1. Solutions Unit
Honors Chemistry

2. Solution
Definition: a homogeneous mixture of 2 or more substances in a single physical state
Parts: solute and solvent (usually water)
Types:
Physical states: solid (alloys), liquid, gas
Miscible vs. Immiscible
Miscible: Liquids that dissolve freely in one another in any proportion
Immiscible :Liquid solutes and solvents that are not soluble
Saturated, Unsaturated and Supersaturated
Dilute vs. Concentrated
Electrolyte vs. Nonelectrolyte

4. Saturated – soln containing the max amt of solute
Unsaturated – soln containing less solute than a sat soln under the existing conditions
Supersaturated – contains more dissolved solute than a saturated solution under the same conditions
Solubility Curves

5. supersaturated solution (stirred)
Supersaturated Solution of Sodium Thiosulfate

6. Solubility (physical change)
Definition: mass of solute needed to make a saturated solution at a given temperature
solution equilibrium in a closed system
dissolution ↔ crystallization
Unit = g solute/100 g H2O

7. Solubility of solids in liquids
For most solids, increasing temperature, increases solubility.
In general, “like dissolves like”. Depends on
Type of bonding
Polarity of molecule
Intermolecular forces between solute and solvent

8. What is the solubility at 70oC?
At 20oC, a saturated solution contains how many grams of NaNO3 in 100 g of water?
Saturated sol’n
90 g
What is the solubility at 70oC?
Supersaturated solution
135 g/100 g water
What kind of solution is formed when 90 g NaNO3 is dissolved in 100 g water at 30oC?
Unsaturated solution
unsaturated
What kind of solution is formed when 120 g NaNO3 is dissolved in 100 g water at 40oC?
supersaturated

9. Solubility of Gases
Gases are less soluble at high temperatures than at low temperatures
Increasing temperature, decreases solubility.
Increasing pressure, increases solubility.

10. Increasing pressure, increases solubility.
The quantity of gas that dissolves in a certain volume of liquid is directly proportional to the pressure of the gas (above the solution).

11. Effervescence – rapid escape of gas dissolved in liquid

12. Factors Affecting Solubility
Increase surface area of solute (crushing)
Stir/shake
Increase temperature

13. Dissolution Process Ionic Compounds NaCl(s)  Na+1(aq) + Cl-1(aq)
For dissolution to occur, must overcome solute attractions and solvent attractions.
Dissociation Reaction: the separation of IONS when an ionic compound dissolves (ions already present)
Try calcium chloride
electrolyte
nonelectrolyte
Dissolving NaCl in water
hexahydrated for Na+1;
most cations have 4-9 H2O molecules
6 is most common
Solvation: process of solvent molecules
surrounding solute
Hydration: solvation with water

15. Dissolution Process Molecular Compounds
Nonpolar molecular solids do not dissolve in polar solvents
naphthalene
Polar molecule
C12H22O11(s)  C12H22O11(aq)
Molecular solvation
Nonelectrolyte
HCl(g)  H+1(aq) + Cl-1(aq) or
HCl(g) + H2O  H3O+1(aq) + Cl-1(aq)
Ionization: ions formed from solute molecules by action of solvent (no ions initially present)
Nonelectrolyte (HCl)  electrolyte (ions)

16. Electrolyte vs. Nonelectrolyte

17. Concentration Percent concentration by mass (mass % Molarity (M)
(solute/solution) x 100% = % Concentration
Molarity (M)
Moles of solute/Liters of solution = mol/L
Molality (m)
Moles of solute/mass of solvent = mol/kg
ppm and ppb
Used for very dilute solutions
Dilution – a process in which more solvent is added to a solution
How is this solution different?
Volume, color, molarity
How is it the same?
Same mass of solute, same moles of solute
In Dilution ONLY – M1V1 = M2V2

18. Net Ionic Equations
Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and omit the spectator ions, which go through the reaction unchanged.
Substances that are aqueous break down into ions.
Substances that pure solids, liquids, or gases do not break down in solution.
Hint: Remember to check solubility rules to determine if a precipitate forms.

19. Energy Changes Heat of solution = Hsoln Endothermic Exothermic
Solute particles separating in solid
Solvent particles moving apart to allow solute to enter liquid
Energy absorbed
Exothermic
Solvent particles attracted to solvating solute particles
Energy released

20. Colligative Properties
Definition: physical properties of solutions that differ from properties of its solvent.
Property depends upon the number of solute particles in solution.
Types:
Vapor Pressure
Boiling Point ELEVATION
Freezing Point DEPRESSION

21. Vapor Pressure A measure of the tendency of molecules to escape from a liquid
For nonvolatile liquids or solid solutes
A nonvolatile solute will typically increase the boiling point and decrease the freezing point.
Adding a nonvolatile solute lowers the concentration of water molecules at the surface of the liquid.
This lowers the tendency of the water molecules to leave the solution and enter the gas phase.
Therefore the vapor pressure of the solution is LESS than pure water.

22. Same Temperature H2O H2O Sugar H2O Temperature (ºC)
Vapor Pressure (kPa)
100 40 20 60 80
H2O
solution

23. Boiling Point Elevation
tb = boiling point elevation
tb = iKbm
i = molality conversion factor; for ionic compounds adjust for # of ions actually present in solution (dissociation process)
Kb = molal bp elevation constant
Kb = 0.512°C·kg H2O
moles of solute (ions or molecules)
m = molality = moles solute
kg of solvent
bp of solution = bp of solvent + tb

24. Freezing Point Depression when a solution freezes, the solvent solidifies as a pure substance; deviates for more concentrated solutions
tf = freezing point depression
tf = iKfm
i = molality conversion factor; for ionic compounds adjust for # of ions actually present in solution (dissociation process)
Kf = molal freezing point depression constant
Kf = 1.858°C·kg H2O
moles of solute (ions or molecules)
m = molality = moles solute
kg of solvent
fp of solution = fp of solvent - tf

25. Boiling Point Elevation and Freezing Point Problems
At what temperature will a solution begin to boil if it is composed of 1.50 g potassium nitrate in 35.0 g of water?
Solute:
At what temperature will a solution begin to freeze when 18.0 g ammonium phosphate is dissolved in g water?

26. Naming Acids Review: Ex) HCl hydrochloric acid
A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid
Ex) HCl hydrochloric acid
Ex) H3P hydrophosphoric acid
B. Tertiary – H + polyatomic anion no Prefix “hydro”
(oxo) end “ate” = “ic” acid end “ite” = “ous” acid
Ex) H2SO4 sulfuric acid
Ex) H2SO3 sulfurous acid

27. Properties of Acids and Bases:
Taste
Touch
Reactions with Metals
Electrical Conductivity
Acid
sour
looks like water, burns, stings
Yes-produces H2 gas
electrolyte in solution
Base
(alkali)
bitter
looks like water, feels slippery
No Reaction

28. Indicators: Turn 1 color in an acid and another color in a base.
Litmus Paper: Blue and Red
An aciD turns blue litmus paper reD
A Base turns red litmus paper Blue.
Phenolphthalein: colorless in an acid and pink in a base
pH paper: range of colors from acidic to basic
pH meter: measures the concentration of H+ in solution

30. Reactions
Neutralization: A reaction between an acid and base. When an acid and base neutralize, water and a salt (ionic solid) form.
Acid Base → Salt Water
Ex) HCl NaOH → NaCl HOH

31. Arrhenius Definition (1884):
A. An acid dissociates in water to produce more
hydrogen ions, H+.
HCl  H Cl-1
B. A base dissociates in water to produce more hydroxide ions, OH-.
NaOH  Na OH-1
C. Problems with Definition:
• Restricts acids and bases to water solutions.
• Oversimplifies what happens when acids
dissolve in water.
• Does not include certain compounds that have
characteristic properties of acids & bases.
Ex) NH3 (ammonia) doesn’t fit

32. Bronsted-Lowry Definition (1923):
A. An acid is a substance that can donate hydrogen ions. Ex) HCl → H+ + Cl-
Hydrogen ion is the equivalent of a proton.
Acids are often called proton donors.
Monoprotic (HCl), diprotic (H2SO4) , triprotic (H3PO4)
B. A base is a substance that can accept hydrogen ions. Ex) NH3 + H+ → NH4+
Bases are often called proton acceptors.
C. Advantages of Bronsted-Lowry Definition
•Acids and bases are defined independently of how
they behave in water.
•Focuses solely on hydrogen ions.

33. Hydronium Ion:
Hydronium Ion – H3O+ This is a complex ion that forms in water.
H H2O  H3O+1
To more accurately portray the Bronsted-Lowry, the hydronium ion is used instead of the hydrogen ion.

34. STRONG Acid/Base versus WEAK Acid/Base
Strength refers to the % of molecules that form IONS.
A strong acid or base will completely ionize (>95% as ions). This is represented by a single () arrow.
HNO3 + H2O  H3O+ + NO3-
A weak acid or base will partially ionize (<5% as ions). This is represented by a double (↔) arrow.
HOCl + H2O ↔ H3O+ + ClO-

35. HF < HCl < HBr < HI increasing strength
7 Strong Acids
HNO3 H2SO4 HClO3
HClO4 HCl HBr HI
8 Strong Bases
LiOH NaOH KOH
RbOH CsOH Ca(OH)2
Sr(OH)2 Ba(OH)2

36. Strength vs. Concentration
Strength refers to the percent of molecules that form ions
Concentration refers to the amount of solute dissolved in a solvent. Usually expressed in molarity.

37. Ionization of Acids & Bases
H2SO4  2 H+ + SO4-2
Sulfuric acid
H3PO3 
Phosphorous acid
Ca(OH)2 
Calcium hydroxide
3 H+ + PO3-3
Ca OH-1

38. Conjugate Acid-Base Pairs: A pair of compounds that differ by only one hydrogen ion
Acid donates a proton to become a conjugate base.
Base accepts proton to become a conjugate acid.
A strong acid will have a weak conjugate base.
A strong base will have a weak conjugate acid.

39. Acid (A), Base (B), Conjugate Acid (CA), Conjugate Base (CB)
NH3 + H2O ↔ NH OH-
HCl + H2O ↔ Cl H3O+
Base and Conjugate Acid are a Conjugate Pair.
Acid and Conjugate Base are a Conjugate Pair. B A CA CB A B CB CA

40. AciDonates & Bases accept
H2O H2O ↔ H3O OH−
B A CA CB
H2SO OH− ↔ HSO4− H2O
A B CB CA
HSO4− H2O ↔ SO4− H3O+
A B CB CA
OH− H3O+ ↔ H2O H2O
B A CA CB

41. The Self-ionization of Water & pH
1. Water is amphoteric, it acts as both an acid and a base in the same reaction.
Ex) H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)
Keq = equilibrium constant = [H3O+] [OH-]
Because reactants and products are at equilibrium, liquid water is not included in the equilibrium expression
@ 25C, [H3O+] = 1 x 10-7 M and [OH-] = 1 x 10-7 M
Kw = ion product constant or equilibrium constant for water
Kw = [H3O+] [OH-] = 1 x M2
1.0 x M2 = [1.0 x 10-7 M] [1.0x10-7 M]
1.0 x = [H3O+] [OH-]

42. Acids: [H3O+] > 1 x 10-7 M Bases: [OH-] > 1 x 10-7 M
Using Kw in calculations: If the concentration of H3O+ in the blood is 4.0 x 10-8 M, what is the concentration of OH­ ions in the blood? Is blood acidic, basic or neutral?
Kw = [H3O+] [OH-]
1.0 x M2 = [4.0 x 10-8 M] [OH-]
2.5 x 10-7 M = [OH-]
slightly basic

43. The pH scale (1909): the power of Hydrogen
Measure of H3O+ in solution.
pH = -log[H3O+]
Range of pH: 0-14
pH < 7: acid
pH > 7: base
pH = 7: neutral
pOH = -log[OH-]
pH + pOH = 14

44. OH- H+ 14 1 pH [H3O+] [OH-] 14 1x10-14 1x100 13 1x10-13 1x10-1 1211
1x10-11
1x10-3 10
1x10-10
1x10-4 9
1x10-9
1x10-5 8
1x10-8
1x10-6 7
1x10-7 6 5 4 3 2 1 14 H+
OH- 1

45. Significant Digits Rule
The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number
Ex -log[8.7x10-4M]
Calc Answer =
Sig Fig pH = 3.06