1. Solutions/ Concentrations
Georgia Performance Standard SC7: Students will characterize the properties that describe solutions and the nature of acids and bases.

2. Polar bonds & polar overall
The Water Molecule
Triatomic
Covalent
Polar bonds & polar overall
Bent at 105° angle

3. Intermolecular Forces
Polar molecules are attracted to one another by dipole forces
Water is attracted to other water molecules by a special dipole force, a hydrogen bond

4. Solution=Homogeneous Mixture
Water Solutions
“Chemically pure water never exists in nature because water dissolves so many substances.” textbook
Universal solvent
Aqueous solutions
Solution=Homogeneous Mixture

5. What happened to the ionic compound?
Water Solutions
Ionic compounds most readily dissolve in water due to extreme polarity
Polar covalent compounds also dissolve in water
Nonpolar compounds don’t
What happened to the ionic compound?

6. Solvation of Ionic Compounds

7. A Few Exceptions... Calcium carbonate Remember the solubility rules...
In some ionic compounds, the ions are so attracted to each other that they won’t break apart and dissolve.
These are INSOLUBLE ionic compounds.
Calcium carbonate

8. Solvation of Covalent Compounds
Covalent compounds do NOT break apart in water when dissolving.
Solvation of covalent compounds means that each solute molecule is surrounded by water molecules.

9. Conductors
In general, aqueous solutions of ionic compounds are electrolytes.
Generally, aqueous solutions of covalent compounds are nonelectrolytes.

10. Solution Vocabulary
Solute: Dissolves in the solvent Soluble: Able to be dissolved in the solvent (applicable to any states of matter) Insoluble: Unable to be dissolved in the solvent (applicable to any states of matter) Miscible: Able to be dissolved in the solvent (applicable to liquid/liquid solutions) Immiscible: Unable to be dissolved in the solvent (applicable to liquid/liquid solutions)

11. Solubility Vocab
Unsaturated: less than maximum amount of solute is dissolved in the solvent
Saturated: maximum amount of solute is dissolved in the solvent
Supersaturated: special conditions have been created to dissolve more than maximum amount of solute in the solvent

12. Determines IF Solute Will Dissolve...
The nature of the solvent and solute governs whether a solute will solvate in a particular solvent. Specifically, the nature of the intramolecular bond. Polar molecules will solvate with polar molecules. Nonpolar molecules will solvate with nonpolar molecules. BUT, polar and nonpolar will not form solutions together.

13. Determines Speed of Dissolving
BRING SOLUTE IN CONTACT WITH SOLVENT
Agitation: Create more collisions mechanically
Temperature: More kinetic energy creates more collisions
Surface Area: Dissolving process is a surface phenomenon, the more surface of the solute that is exposed the faster the solvation

14. Determines How Much Will Dissolve
SOLUBILITY: HOW MUCH WILL DISSOLVE
Temperature: solubility of solid solute increases as the temp. increases; solubility of gaseous solute decreases as temp increases ex. Hot water bubbles, thermal pollution
2. Pressure: solubility of gaseous solute increases as the pressure increases
-Henry’s Law ex. Soft drinks

15. Temperature & Solubility
Higher the temperature, the more solid will dissolve in a liquid
Higher the temperature, the less gas will dissolve in a liquid
Why do soft drinks taste better cold?

16. Henry’s Law: Pressure & Solubility
At a given temperature, the solubility of a gas is proportional to the pressure of the gas above the liquid.
Page 506
The higher the pressure, the more carbon dioxide will dissolve in the syrup giving a less “flat” taste.

17. Solubility Graph

18. Concentration of Solutions
The concentration of a solution is a measure of how much solute is dissolved in a specific amount of solvent or solution. Molarity: most common units of solution concentration; # of moles solute dissolved in one liter of the solution

19. Molarity: moles of solute liter of solution
If given grams, remember to change it to moles.

20. Example 15.3 (page 531) Example 15.4 (page 532)
Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50L of solution.
Example 15.4 (page 532)
Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl into enough water to make 26.8 mL of solution.

21. molality: moles of solute kg of solvent
The most common concentration term in chemistry is Molarity (M), but chemists also report concentration in molality (m) sometimes.
molality: moles of solute
kg of solvent
If given grams, remember to change it to moles.

22. Practice Molality
The front of the last page in the calculations packet refers to Molality. Complete 1 (a), 2 (a), 3 (b), and 4 (b).
NOTE: Questions 3 and 4 will require you to manipulate the equation. Question 4 might even require that you change the units of your answer at the end of the calculation.

23. Preparing a Solution in the Lab
You’ve learned about concentration (Molarity and molality). Therefore, you should be able to create your own solutions for use in the lab from now on.
Example:
0.5 M HCl reacts with Mg(s)
If I gave you HCl powder, what would you do to make the solution?
Watch these kids, and assess your plan.
Choose a lab group and station. Draw an assignment from the cup, and make the solution. Write your steps (including materials) as you go.
ALSO: Solve 15-2 Practice Problems WS 6, 8, 10-12,

24. Diluting Solutions
If you already have a solution molarity, but you want a different molarity: Example: Your lab asks you to use 250 mL of 0.25M HCl, but you only find a jug of 6M HCl in the stockroom. USE THIS EQUATION: M1V1 = M2V2

25. Example 15.8 (page 539)
What volume of 16 M sulfuric acid must be used to prepare 1.5 L of a 0.10 M H2SO4 solution?
*Be sure to solve the Dilution Worksheet in your practice packet!

26. Using Molarity as a Conversion Factor
You have learned to calculate molarity using the equation:
Molarity = moles of solute / liter of solution
Did you realize that the calculation is simply a ratio of solute to solution?
Ratios (ie mole ratios, energy to mole ratios) can be used as conversion factors in stoichiometry.
THEREFORE, molarity can be written into our dimensional analysis charts to solve stoichiometry problems.

27. Solutions Stoichiometry
What volume of 1.5 M HCl is needed to react with 21.5 grams of NaOH?
What is the molarity of a solution of H2S if 48.5mL are required to titrate 35.6mL of 0.35M Fe(OH)3 solution?
A white precipitate forms when 200 mL of 0.200M K3PO4 solution is mixed with 300 mL of M CaCl2 solution. What mass of precipitate will form?
BE SURE TO SOLVE THE PRACTICE PROBLEMS IN YOUR PACKET!

28. Diluting A Solution Lab
MUST HAVE NOTEBOOK APPROVED BEFORE MOVING TO STATION
Purpose: See group assignment
Materials: What else would you need?
Safety: What should you think about?
Procedure: Part One: Making the solution from “scratch”. Part Two: Dilution. Write it out step by step!
Results & Conclusions: How do you need to revise your procedure?

29. Colligative Properties
Vapor Pressure lowering
Freezing Point Depression
Boiling Point Elevation
Physical properties of solution are different from the physical properties of the solvent. Some properties are different simply because there are “foreign” particles (solute) in the solvent. Colligative properties of solutions depend only on the number of solute particles.
Colligative Properties

30. More solute particles means less solvent particles able to escape
to become vapor!

31. Boiling Point Elevation
More solute particles means that vapor pressure is lower which means that more kinetic energy is needed to make vapor pressure equal atmospheric pressure!
Change in temperature is calculated:
∆Tb = Kb m i
Since boiling point increases...ADD the change.
Kb will be given.

32. Freezing Point Depression
Liquid particles get into an orderly pattern to become a solid. The solute particles disrupt the orderly pattern causing more kinetic energy to be drawn from the solution for it to freeze!
∆Tf = Kf m i
Since freezing point decreases...SUBTRACT the change.
Kf will be given

33. About.com Chemistry
Ice has to absorb energy in order to melt, changing the phase of water from a solid to a liquid. When you use ice to cool the ingredients for ice cream, the energy is absorbed from the ingredients and from the outside environment (like your hands, if you are holding the baggie of ice!). When you add salt to the ice, it lowers the freezing point of the ice, so even more energy has to be absorbed from the environment in order for the ice to melt. This makes the ice colder than it was before, which is how your ice cream freezes. Ideally, you would make your ice cream using 'ice cream salt', which is just salt sold as large crystals instead of the small crystals you see in table salt. The larger crystals take more time to dissolve in the water around the ice, which allows for even cooling of the ice cream.

34. Molality Another unit for concentration
m=moles of solute per kilogram of solvent
page 520

35. Solutions Stoichiometry
A white precipitate forms when 200 mL of 0.200M K3PO4 solution is mixed with 300 mL of M CaCl2 solution. What mass of precipitate will form?
What volume of 1.5 M HCl is needed to react with 21.5 grams of NaOH?
What is the molarity of a solution of H2S if 48.5mL are required to titrate 35.6mL of 0.35M Fe(OH)3 solution?