1. By Logan Danielson

2. Titrations
Titrations are done to determine the unknown
concentrations of unknown acids.
It is based of off determining how much known acid or
base is needed to reach the equivalence point of the
unknown liquid and at what pH the equivalence point
exists.
The equivalence point is the middle point of the fastest
changing pH area. Usually a color indicator is used to help
find the equivalence point as well as a pH indicator.

3. Solubility Equilibria
Adding an ionic solid to water
It will be in equilibrium so long as some solid exists in
solution
Mg(OH)2(s)<-> Mg2+(aq) + 2 OH-(aq)
Ksp= [Mg2+][OH-]2
 This is the solubility product constant for the above equation
 Note, the solid is not part of the equation

4. Predicting precipitation
For Mg(OH)2(s)<-> Mg2+(aq) + 2 OH-(aq)
If Q= [Mg2+][OH-]2 < Ksp then the there will be no
precipitate.
If Q > Ksp then the there will be a precipitate.
Note, if Q was < Ksp then only Mg2+(aq) or OH-(aq) would need
to be added to cause precipitation.

5. Thermodynamics
Is the flow of energy
1st law of thermodynamics
 The total energy of the universe is constant, energy can not
be created nor destroyed
 If you decrease the energy in the system, then you increase the
energy of the surrounding
 If you increase the energy in the system, then you decrease the
energy of the surrounding
 U = internal energy = all potential energy + all kinetic energy
 dU universe = dU system + dU surroundings
 dU system = - dU surroundings

6. Thermodynamics 2: heat and work
dU system = Heat (q) + Work (w)
 q= Heat = energy flow from the change in temperature
 w= Work = energy flow from movement against a force
Closed system- do not allow matter in or out of the system
q<0 – heat flows out of the system
q>0 – heat flows into the system
w<0 – the system does work on the surrounding
w>0 – the surrounding does work on the system

7. Thermodynamics 3: enthalpy
The most common work in chemical systems
w= -P dV
 Work= - Pressure * change in Volume
dU system = q – P * dV
dH (enthalpy) = dU + d(PV)
dH = dU + P*dH @ constant P
Plugging in “dU system = q – P * dV” you get:
dH = q - P*dH + P*dH = q @ constant P
dH = q @ constant P

8. 2nd law of thermodynamics
All spontaneous processes increase the entropy of the
universe
dS (entropy) > 0 for spontaneous processes
Entropy is a measure of the dispersive-ness of energy
A larger value of entropy is a more dispersed energy

9. Guessing thermodynamics
C6H12O6(s) -> 2 C2H3OH(s) + 2 CO2(g)
dH>0 because we are breaking more bonds then we are forming
dS>0 because there are more moles on the right
H2O(s) -> H2O(l)
dS>0 because liquid is more dispersible than solid
H2O(l) -> H2O(g)
dS>0 because gas is more dispersible than liquid

10. Entropy guessing rules
Increase Entropy
Breaking bonds without making new ones
Change to a favored phase
More moles on the product side

11. Entropy
dS universe = (dS system + dS surroundings) > 0
dS surroundings = -q / T
dS system - q / T > 0
At constant P, q=dH so:
 dS - dH / T > 0
system
For phase changes:
dS system - dH / T is approximately 0, so:
 dS system = dH / T

12. Calculating entropy
aA + bB -> cC + dD
dS reaction = c*SCo + d*SDo – a*SAo – b*S Bo
SAo is the standard entropy of A

13. Gibbs free energy
 G = free energy
 dG = dH – TdS @ constant P and T
 dG < 0 for a spontaneous process
 Reactions that are dG>0 won’t happen, but their reverse reaction will
 Reactions that are dG=0 will have nothing happen because the system is at
equilibrium
 The free energy and the equilibrium constant are related:
 dGo= - RT ln K
 dGo is the dG at standard temperature and pressure
 R is the gas law constant, R=8.314 J/(K*mol)
 T is the temperature
 ln K is the natural logarithm of the equilibrium constant

14. Gibbs free energy
dG = dH – T*dS @ constant P and T
dHo dSo dGo
>0, endothermic >0, increase in energy (+)-T(+), more negative at higher temperatures
<0, exothermic <0, decrease in energy (-)-T(-), more negative at lower temperatures
>0, endothermic <0, decrease in energy (+)-T(-), always positive, never spontaneous
<0, exothermic >0, increase in energy (-)-T(+), always negative, always spontaneous
At low T: dG is approximately= dH
At high T: dG is approximately= – T*dS