1. Unit 9: Solutions

2. H2O What properties of water make it essential to life of earth?
Good solvent
High Surface tension
Low vapor pressure
High boiling point

3. Water is a polar molecule
Water is a polar molecule. It experiences hydrogen bonding due to its polar nature.

4. After many hydrogen bonds are formed, you have a weak force holding all the water molecules to each other.

5. Hydrogen bonding is the reason water freezes into ice crystals of a certain repeating shape (hexagons)
Volume increases.

6. Hydrogen bonding is what gives proteins and nucleic acids their three dimensional shape.

7. Hydrogen bonding also causes water to have….
High Surface Tension (cohesion)
A large amount of energy to boil water into vapor because of the bonds holding the molecules together. The bonds must be broken in order for the liquid to change to a gas. Water molecules on the surface are 'holding' each other together (cohesion of water molecules)

8. Surfactant
Surface Tension can be decreased by adding a surfactant – this type of substance interferes with hydrogen bonding.
Surfactants are used in soaps, detergents, paints, adhesives, inks…

9. Hydrogen bonds also cause water to have…
A high boiling point and a low vapor pressure.

10. Water is also attracted to other substances like glass.
Water can be drawn up into a thin glass tube with no effort (capillary action) because of the attraction between the water and glass molecules.

11. Classification of Matter
anything composed of atoms
Two types of mixtures!

12. Heterogeneous Mixtures
Two or more pure substances mixed unevenly (you can see the different components).
Example: fruit salad, pizza, granite

13. Suspensions
A suspension is a heterogeneous mixture that has large particles that will stay suspended as long as the mixture is in motion Once the motion stops, the particles will fall to the bottom or settle out Can be separated through filtering

14. Colloids
Homogeneous mixtures where the medium sized particles are dispersed throughout but are not heavy enough to settle out. Appear cloudy and opaque Cannot be separated by filtering Can be separated by a centrifuge

15. How can light be used to tell the difference between colloids and solutions?
Tyndall effect A beam of light passing through a solution, such as air, is not visible. Light passing through a colloid, such as fog, will be scattered by the larger particles and the light beam will be visible.

16. Homogeneous Mixtures
Two or more pure substances mixed evenly. When you look at it, you can’t see separate parts.
Also called solutions (can be liquid, gas or solid solutions)
Examples: salt water, soda, coffee,

17. Parts of a solution Solute – the substance being dissolved.
The solute dissolves INTO THE solvent to form a solution.
Solute – the substance being dissolved.
Solvent – the substance that is doing the dissolving (usually present in a greater amount).
Solute
Solvent

18. Miscible – two substances that will mix and dissolve together in any proportion
Immiscible – two substances that are insoluble in one another.
Water and oil are _______________.

19. Factors that affect the rate of dissolution
Temperature
Dissolve faster at higher temp
Agitation
Dissolve faster when the mixture is shaken or stirred.
Surface Area
Dissolve faster if you increase their surface area (broken up into smaller pieces)
VS.

20. Solubility in Water :
The simple rule to remember is ‘like dissolves like’ +
Water is a polar solvent and will dissolve polar molecules and substances that contain charged particles. - +
Substances that dissolve in H2O are said to be soluble
Ex: Sugar, ethanol (type of alcohol) which are polar, most ionic compounds
Site of polarity can form H-bonding
sucrose

21. When solid sodium chloride (NaCl) dissolves in water it breaks up into its individual ions.

22. Water is polar, the different ends are attracted to the charged ions
Water is polar, the different ends are attracted to the charged ions. The positive and negative ions become surrounded by solvent- solvation.
 + -
 +
Cl-
Na+

23. Electrolytic Solutions
Ionic Compounds dissociate in aqueous solutions and create positive and negative ions in solution. Their ability to move nearly independently through the solution permits them to carry positive or negative electrical charges from one place to another. Hence the solution conducts an electrical current and are called electrolytes.

24. Electrolytic Solvation- +
sugar - +
acetic acid - +
salt
Non-
Electrolyte
Weak
Electrolyte
Strong
Electrolyte
solute exists as
molecules only
solute exists as
ions and
molecules
solute exists as
ions only

25. Electrolytic Solvation- +
sugar - +
acetic acid - +
salt
Non-
Electrolyte
Weak
Electrolyte
Strong
Electrolyte
covalent compounds
weak acids and bases
ionic compounds/ strong acids and bases

26. Which beakers contain an electrolytic solution:
KCl Solution
HNO3 Solution
Sugar Solution
Acetic Solution
AmmoniaSolution

27. Substances that don’t dissolve are called insoluble
E.g. Petroleum (crude oil), which are non-polar
So if you want to dissolve grease which is non-polar, you need to use a non-polar solvent.
Petroleum in a non-polar organic molecule

28. How is solubility of a compound denoted in a chemical equation?
, yields
A precipitate!!

29. How do we know if the products of a chemical reaction are (aq), a solution or (s), a suspension or a precipitate?

30. Double Replacement Reactions
AB + CD  AD + CB
Which of the products produced will be soluble?
potassium iodide + lead (II) nitrate
barium nitrate + calcium carbonate

31. Saturation

32. MAXIMUM Amount of solute dissolved
Saturation – The point when no more solute can dissolve into a solvent at a given temperature ( a dynamic equilibrium exists between the solution and undissolved solute).

33. Unsaturated Solution
An unsaturated solution is like a sponge that can hold more water.
More solute can be dissolved at that temperature.
I’m still thirsty!

34. Supersaturated Solution
A supersaturated solution is a solution that contains more of the solute than it can normally hold at a given temperature.
Can be prepared by changing the conditions of a saturated solution (temperature, volume, or pressure).
Ex) carbonated water – CO2 in water under pressure
Ex) sugar in heated water, after it cools again it still holds the dissolved sugar (this is how to make rock

35. Solubility UNSATURATED SOLUTION more solute dissolves
no more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable, crystals form
concentration

36. Solvation & solubility curves

37. Solubility Solubility Curve
shows the dependence of solubility on temperature

38. Solubility Curve
Most solubility curves show the solubility of more than one solute. These graphs show comparisons. Which solute is most soluble at 100˚C?
KNO3 is most soluble at 100 degrees.

39. Practice: What is the solubility of KBr at 80ºC?
Which salt is the most soluble at 50ºC?
At which temperature can you dissolve 160 g of KNO3 in 100 g of water?

40. Unsaturated, Saturated, and Supersaturated solution
Solutions that are below the line are always UNSATURATED SOLUTIONS.
Solutions that are described by the line are always SATURATED SOLUTIONS.
Solutions that are above the line are always SUPERSATURATED SOLUTIONS.

41. Practice:
How many grams of sodium chloride can be dissolved in 100 g of water at 80ºC?
If 40 g of potassium bromide is dissolved in 100 g of water at 80ºC, is the solution saturated, unsaturated, or supersaturated?

42. Solubility Curve What if we change the amount of water?
In order to reach saturation, if we double the amount of water that is being used to prepare the solution, we would need to also double the solute.
20 g / 100g of water at 25 °C
40 g / 200g of water at 25 °C

43. Example:
How many grams of KBr would be needed to saturate 50g of water at 90°C?
It takes 100g of KBr to
saturate 100g of water,
so it would take 50g of KBr
to saturate 50g of water.

44. What do you notice about the relationship between temperature and solubility of solids and gases in water?
Generally, as temperature increases, the solubility of solids increase.

45. Solubility and Temperature
As the temperature increases, the solubility of a gas decreases. Why?

46. Solubility of gases and 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.
The solubility of gases decreases as the temperature increases.
Ex: room temp soda tastes more “flat” than cold soda

47. Gas pressure and solubility
Gases increase in solubility as pressure increases (Henry’s law) The gas molecules are "forced" into the solution since this will best relieve the pressure

48. Problems
A solution is prepared with 10g of NaCl at 120°C. What type of solution is it? A. Unsaturated B. Saturated C. Supersaturated D. Heterogeneous

49. Problems
How many grams of KBr are needed to saturate 100g of water at 90°C? A. 80g B. 90g C. 100g D. 110g

50. Problems
How many grams of KBr are needed to saturate 200g of water at 90°C? A. 160g B. 180g C. 200g D. 220g

51. Concetrations & Dilutions

52. Concentration
Measure of how much solute is dissolved in solution, or “strength” of solution.
Concentrated: A solution with a relatively high concentration of solute
strong coffee
Dilute: very little solute in the solution
weak coffee
These statements are vague so lets get more specific…..

53. Concentration
Molarity(M) – is used by chemist to describe concentrations. It’s a measure of how many moles of solute are present for each liter of solution.
A 9M solutions is more concentrated than a 3M solutions.

54. Suppose you have 2.00 moles of sugar and you mix it with enough water to make 1.00 liter of solution- what is the molarity of the solution?
2M or “two molar”

55. Practice
Find the molarity of a solution containing 75 g of MgCl2 in 250 mL of water.
Given:
Mass = 75g
Volume = 250 mL
Unknown:
M=?

56. Practice
How many grams of NaCl are required to make a 1.54M solution using L of water?
Unknown:
mass=?
Given:
Molarity= 1.54 M
Volume = 0.500L

57. Other ways to express concentration…
Molality (m)
Concentration calculated by dividing the moles of solute by the kilograms of solvent (usually water) used to dissolve.
m=mols solute/kg solvent
Suppose you dissolve 18.2 grams of sucrose (C12H22O11) into 200. grams of water. What is the molality of the solution?

58. Other ways to express concentration…
Mole fraction (X)
X= moles of component/total moles of all components
You mix 1.7 moles of sodium chloride with 6.7 moles of water. What is the mole fraction of the sodium chloride?

59. Other ways to express concentration…
Percent (%) mass
% mass = grams of component/total mass of mixture
16 grams of NaBr are mixed with 32 grams of water. What is the % mass of sodium bromide?

60. Dilution
Preparation of a desired solution by adding water to a concentrate. Moles of solute remain the same.

61. What volume of 15. 8M HNO3 is required to make 250 mL of a 6
What volume of 15.8M HNO3 is required to make 250 mL of a 6.0M solution?

62. Practice
How much 0.05 M HCl solution can be made by diluting 250 mL of 10 M HCl?
How much water would I need to add to 500 mL of a 2.4 M KCl solution to make a 1.0 M solution?