C.T PILLAY
21910221
GROUP 3A
M SABELA
5TH SEPTEMBER 2019
Contents
Introduction ............................................................................................................................................ 2
Methodology........................................................................................................................................... 3
Apparatus............................................................................................................................................ 3
Reagents ............................................................................................................................................. 3
Procedure ............................................................................................................................................ 3
Instrument parameters ....................................................................................................................... 4
Calculations ............................................................................................................................................. 4
Table 1: Calculation of masses to be used .......................................................................................... 4
Table 2: Calculation for equal quantities of the copper solution to be used ...................................... 4
Table 3: Separate calculations ............................................................................................................ 5
Results ..................................................................................................................................................... 6
Table 4: Measured results ................................................................................................................... 6
Table 5: Calculated results .................................................................................................................. 6
Table 6: Required information for the graph below ........................................................................... 7
Graph 1: A graph comparing Ecell and log[Cu2+] for cells 7,8,9 and 10. ............................................... 7
Discussion................................................................................................................................................ 8
Conclusion ............................................................................................................................................... 8
Safety/Precautions .................................................................................................................................. 9
References .............................................................................................................................................. 9
1
Introduction
The aim of this experiment is to determine the electrode potentials of some metal electrodes.
Electrochemistry is a branch in chemistry that deals with the interconversion of electrical energy and
chemical energy. There are two types of cells. A galvanic or voltaic cell, which converts chemical
energy to electrical energy and an electrolytic cell, which converts electrical energy to chemical
energy.
Electrochemical cells contain two electrodes, namely the anode (where oxidation takes place) and
the cathode (where reduction takes place), these are the basis for this process. The metal electrodes
must be placed in an electrolyte containing the same type of metal that makes up the electrodes.
When the metal reacts, they give away electrons and form positive ions. The electrons are
transferred through an external circuit, by an electromotive force, while ions are transferred through
a salt bridge. The salt bridge maintains the overall charge balance for the two compartments.
The potential that develops in a cell is the measure of the tendency for a reaction to proceed toward
equilibrium. E0 values give you a way of comparing the positions at equilibrium, as well as predicting
if it is spontaneous or not.
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Methodology
Apparatus
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A Balance
50 ml volumetric flasks
100 ml beakers
Funnels
Plastic droppers
Pipettes
Pipette pump
Small plastic dish
Copper, Zinc, Iron and Lead electrodes
Digital Multimeter
Salt bridge
Reagents
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Copper(II) nitrate
Zinc nitrate
Lead nitrate
Ammonium iron(II) sulphate
Procedure
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The mass that was needed to prepare a 1 molar concentration of each reagent was
calculated.
The empty plastic dish was weighed and recorded. The required mass of each reagent was
then weighed.
Each reagent was transferred into separate 50 ml volumetric flasks using a funnel, ensuring
the stem does not get clogged. The funnel was removed, and each sample was diluted to the
calibration mark with deionised water. A stopper was added to the flasks and then swirled
till the reagent was dissolved. Each flask was labelled.
By serial dilutions, four copper(II) solutions of concentrations 0.1M, 0.01M, 0.001M and
0.05M were prepared in 50 ml volumetric flasks, using the 1M copper solution that was
previously prepared. The flasks were then filled with deionised water to the calibration
mark, and then swirled, to evenly dilute it. The flasks were labelled accordingly.
Equal quantities of each of the metal ion solutions were calculated and then poured into
four different 100 ml labelled beakers.
The potentials were obtained, using a Digital Multimeter. The zero on the Multimeter was
checked by short-circuiting across the positive and negative leads. The leads were then
connected to the metal electrodes, so that a positive reading was displayed. A salt bridge
taken from a solution was also connected to each half cell. The highest potential shown in
the first 30 seconds was recorded, as well as the electrode which produces a positive
reading. The procedure was repeated for each cell. It was made sure that the salt bridge was
rinsed with deionised water before placing in a different solution.
3
Instrument parameters
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Make sure equipment is clean, if not, rinse with tap water then deionised water.
Clean the electrodes by dipping them in 6M nitric acid for a few seconds, then rinse with
deionised water.
When using the Digital Multimeter, make sure a positive sign is shown on screen after the
electrodes are connected, if not, switch the leads.
Calculations
Table 1: Calculation of masses to be used
Mass
calculated(g)
=Molarity x
volume x
RAM
Mass
weighed(g)
Cu(NO3)2 .
3H2O
Zn(NO3)2 .
6H2O
Pb(NO3)2
FeSO4 .
(NH4)2SO4.
6H2O
= 1 x 0.05 x
241.60
= 12.0800
= 1 x 0.05 x
297.46
= 14.8730
= 1 x 0.05 x
331.21
= 16.5605
= 0.5 x 0.05 x
392.14
= 9.8035
= 12.0927
= 14.8679
= 16.5613
= 9.8068
Table 2: Calculation for equal quantities of the copper solution to be used
Cu2+(0.1M)
Cu2+(0.01M)
Cu2+(0.001M)
Cu2+(0.05)
C1V1=C2V2
0.1 x 0.05
V1=
1
V1= 5 x 10-3 L
= 5 ml
C1V1=C2V2
0.01 x 0.05
V1=
1
V1= 5 x 10-4 L
=0.5 ml
C1V1=C2V2
0.001 x 0.05
V1=
1
V1= 5 x 10-5 L
=0.05 ml
C1V1=C2V2
0.05 x 0.05
V1=
1
V1= 2.5 x 10-3 L
=2.5 ml
4
Table 3: Separate calculations
Cell
No.
E0cell = Ecathode - Eanode
(Volts)
Ecell = E0cell (Volts)
𝟎.𝟎𝟓𝟗𝟐
𝒏
log[Q]
emf = |Ecell|
0.0592
1.0
log [1.0]
2
= |0|
=0
0.0592
1.0
log [ ]
2
1.0
=|0|
=0
0.0592
1.0
log [0.5]
2
=|0.226|
=0.226
0.0592
1.0
log [1.0]
2
=|0|
=0
0.0592
0.5
log [1.0]
2
=|-0.093|
=0.093
0.0592
2
=|0.196|
=0.196
1
= 0.34-(-0.76)
=1.10
= 1.10 = 0.00
2
= 0.34-(-0.13)
=0.47
= 0.47 =0.00
3
= 0.34-(-0.44)
=0.78
= 0.78 = 0.226
4
= -0.13-(-0.76)
=0.63
= 0.63 = 0.00
5
= -0.44-(-0.76)
=0.34
= 0.34 = -0.093
6
= -0.13-(-0.44)
=0.31
= 0.31= 0.196
7
= 0.34-(-0.76)
=1.10
= 1.10 = -1.070
8
= 0.34-(-0.76)
=1.10
= 1.10 = -1.393
9
= 0.34-(-0.76)
=1.10
= 1.10 = -2.141
10
= 0.34-(-0.76)
=1.10
= 1.10 = -3.211
11
= 0.34-0.34
=0
= 0.78 = -0.75
0.5
log [0.1]
0.0592
0.1
log [1.0]
2
=|-1.070|
=1.070
0.0592
0.05
log [ 1.0 ]
2
=-1.393|
=1.393
0.0592
0.01
log [ 1.0 ]
2
=|-2.141|
=2.141
0.0592
0.001
log [
]
2
1.0
=|-3.211|
=3.211
0.0592
0.1
log [1.0]
2
=|-0.75|
=0.75
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Results
Table 4: Measured results
Cell
No.
Anode
+
Or
Or
Cathode −
Anode
+ Emf/V
Or
Or
Cathode -
Cell Notation
1
Anode
−
Zn | Zn2+ (1.0M) || Cu2+ (1.0M) | Cu
Cathode
+
0.876
2
Anode
−
Pb | Pb2+ (1.0M) || Cu2+ (1.0M) | Cu
Cathode
+
0.524
3
Anode
−
Fe | Fe2+ (0.5M) || Cu2+ (1.0M) | Cu
Cathode
+
0.753
4
Anode
−
Zn | Zn2+ (1.0M) || Pb2+ (1.0M) | Pb
Cathode
+
0.4136
5
Anode
−
Zn | Zn2+ (1.0M) || Fe2+ (0.5M) | Fe
Cathode
+
0.1411
6
Cathode
+
Pb | Pb2+ (1.0M) || Fe2+ (0.5M) | Fe
Anode
−
0.2493
7
Anode
−
Zn | Zn2+ (1.0M) || Cu2+ (0.1M) | Cu
Cathode
+
0.4591
8
Anode
−
Zn | Zn2+ (1.0M) || Cu2+ (0.05M) | Cu
Cathode
+
0.771
9
Anode
−
Zn | Zn2+ (1.0M) || Cu2+ (0.01M) | Cu
Cathode
+
0.7627
10
Anode
−
Zn | Zn2+ (1.0M) || Cu2+ (0.001M) | Cu
Cathode
+
0.6154
11
Cathode
+
Cu | Cu2+ (1.0M) || Cu2+ (0.1M) | Cu
Anode
−
0.821
Table 5: Calculated results
Cell
No.
Spontaneous Cell Reaction
Calculated (Theoretical)
Cell
Potential/V
1
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
2
Pb(s) + Cu2+(aq) → Pb2+(aq) + Cu(s)
3
Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)
4
Zn(s) + Pb2+(aq) → Zn2+(aq) + Pb(s)
5
Zn(s) + Fe2+(aq) → Zn2+(aq) + Fe(s)
6
Pb(s) + Fe2+(aq) → Pb2+(aq) + Fe(s)
= -0.13-(-0.44)
=0.31
7
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
8
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
9
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
= 0.34-(-0.76)
=1.10
= 0.34-(-0.76)
=1.10
= 0.34-(-0.76)
=1.10
= 0.34-(-0.76)
=1.10
= 0.34-(-0.13)
=0.47
= 0.34-(-0.44)
=0.78
= -0.13-(-0.76)
=0.63
= -0.44-(-0.76)
=0.34
6
10
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
11
Cu(s) + Cu2+(aq) → Cu2+(aq) + Cu(s)
= 0.34-(-0.76)
=1.10
= 0.34-0.34
=0
Table 6: Required information for the graph below
Cell
Ecell (measured)/V
[Cu2+]
No.
0.1
7 = -1.070
0.05
8 = -1.393
=
-2.141
0.01
9
0.001
10 = -3.211
log [Cu2+]
=log[0.1] = -1
=log[0.05] = -1.301
=log[0.01] = -2
=log[0.001] = -3
Graph 1: A graph comparing Ecell and log[Cu2+] for cells 7,8,9 and 10.
A graph comparing Ecell and log[Cu2+]
0
-0,5
-1
-1,5
-1
-1,07
-1,301
-1,393
-2
-2
-2,141
-2,5
-3
-3
-3,211
-3,5
Ecell
log[Cu2+]
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Discussion
In this experiment a Digital Multimeter was used to determine the electrode potentials for four
metal electrodes (Cu, Zn, Pb and Fe). The electrodes were placed in aqueous solutions, with a salt
bridge connected in between. The results obtained were then used as a comparison to the
calculated potentials.
It is known that if the concentrations of the solutions increase, the cell reaction becomes more
spontaneous and the emf increases (Anon. 2006). A positive voltage means the reaction is
spontaneous. The concentration trend is seen in the results, for cells 7, 8, 9 and 10. The Nernst
equation was used to calculate the cell emf at nonstandard conditions. It helps relate the cell emf
(Ecell)to its standard emf(E0cell).
In table 3, it shows that the E0 value for copper is 0. This means that the products being made, and
reactants being dissociated are happening at the same rate, therefore there is no net potential since
the system is already in equilibrium (Tiwari. 2015).
Graph 1 was plotted according to the calculations in table 6. It shows that there is a relationship
between Ecell and log[Cu2+]. Both line graphs decrease proportionally, in relation to the decrease in
concentration. A higher concentration makes the electrode potential more positive and gives it a
greater ability to move to equilibrium (Charco. 2007). This is seen at the beginning of the graph,
when it starts to curve or level out in a horizontal line. This could have been shown more clearly if
more values were taken into consideration.
The discrepancies in the results may have been caused by the difference in the calculated mass and
the actual mass weighed in table 1, the inaccurate measuring of the solutions using the pipette and
when transferring the solutions into different flasks. The difference in the reading time on the Digital
Multimeter could have been greater than 30 seconds, and this will cause a great difference since the
values on display change constantly.
While doing the experiment, it was made sure that all reagents were handled with care and away
from our eyes and skin, the salt bridge was washed thoroughly, and all apparatus was clean.
Conclusion
In conclusion, the experiment was successful and quite accurate because the results measured, and
the calculated results are similar. It was also proven that there is a relationship between the Ecell
values and concentration. The results have also shown that the concentration effects the emf of a
cell and what the emf of a cell is, when it has reached equilibrium.
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Safety/Precautions

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Make sure all beakers and flasks are correctly labelled.
Dispose all used solutions in an approved manner, as instructed.
Make sure the apparatus is cleaned properly after use, to avoid contamination for the next
users.
Make sure you are dressed appropriately, wearing safety goggles and closed shoes, to avoid
chemicals harming your skin.
Do not taste of sniff chemicals.
Always read the labels on containers, to make sure you are working with the correct
reagent.
Handle glassware with care.
Always wash your hands before leaving the lab.
References

[1] Anon. 2006. Effect of concentration on cell emf. Available:
http://wps.prenhall.com/wps/media/objects/3313/3392587/blb2006.html (Accessed 18
September 2019)

[2] Tiwari, C. 2015. What is the emf of the cell when the cell reaction attains equilibrium.
Available: https//www.quora.com/what-is-the-EMF-of-the-cell-when-the-cell-reactionattains-equilibrium (Accessed 18 September 2019)

[3] Charco. 2007. Effect of concentration on electrode potential. Available:
https://www.thestudentroom.co.uk/showthread.php?t=360208 (Accessed 18 September
2019)
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