South Pasadena • Honors Chemistry
Name
4 • Salts and Solutions
Period
4.1
NOTES
Date
–
SOLUBILITY
Defining a Solution
Solution – a homogeneous mixture. Two or more substances combined together, but are visibly
indistinguishable. An aqueous solution – a substance dissolved in water.
NaCl
Solute – what’s being dissolved (the lesser part)
Solvent – what the solute is dissolved in (the greater part)
We represent NaCl (s) dissolved in H2O (ℓ) as NaCl (aq).
H2O
Saturated vs. Unsaturated Solutions
Saturated Solution – a solution that contains the maximum amount of solute that the solvent can “hold”.
Excess solute will “precipitate” out, but the solution remains saturated.
Unsaturated Solution – a solution that contains less solute than in a saturated solution.
Supersaturated Solution* − It is possible to dissolve more solute than what the solvent can “hold.” When the
solvent of a saturated solution is evaporated, the solute stays dissolved (i.e., it does not precipitate), even
though it is beyond “saturated”, and is unstable until it a “seed crystal” is added and the solute crystalizes.
Factors that Affect Solubility
Substances have differing ability to dissolve in water. Some substances (e.g. NaCl) dissolve well in water;
others don’t dissolve very well.
Solubility – a solute’s ability to dissolve in water. A substance that is very soluble dissolves well.
A substance that is insoluble does not dissolve well, and tends to precipitate.
g of solute
mol solute
Solubility can be expressed as volume of solution or L solution
1. Nature of the Substances: “Like Dissolves Like”

Solubility of Ionic Substances (or salts)
dissociates into ions.
A salt is any ionic compound. A salt is soluble if it easily
Soluble salts – dissolve well because they easily
break up into ions when polar water molecules pull
the ions apart. The ionic bonds are not very strong.
NaCl (aq)  Na+ + Cl−
Insoluble salts – do not break up into ions much
because ionic bonds are very strong. Insoluble salts
form precipitates.
AgCl (s)  Ag+ + Cl−
Solubility Rules for Salts*
Always soluble:
 alkali ions, NH4+, NO3−, ClO3−, ClO4−, C2H3O2−, HCO3−
Generally soluble:
 Cl−, Br−, I−
Soluble except with Ag+, Pb2+, Hg22+
 F−
Soluble except with Pb2+, Ca2+, Ba2+, Sr2+, Mg2+
2−
 SO4
Soluble except with Pb2+, Ca2+, Ba2+, Sr2+
Generally insoluble:
 O2−, OH−
Insoluble except with Ca2+, Ba2+, Sr2+, alkali ions, NH4+
2−
3−
2−
 CO3 , PO4 , S , SO32−, CrO42−, C2O42−
Insoluble except with alkali ions and NH4+

Solubility of Covalent Substances
Polar compound – a compound that has an uneven distribution of charges. (One side of the molecule is
more positive, and another side of the molecule is more negative.) A non-polar compound has charges
relatively evenly distributed throughout the molecule.
Solute
Polar
Nonpolar
Polar
Nonpolar
Solvent
Polar
Nonpolar
Nonpolar
Polar
Soluble?
Yes!
Yes!
No.
No.
Liquids that mix well (are soluble in each other) are said to be miscible. If liquids are immiscible, they
form layers (e.g. oil and water).

Acids and Bases*
Acids (HA) are covalent compounds and are often soluble in water, because they are polar substances.
However, the H−A bond is weak, so H+ often breaks off in solution (similar to how salts behave), and the
ions can be soluble.
Strong
Acids/Bases:
Weak
Acids/Bases:
Acids
Bases
An acid that dissociates 100%.
HA (aq)  H+ (aq) + A− (aq)
A base that dissociates 100%.
BOH (aq)  B+ (aq) + OH− (aq)
Strong acids: HCl, HBr, HI, HNO3,
H2SO4, HClO3, HClO4, HIO4
(memorize these!)
Strong bases: LiOH, NaOH, KOH, RbOH,
CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
(Family 1 and 2 hydroxides)
An acid that dissociates <5%
HA (aq)  H+ (aq) + A− (aq)
A base that does not dissociate 100%.
BOH (s)  B+ (aq) + OH− (aq)
(Any acid that is not strong.)
(Any base that is not strong.)
2.
Temperature
 Solid solutes (think NaCl in water)
Increasing temperature – increases solubility
Decreasing temperature – decreases solubility

Gas solutes (think CO2 dissolved in water in a carbonated
drink)
Increasing temperature – decreases solubility
Decreasing temperature – increases solubility
Solubility Curve*:
(Notice slopes of solid vs. gas solutes on solubility curve)
Example:
 How many grams of KNO3 can be dissolved in 300 g of
water at 50°C?
85 g KNO3
At 50°C: 300 g H2O  100 g H O  = 255 g KNO3

2 

If this solution is cooled to 10°C, how many grams of KNO3 will precipitate out of the solution?
22 g KNO3
At 10°C: 300 g H2O  100 g H O  = 66 g KNO3 dissolved at 10°C

2 
255 g – 66 g = 189 g KNO3 will precipitate out of the solution.
3. Pressure (for gaseous solutes)*
Henry’s Law
P = kC
P = partial pressure of the gas above the solution (atm)
C = concentration of dissolved gas (mol/L)
The concentration of a dissolved gas is directly related to the partial pressure of the gas above the solution.
Examples: At 0.750 atm and 25.0°C, 3.36 × 10−3 g of N2 can be dissolved in 0.250 L of water.
 What is the value of Henry’s Law constant (in atm·L/mol) for N2 at 25°C?
1 mol N2
nN2 = 3.36 × 10−3 g 28.02 g N  = 1.20 ×10−4 mol N2

2
−4
mol N2
1.20 × 10 mol N2
CN2 = L solution = 0.250 L water = 4.8 × 10−4 mol/L
P
0.750 atm
K = C = 4.8 × 10−4 mol/L = 1560 atm·L/mol
 How many grams of N2 can be dissolved in 0.500 L of water at 1.00 atm pressure?
P
1.00 atm
C = k = 1560 atm·L/mol = 6.40 × 10−4 mol/L
6.40 × 10−4 mol 28.02 g N2
0.500 L 
= 0.00897 g N2
L

  1 mol N2 
1000 × Msolute × Kb
∆T = i  D
 solution × %solvent 