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A Cu-Ni galvanic cell

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Q1)  In an experiment, if you are using  highly toxic chemical (eg lead a NEUROTOXIN) and a relatively harmless chemical and you have to choose which chemical to write about, make sure you focus the risk assessment on the dangerous chemical. For three marks you should consider three different 'things' that were done (NB gloves, glases, coat all come under the same categoty of 'protection equipment'). HINT You must mention storage (for proper disposal) all heavy metal solutions as part of your safety (i.e. do not tip down the sink). Note specifically why lead is a concern.

 

Q2ab)  Explain means give a reason – so in the Chemistry context, that means give a reason based on some Chemistry concept. E.g. it is not good enough to say ‘The cell needed a salt bridge’. You need to say ‘The cell needed a salt bridge because...’ where you go on to give a reason based on chemistry. Similarly, the electrodes needed to be separated so the reduction and oxidation reactions are separated (ie don't react directly) so a potential difference (voltage) between the electrodes is establihed and current has to flow between them (through the circuit - doing useful work).   If you are going to say ‘the salt bridge completes the circuit’ make sure you follow up with ‘by maintaining charge neutrality in the half cells’.

 

Q2c)      KIS – Keep it simple . Answer the question and don’t over explain. Some of the explanations we use in class are gross-oversimplifications of very complex processes. I use these so that you get a better understanding of the basic process (so you will more easily comprehend & remember the theory) – but they aren’t so good in written reports. Eg for Q3c – just say ‘Anions migrate to the …. through the …. while cations….

 

Q2d)      While it is true that the greater difference in standard reduction potential, the greater the voltage (and you should say this) the Q refers you to the activity series so you should also relate the voltage to the relative reactivity of electrodes. Basically, the bigger difference in reactivity, the greater the difference in reduction potential and thus the greater the cell voltage.

 

Q3)     Standard conditions are 25^oC and 1mol/L solutions (and gases at 100kPa). If these conditions are not met the E^o values for the cells will be different to the data sheet and thus the cell potential will be different to that predicted by the sheet (and possibly anode and cathode will be switched) as different cells are effected differently by changing conditions will be different . Corrosion or damage to the electrode surface can also change the transport of ions to the surface and effect E^o values. Impure samples (PURE samples of metals are very hard to obtain) also lead to a different reduction potential for the electrodes.

 

Q3b)  Beware of the Cu⁺ vs Cu²⁺ problem identified below.

 

Q4)     Extracting data from graphs and tables is not just vital to extract data from standard potential tables to assess the validity of our experiment, it is also vital as that is how most numerical data is reported and scientists need to use this data to make calculations, judgements, conclusions and generalisations (i.e. about tends and patterns). What can YOU learn from the Standard reduction Potential table?

 

Q7)       Your cell diagram needs to have ALL the features and calculations specified in the question

Worksheet 16 and Compulsory Activities

But before I look at the most troublesome questions here is my Rant of the Week

1)      Electrode potentials MUST include the unit VOLTS. Without Volts (or at least a capital V) the answer is meaningless. Don’t let all that hard work calculating the answer be worth nought because you forgot to add a V at the end of your answer.

2)      Cu²⁺ and Cu⁺ are NOT the same thing. Use the correct cell potential!

Cu²⁺ + 2e^- ® Cu _(s)  E^o = 0.34V

Cu⁺ + e^- ® Cu _(s)  E^o = 0.52V

If you are ever doing a copper cell then it is mostly likely to be a 2+ solution and the reduction potential is 0.34V. Learn to love that number and learn to be wary of 0.52V.

3)      We learn these things for a reason: OIL RIG, RED CAT, ANGRY OX, ABC’s. At the very least check what you have written against these rules. e.g. you have written a half equation where electrons are being lost then (OIL RIG) that means it is oxidation and (ANGRY OX) that means it is at the anode which is –ve.

4)      USE THE REDUCTION POTENTIAL SHEET!. E.g., If you are told that there is a chlorine half cell then look at the data sheet to, not just figure out ABC’s but to also to figure out what reactions are possible. All of these examples below are wrong for many reasons, but that didn’t seem to stop anyone – you should be able to explain what is wrong with each one of these:

Cl₂ ® Cl^- + e^-                       FAIL

½Cl₂ ® Cl₂ + e^-                   WT FAIL

Cl^- + e^- ® Cl₂                       SO FAIL

Cl^- + e^- ® Cl⁺                       SO VERY FAIL

Cl⁺ + e^- ® ½ Cl₂                  FAIL

½ Cl₂ + e^- ® Cl⁺                  EPIC FAIL

½ Cl₂ ® Cl⁺ + e^-                  FAIL

Q2) Just do it – people must have thought this was part of Q3 as many skipped it (and it was easy)

Q4) When converting from half equations to a full cell equation you must make the electrons balance before you add the half equations together. But when converting a full equation into half equations you should simplify the half cell equations at the end.

Q5,7, 8) Main issue is reading the question and doing all the things the question asked.

Q6) See Rant point #4 above. Also in shorthand notation it should go: ion|related gas,Pt

E.g.  ||Cl^- _(aq)|Cl_2(g),Pt

Q9) All pretty good except the overall cell equations. Just like any galvanic cell overall equation you: a) make the # of electrons the same in both half equations (i.e. for the dry cell you need to double that whole monkey-nuts equation.  Also - simplify at the end (i.e. for the silver cell, the fact that electrolyte concentration stays the same means you should be able to cancel out the OH and H₂O on both sides)

Q10) Just be careful - ‘ decreasing tendency’ means gong from the most likely to least likely

Q11) Similar to rant point 2 above. You need to be aware that there is a Fe²⁺/Fe half equation and a Fe²⁺/Fe³⁺ equations. Read the question carefully to know which one you need to use

Q13) To prove a reaction occurs ’as written’ means you need to prove the reaction is spontaneous. Spontaneous redox reactions give a positive E^o value. If you look at the equation provided, one reaction taking place is the oxidation of Iron(II) ions:  Fe²⁺ ® Fe³⁺ + e^-. Take out the Fe²⁺ and Fe³⁺ terms from the equation and you are left with the other half equation. That is on the data sheet too. Just add the Fe²⁺ oxidation E^o to the E^o of the reduction process and see if it is spontaneous.

Q14) First - figure out which is oxidised and which is reduced from the equation, then just add the two potentials together (remembering to reverse the sign of the oxidation E^o)