|Lecture Notes::Lec 8_8 June||
In Chem 109 we used le Chatelier's Principle to predict the effects of changes in temperature on reaction equilibria. Thus, for example, for an exothermic reaction, if T increases heat must have been added, therefore the reaction will use up some of the extra heat, driving the reaction to the left and lowering the temperature back toward its original value. With our new thermodynamic knowledge we can determine this relationship quantitatively.
Substituting we can say
dividing by - RT:
Assuming deltaH° and deltaS° are constants (a reasonable approximation if deltaT is not large) this is now in the form of the equation of a straight line:
So we can find both deltaH and deltaS (as seen in the lab).
We can also find these values non-graphically by looking at a system at two different temperatures and doing a bit of algebra:
Subtracting (2) from (1) we get:
And using log rules:
Another common form of the van't Hoff Equation is:
Reversible vs. irreversible processes. In order for a process to be completely efficient we say it must be reversible - that is we can get back to the starting point without energy loss. At the microscopic level (atoms and molecules or smaller) processes are frequently reversible. However at the macroscopic level there are inevitably heat losses due to friction etc.
For real processes a further complication arises in that the only energy available is the Free Energy, so we measure the efficiency of a process by how much of the free energy (deltaG) is captured.
As state above, in real processes some energy is lost to the environment as heat. Folks often lament this loss because it appears to be wasted - wouldn't it be better to capture all the deltaG to do work? If so why hasn't nature done so? After all life on Earth has had somewhere around 3.5 x 109 years to get it right, and glycolysis still "wastes" about 40% of the available deltaG. Why? Unfortunately in order to capture the maximum amount of deltaG we have to accomplish things in very small equilibrium steps. Very highly simplified, if you're very efficient, its hard to generate enough working energy, and you get eaten! For human designed processes you have to be somewhat inefficient to get anything done during your lifetime.
So like many things we have to compromise and try to optimize the process between speed and completion versus efficiency. The challenge is to reach this optimum, which generally changes with circumstance -life's tough!
Why do we care? Examples of electrochemistry all around us (batteries, plating, corrosion, fuel cells, manufacture etc.)
Electrochemistry is the interchange of electrical and chemical energy.
First let's review Redox chemistry.
Terms: Look briefly at some terms used in redox/electrochemistry.
Note that in any chemical redox reaction that oxidation and reduction are always coupled getting an exchange of status:
Balancing Half-Reactions. In electrochemistry we separate a redox reaction physically into half-reactions. Thus we need to be able to separate a redox reaction into half-reactions and to balance them, as we did in Chem 109. So let's review redox balancing by the half-reaction method. In the half-reaction method what we do is first break an equation into two parts and then balance the parts individually. We will just review the method for reactions in acid solution. But remember the same method works for basic solutions as well with a few additional steps.
Presented step wise for Acid Solution:
Example. Balance the following equation as it occurs in acid solution:
MnO4- + Cl- Mn2+ + Cl2
First break the equation into two half reactions, one for Mn and one for Cl, then follow the steps above for each:
MnO4- Mn2+ + 4 H2O
8 H+ + MnO4- Mn2+ + 4 H2O
5 e- + 8 H+ + MnO4- Mn2+ + 4 H2O
10 e- + 16 H+ + 2 MnO8- 2 Mn2+ + 8 H2O
2 Cl- Cl2
2 Cl- Cl2 + 2 e-
10 Cl- 5 Cl2 + 10 e-
Finally, combine half reactions and cancel terms:
10 e- + 16 H+ + 2 MnO4- + 10 Cl- 2 Mn2+ + 8 H2O + 5 Cl2 + 10 e-
Consider a simple Redox system consisting of a beaker full of copper sulfate and a zinc strip as an example:
When we place the zinc strip into the Cu2+ solution we will observe a gradual darkening of the zinc strip while the solution gradually goes from light blue to colorless. (you may recall doing this this reaction in the Ionic Reactions lab in Chem 109).
So what's going on?
What we'd really like to do is capture the energy of this redox reaction as a flow of electrons. After all our half-reactions tell us electrons are being exchanged. So how do we do this? We can separate the reaction into its two half-reactions by putting the components of each half reaction into its own container. We then have:
If we now connect the two metal strips in our set-up with a wire with a meter on it what will occur?
Before we go further, we need to understand the units and instruments for measuring electricity. These are discussed in the box below.
Let's go back to our two containers each with its own half-reaction and connect them with a wire. What happens? If we watch very carefully with a sensitive current meter we should see a pulse of electricity followed by a return to zero.
Why does the current not continue? The problem is that the charges in the two beakers quickly build up until the free energy of the redox reaction is not sufficient to over come the work needed to move the charges against the potential gradients.
So what can we do to allow the flow of electrons to continue? We need to connect the solutions in the beakers so that the charges can be neutralized with a counter flow of ions. The connecting ionic fluid is referred to as a salt-bridge, as seen in the figure. Other arrangements are possible such as semi-permeable membranes etc. Such an arrangement is called a Galvanic cell.
Galvanic cells (or Voltaic cells) are cells in which the overall redox reactions occur spontaneously (equilibrium favors products) as written. They can serve as a source of electric power (as a battery).
© R A Paselk
Last modified 8 June 2006