1)

41. + Video Review Boiling Point Elevation Freezing Point Depression Osmotic Pressure Video 1 Video 2 Vapor Pressure Lowering The Chemistry of Cars

42. + Individual Practice Unit Packet

43. Mini-Lab20 Minutes! Ice Cream Anyone?

44. + Here's what happens. When ice cream is made the old-fashioned way, rock salt (big chunks of salt crystals) is mixed with ice. Only a little water melts before some of the dissolved salt lowers its freezing point. Now when the ice wants to melt, it can absorb lots of heat from the water to do it, and the water still will not freeze. (In the recipe you will try, the temperature of the water may decrease to almost minus 20 degrees C and still be liquid! Pretty cool!) The water is very much colder than the cream mixture. Because heat flows from hot things to cold things, the cream now loses its heat to the water and rapidly cools down. http://www.eduplace.com/kids/hmxs/g6/weatherwater/cricket/s ect2cc.shtml http://www.eduplace.com/kids/hmxs/g6/weatherwater/cricket/s ect2cc.shtml Copyright © Houghton Mifflin Company. All rights reserved.