Unit 11- Solubility

1
Unit 11- Solubility Solutions, Ch. 17 18
2
I. Water

• A. The Molecule
• 1. OH bond is highly polar
• 2. Bond angle 105 making it Bent shaped
• 3. Water Molecule as a whole is polar
• 4. Attracted to each other by intermolecular
  hydrogen bonds

Greater electronegativity
3
I. Water (cont.)

• B. Important Properties
• 1. High surface tension
• 2. low vapor pressure
• hydrogen bonds hold molecules to one another,
  tendency to escape surface is low
• 3. high specific heat capacity
• 4.184 J/gC
• 4. high melting and boiling points
• 0C and 100C

4
I. Water (cont.)

• C. Surface Tension inward force, or pull, that
  tends to minimize the surface area of a liquid
• Surfactant wetting agent such as soap or
  detergent that decreases the surface tension by
  interfering with hydrogen-bonding

Responsible for high surface tension
5
I. Water (cont.)

• D. Atypical Ice
• 1. As a typical liquid cools, density increases
  b/c Volume decreases as the mass stays constant
• 2. As water cools it first behaves like a
  typical liquid until it reaches 4C
• 3. Below 4C the density of water starts to
  decrease
• Ice is one of only a few solids that float in
  their own liquid.

6
Atypical Ice Do not need to write.
Density of Liquid Water and Ice Density of Liquid Water and Ice
Temperature (C) Density (g/cm3)
100 (liquid water) 0.9584
50 0.9881
25 0.9971
10 0.997
4 1.000 (most dense)
0 (liquid water) 0.9998
0 (ice) 0.9168

7
Atypical Ice
Why does ice behave so differently?
Open framework arranged like a honeycomb.
Framework collapses, molecules packed closer
together, making it more dense
8
II. The Solution Process

• A. Solution - homogeneous mixture consisting of
  the same properties throughout (exists in a
  single phase)

Solute - substance being dissolved - smaller
quantity substance
Solvent - dissolving medium - larger quantity
substance - usually water
9
II. The Solution Process (cont.)

• Aqueous Solution (aq) a solution in which the
  solvent is water

HCl(g) H2O(l) ? H3O(aq) Cl-(aq)
potassium chromate, K2CrO4 (yellow)
cobalt(II) nitrate, Co(NO3)2 (red)
copper(II) sulfate, CuSO4 (blue)
potassium permanganate, KMnO4 (purple)
potassium dichromate K2Cr2O7 (orange)
nickel(II) chloride, NiCl2 (green)
10
II. The Solution Process (cont.)
B. Solvation the process of dissolving
1st solute particles (salt) are surrounded by
solvent particles (water)
First...
2nd solute particles (salt) are separated and
pulled into solution (salt water)
Then...
11
II. The Solution Process (cont.)
C. Like Dissolves Like
Polar solvents dissolve polar molecules and ionic
compounds
Nonpolar solvents dissolve nonpolar compounds
12
Polar Dissolves Polar
13
Polar vs. Nonpolar

• Polar Molecule uneven electron forces acting on
  the central atom.
• Water

• Nonpolar Covalent Bond when 2 atoms are joined
  by a covalent bond and the bonding electrons are
  shared equally

14
III. Water of Hydration

• Water molecules that are integral part of the
  crystal structure
• Compound that contains water of hydration is
  called a hydrate
• (ex. CuSO4.5H2O)
• Effloresce losing the water of hydration
• Deliquescent removes water from air to form a
  solution

15
III. Water of Hydration

• Calculate the percent by mass of water in sodium
  carbonate decahydrate
• (Na2CO3 10H2O)
• Percent H2O mass of water x 100
• mass of hydrate
• Mass of 10H2O 180 g
• Mass of Na2CO3 10H2O 286g
• Percent H2O 180 g x 100 62.9 or
• 286 g 63

16
A. Electrolytes compounds that conduct an
electric current in solutions
IV. Electrolytes

• All ionic compounds are electrolytes
• Compounds that dont conduct an electric current
  are called

nonelectrolytes not composed of ions, includes
many molecular compounds
(covalent compounds)
17
IV. Electrolytes (cont.)
Non- Electrolyte
Weak Electrolyte
Strong Electrolyte
solute exists as ions and molecules
solute exists as molecules only
solute exists as ions only
18
Some Examples of Strong Electrolytes, Weak Electrolytes and Nonelectrolytes Some Examples of Strong Electrolytes, Weak Electrolytes and Nonelectrolytes Some Examples of Strong Electrolytes, Weak Electrolytes and Nonelectrolytes
Strong Electrolyte Weak Electrolyte Nonelectrolyte
Acids (HCl, HBr, HI, HNO3, HClO4) Heavy metal halides HgCl2, PbCl2 Most organic Compounds
Bases (NaOH, KOH) Water (very weak) Glucose
Ions, Ionic compounds (NaCl, KCl, CaCl2, KClO3, MgSO4) Also called salts Organic (acids bases) Acetic acid, aniline Glycerol
19
V. Heterogeneous Mixtures

• A. Suspensions mixtures from which particles
  settle out upon standing and the average particle
  size is greater than 100 nm in diameter.
• Clearly identified as two substances
• Gravity or filtration will separate the particles
• B. Colloids heterogeneous mixtures containing
  particles that are between 1 nm and 100 nm in
  diameter
• Appear to be homogeneous but particles are
  dispersed through medium
• Ex paint, aerosol spray, smoke, marshmallow,
  whipped cream

20
V. Heterogeneous Mixtures

• C. Tyndall Effect - phenomenon observed when beam
  of light passes through a colloid or suspension
• Colloids exhibit the Tyndall effect

Solution
Colloid
21
VI. Solubility

• defined as the maximum grams of solute that will
  dissolve in 100 g of solvent at a given
  temperature
• based on a saturated solution

22
Solubility
SATURATED SOLUTION max amount no more solute
dissolves
SUPERSATURATED SOLUTION over max amount becomes
unstable, crystals form
UNSATURATED SOLUTION capable of dissolving more
solute
Concentration Increasing

• Supersaturated solutions are not in equilibrium
  with the solid substance and can quickly release
  the dissolved solids.
• Saturated solution is one that is in equilibrium
  with respect to the dissolved substance. These
  conditions can quickly change with temperature.

23
VI. Solubility (cont.)

• B. Factors Affecting Solubility
• 1. Stirring (agitation)
• Increases solubility b/c fresh solvent is brought
  in contact with the surface of the solute
• Surface phenomenon

24
VI. Solubility (cont.)

• 2. Temperature
• Increases solubility by increasing kinetic
  energy, which increases the collisions b/w
  molecules of solvent and the surface of the
  solute
• Increases amount and rate of solute dissolved

25
VI. Solubility (cont.)

• 3. Surface Area
• A smaller particle size dissolves more rapidly
  than larger particles size
• Surface phenomenon
• More surface area exposed, faster rate of
  dissolving

26
VI. Solubility (cont.)

• A. Solubility Curve
• shows the dependence of solubility on temperature
• Note the solubility of the gases are greater in
  cold water than in hot water
• Out of the solids, which has the lowest
  solubility at 40C?
• KClO3
• Which has the highest solubility at 20C?
• KI

27
VI. Solubility (cont.)

• How many grams of KNO3 can be dissolved in 100g
  of water at 60C?
• 110g
• What is the solubility of NH4Cl at 50C?
• 50 g/L
• Which salt shows the least change in solubility
  from 0C to 100C?
• NaCl

28
Output side

• 1. Which of the following compounds dissolved the
  highest at 20C?
• 2. The lowest at 20C?
• 3. Overall which compound dissolved the fastest
  from 0C to 100C?
• 4. Name a compound in the graph that is a gas?
  How do you know its a gas?

29
VII. Concentrations of Solutions

• Concentration of a solution is a measure of the
  amount of solute that is dissolved in a given
  quantity of solution.
• Dilute solution contains a low concentration of
  solute.
• Concentrated solution contains a high
  concentration of solute.
• Molarity (M) number of moles of a solute
  dissolved per liter of solution
• a.k.a. molar concentration

30

• A. Molarity
• Calculates the number of moles in
• 1 L of the solution
• Molarity (M) moles of solute
• liters of solution
• Ex 1
• Calculate the molarity when 2 mol of glucose is
  dissolved in 5 L of solution.
• 2 mol glucose
• 5 L solution

M n L
M n L
0.4 mol/L 0.4 M
31

• Ex 2
• How many moles of solute are present
• in 1.5 L of 0.24M Na2SO4?
• M n of solute
• L of solution

Triangle Trick
0.24M n 1.5 L
1.5
1.5
Divide
n
1.5L
0.24M
0.36 mol n
Multiply
32

• Ex 3
• A saline solution contains 0.90g NaCl in exactly
  100.0 mL of solution. What is the molarity of the
  solution?
• 1. CONVERT GRAMS TO MOLES!!
• 2. CONVERT mL to L.

What are you solving for?
Molarity (M)
0.90g NaCl
1 mol NaCl
0.016 mol
58g NaCl
100.0 mL 1L 0.1000 L 1000mL
Molarity 0.16 M
K h D M d c m L
33

• Ex 4
• How many grams of solute are in 2.40L of 0.650M
  HClO2?
• M n
• L

2.40 L
2.40 L
1.56 mol n
68g HClO2
1.56 mol HClO2
106.08g HClO2
1 mol HClO2
34
Title Molarity ProblemsYou do not have to write
the problem. You MUST show your work. BOX in your
answer!

• 1. A solution has a volume of 2.0L and contains
  36.0g of glucose. If the molar mass of glucose is
  180g, what is the molarity of the solution?
• 2. A solution has a volume of 250 mL and contains
  0.70 mol NaCl. What is its molarity?
• 3. How many moles of ammonium nitrate are in 335
  mL of 0.425M NH4NO3?
• 4. How many moles of solute are in 250 mL of 2.0M
  CaCl2? How many grams of CaCl2 is this?(molar
  mass 110g)

35
VII. Concentrations of Solutions (cont.)

• B. Making Dilutions
• You can make a solution less concentrated by
  diluting it with a solvent.
• A dilution reduces the moles of solute per unit
  volume, however, the total moles of solute in
  solution does not change
• M1 V1 M2 V2
• Volumes can be in L or mL, as long as the same
  units are used for both V1 V2

36

• Ex 1
• You need 150 mL of 0.40M NaCl and you have a 1.0M
  of NaCl solution. Calculate the volume of the
  NaCl solution.
• V1 150 mL of 0.40 NaCl
• M1 0.400M NaCl
• M2 1.0M NaCl
• Unknown V2

M1 V1 M2 V2
0.400M
150 mL
V2
1.0M
60. mL V2
37

• Ex 2
• What volume of 5.00M sulfuric acid is required to
  prepare 25.00L of 0.400M sulfuric acid?
• M1 5.00M
• V2 25.00L
• M2 0.400M

M1 V1 M2 V2
5.0M
V1
0.40M
25 L
V1
2.0 L
38

• Ex 3
• How many milliliters of a stock solution of
  2.00M MgSO4 would you need to prepare 100.0 mL of
  0.400M MgSO4?
• M1 2.00M MgSO4
• M2 0.400M MgSO4
• V2 100 mL of 0.400M MgSO4
• Unknown V1

M1 V1 M2 V2
0.400M
V1
100 mL
2.00M
V1 20 mL
39
LEFT SIDETitle Dilution ProblemsYou do not
have to write the problem. You MUST show your
work. BOX in your answer.

• 5. In making a dilution how many mL of 4.0M HCl
  are needed to make 250 mL of 0.30M HCl?
• 6. In making a dilution how many mL of a 3.00M
  NaBr solution are needed to make 175 mL of 0.400M
  NaBr?
• 7. How many milliliters of a stock solution of
  2.00M MgSO4 would you need to prepare 100mL of
  0.400M MgSO4?
• 8. How many milliliters of a stock solution of
  4.00M KI would you need to prepare 250.0mL of
  0.760M KI?

40
Adding to the mole road map!!
Solutions
? mol 1 L
1 L ? mol
Molarity 5.0 M 5.0 mol 1 L
Molecule
Formula unit
Atoms (ions)
41
VIII. Using Molarity in Stoichiometry

• Zn 2 HCl ? ZnCl2 H2
• How many milliliters of 3.00M HCl are required
  to react with 17.35 g of zinc?

17.35 g Zn
1 mol Zn 65 g Zn
1 L HCl 3.00 mol HCl
1000 mL HCl 1 L HCl
2 mol HCl 1 mol Zn
177.95 mL HCl
42
Using Molarity in Stoichiometry

• Cu 2 AgNO3 ? 2 Ag Cu(NO3)2
• How many grams of copper are required to react
  with 40.0 mL of 9.0M AgNO3?

• 9.0 mol AgNO3
• 1 L AgNO3

1 mol Cu 2 mol AgNO3
1 L AgNO3 1000 mL AgNO3
64 g Cu 1 mol Cu
40.0 mL AgNO3
11.52 g Cu
43
IX. Percent Solutions
If both solute and solvent are liquids the
concentration of the solute is expressed as a
percent. Percent by volume ((v/v)) volume
of solute X 100 solution
volume For solutions of solids dissolved in
liquids is percent (mass/volume). Percent
(mass/volume) ((m/v)) mass of solute (g) X
100 solution volume (mL)
44
IX. Percent Solutions

• Ex 1
• What is the percent by volume of ethanol (C2H6O),
  or ethyl alcohol, in the final solution when 85
  mL of ethanol is diluted to a volume of 250 mL
  with water?
• (v/v) 85 mL ethanol X 100
• 250 mL solution
• 34 ethanol (v/v)

45
IX. Percent Solutions

• Ex 2
• A solution of glucose (C6H12O6) contains 2.8 g
  in 20.0 mL of solution. What is the percent (m/v)
  of the solution?
• (m/v) 2.8g X 100
• 20.0 mL
• 14 glucose (m/v)

46
Why do we get to make ice cream?

• X. Freezing-Point Depression
• the difference in temperature between the
  freezing point of a solution and that of the pure
  solvent.
• The salt depresses the freezing point of water
  because it disrupts the crystal formation of the
  water

47
Why do we get to make ice cream?

• 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
• Compounds that break into ions upon dissolving,
  like NaCl breaks into Na and Cl-, are better at
  lowering the freezing point than substances that
  don't separate into particles because the added
  particles disrupt the ability of the water to
  form crystalline ice.
• The more particles there are, the greater the
  disruption and the greater the impact on
  particle-dependent properties (colligative
  properties) like freezing point depression.

48
Ice Cream Lab

• 1. Put all ingredients on my desk.
• 2. While I am mixing the ingredients. You need to
  
• prepare your coffee can and freezer bag.
• 3. To prepare your coffee can, put a 5cm layer of
  ice then sprinkle 1cm of rock salt, then add
  another layer of ice then rock salt. (3/4th full,
  leave room to put in the ice cream bag)
• 4. Then you will get your freezer zip lock bag
  and duct tape the 3 side edges.
• 5. I will then put ice cream in your bag and you
  will seal the opening with duct tape.
• 6. Place the taped bag of ice cream inside the
  middle of the coffee can and duct tape the lid
  onto the coffee can.
• 7. Roll your coffee can outside for 20 minutes!
• 8. When complete, take bag out of can and rinse
  off bag.
• 9. Then enjoy ice cream!