Hi Everyone,

Anyone want to try giving these a shot?

The right electrode in a cell is a piece of copper immersed in a 1.00 L solution of 0.0701 F copper sulfate. The left electrode is a galvanized nail (i.e., an iron nail that is completely coated with a layer of zinc), immersed in a solution containing 1.00 L 0.0224 F zinc sulfate.
As the reaction proceeds, the zinc coating of the nail will gradually get thinner and thinner as zinc is consumed in the reaction. Assume that the iron nail (excluding the zinc coating) was 1/8" in diameter, and that the zinc coating for the nail was 1/32" thick. What is the thickness of the zinc coating that remains once the reaction goes to completion? Provide your answer in centimeters to 3 sig. figs. (don't enter units). You will need to know that the nail was immersed in the zinc solution to a depth of 10.5 cm.

I have found the Area of the Nail and the Area of the zinc and subtracted the two to find the area difference. I then tried to figure out which one is limiting. I am stuck in trying to convert from area into concentrations. I realize I need to use the density of zinc to do this, but can someone please help?

Q2: A cell consisting of a Pt indicator electrode and an SCE reference electrode is used to follow the titration of 40.00 mL of a 0.0200 F Br2 (the unknown) with 0.0200 F Fe2+, each in the presence of 1 mol/L HCl (the titrant). Calculate the voltage of the cell following the addition of 60.0 mL of the titrant to the unknown.

I found the moles of iron that I have and tried to write an overall reaction between Fe and Bromine but I am having trouble doing that. Can anyone help walk me through that?

For #1.
I would try this. The problem states, implicitly at least, that copper is the limiting reagent (otherwise, there would be no Zn coating left on the galvanized iron nail at the end of the experiment). Therefore, you will have used 1.00 L x 0.0701 M = 0.0701 mols Cu^+2. Since the reaction is
Zn + Cu^+2 ==> Zn^+2 + Cu, there is a 1:1 ratio, and you must have used up 0.0701 mols Zn that came from the galvanized iron nail. Mols should convert into grams. You will need to calculate the grams Zn at the beginning, subtract the grams Zn used in the reaction to arrive at grams Zn remaining, convert that to thickness at the end knowing the original thickness was 1/32". I think you will need to work in volume in order to obtain final thickness but there may be other ways to get at it. This should get you started.

I'll think about #2 but don't hold your breath. I've done these BUT the procedure I use is different than is being taught today and I don't know if I can convert to the present standards or not. However, the reduction half reactions you need are as follows:
Br2 + 2e ==> 2Br^-
Fe^+3 +e ==> Fe^+2
I presume the ferrous ion is being oxidized to ferric ion and the bromine is being reduced to bromide ion. Note that the iron equation must be multiplied by 2 before adding to the bromine equation to obtain the full reaction. That is, the reaction, as it occurs spontaneouslyhalf as an oxidation and half as a reduction) is
2Fe^+2 + Br2 ==> 2Fe^+3 + 2Br^-

If the equation is all that's holding you up that should take care of it. I hope this helps.

differentriate the reaction between metals and non-metals based on the result of experiment

For question 1, you have correctly determined that copper is the limiting reagent and you have calculated the moles of copper sulfate used. Now, you need to calculate the grams of zinc used by converting the moles of copper sulfate to moles of zinc and then to grams of zinc.

To convert moles of copper sulfate to moles of zinc, you will use the stoichiometry of the reaction. The balanced equation is Zn + Cu^+2 -> Zn^+2 + Cu, which shows that for every 1 mole of copper sulfate, 1 mole of zinc is used. Since you have 0.0701 moles of copper sulfate, you will also have 0.0701 moles of zinc.

Next, you will convert moles of zinc to grams of zinc. To do this, you need to know the molar mass of zinc. The molar mass of zinc is 65.38 g/mol. Multiply the number of moles of zinc by the molar mass to get the grams of zinc used.

Once you have the grams of zinc used, subtract this amount from the initial mass of the zinc coating to find the remaining mass of zinc. Convert this mass to thickness by considering the density of zinc. The density of zinc is 7.14 g/cm^3. Divide the mass of zinc by the volume occupied by the remaining zinc coating to get the thickness in centimeters.

For question 2, you are trying to determine the voltage of the cell following the addition of the titrant to the unknown. To do this, you need to consider the half-reactions and the standard electrode potentials.

The half-reactions involved are:
Br2 + 2e^- -> 2Br^- (reduction at the Pt indicator electrode)
Fe^2+ -> Fe^3+ + e^- (oxidation at the SCE reference electrode)

In order to write the overall reaction, you need to balance the number of electrons transferred. Multiply the iron oxidation half-reaction by 2 and add it to the bromine reduction half-reaction. This gives you the overall reaction:
2Fe^2+ (aq) + Br2 (aq) -> 2Fe^3+ (aq) + 2Br^- (aq)

Now, use the Nernst equation to calculate the cell voltage. The Nernst equation is:
Ecell = E°cell - (RT/nF)ln(Q)

Where Ecell is the cell voltage, E°cell is the standard cell potential, R is the gas constant, T is the temperature in Kelvin, n is the number of electrons transferred in the balanced reaction, F is Faraday's constant, and Q is the reaction quotient calculated using the concentrations of the species involved.

Given the concentrations of the species involved and the standard electrode potentials (which you can look up in a table), you can calculate the cell voltage following the addition of the titrant.