Problem Set 2

There's a key posted for the problem set we did in class that was titled "Problem Set #2" on my Gen Chem webpage:
http://www.drbodwin.com/teaching/genchem.php
Direct link to the key:  http://www.drbodwin.com/teaching/problemsets/c150gps02k.pdf

Don't forget, there are also old exams posted at:
http://www.drbodwin.com/teaching/examarchive.php

Other questions, let me know…

Redox Lab question...

A few people have asked about the chemical reactions for the Redox lab hand-in {Redox Hand-in}, and this is often a source of confusion for students, so let me answer it here. First of all, I think a number of people get confused because you're trying to over-think the question and make it more complex than it really is.

So you need to draw a voltaic cell... you can draw this by hand, or you can draw it electronically. DO NOT just find one online and copy-paste it into your hand-in, if that's all I wanted I would have pasted this {voltaic cell} into the hand-in myself before I posted it.

The part that causes some confusion is the "...write a correctly balanced net ionic equation for the spontaneous process..." Don't over-read that! The metal cation solutions you were using in lab were probably nitrate salts, but nitrate (or whatever anion might have been present) was a spectator in all of your reactions. Net ionic equations are actually easier than full-formula equations because they're not cluttered up with with a bunch of extra stuff, and net ionic equations actually just describe the CHEMISTRY that's happening rather than distracting you with a bunch of spectator ions and species.

Now, I'm not going to write out a net ionic equation that's the exact answer to one that you have to write, but here's an example. Let's say I made up the voltaic cell Fe|Fe⁺³||Cd⁺²|Cd with the black/negative lead of my meter hooked up to the Fe(s) electrode and the red/positive lead connected to the Cd(s). The potential I measure is -0.32V. Because the measured potential is negative, the cell is running backwards, so the spontaneous cell reaction is Cd|Cd⁺²||Fe⁺³|Fe. Translating that into a reaction, we can write the two half-reactions as:

Cd(s) <=> Cd⁺²(aq) + 2 e^- 

3e^- + Fe⁺³(aq) <=> Fe(s)

Adding those up gives the overall (or "net") reaction:

3Cd(s) + 2 Fe⁺³(aq) <=> 2 Fe(s) + 3 Cd⁺²(aq)

{Remember to multiply each half-reaction by an appropriate integer to make all the electrons cancel...}

That's a "correctly balanced net ionic equation for the spontaneous process" in this case. Now go do that for all the cells you measured in the experiment.

Redox lab question

Quite a few people have questions about the redox lab, so let me give a hint/some guidance to everyone...

In the first part, you looked at reactivity and found Zn to be the most active metal, followed by Pb, then Cu.  I'll just use those three as an example, you will also need to include Ni and Ag in your assignment.  You measured the potential for a Zn/Pb cell, a Pb/Cu cell, and a Zn/Cu cell.  Is there a relationship between those measured potentials?  There's a relationship between those reactions, but how are cell potentials related to one another?

Zn(s) + Pb²⁺(aq)  ⇄  Pb(s) + Zn²⁺(aq)

Pb(s) + Cu²⁺(aq)  ⇄  Cu(s) + Pb²⁺(aq)

Zn(s) + Cu²⁺(aq)  ⇄  Cu(s) + Zn²⁺(aq)

Are cell potentials like kinetics (the cell potential for the overall process is determined by the lowest potential)?  Are cell potentials like equilibrium (the cell potential for the overall process is the product of the step-wise potentials)?  Or is there another relationship between the step-wise potentials and the overall potential?  When you think you see (observe) a relationship with the Zn/Pb/Cu system (hypothesis), check to see if the same relationship is true with some of the other cell combinations you measured.(test/experiment)

{Hmm, it looks like we could use the scientific method to analyze the data and results from this experiment.  Who would have guessed?!}

On your hand-in assignment for lab, the "calculated" cell potentials for pairs that are not next to each other refers to the treatment you see above.  You have measured all of the potentials for cell constructed from metals that are adjacent to each other in your activity series (step-wise potentials), so if there is a relationship between step-wise potentials and the potentials for cells constructed from metals that are not adjacent to each other in your activity series, you should be able to calculate the expected cell potential for those non-adjacent cells.

Voltaic Cells

Not a lot of information right now, but I just finished drawing a voltaic cell diagram.  Not perfect, but I'm pretty satisfied with it.

Balancing Redox Reactions

Balancing redox reactions can be pretty simple for some system, but some redox reactions can be exceptionally challenging. To balance any redox reaction, we can follow a systematic set of steps, and every Gen Chem book happily provides a set of steps. In my experience, most books use rules that require a little bit of faith and function like a black-box. There are a LOT of places that mistakes can be made when balancing redox reactions, so I prefer to use rules that have built-in places to check my answer before I go through the whole process. Here they are:

Balancing redox rules (in acidic or neutral aqueous solutions):

1. Assign oxidation numbers to all atoms in the equation

2. Identify the oxidation and reduction half reaction

3. In each half reaction, balance all atoms except hydrogen and oxygen

4. In each half reaction, add electrons to the reactant or product side to balance the change in oxidation state. Note: you are not adding electrons to balance charge

5. In each half reaction, add water molecules to balance any oxygen atoms

6. In each half reaction, add H⁺(aq) to balance any hydrogen atoms

7. At this point, the half reactions should be balanced, check the charge balance to confirm

8. Multiply each half reaction by an appropriate integer to balance the electrons involved in the oxidation and reduction processes

9. Add the half reactions together

10. Again, the resulting reaction should be balanced, check the charge balance to confirm

11. Cancel out any spectator species, sit back and pat yourself on the back for writing such a lovely balanced redox equation.

These rules work well, BUT rely upon some assumptions. First, since we're using water and H⁺(aq), the reaction must be taking place in aqueous solution that is neutral or acidic. That's OK for most redox reactions in Gen Chem, but once in a while we'll run into a rxn that takes place in basicaqueous solution. How do we handle that? Think about it... basic solutions have excess (relatively speaking) hydroxide ions. Hydroxide ions react with H⁺(aq) ions to form water. If a reaction is taking place in basic aqueous solution, balance it according to the above rules and then add:

12. Add enough OH^-1(aq) to each side to react with all the H⁺(aq) that is present

13. Check charge balance

14. Cancel any excess water

As with any process, practice is the key, so practice balancing redox rxns, then practice a little more, and when you think you have it all figured out, practice a couple more times. We'll do some of that in class...

Oxidation Numbers

Oxidation numbers describe the balance between electrons and protons on an atom, whether that atom is happily floating around all by itself or part of a massive molecule. Oxidation numbers can be determined two different ways: by using rules, or by looking at structure. Let's look at the rules first.

Oxidation Numbers by the Rules:
1. For neutral, uncombined elements, Ox# = 0. Examples: Fe(s), H₂(g), Hg(l), Ne(g)
2. For monoatomic ions, Ox# = charge. Examples: Fe²⁺(aq) {Ox# = +2}, P^3-(g) {Ox# = -3}
3. Oxygen is almost always Ox# = -2, except in O₂ {Ox# = 0, Rule #1} and peroxides {Ox# = -1}
4. Hydrogen is almost always Ox# = +1, except in H₂ {Ox# = 0, Rule #1} and hydrides {Ox# = -1}
5. The sum of the Ox#s on all the atoms in a polyatomic molecule or ion is equal to the charge on the whole polyatomic molecule or ion.

Let's look at a redox reaction and assign Ox#s by the rules:

Cl₂(g) + 2 O₂(g) ↔ 2 ClO₂(g)

Cl₂(g) : Rule #1, Ox# = 0

O₂(g) : Rule #1, Ox# = 0

ClO₂(g) : Rule #3, oxygen is Ox# = -2.

ClO₂(g) : Rule #5, (Ox# Cl) + 2(Ox# O) = 0 (the charge on a neutral molecule)

(Ox# Cl) + 2(-2) = 0

(Ox# Cl) = +4

So in this redox reaction, each Cl is going from 0 to +4, losing 4 electrons, Losing Electrons is Oxidation; and each O is going from 0 to -2, gaining 2 electrons, Gaining Electrons is Reduction.

For many substances, it's actually easier to assign Ox#s by looking at the structure. The process is very similar to finding Formal Charge, the electrons are just assigned a little differently. Formal Charge assigns electrons as if all bonds are purely covalent, meaning that all bonding pairs of electrons are split with one electron given to each atom in the bond. Oxidation Number assigns electrons as if all bonds are purely ionic, meaning that all of the bonding electrons go to the more electronegative element in the bond.

Oxidation Number by the Structure:
1. Draw a good Lewis Structure
2. Assign all bonding electrons to the more electronegative element in the bond
3. Compare the electrons assigned to each atom to the valence electrons of the neutral element

Again, let's look at an example, in fact, let's look at the sameexample as above, ClO₂(g). Drawing a good Lewis Structure:

This is a very interesting molecule, it violates the octet rule andit has an unpaired electron. Looks like it would be pretty reactive. Looking at the electronegativities, oxygen is more electronegative than chlorine, so all of the bonding electrons will be assigned to oxygen, giving each oxygen 8 assigned electrons and the chlorine 3 assigned electrons.

Neutral oxygen has 6 valence electrons, we've assigned 8 electrons to oxygen, so the oxidation number for oxygen in this molecule is -2, just like we predicted using the rules. Neutral chlorine has 7 valence electrons, we've assigned 3 electrons to chlorine, so the oxidation number for chlorine in this molecule is +4, again, just like we predicted using the rules. Both methods work. Why would you ever use structures when the rules work? Try looking at hydrogen peroxide, H-O-O-H, or a more complex molecule like glucose. The rules don't always give the best picture of what's happening in a molecule, and this can ultimately make it harder to predict reactivity or other behaviors.

Redox - Definitions

If you want to understand chemistry, you have to follow the electrons. If electrons are transferred during a chemical reaction (as opposed to just being rearranged...), then a reduction-oxidation process is taking place. To help keeping track of the electron transfer in redox processes, we can use a couple acronyms/mnemonics related to the definitions of reduction and oxidation.

OIL / RIG - “Oxidation Is Losing electrons” / “Reduction Is Gaining electrons”

LEO / GER - “Losing Electrons is Oxidation” / “Gaining Electrons is Reduction”

Hmm, why is “reduction” associated with gainingelectrons? Remember, electrons are negatively charged, so gaining electrons increases the number of negatively charged particles associated with an atom which reduces its net charge. But are we really looking at “charge” to determine redox chemistry? Sometimes it seems like it, but other times the charge doesn't seem to line up with the processes. We really have to look at oxidation number, which is related to charge, but a with some subtle differences. One way to distinguish charge and oxidation number is that “charge” can be used to describe the net overall balance between electrons and protons in a system that might contain multiple atoms, but “oxidation number” describes the balance between electrons and protons for each individual atom in a structure regardless of its size. Oxidation numbers sound important enough for their own post, look for it soon.

Another thing that can cause some mix-ups is the term “oxidation” or {in verb form} “oxidize”. {Or for the British English spellers in the crowd, “oxidise”.} The element oxygen is very often involved in redox reactions. Is oxygen usually undergoing oxidation or reduction? You know you want to say oxidation, the words look so similar... But if we start with molecular oxygen, O₂(g), it's almost always going to gainelectrons. Gaining Electrons is Reduction. GER, indeed! The important thing to remember here is that reduction and oxidation are ALWAYS coupled processes. You can't have one without the other. This leads to some other terminology...

The process of one substance undergoing oxidation causes something else in the system to be reduced. The substance that is being oxidized is the reducing agent or reductant because it is causing reduction to take place.
The process of one substance undergoing reduction causes something else in the system to be oxidized. The substance that is being reduced is the oxidizing agentor oxidant because it is causing oxidation to take place.

So if oxygen is usually being reduced, it is a good oxidizing agent.