1. Unit 5:
Solutions,
Kinetics
and
Equilibrium

2. Solutions Parts of a Solution -- any homogeneous mixture
Not to be confused with:
a. suspensions: small particles of one
substance suspended in another
b. colloids: a suspension that doesn’t settle
Parts of a Solution
1. Solute : the material that gets dissolved
makes up less than 50%
2. Solvent : material that dissolves the solute
makes up more than 50%
Often the solvent is a liquid
--but not always

3. Solution Concentration
Qualitative:
lots of solute per unit solvent = concentrated
little solute per unit solvent = dilute
(how much Mio you add to your water- a little or a lot)
Quantitative: This is an actual measured quantity
Molarity (M) moles per liter
solute solution
moles
M =
Liters
(Table T)

4. Example: 100 g of MgO are dissolved in enough water to make
5.0 L of solution. What is the molarity?
100
Solution: first, find the moles of solute
moles =
= 2.5 moles 40
2.5
next, use molarity formula
M =
= 0.5M
5.0
Example: How many grams of BaCl2 are needed to make mL
of a 0.70M solution?
moles
Solution: find the moles of solute required
= 0.21
moles
0.7 =
0.3
grams
now, change moles to grams =
0.21 =
44 g
207

5. ) 2. Percent (%) (Table T) % = ( amount of solute . x 100
% = (
amount of solute . )
x 100
amount of total solution
amounts can be volumes or masses, but units must agree
-- except it’s OK to mix g and mL because 1mL has a mass of 1 g
Example: 100 mL alcohol is diluted with enough water to form
a final volume of 250 mL. % alcohol = ?
100
Solution:
x 100 =
40%
250
Example: 100 mL alcohol is mixed with 250 mL water. % alcohol=?
100
29%
Solution:
x 100 =
350

6. ( ) ( ) ( ) 3. Parts per Million (ppm) (Table T) amount of solute .
x 1,000,000
amount of total solution
Example: a 50.0 gram sample of apple contains 3.2 x 10-4 g of
pesticide. ppm = ? (
3.2 x 10-4 ) ( )
Solution:
1,000,000 =
6.4 ppm
50.0
When solutes dissolve in liquids, they change some physical
properties of the liquid, such as:
1. The freezing point is lowered
-- salt used to melt ice on roads
-- salt used to freeze home-made ice cream
-- antifreeze used to prevent car freeze-ups

7. 2. The boiling point is raised
--antifreeze also prevents boiling
These changes depend on the concentration
of solute particles, not what kind
Some substances cause larger-than-expected changes
-- ionic compounds (salts)
-- acids (H+) and bases (OH-)
These substances dissociate, or ionize, when dissolved in water
(separate into + and – ions)
Example: NaCl (s) 
Na+(aq) + Cl-(aq)
1 mole mole mole
2 moles
So, 1M NaCl affects the fp and bp as if it were a 2M solution
Examples: CaCl2(s) 
Ca+2(aq) + 2Cl-(aq)
Na2CO3(s) 
2Na+(aq) + CO3-2(aq)

8. These ions behave independently in solution
Example: Na2CO3 + CaCl2  2NaCl + CaCO3
2Na+(aq) + CO3-2(aq) + Ca+2(aq) + 2Cl-(aq) 
2Na+(aq) + 2Cl-(aq)
+ CaCO3(s)
This is an Ionic Equation
The solid that forms from a solution (CaCO3) is a precipitate
The Na+ and Cl- do not participate in the reaction --
they are spectator ions
The Net Ionic Equation is:
Ca+2(aq) + CO3-2(aq)  CaCO3(s)

9. Reactions between ions in solution will occur only if one of
the products is insoluble (gas or solid), or water
Example: NaCl + KNO3  ?
Solution: the possible products, NaNO3 and KCl, are both soluble
so, no reaction
Use reference table F to predict solubilities
Pb I- = PbI2

10. Kinetics All substances contain chemical potential energy A  B
high PE low PE
Energy released -- Exothermic
low PE high PE
Energy absorbed -- Endothermic
Most chemical reactions require some energy to get started --
activation energy

11. Potential Energy Diagrams
A – PE of reactants 50 40 30 20 10 PE
B – Activation energy B kJ
C – PE of products
D – Heat of reaction; ΔH D A C
Progress of Reaction
In this diagram, PE decreased
energy released; exothermic
ΔH is negative

12. Exothermic reactions have
a negative ΔH
Endothermic reactions have
a positive ΔH
See reference Table I Notice that ALL values of ΔH
have a sign (+/-)!

13. How do we tell if a reaction will occur spontaneously?
Consider 2 factors:
1. Heat of Reaction (ΔH)
--low energy conditions are more stable
-- exothermic reactions (- ΔH ) are favored
2. Change in entropy (ΔS)
--entropy is a measure of disorder, or randomness
gas
liquid
solid
disordered
neat, orderly
high entropy
low entropy
--increasing entropy (+ ΔS ) is favored

14. Favorable conditions are:
Exothermic -ΔH
Increasing entropy +ΔS
Any reaction with -ΔH, +ΔS is always spontaneous
+ΔH, -ΔS is never spontaneous
Reactions where only one factor is favorable are spontaneous under some conditions – depends on temperature.
Example: 2C(s) + 3H2(g) → C2H6(g)
ΔH = kJ
Exothermic, favorable (Table I)
ΔS solid to gas, entropy increases, favorable
So this reaction is spontaneous

15. Example : H2O (solid)  H2O (liquid)
Types of Reactions
heat of reaction(ΔH) [favorable?] entropy (ΔS) [favorable?]
1. decrease ( - )
[yes]
increase ( + )
[yes]
YES
[no]
decrease ( - )
2. increase ( + )
[no] NO
3. decrease ( - )
decrease ( - )
Only at
low temp
[yes]
[no]
4. increase ( + )
[no]
increase ( + )
[yes]
Only at
high temp
The result in cases 3 and 4 depends on temperature
Example : H2O (solid)  H2O (liquid)
-- solid to liquid, energy increases, energy absorbed
ΔH positive, unfavorable
-- solid to liquid, entropy increases
ΔS positive, favorable
So,
ice melts spontaneously only at high temperatures

16. Reaction Rates Collision Theory
--reactions occur when molecules physically collide
-- the collision must be effective – it must have
1. enough energy
2. proper orientation
-- any change that increases the number of effective collisions
will increase the reaction rate

17. Factors affecting reaction rate
1. Temperature --
faster motion of molecules means more frequent
and more effective collisions
2. Surface Area --
more contact between reactants means more
frequent collisions
(increase surface area by dividing a solid into many small pieces)
Concentration of Reactants –
higher concentration = more frequent collisions
in solutions – raise concentration by dissolving more reactant, or by reducing water
in gases – raise concentration by raising pressure

18. 4. Nature of Reactants --
--if bonds must be broken first, reaction is slow
--if not, reaction is fast ex., reactions between solutions
Catalysts– a catalyst is a substance that
increases the rate of a chemical reaction
catalysts are not used up during the reaction
catalysts do not affect the ΔH or the
spontaneity of the reaction
catalysts work by lowering the
activation energy

19. ineffective
no catalyst
low energy collision
catalyst
effective

20. Equilibrium Kinds of Equilibrium
-- a state of balance between opposing forces
-- no overall changes occur at equilibrium
-- in a dynamic equilibrium, opposing changes occur at equal rates
and cancel each other
Kinds of Equilibrium
1. Phase equilibrium:
ex.) ice and water at 0ºC
ex.) water and water vapor in a closed jar
2. Solution equilibrium
3. Chemical equilibrium

21. Solution Equilibrium
When solute is added to solvent, it dissolves until an equilibrium
is reached:
solid dissolved
The rate of dissolving = the rate of precipitation
when the solution is saturated.
A saturated solution contains the maximum amount of solute
that can be dissolved under specific conditions, like temperature.
By comparison, an unsaturated solution contains less, and a supersaturated solution contains more. (Table G)
The solubility of a
substance is its
concentration
at equilibrium
(at saturation with
given conditions)
at STP

22. Factors affecting solution equilibrium:
A. solid/liquid solutions
-- Temperature
almost all solids are more soluble at high temps

23. B. gas/liquid solutions
-- Temperature
gases are less soluble
at high temperatures
See reference
Table G
The lines that
go DOWN are
gases
-- Pressure
gases are more soluble
at high pressures. Think
of soda. What happens
when you open it?

24. Chemical Equilibrium A  B Many chemical reactions are reversible
forward reaction A  B
reverse reaction B  A
Both reactions use the same PE diagram
forward: ΔH +; endo
act. E
act. E B
reverse: ΔH ; exo ΔH ΔH
ΔH’s are equal but
opposite A

25. A  B Equilibrium occurs when:
1. the rate of the forward and reverse reactions are equal
2. the concentration of the product and reactant remain constant
3. no visible changes are occurring

26. Equilibrium Shifts
LeChatelier’s principle: When a system at equilibrium is subject
to a change, the equilibrium will shift to undo the change
Three factors can affect chemical equilibrium:
1. Concentration of product or reactant
Example: A + B  C + D
is at eq.; then more D is added
Result:
the equil. will shift to reduce the amount of D
more A and B will form, and C will be used up to combine
with the excess D.
the equilibrium shifts left
Example:
AgCl(s)  Ag+(aq) + Cl-(aq)
at eq.; then NaCl is added
Result:
adding NaCl would add more Cl-; therefore, the equil.
shifts left, and AgCl(s) will precipitate

27. 2. Pressure: affects equil. involving gases
If pressure increases,
the equil. shifts to the side with
fewer moles of gas
Example: N2(g) + 3H2(g)  2NH3(g) at eq., then pressure
1 mole + 3 moles  2 moles
Result: to cause the pressure to go back down, the eq. shifts right
Example: CaCO3(s) + HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
What happens to the amount of CaCl2 if pressure decreases?
Result: moles  mole
Eq. shifts right to restore pressure,
producing more CaCl2
If gases are equal on both sides, equil. does not shift

28. 3. Temperature: consider heat as a product or reactant
Example: exothermic reaction
NaOH + HCl  NaCl + H2O + heat
at eq., then temp
Result: Eq. shifts left to get rid of excess heat
Example: endothermic reaction
N2 + O2 + heat  2NO
at eq., then temp
Result: Eq. shifts right
Catalysts do not shift the equilibrium --
they speed up the forward and reverse reactions equally