1.
Solutions- A solution is a homogeneous mixture in which one substance disperses uniformly throughout another.

Solutions may exist in any of the three states of matter: gas, liquid, solid, or combinations: carbonation (CO
2
dissolved in H
2
O)

Solution consists of two parts: i.
Solvent- dissolving medium, usually the majority substance, can be any phase. ii.
Solute- that which is dissolved in a solvent (usually minority)
2.
Solution formation- The driving force of most solution processes is the natural tendency for systems to increase in Entropy, (state of randomness or disorder of a system.) Steps in solution formation. a.
Solute particles separate from each other.
 intermolecular forces of attraction must be broken.
 𝑠𝑜𝑙𝑢𝑡𝑒 + 𝑒𝑛𝑒𝑟𝑔𝑦 ⟶ 𝑠𝑜𝑙𝑢𝑡𝑒(𝑠𝑒𝑝𝑎𝑟𝑎𝑡𝑒𝑑)
 𝑏𝑟𝑒𝑎𝑘 𝑏𝑜𝑛𝑑𝑠 = 𝑒𝑛𝑑𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 = 𝑒𝑛𝑒𝑟𝑔𝑦 𝑎𝑏𝑠𝑜𝑟𝑏𝑒𝑑

∆𝐻 = 𝑝𝑜𝑠𝑖𝑡𝑖𝑣𝑒(+) b.
Solvent particles separate.
 intermolecular forces of attraction must be broken.
 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 + 𝑒𝑛𝑒𝑟𝑔𝑦 ⟶ 𝑠𝑜𝑙𝑣𝑒𝑛𝑡(𝑠𝑒𝑝𝑎𝑟𝑎𝑡𝑒𝑑)
 𝑏𝑟𝑒𝑎𝑘 𝑏𝑜𝑛𝑑𝑠 = 𝑒𝑛𝑑𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 = 𝑒𝑛𝑒𝑟𝑔𝑦 𝑎𝑏𝑠𝑜𝑟𝑏𝑒𝑑

∆𝐻 = 𝑝𝑜𝑠𝑖𝑡𝑖𝑣𝑒(+) c.
Solute and solvent particles mix and form a solution

 particles attract and new bonds are formed 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 + 𝑠𝑜𝑙𝑢𝑡𝑒 ⟶ 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 + 𝑒𝑛𝑒𝑟𝑔𝑦
 𝑓𝑜𝑟𝑚𝑎𝑡𝑖𝑜𝑛 𝑜𝑓 𝑏𝑜𝑛𝑑𝑠 = 𝑒𝑥𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 = 𝑒𝑛𝑒𝑟𝑔𝑦 𝑟𝑒𝑙𝑒𝑎𝑠𝑒𝑑
 ∆𝐻 = 𝑛𝑒𝑔𝑎𝑡𝑖𝑣𝑒(−) d.
Overall solution formation:
∆𝑯 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏
= ∆𝑯 𝒔𝒐𝒍𝒖𝒕𝒆
+ ∆𝑯 𝒔𝒐𝒍𝒗𝒆𝒏𝒕
+ ∆𝑯 𝒎𝒊𝒙𝒕𝒖𝒓𝒆
3.
Solubility- amount of solute that can dissolve in a fixed amount of solvent. a.
At a given temperature when the solid is dissolving at the same rate as recrystallization is happening an equilibrium is established: 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑢𝑛𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑) ⟷ 𝑠𝑜𝑙𝑢𝑡𝑒(𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑) i.
At equilibrium no additional solute will be dissolved and the solution is saturated. ii.
An unsaturated solution contains less than the equilibrium concentration of dissolved solute.
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iii.
A supersaturated solution contains more than the equilibrium concentration
4.
Factors affecting solubility- the extent to which one substance dissolves in another. a.
Solute-solvent interactions- if solute-solvent IMF are comparable then solution will form. Paraphrased as “likedissolves-like.”

Depends on strength of the Intermolecular Forces of attraction: H-bonding, dipole-dipole, dispersion forces, etc. b.
Temperature-

Solids- increased temperature shows increasing solubility.
 Gases- decreased solubility with increasing temperature. Gases have weak IMF with water, as T ↑ ,
KE ↑ , more gas molecules escape solution. c.
Pressure- solids and liquids not affected by pressure
 changes, but gases are. solubility of gases increases in direct proportion to it’s partial pressure above the solution.
 Henry’s Law 𝑆 𝑔
𝑆 𝑔
𝑃 𝑔
= 𝑘 ∙ 𝑃 𝑔
= 𝑠𝑜𝑙𝑢𝑏𝑖𝑙𝑖𝑡𝑦 𝑜𝑓 𝑔𝑎𝑠 𝑖𝑛 𝑚𝑜𝑙𝑒𝑠
𝐿
= 𝑝𝑎𝑟𝑡𝑖𝑎𝑙 𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 𝑜𝑓 𝑔𝑎𝑠
, 𝑘 = 𝐻𝑒𝑛𝑟𝑦 ′ 𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡
5.
Solution concentration- amount of solute per given o amount of solution.
𝑀𝑎𝑠𝑠% = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
× 100% o 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 = 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 o 𝑀𝑜𝑙𝑒 𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 = 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 o 𝑃𝑃𝑀 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
× 10 6 o 𝑉𝑜𝑙𝑢𝑚𝑒% = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓𝑠𝑜𝑙𝑢𝑡𝑒 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
× 100 o Making a molar solution from stock- o 𝑀 𝑠𝑡𝑜𝑐𝑘
𝑉 𝑠𝑡𝑜𝑐𝑘
= 𝑀 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑
𝑉 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 o Converting between concentration units.
 to convert 𝑚𝑜𝑙𝑒𝑠 ⟶ 𝑚𝑎𝑠𝑠 , use MM
 to convert 𝑚𝑎𝑠𝑠 ⟶ 𝑣𝑜𝑙𝑢𝑚𝑒 , use solution density
6.
Solvent-solute interactions- a.
Charged ends of polar water molecules attract to charged ions and separate and surround ions from crystal lattice. b.
Polar water molecules also can solvate molecular substances, but don’t form ions.
Type of solute vs. solubility of solvent
SOLUTE
POLAR
SOLVENT
NONPOLAR
SOLVENT
2

B.
7.
Electrolytic properties a.
Aqueous solution in which ions are formed is an
electrolyte. b.
Strong electrolyte- solutes that exist in solution completely or nearly completely as ions. Good conductor of electricity. c.
Weak electrolyte- solutes exist in solution mostly as neutral molecules, only small fraction as ions. Poor conductor of electricity d.
Nonelectrolyte- no ions, doesn’t conduct electricity.
reaction that results in the formation of an insoluble product.
1.
form insoluble salt (precipitate) from aqueous ions
2.
solubility rules for predicting MX(s) (memorize)

All group 1 and NH
4
+
Mostly Soluble Salts

All NO
3
and C
2
H
3
O
2
–

Most ClO
3
– , ClO
4
–

Most Cl , Br , I except Ag + , Pb 2+ , Hg
2
2+

Most SO
4
2 except Ca 2+ , Sr 2+ , Ba 2+ , Ag + , Pb 2+ , Hg
2
2+
Mostly Insoluble Salts

Most OH , S 2 except group 1, NH
4
+ , Ca 2+ , Sr 2+ , Ba 2+
If OH- bonds, cmpds are slightly soluble

CO
3
2- and PO
4
3- except group 1, or NH
4
+
 All other anions except with column 1 and NH
4
+
3.
Writing equations for reactions in aqueous solutions

Molecular equation- shows complete chemical formulas of reactants and products.
 Ionic equation- shows all species as they are actually present in solution.

Net ionic equation- shows only the species that actually change during the reaction.

Spectator ions- ions that remain dissolved in solution, don’t participate in reaction.
C.
Colligative Properties-
1.
On adding a solute to a solvent, the properties of the solvent are modified. a.
Vapor pressure lowering- addition of solute particles reduces vapor pressure. b.
Freezing point decreases- solute particles interfere with the intermolecular forces thus preventing the solvent from entering the solid state.
•
•
K f
is the freezing point constant, for water it is 1.86 ºC/m
m is the molality of the solution. c.
Boiling point increases- addition of solute particles lowers vapor pressure and therefore increases boiling point.
3

•
K b
is the boiling point constant, for water it is 0.512 ºC/m
• m is the molality of the solution. d.
Osmosis- adding additional solute particles to a solution increases the osmotic pressure.
2.
These changes are called COLLIGATIVE PROPERTIES. a.
They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.
D.
Volumetric Analysis-
1. stoichiometry flow chart grams A ⟵ MM
A
⟶ moles A
⟵ M
A
⟶
⟵ V
A
⟶
V
A
M
A
↕
⟵ M
B
⟶ V
B grams B ⟵ MM
B
⟶ moles B
⟵ V
A
⟶ M
B
2. titration
 add standard solution (titrant) to buret o rinse buret with standard solution o clear air pockets o record initial volume (bottom of meniscus)
 add unknown and indicator to flask
 add standard solution until color change (end point) o touch tip to flask to release hanging drop o record final volume (bottom of meniscus)
 calculate moles of unknown X (see flow chart)
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