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Final

Chapter 1 Summary This is a complete summary of chapter 1 that includes examples and all the important formulas.


Department
Chemistry
Course Code
CHEM 1027A/B
Professor
Felix Lee
Study Guide
Final

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Chemistry 1200b Notes
Chapter 1 Chemical Equilibrium
1.1 The Equilibrium Constant

Note: omit solids and liquids from K
Reverse equation, K gets flipped.
o

Multiply by a constant
o K#
Add equations
o 
Homogeneous means that all the substances in the reaction are in the same phase
Heterogeneous means that all the substances in the reactions are in different phases
o Note: Aqueous is considered the same phase as water
The Reaction Quotient Q
Where Q > K then a shift to the left will occur
Where Q < K then a shift to the right will occur
Q = k at equilibrium
Q is at anytime in the reaction, not necessarily at K
Le Châtelier’s Principle
After equilibrium is reach, if a disturbance of temperature or pressure or concentration occurs,
equilibrium will shift to oppose the action.

In the above example, 4 moles of substance are on the right versus the 2 moles of substance on
the left, which mean that the equilibrium will shift to the right.
Note: liquids/solids are not affected by volume change
Concentration:
o An increase of reactant/decrease in product : Q < K shift right
o A decrease of reactant/decrease in product : Q > K shift left

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Pressure: an increase in pressure will shift to fewer moles and vice versa as in the above
example.
Temperature: The system will shift opposite to heat added.
Gibbs Free Energy and the Equilibrium Constant



At equilibrium:
 
Where:
o R = 8.314

o T = 298.13K
o 
o K = atm




 

Solubility and Precipitation
Solubility: the amount of substance that will dissolve in a certain volume of a solvent
o Soluble: 10g or more can dissolve in a L
o SS: 0.1 to 10g can dissolve in a L
o Insoluble: less than 10g can dissolve in a L
This applies to room temperature reactions. In general, ionic compounds become more soluble
as temperature increases.
1.2 Solubility Table
Soluble Ions
Exceptions
Nitrates 

NONE
All Alkali metals
NONE
All salts of 
NONE
Halides (Cl-, Br-, I-)
Ag+, Hg2+, Cu+, Pb2+
Sulphates 

Cu2+, Sr2+, Ba2+, Pb2+

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Insoluble Ions
Exceptions
Sulphides S2-
Group 1, 2 and 
Carbonates 

Group 1 and 
Phosphates 

Group 1 and 
Hydroxides (OH-)
Group 1 and 
and Ba2+
Note: The common ion effect can significantly reduce the solubility of the solid relative to that in
pure water. This means that if a substance dissolves in a solution with ions already in it, the
uch since some is already present.
1.3 Weak Acids and Bases




 

Amphiprotic anion: a species that can act as both an acid and a base, eg. Water.
Weak acid and weak base reactions proceed until an equilibrium is achieved between the
reactants and products
In the case of weak acids and bases, their K values are much less than one.
The relationship between equilibrium constants shows that the stronger the weak acids is the
weaker the weak conjugate base will be
In strong acids, their bases are so weak that they are spectator ions and do not affect pH at all.
Arrhenius Theory
An acid produces ions in water
A base produces OH- ions in water
Bronsted-Lowry Theory:
An acid is a proton donor
A base is a proton acceptor
Lewis Acids and Bases
A Lewis acids is a substance that can accept a pair of electrons from another atom, this will
result in the formation of a coordinate solvent bond (i.e. both electrons shared in the bond from
the same atom).
A Lewis base is a substance that donates a pair of electrons to another atom. The species
formed in this reaction is called n acid-base adduct.
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