Review Sheet - Exam 3 - CH302 - McCord

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Textbook Chapters for Exam 3

Chapter 8, sections 8-10. Chapter 11, all sections (1-8). You’re responsible for any material covered in class whether it was on the homework or not. This review sheet covers/mentions most everything we have done. Also remember that no equations are given on the exams so you need to know them. Values for equilibrium constants and any standard potentials are given where needed.

Chapter 8 - Solubility Equilibria

Know how to calculate the molar solubility (\(x\)) from \(K_{\rm sp}\)

Know how to calculate \(K_{\rm sp}\) from molar solubility. Do realize that for many problems you must first convert the plain ol’ solubility (g/L, mg/L, g/100mL, ppm, ppb, etc...) into molar solubility first, then convert that to \(K_{\rm sp}\).

Know also how to calculate the apparent solubility in the presence of a common ion. This is where one of the ions concentrations is already SET in the solubility product expression. My example in class was dissolving AgCl into water (no common ion) or into 0.001 M NaCl (common ion of Cl^-).

I already said this above but I’ll say it again: be able to calculate molar solubility (\(x\)) from “plain” solubility which can be expressed as grams per liter or, in general, mass per unit volume like ppm (mg/L) or even g/100 mL..

Know how to calculate all final concentrations in saturated solutions. OK, so you can calculate \(x\) from \(K_{\rm sp}\). Now tell me what the actual concentration of Mg²⁺ is, or Cl^-, or OH^-, or etc... sometimes it is \(x\), sometimes it’s \(2x\), or \(3x\), and so on. Remember, \(x\) isn’t always the answer.

REMEMBER: Not all salts have the same solution for \(x\) from \(K_{\rm sp}\). There are 1:1 salts, 1:2 salts, 1:3 salts, 2:3 salts and so on. Also remember that 3³ = 27 and not 9.

Be able to predict whether a precipitation will occur. This is just comparing \(Q_{\rm sp}\) to \(K_{\rm sp}\). If a precipitation does occur, how much does precipitate? What are the concentrations of the ions in solution after the precipitation?

Fractional or Selective Precipitation

Know what will precipitate 1st, 2nd, etc.. in solutions containing many different ions that can precipitate. Remember that to do this you must calculate the minimum concentration of the added ion in order to reach the saturation point (solve the \(K_{\rm sp}\) expression). The one that precipitates first is the one with the lowest concentration needed of the added ion.

Know how to solve questions about fractional precipitation or selective precipitation. What % of the 1st precipitate is precipitated (or % NOT precipitated) when the 2nd precipitate first starts to precipitate?

Dissolving Precipitates

What do you add/use to get certain insoluble compounds to dissolve? Usually this involves some sort of chemical reaction with one of the ions of the salt. As the ion from the salt reacts, its concentration starts to drop and Le Chatlier’s Principle takes over shifting the reaction to the right thus dissolving more salt.

Complex Ions

Know what a complex ion is (metal cation + ligands) and how they dissociate. Know how to write a \(K_{\rm f}\) or \(K_{\rm d}\) expression and how to SOLVE it. Remember that \(K_{\rm f}\) is a formation constant and \(K_{\rm d}\) is a dissociation constant and they have an inverse relationship:

\[K_{\rm d} = {1 \over K_{\rm f}}\]

Chapter 11 - Electrochemistry

Remember your redox chemistry - Chapter 4. Know the basic concepts involved in redox reactions.
Know how to assign oxidation numbers to all the elements in a given formula.
Know how to balance a redox equation in acid.
What is the difference in oxidation and reduction. KNOW your definitions here!
Know what a redox reaction is. What is being oxidized and what is being reduced? What is the oxidizing agent and what is the reducing agent?
What is a half-reaction? Why is it so useful?
Know how to write half-reactions for both oxidations and reductions.
Know how to combine 2 half-reactions to give an overall cell reaction (make sure the electrons balance). This is essentially the same as what was stated above except that in electrochemical cells we run the 2 half-reactions in 2 different locations - that is we have them physically separated.
Know the difference in voltaic vs electrolytic cells
Know how to draw a picture of both a voltaic and an electrolytic cell if given the shorthand notation.

Electrochemical Cell Shorthand

Given below is the generalized cell notation. Note what is on the left side and what is on the right side.

anode | anodic solution || cathodic solution | cathode

where the "|" are phase changes and "||" is a salt bridge.

What's the sign convention used for the cathode and anode?
What's the significance of the SHE?
What's an inert electrode? Used for what? Pt in the SHE, graphite in a dry cell
Know what standard conditions are and how it is symbolized (same as always).

Calculations

Know how to get:

\({\cal E}^\circ \) for any half-reaction (look it up in a table if it is not given in the question itself).
\({\cal E}^\circ \) for any overall cell reaction. remember that Zumdahl uses cathode potential minus the anode potential which is the same as above. My oxidizing potential is the “minus the anode potential”. Bottom line is that you always leave the reduction reaction as is off the table but you must FLIP the table value for an oxidation and when you FLIP, you must change the sign.
How to convert between \(\Delta G^\circ \), \({\cal E}^\circ \), and \(K\).

\[\Delta G^\circ = -nF{\cal E}^\circ \hskip36pt \Delta G^\circ = -RT\ln K \hskip36pt nF{\cal E}^\circ\ = RT\ln K \]

"n" here is the # of total e⁻ cancelled out in the whole rxn.

Non-Standard Conditions : The Nernst Equation

Know how to use the Nernst Equation to get cell potentials at non standard conditions. Also know how to calculate an unknown concentration from a given non-standard potential. Here are 3 forms of the Nernst Equation:

\(\displaystyle{{\cal E} = {\cal E}^\circ - {RT\over nF} \ln Q \hskip36pt {\cal E} = {\cal E}^\circ - {0.0257\over n} \ln Q \hskip36pt {\cal E} = {\cal E}^\circ - {0.05916\over n} \log Q} \)

Electrolysis

When you pass current through an electrolytic cell you are pushing and pulling electrons. You use your power source to PULL electrons out of the anodic solution via the anode (oxidation) and then PUSH electrons into the cathodic solution via the cathode (reduction). Here, the charge on the electrode is positive (+) for the anode and negative (−) for the cathode. This is the opposite sign convention as for voltaic cells.

You can count moles of electrons via a calculation using current (I), time (t), and the faraday constant (F). Unit analysis will get you there although the following equation works well also.

\[{I\;\;t\over n\;F} = mol {\rm \;of\;reaction}\]

From here you can easily calculate the number of grams of metal that are electroplated during the process. You should also be able to calculate the number of liters of gas formed during the process if there is a gas forming reaction. Remember, you’ll get your moles of gas from the previous equation and then you’ll plug that answer into the Ideal Gas Law for your answer. Know how to solve for any of those variables (I, t, n, and mol) in that equation above. Remember, sometimes we work backwards from the amount to time, current, or even number of electrons.

How to calculate K (or K_sp) for a reaction using electrochemical data. The key is to find the 2 half-reactions that add up to equal the overall K_sp reaction. Remember, all K_sp reactions are just dissociations: MX(s) ⇌ M⁺ + X⁻

Corrosion

What is going on when iron corrodes? What are some ways to prevent or slow down the rate of corrosion? What is a sacrificial electrode? Why doesn't aluminum corrode more than iron - it's has a much bigger oxidizing potential!

Batteries

Know the general concepts of how batteries work. Read book and refer to class notes. What's the difference in a primary, secondary, and fuel cell? Know examples of each type as mentioned in the textbook and from my lecture - names only here. However, you DO need to know the specific reactions that drive a lead storage battery. Know how the reactions work during discharge and during recharge.

What would make one battery capable of delivering more current than another if they have the same potential (voltage)?