WORKSHEET 14 FEEDBACK
Generally good responses but a disappointing amount of errors relating to oxidation #’s, oxidation and reduction.
Q1f) For positive metal ions the oxidation state is always equal to the charge on the ion. Such ions can range from +1 (eg Na+) to +8 (eg Mn8+).
Q3b) For any redox reaction the atoms and the charge must balance
Eg Ag+ (aq) + Pb (s) → Ag (s) + Pb2+ (aq) – the atoms are balanced but the charge is not (+1 in the
reactants and +2 in the products)
reactants and +2 in the products)
If you think about the half equations for this reaction
Ag+ (aq) + e-→ Ag (s)
Pb (s) → Pb2+ (aq) ++ 2e-
It is clear that 2 silver reductions are needed for every lead oxidation, so the balanced equation is
2Ag+ (aq) + Pb (s) → 2Ag (s) + Pb2+ (aq) – now atoms and charge are balanced
Q5) writing ionic equations means that there should be no spectator ions. For 5c) according to metal displacement only the most unreactive metal ions will be reduced and form a solid. All other metal ions will end up in solution. Even still, do not include spectator ions in your equation
Q7) People who drew the grid typically got this correct. If you didn't and you got it wrong – DRAW THE GRID. If you did draw the grid and got it wrong it was typically because you i) didn't complete the grid for all species or ii) forgot what oxidation and reduction mean.
If the oxidation # increases it has been oxidised (ie, it is the reductant).
Eg if lead goes from 2+ to 4+, because 4>2 it has been oxidised.
Eg If chlorine goes from -3 to -1, because -1>-3 it has been oxidised (don’t forget the basics of the number line!)
If the oxidation number decreases (ie reduces) it has been reduced (ie, it is the oxidant).
Eg if manganese goes from +3 to +2, because +2 < +3 it has been reduced.
Eg If sulfur goes from +3 to -2, because -2 < +3 it has been reduced
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SUMMARY FEEDBACK
Summaries were very good, just a few errors kept appearing.
A) Remember that we had to clean the electrodes between each different solution tested & at the end we tested the solution that gave a reading with two Cu electrodes and then two Zn electrodes.
B) You need to put the units for all measurements in the title-row of tables and you need to use units when referring to quantities in your text.
C) This experiment was about the BASIC requirements of a galvanic cell. Thee are only three things needed (four if you count each electrode separately). Even though the cells we analyse in problems have two different electrolytes and a salt bridge that is not really necessary (just easy to mathematically analyse). Also - remember that the electrolyte AND conducting path need to connect both electrodes.
D) Two different metals are needed so they have different reactivity. One tends to be oxidised more (it is a stronger reductant) than the other. This means there will be a potential (voltage) established between them that will allow current to flow when connected in a galvanic circuit.
E) You must mention that the flow of ions in the salt bridge maintains charge neutrality and allows the current to keep flowing (“completes the circuit” is OK but not sufficient).
F) Keep a) and b) simple – one is oxidation, one is reduction… For c) you have to think carefully
This is the galvanic cell we set up in the experiment
i.e. a zinc and copper electrode, separated by a salt bridge with an electrical contact (I’m going to refer to it as the experiment cell).
This is SIMILAR to the galvanic cell below.
Zinc is the anode because it is more reactive, copper the cathode, zinc gets oxidised at the anode (ie Zn ® Zn2+ + 2e-) & copper ions get reduced at the cathode (Cu2+ + 2e- ® Cu)
It is easy to assume that this is exactly what happens in our experiment cell, BUT -
While the Zinc in the experiment cell can be oxidised (so the anode reaction is the same) we have a problem at the cathode. This is because we have no copper ions in our experiment cell. So there must be some other cathode (reduction) reaction occurring.
This is where we have to ask yourself: “What else could be being reduced if not copper ions…?”
We do know that we have water in our salt bridge AND because it is exposed to air there will be a small amount of dissolved oxygen in the salt bridge too.
Hmmm…
Perhaps there is a reduction equation on the reduction potential list that has oxygen and water as the reactants in the reduction equation…
G) The results table needed a simple summary e.g. “The only cell that gave a reading ….”
H) You need to say that the more reactive metal (has a lower reduction potential &) will be oxidised, thus act as the anode which by definition has negative polarity. The less reactive metal will be the site of reduction…. Don’t forget to always apply OIL RIG, RED CAT and ANGRY OX to remember you polarities, sites of oxidation/reduction and which electrode gains/loses electrons.
I) I can no longer remember what the ‘I’ was supposed to indicate…If you figure it out please let me know.
J) Always be specific, don’t just say electrons (or ions) flow between electrodes, say exactly where they come from and go to (just like we never say ‘colour changed’ we say exactly what colour changed to what other colour).
K) The multimeter advantages are very simple. They relate to accuracy, ease of use and reduction in human error – just think to when you used the multimeter - what made it better than trying to use a voltmeter?
L) Don’t forget that you need to include your original hypothesis.
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