1. Unit 6 – Solutions and Solubililty
Text – Chapters 6 and 7

2. Solutions and Solubility - Definitions
aqueous solutions
solute
solvent
electrolyte
cation
anion
ionization
dissociation
saturated solution
non-saturated solution
supersaturated solution

3. Dynamic Equilibrium
For salts, when a solution is saturated, there is a point when the production of ions (dissolving) is in equilibrium with the ions coming back together to form the salt.
This is called dynamic equilibrium.
The more ions that form means the more soluable the substance is.
The less ions that form means the substance is less soluable.

4. Properties of Liquids
Liquids have their particles further apart than a solid but closer than a liquid
The spaces allow the movement of particles
As a result, liquids can take the shape of their container
Since the particles are further apart, the intermolecular forces are weaker than most solids
This affects the boiling point and vapour pressure of a liquid.

5. Properties of Liquids Boiling point Vapour Pressure
The temperature in which the vapour pressure of the substance must equal or be greater than the surrounding pressure on the surface of the liquid and the substance changes from a liquid to a gas
Vapour Pressure
The pressure exerted by a vapour at equilibrium with its liquid phase in a closed system

6. Properties of Liquids
Substances with high vapour pressures have low boiling points
If the vapour pressure is high, this means that a lot of the liquid is turning into a gas. The intermolecular forces are weak and the energy needed to boil is low.

7. Properties of Liquids
Substances with low vapour pressures have higher boiling points
If the vapour pressure is high, this means that a lot of the liquid is turning into a gas. The intermolecular forces are weak and the energy needed to boil is low.
Ex. Vapour pressure of pure water
0oC = 0.6 kPa
100oC = kPa

8. Properties of Liquids
In order for a substance to boil, energy is added to the substance and the particles move faster.
If the energy added allows the particles’ vapour pressures to be equal to the surrounding pressure, the particles will escape the surface of the liquid
If the energy added does not increase the energy enough, the substance will not boil

9. Properties of Liquids
Why can you not cook an egg on a very high mountain?
How does a pressure cooker work?

10. Factors that Affect Solubility
Intermolecular Forces
The attraction between molecules is increased as the polarity increases.
The most polar substances are ionic substances
When an ionic substance comes in contact with water, the polar ends of the water attract the ions, surrounding the ions and separating them.
The process of surrounding the ions is called solvation and the breaking apart of the ions is called ionization or dissociation.

11. Factors that Affect Solubility
For non polar substances, the slight polar ends attract to each other (called London Dispersion Forces)
This is why we can make solutions of things like gasoline and oil, for two stroke engines
Therefore the rule is “Like dissolves like”, as polar solutes and solvents will dissolve each other and non-polar will dissolve non-polar.
How about something, like alcohol that has both a polar and non-polar end?

12. Factors that Affect Solubility
Temperature
Increase in Temperature increases the solubility of solids (more solute can be dissolved)
If the heat given off in the dissolving reaction is less than the heat required to break apart the solid, the net dissolving reaction is endothermic (energy required).
The addition of more heat facilitates the dissolving reaction by providing energy to break bonds in the solid.

13. Factors that Affect Solubility
Temperature (cont)
Decrease in Temperature increases the solubility of gases (more solute can be dissolved)
The reason for this gas solubility relationship with temperature is very similar to the reason that vapor pressure increases with temperature.
Increased temperature causes an increase in kinetic energy. The higher kinetic energy causes more motion in molecules which break intermolecular bonds and escape from solution.
Decreasing the temperature lowers the KE and slows the gas particles, lowering the vapour pressure

14. Factors that Affect Solubility
Nature of the solute
Intermolecular force of attraction will determine the degree of solubility
Water = more polar more soluble
Non-polar solvent = less polar, more soluable

15. Factors that Affect Solubility
Pressure (Gases only)
Liquids and solids exhibit practically no change of solubility with changes in pressure. Gases increase in solubility with an increase in pressure.
Henry's Law states that: The solubility of a gas in a liquid is directly proportional to the pressure of that gas above the surface of the solution.
If the pressure is increased, the gas molecules are "forced" into the solution since this will best relieve the pressure that has been applied. The number of gas molecules dissolved in solution has increased as shown in the graphic on the left.

16. Factors that Affect Solubility
Examples
Tea
Chemical cold packs
Soft drinks
Scuba diving and the bends

17. Properties of Aqueous Solutions
One phase
Aqueous solutions are homogeneous mixtures
The solute and solvent are indistinguishable from each other (one clear phase is observed)
The solute particles are surrounded and separated by the water molecules and are in solution

18. Properties of Aqueous Solutions
Conduct electricity
Aqueous solutions are electrolytes
This means that when a solute is put into solution, it ionizes into a cation and an anion
When positive and negative poles are put into a solution, electricity travels, as the cation migrates to the negative pole and the anion migrates to the positive pole

19. Properties of Aqueous Solutions
Acidity and basicity
Acids and bases ionize in water.
When acids ionize, H+ is put into solution, making the solution acidic
When bases ionize, OH- is put into solution, making the solution basic
Some salts are acidic, basic or neutral (more in Grade 12)

20. Properties of Aqueous Solutions
Changes in the Freezing Point (Freezing Point Depression)
Freezing Point Depression
The temperature at which water freezes or melts is lower than 0oC
Solutions have a lower freezing point than pure water

21. Properties of Aqueous Solutions
Reason
For water to freeze, the overall energy of the water molecules needs to be low
At 0oC the particles move slower and the intermolecular forces allow the molecules to go close together, forming a solid
When solute is added it gets in between the water molecules and does not allow the intermolecular forces to attract to bring the molecules together

22. Properties of Aqueous Solutions
Therefore the temperature has to be lowered more to slow the water molecules and solute particles to allow the intermolecular forces to interact between molecules
The solute does not melt ice, but does not allow it to freeze

23. Properties of Aqueous Solutions
Changes in the Boiling Point (Boiling Point Elevation)
Boiling Point Elevation
The temperature at which water boils or condenses is higher than 100oC
Solutions have a higher boiling point than pure water

24. Properties of Aqueous Solutions
Reason
For water to boil, the overall energy of the water molecules needs to be high for the vapour pressure to equal the atmospheric pressure
At 100oC the particles move faster and the intermolecular forces are broken allowing the molecules to move further apart to form a gas
When solute is added it gets in between the water molecules and the surface of the water
The particles cannot escape the surface and move further apart

25. Properties of Aqueous Solutions
Therefore the temperature has to be raised more to speed up the water molecules and increase the vapour pressure of the water molecules
As a result, the water molecules get past the solute particles and escape from the surface of the water

26. Properties of Aqueous Solutions
Examples
“Salt” used in the winter time
Making ice cream
Antifreeze

27. Solution Concentrations
Of all the concentration units that are in general use, only four are likely to be encountered in an introductory chemistry course:
mass percent
volume percent
mole fraction
moles per liter (mol /L) Moles per liter, or molar concentration, is called molarity (M).
Molarity is the most useful

28. Solution Concentrations
Mass Percent
Example - A solution of sodium hydroxide is made up by dissolving 10.0g of sodium hydroxide in g of water. Find the mass percentage of sodium hydroxide in the solution.
Example - Water containing more than 50 ppb of lead is unfit to drink. A certain sample of water is found to contain 8.5 x 10-5 g of lead in 1.0 L of water. Is it safe to drink? (Density of water is 1.0 g / mL)

29. Solution Concentrations
Volume Percent
Example - A total of 25.0 mL of alcohol is placed in a container and sufficient water is added to bring the volume of the solution up to 125 mL. Find the percent alcohol by volume.
A bottle of vodka is 30% by volume. How many ml of alcohol are in a 750 ml bottle?

30. Solution Concentrations
Mole Fraction
Example - A solution is made up of 23.0 g of ethanol (C2H5OH) and 18.0 g of water. What is the mole fraction of ethanol in the solution?

31. Solution Concentrations
Molarity
Example - A solution contains 5.85 g of sodium chloride dissolved in 5.00 x 102 mL of water. What is the concentration of the sodium chloride in mol/L?
Example - What mass of potassium hydroxide is required to prepare 600 mL of a solution with a concentration of mol/L?
Example - A solution of magnesium chloride of mol/L is required. What is the maximum volume of solution that we can prepare if we have only 87.8 g of solid magnesium chloride?

32. Solution Concentrations
Concentrations of Ions
We can figure out the concentrations of the ions that ionize when put in solution.
It is important to know the concentrations to calculate precipitates.
Example – What is the concentration of sodium ion and chloride ion in a 0.5 M solution of sodium chloride?

33. Solution Concentrations
Suppose we have a solution containing 15.6 g of magnesium chloride in 1.25 L of solution. How can we find the concentrations of the magnesium ions and the chloride ions?
Example – g of aluminum sulfate are dissolved in 250 ml of water. What is the [SO4-2]?
Example - A solution of volume 525 mL containing 6.78 g of calcium bromide (CaBr2) is mixed with 325 mL of a solution containing 11.4 g of potassium bromide (KBr). What is the concentration of bromide ions in the resulting solution? Assume the volumes are additive (i.e. the final volume is the sum of the two initial volumes).

34. Solution Concentrations
Dilution
We can prepare solutions by adding solid to a set volume of water.
Occasionally you may find that your laboratory work calls for the use of a dilute solution of a substance, but you discover that the only solutions available are more concentrated than what you need.
This situation calls for a dilution to be carried out.
The thing to remember is that the number of moles of solute does not change as we are only adding more solvent to the existing solution.

35. Solution Concentrations
Example - We want to make 500 mL of a hydrochloric acid solution with a concentration of 3.0 mol/L starting with a concentrated solution with a concentration of 12.0 mol/L. What volume of the concentration solution should we start with and dilute to 500 mL to end up with the correct concentration of 3.0 mol/L? What volume of water needs to be added?

36. Solution Concentrations
Example – You make a stock solution by adding g of sodium chloride to 500 ml of water. You take 50 ml of the stock solution and add it to 100 ml of water. What is the new concentration of the sodium chloride solution?
Example – You have a 0.3 M solution of cobalt (II) nitrate. How many grams were dissolved in 250 ml of water to make the solution. If you add 50 ml of the stock solution to 50 ml of water, what is the new concentration?

37. Solubility Rules
Earlier we said that if a substance is soluable it puts lots of ions in solution. It ionizes 100%. If a substance is insoluable very little ionizes and stays dissociated together.
Examples
Sodium hydroxide
Copper (II) sulfate
Aluminum carbonate
Magnesium hydroxide

38. Solubility Rules How do we know what is soluable?
We use a list of Solubility Rules. (Handout)
Are the following soluable? Use the rules.

39. Solubility Rules Salt Formula Soluable? Rule Sodium iodide
Potassium sulfate
Copper (II) chloride
Lead (II) acetate
Lead (II) Iodide
Mercury (II) nitrate
Sodium Carbonate
Calcium carbonate

40. Precipitates and Net Ionic Equations
When two solutions combine and some of the ions re-associate, the resulting product is a precipitate
Precipitate – insoluable solid that comes out of solution
The precipitate is insoluable based upon our solubility rules
If no precipitate forms, or no new substance forms, there is no reaction. In solutions, the ions stay dissociated and do not re-associate.

41. Precipitates and Net Ionic Reactions
We can predict if a precipitate forms by ionizing the reactants and reforming the products by combining ions. If one of the products is insoluable, determined by the solubility rules then a precipitate forms.
Ex.
Sodium chloride reacts with silver (I) nitrate

42. Precipitates and Net Ionic Equations
A Net Ionic Equation is a chemical equation for a reaction which lists only those species participating in the reaction. We eliminate the spectator ions.
Spectator Ions
ions that stay dissociated in solution and do not participate in the reaction

43. Precipitates and Net Ionic Equations
Steps
Start by simply writing the overall balanced chemical reaction. This is also called the Molecular Equation.
Then, you break apart the soluble ionic compounds into the two ions from which it is formed (one positive and one negative). You will have to use the solubility rules to do this, they can be found online. If something is insoluble, it should not be broken apart. Write the reaction out with all of the separated ions. This is called the Total Ionic Equation.
Then, you simplify by cancelling things out if they appear on both sides of the reaction, resulting in the Net Ionic Equation.

44. Precipitates and Net Ionic Equations
Ex.
Identify the spectator ions and write the net ionic equation for the reaction of sodium chloride with silver (I) nitrate
Write the net ionic equation for the reaction of Copper (II) nitrate and sodium hydroxide

45. Solution Stoichiometry
Solution Stoichiometry (same steps, same stuff)
Quite often the best way to make two chemicals react with one another is to dissolve them in a suitable solvent and to mix the resulting solutions.
It is particularly important to be able to perform stoichiometric calculations when situations such as this arise.
As with the previous examples on stoichiometry, we usually have to determine the amount of reactant (or product) involved.

46. Solution Stoichiometry
Ex.
If 10 ml of a 0.1 M solution of sodium chloride reacts with excess silver nitrate solution, what mass of precipitate is formed? Ex
Calculate the volume of silver nitrate solution [0.800] that is needed to react completely with 12.0 g of copper metal.

47. Solution Stoichiometry
Ex.
If 25 ml of a 0.2 M solution of silver (I) nitrate reacts with 5.00 g of copper wire, what mass of silver would be produced?
Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag (s)