Solutions
Unit
Honors
Chemistry
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
• 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
supersaturated solution
(stirred)
Supersaturated Solution of Sodium Thiosulfate
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
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
At 20oC, a saturated
solution contains how
many grams of NaNO3
in 100 g of water?
90 g
What kind of solution is
formed when 90 g
NaNO3 is dissolved in
100 g water at 30oC?
unsaturated
What kind of solution
is formed when 120 g
NaNO3 is dissolved in
100 g water at 40oC?
supersaturated
180
Saturated sol’n
170
160
150
140
Supersaturated
solution
130
120
Solubility ( g/100 g water )
What is the solubility
at 70oC?
135 g/100 g water
Solubility Graph for NaNO3
110
100
90
80
70
Unsaturated solution
60
50
40
30
20
10
0
0
10
20
30
40
50
60
70
Temperature (deg C)
80
90
100
110
Solubility of Gases
• Gases are less
soluble at high
temperatures than
at low temperatures
• Increasing
temperature,
decreases
solubility.
• Increasing
pressure, increases
solubility.
• 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).
• Effervescence – rapid escape of gas
dissolved in liquid
Factors Affecting Solubility
• Increase surface area of solute
(crushing)
• Stir/shake
• Increase temperature
Dissolution Process
electrolyte
• Ionic Compounds
NaCl(s)  Na+1(aq) + Cl-1(aq)
nonelectrolyte
– 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
Dissolving
NaCl in water
hexahydrated for Na+1;
Solvation: process of solvent molecules
most cations have 4-9 H2O molecules
surrounding solute
6 is most common
Hydration: solvation with water
Dissolution Process
• Molecular Compounds
– Nonpolar molecular solids do not dissolve in
polar solvents
• naphthalene
– Polar molecule
• C12H22O11(s)  C12H22O11(aq)
• Molecular solvation
• Nonelectrolyte
– Polar molecule
• 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)
Electrolyte vs. Nonelectrolyte
Concentration
• Percent concentration by mass (mass %
– (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
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.
Energy Changes
• Heat of solution = Hsoln
• Endothermic
– Solute particles separating in solid
– Solvent particles moving apart to allow solute to enter
liquid
– Energy absorbed
• Exothermic
– Solute particles separating in solid
– Solvent particles attracted to solvating solute particles
– Energy released
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:
1. Vapor Pressure
2. Boiling Point ELEVATION
3. Freezing Point DEPRESSION
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.
H2O
H2O
H2O
Sugar H2O
Same Temperature
Vapor Pressure (kPa)
100
H2O
80
60
40
20
Temperature (ºC)
100
solution
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
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
Boiling Point Elevation and
Freezing Point Problems
1. 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:
2. At what temperature will a solution begin
to freeze when 18.0 g ammonium
phosphate is dissolved in 200.0 g water?
–
Solute:
Naming Acids Review:
A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid
Ex) HCl
Ex) H3P
hydrochloric acid
hydrophosphoric acid
B. Tertiary – H + polyatomic anion
(oxo)
Ex) H2SO4
Ex) H2SO3
no Prefix “hydro”
end “ate” = “ic” acid
end “ite” = “ous” acid
sulfuric acid
sulfurous acid
Properties of Acids and Bases:
Acid
Base
(alkali)
Reactions
Electrical
with Metals Conductivity
Taste
Touch
sour
looks like
water,
burns,
stings
Yesproduces
H2 gas
electrolyte
in solution
bitter
looks like
water,
feels
slippery
No
Reaction
electrolyte
in solution
Indicators: Turn 1 color in an acid and
another color in a base.
A. Litmus Paper: Blue and Red
An aciD turns blue litmus paper reD
A Base turns red litmus paper Blue.
B. Phenolphthalein: colorless in an acid and
pink in a base
C. pH paper: range of colors from acidic to
basic
D. pH meter: measures the concentration
of H+ in solution
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
Arrhenius Definition (1884):
A. An acid dissociates in water to produce more
hydrogen ions, H+.
HCl  H+1 + Cl-1
B. A base dissociates in water to produce more
hydroxide ions, OH-.
NaOH  Na+1 + 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
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.
Hydronium Ion:
Hydronium Ion – H3O+ This is a complex ion
that forms in water.
H+1 + H2O  H3O+1
To more accurately portray the Bronsted-Lowry,
the hydronium ion is used instead of the
hydrogen ion.
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+ + NO3A weak acid or base will partially ionize (<5% as
ions). This is represented by a double (↔) arrow.
HOCl + H2O
↔
H3O+ + ClO-
HF < HCl < HBr < HI
increasing strength
7 Strong Acids
HNO3
H2SO4
HClO4
HCl
HI
HClO3
HBr
8 Strong Bases
LiOH
NaOH
RbOH
CsOH
Sr(OH)2
Ba(OH)2
KOH
Ca(OH)2
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.
Ionization of Acids & Bases
• H2SO4  2 H+ + SO4-2
– Sulfuric acid
• H3PO3 
3 H+ + PO3-3
– Phosphorous acid
• Ca(OH)2  Ca+2 + 2 OH-1
– Calcium hydroxide
Conjugate Acid-Base Pairs: A pair of
compounds that differ by only one
hydrogen ion
A. Acid donates a proton to become a
conjugate base.
B. 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.
Acid (A), Base (B),
Conjugate Acid (CA), Conjugate Base (CB)
NH3 +
H2O
↔ NH4+ +
OH-
HCl +
H2O
↔ Cl-
H3O+
B
A
A
B
CA
CB
+
CB
CA
• Base and Conjugate Acid are a Conjugate
Pair.
• Acid and Conjugate Base are a Conjugate
Pair.
AciDonates & Bases accept
1. H2O
B
+
H 2O
↔
A
B
B
4. OH− + H3O+
B
A
HSO4−
OH−
CB
+
H 2O
CB
3. HSO4− + H2O ↔
A
+
CA
2. H2SO4 + OH− ↔
A
H3O+
SO4−2
CA
+
H3O+
CB
↔
CA
H 2O + H 2 O
CA
CB
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 10-14 M2
1.0 x 10-14 M2 = [1.0 x 10-7 M] [1.0x10-7 M]
1.0 x 10-14 =
[H3O+]
[OH-]
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 10-14 M2 = [4.0 x 10-8 M] [OH-]
2.5 x 10-7 M = [OH-]
slightly basic
The pH scale (1909): the power of Hydrogen
A. Measure of H3O+ in solution.
B. pH = -log[H3O+]
C. Range of pH: 0-14
pH < 7: acid
pH > 7: base
pH = 7: neutral
D. pOH = -log[OH-]
E. pH + pOH = 14
H+
OH-
14
1
pH
[H3O+]
[OH-]
14
1x10-14
1x100
13
1x10-13
1x10-1
12
1x10-12
1x10-2
11
1x10-11
1x10-3
10
1x10-10
1x10-4
9
1x10-9
1x10-5
8
1x10-8
1x10-6
7
1x10-7
1x10-7
6
1x10-6
1x10-8
5
1x10-5
1x10-9
4
1x10-4
1x10-10
3
1x10-3
1x10-11
2
1x10-2
1x10-12
1
1x10-1
1x10-13
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 = 3.0604807474
– Sig Fig pH = 3.06