Schedule1 Schedule2 Schedule3 Schedule4
Schedule to Exam 4
| Date | Day | Topics | 4/6 | Mon | talk about preparing for exam 3 - introduced a little electrochemistry. Electric potential is measured in volts (voltage). |
|---|---|---|
| 4/8 | Wed | Talked about exam 3. Showed bar graphs of score distributions. Showed the most missed problems and discussed them. Said to READ/STUDY Chapter 4 sections 10-12 for a good redox review. Talked about the importance of oxidation numbers and how they help us identify what is being oxidized and reduced in reactions. Discussed RULES for assigning oxidation numbers (p 118 - blue box in your book). Defined oxidizing agent and reducing agent. |
| 4/10 | Fri | Discussed activities of metals. More active means more likely to lose electrons. Will now use a table of standard potentials to determine activities. A more active metal will be on the RIGHT side of the rxn in the table and will be more negative. The ultimate active metal is lithium which is at -3.05 V for the Li-ion + electron goes to lithium metal. Showed how you combine two 1/2-reactions to get an overall or whole reaction. Defined and showed the differences in electric current in wires (e- flow) and electrolytic current in solutions (ions migrating). Electrodes provide the conduit between e- flow and electrolytic current. Defined the 2 types of electrodes. Cathode is where reduction takes place and the anode is where oxidation takes place. |
| 4/13 | Mon | Discussed how to "read" the table of standard potentials for determining the order of oxidizing agents and reducing agents. Drew a voltaic cell for the Zn/Cu half reactions. Labeled the parts and showed electron flow. Note: electric current (I) is opposite of electron flow. Showed the shorthand notation for cells --- anode on the left and cathode on the right. Standard potential for an overall rxn or cell is the cathodic std potential minus the anodic std potential. Remember you SUBTRACT the anodic standard potential. Or in other words, you "flip" the reaction from the table for the anode and show it as an oxidation, you then change the SIGN on the potential and add it to the reduction potential from the cathode. Also showed the relationship between ΔG and E, which is: ΔG = -nFE. |
| 4/15 | Wed | Showed the concentration dependence of cell potential which of course leads to the Nernst equation. KNOW the Nernst equation for the exam. It is basically this: E = E° - RT/nF ln Q . The RT/F is used so often at 25°C (298.15 K), that we just calculate it and remember it as 0.0257. Then we can also switch to good ol' "log" instead of "ln" by using 0.05916 . So most of us use this version: E = E° - 0.05916/n log Q. |
| 4/16 | Thu | homework 12 due by noon |
| 4/17 | Fri | Went over in more detail the shorthand notation we use for cells which is : anode | anodic solution || cathodic solution | cathode . I then used this notation to show the S.H.E. and the Ag/AgCl standard electrodes. Also showed how potential will DECREASE as you draw current from a voltaic cell when the active species are solution species. Why? The product concentrations will increase and the reactant concentration will decrease as the cell is run. This will increase the value of Q and using the Nernst equation you can easily conclude that the potential will drop. This is NOT what we want for battery use. We need a battery to provide a fairly steady potential throughout its useful lifetime. The solution for this is to use more solids as the active parts (reactants and products) in the battery. Pure solids will maintain their activities of 1 (ONE) no matter how much you have. As long as you have some solid present, you'll have the same activity. So Q will remain steady throughout the lifetime of the battery. Batteries: Primary = non-reversible rxn and therefore NOT rechargeable (dry cells, heavy duty, alkaline, Li). Secondary = reversible rxn and therefore IS rechargeable (NiCad, NiMH, Li-ion, Pb-storage). Fuel cell = refillable reactants (H2/O2, and Zn-air). You are responsible for knowing the reactions for the Pb-storage battery - see page 481 in our book. Realize what is shown there is for discharge of the battery. You'll need to reverse the reactions for recharge. |
| 4/19 | Sun | homework 13 due by 3 PM |
| 4/20 | Mon | Talked a bit more about batteries. Another mention of
the lead-storage battery and its reactions. A "true" batteries is a series of cells... the 12V
car battery has six 2 volt cells in a series. Asked what limits current for batteries? Answer
was the surface area of the electrodes. More surface area will allow for more current density and
more electrons passed in a given amount of time. Different sized cells (D, C, AA, and AAA) each
have different maximum current capabilities although all of them have the same voltage. Then
discussed the faraday constant in calculating molar amounts in a 1/2 rxn from current/time information
from the electricity side of things. I gave a formula that tends to work for most of these type
calculations: I t / n F = mol of stuff.
I then introduced chemical kinetics (chapter 15) as the study of reaction rates. Thermodynamics does NOT tell you anything about the rates of the given reactions. You MUST measure rates in the laboratory. At the end of class I wrote down the 4 factors that affect reaction rates: nature of the reactants, concentration, temperature, and presence of a catalyst. |
| 4/21 | Tue | homework 14 due by noon |
| 4/22 | Wed | Showed how rate can be measured via ANY species in the reaction, both reactants and products. The overall rate of the reaction is doubled, tripled, etc... when there are coefficients. So for aA + bB --> cC + dD you get the following rates for the overall reaction: rate overall = -Δ[A]/aΔt = -Δ[B]/bΔt = +Δ[C]/Δdt = +Δ[D]/dΔt . Note the reactants have negative signs and the products have positive signs. |
| 4/24 | Fri | Showed how to determine the rate law via the method of initial rates. This is where you evaluate data from a table with starting concentrations and the rates. You try to make the rate ratio equal the concentration ratio via an exponent on the conc ratio. The exponent is the order in that component. I then integrated the rate law to get the time based equation. You get different equations for different orders. First order is ln([A]0/[A])=akt, second order is 1/[A]-1/[A]0=akt, and zero order is [A]0-[A]=akt. |
| 4/27 | Mon | Pointed out again the THREE integrated rate laws for zero, first, and second order reactions. Also showed the appropriate straight line plots for each type. First order was special in that you could use ANY concentration term for [A]0 and [A] because the units will cancel. This allows one to use straight up percentages as concentration terms in the first order equation. I then defined half-life and got an equation for half-life for each of the three orders. I then gave the Arrhenius equation which relates temperature with rate constant and activation energy (Ea). |
| 4/29 | Wed | LOTS of good stuff covered today! Collision theory: You must
have reactants collide in order to form products BUT not all collisions will lead to products.
Only Effective Collisions will lead to products. 2 Criteria for Effective Collisions are (1)
Must have enough ENERGY. You control this via temperature, (2) Orientation - molecules must
align properly in most cases in order to react. Collision Frequency is proportional to the rate.
You can INCREASE collision frequency by increasing the temperature or concentration (or both).
Reaction Mechanisms: this are a series of elementary step reactions that combine to give the overall reaction. The mechanism should be consistent with experimental data. An elementary step will just be ONE or TWO reactants and that is all. A combo of 3 reactants is very very improbable. Any elementary step can immediately have its rate law written. The overall rate is going to be the same as the slowest STEP in the mechanism. You cannot have intermediate species in the overall rate law though. So you must substitute in for intermediates with other reactants that were used in previous steps. Previous steps to a slow step will generally reach equilibrium. The KINETIC definition of equilibrium is the point in a reaction at which the forward rate equals the reverse rate. Writing this out one can see that the equilibrium constant for a reaction is equal to the ratio of the forward rate constant to the reverse rate constant. That is: K = kf/kf. |
| 4/30 | Thu | homework 15 due by noon |
| 5/1 | Fri |
More on rxn mechanisms. Elementary steps are reactions with KNOWN molecularities. Unimolecular is a one-body reaction (one reactant) and is first order in that reactant. Bimolecular is a two-body reaction and is either first order in each of two different reactants (like A + B ) or is second order in a reactant colliding with itself (like A + A or 2A). Termolecular reactions are rare because three-body collisions are very very low probability. Reaction Energy Diagrams. The reactants are shown at their energies and the products at their energies. In between is an energy maximum known as the transition state. To get from reactants to the transition state you must increase the energy by an amount known as the activation energy. There is a forward and reverse activation energy for any reaction (step). After showing a single step reaction, I showed a 3-step reaction with 3 "humps" for each step. LOTS of quantitative data can be extracted from a diagram such as this. |
| 5/4 | Mon | Chapter 15 |
| 5/5 | Tue | homework 16 due by noon |
| 5/6 | Wed | talk about exam - finish Chapter 15 |
| 5/7 | Thu | EXAM 4, 7-9 PM |
| 5/8 | Fri | Last Class day - discuss Final Exam and Grading |