1 - Galvanic Cell.docx

INTRODUCTION Galvanic Cells allow us to control the electron flow in a redox reaction to perform useful work. Such cells find common use as batteries, pH meters, and as fuel cells. The setup of the cell requires that the oxidation and reduction half-reactions are connected by a wire and by a salt bridge. Electrons will flow through that wire creating an electrical current. The salt bridge or porous disk allows the passage of ions in solution to maintain charge neutrality in each halfcell. The direction of the current in a cell is determined by the standard reduction potential of each half-cell. For a reaction to be spontaneous, the overall cell potential must be positive. Therefore, the half-reaction with the greater reduction potential will be a reduction and the other half-reaction will be an oxidation. The electrode in the oxidation half-reaction is called the anode. The electrode in the reduction half-reaction is called the cathode The objective of this experiment is to measure the relative reduction potentials for several half cell (redox) couples in galvanic cells, to develop an understanding of the movement of electrons, anions and cations, to study factors affecting cell potentials and to estimate the concentration of ions in solution using Nernst Equation.

MATERIALS AND METHOD Copper, zinc and iron metal strip, 0.1M Cu(NO3)2 solution, 0.1M Fe(NO3)2, 0.1M FeSO4, 1M CuSO4, 0.001M CuSO4, ammonia solution, voltmeter, wire, filter paper and beakers.

PROCEDURE Refer to jotter attached.

RESULTS A. Galvanic Cell – Reduction Potentials of Several Half Cell Couples Galvanic Measured Anode Equation for anode Cathode Equation for cathode Cell Ecell reaction reaction 2+ Cu – Zn 0.86 V Zn Zn Zn + 2e Cu Cu2+ + 2eCu 2+ 2+ Cu – Fe 0.45 V Fe Fe Fe + 2e Cu Cu + 2e Cu Zn – Fe 0.18 V Zn Zn Zn2+ + 2eFe Fe2+ + 2eFe Overall equation of each reactions: Cu – Zn Zn + Cu2+ Zn2+ + Cu 2+ Cu – Fe Fe + Cu Fe2+ + Cu Zn – Fe Zn + Fe2+ Zn2+ + Fe Redox couple Cu – Zn Cu – Fe Zn – Fe

Reduction potential (measured) 0.86 V 0.45 V 0.18 V

Reduction potential (calculated) 1.13 V 0.78 V 0.35 V

Reduction potential (calculated)

Percentage error

E°cell = Ecathode – Eanode

% error =

Ecell = E°cell -

0.0592 [𝑝𝑟𝑜𝑑𝑢𝑐𝑡] log [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡] 𝑛

% error 23.89 42.31 48.57

𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑−𝐸𝑐𝑒𝑙𝑙 𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑

Example calculation: 1)

Cu – Zn : E°cell = ECu – EZn = 0.34- (-0.79) =1.13 V Ecell = 1.13 -

0.0592 0.1𝑀 log 0.1𝑀 2

= 1.13 V 2)

Cu – Fe : E°cell = ECu - EFe = 0.34- (-0.44) =0.78 V Ecell = 0.78 -

0.0592 0.1𝑀 log 0.1𝑀 2

= 0.78 V 3)

Zn – Fe : E°cell = ECu - EFe = -0.44- (-0.79) =0.35 V Ecell = 0.35 = 0.35 V

0.0592 0.1𝑀 log 0.1𝑀 2

1) Cu-Zn =

1.13−0.86 1.13

= 23.89%

x 100%

𝑥 100%

B. Effect of concentration changes on cell potential 1. Cell potential of “concentration cell” : 0.0083V Anode reaction : Cu Cu2+ + 2eCathode reaction : Cu2+ + 2eCu A potential is recorded because the solution used has difference concentration. 2. Cell potential from complex formation : 0.016V Observation of solution in half-cell : the colour of solution 0.001M CuSO4 change from colourless to slightly light blue colour The potential changes with the addition of NH3 because the number of Cu2+ ions flow form cathode and the increase of the Cu2+ concentration. 3. The cell potential decreases when NH3 added to 1M CuSO4 instead of 0.001M CuSO4 because the mole concentration of 1M CuSO4 are higher than 0.001M CuSO4. C. The Nernst equation and an unknown concentration 1. Solution number 1 2 3 4

Concentration of Cu(NO3)2 0.1 mol/L 0.001 mol/L 0.00001 mol/L 0.0000001 mol/L

log [Cu2+]

Ecell (measured) 0.945 V 0.906 V 0.780 V 0.753 V

-1 -3 -5 -7

Ecell (calculated) 1.100 V 1.041 V 0.982 V 0.922 V

Dilution of Cu(NO3)2 : M1V1 = M2V2 0.1(1) = M2(100) M2 = 0.001M

M1V1 = M2V2 0.001(1) = M2(100) M2 = 0.00001M

Solution 1 E°cell = ECu – EZn = 0.34- (-0.76) = 1.10 V Ecell = 1.10 -

0.0592 2

log

Solution 2 E°cell = ECu – EZn = 0.34- (-0.76) = 1.10 V 0.1𝑀 0.1𝑀

= 1.10 V

Ecell = 1.10 -

0.0592 2

log

0.1𝑀 0.001𝑀

= 1.041 V

Solution 3 E°cell = ECu – EZn = 0.34- (-0.76) = 1.10 V Ecell = 1.10 -

M1V1 = M2V2 0.00001(1) = M2(100) M2 = 0.000001M

0.0592

= 0.982 V

2

log

Solution 4 E°cell = ECu – EZn = 0.34- (-0.76) = 1.10 V 0.1𝑀 0.00001𝑀

Ecell = 1.10 -

0.0592

= 0.922 V

2

log

0.1𝑀 0.0000001𝑀

2. E cell versus log [Cu2+]

ECELL VERSUS LOG [CU2+] Ecell (measured)

Ecell (calculated)

Linear (Ecell (measured))

Linear (Ecell (calculated)) 1.5

1

ECELL (V)

0.5

-60

0 -50

-40

-30

-20

-10

0 -0.5

-1

LOG [CU2+]

-1.5

3. E cell for the solution of unknown concentration : 0.018V Molar concentration of Cu2+ in the unknown : 4. The concentrations of Cu2+ and Zn2+ for the Cu-Zn cell to maximize the cell potential can be adjusted by increase the concentration one of the solution.

DISCUSSIONS In electrochemistry, redox reactions are a common term that used for oxidation and reduction process that happen in a same time. Oxidation is when electron is lost and take places in anode electron while reduction is when electron is gain at cathode electrode. In this experiment, the interest of redox reaction was carried out for galvanic cell which consist of voltmeter to measure it reduction potential with the presence of salt bridge as medium to transfer electron from one solution to another solution. From part A of this experiment, three metal strips were used as electrode with three different solution with 0.1M respectively. The Ecell that measured by the voltmeter were compared with the calculated Ecell by using formula, Ecell = Ecathode – Eanode The arrangement of the redox couples in decreasing order based on the results are: Cu2+> Fe2+ > Zn2+ For part B, the concentration used for the galvanic cell are different where CuSO4 solution with 0.1M and 0.001M are used as electrolyte. The electron will move from anode to cathode as solution containing 0.001M Cu2+ will transfer electron to the one containing 0.1M Cu2+. Thus, the electron transfer will produce more Cu2+ in the less concentrated solution and consume Cu2+ in the more concentrated solution. Electron will continue to transfer until it reaches the equilibrium concentration. On the other hand, the addition of NH3 to the solution makes Cu2+ ions flow from cathode to anode and the concentration of Cu2+ ions tend to decrease. Nernst equation can be used to find the cell potential at any moment during a reaction or at condition other than standard state. Ecell = E°cell -

0.0592 [𝑝𝑟𝑜𝑑𝑢𝑐𝑡] log 𝑛 [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡]

In part A, this equation was applied, however the differences can be seen clearly in part C since the molar concentration of electrolyte used were different. From Nernst equation, the unknown concentration can be determined by plotting graph of Ecell versus log [Cu2+]. However, the expected Ecell from the experiment cannot be seen from the graph because the intercept between line of Ecell measured and Ecell calculated is larger than expected (0.018V). This may due to some errors during the experiment. First, the solution used may be contaminated with other solutions which affect its molarity. This can be overcome by use a clear and dry beaker to transfer the solution from sample bottle. Next, the crocodile clippers use may not be stable as it loses its sensitivity to the electric voltage and can be avoid by use the clippers that in a good condition. Last but not least, the dilution of the electrolyte could wrongly dilute from standard solution.

CONCLUSIONS As conclusion, galvanic cell is one of the examples of cell in electrochemistry. The energy conversion involving from chemical energy to electrical energy that happen spontaneously as electron move from higher potential energy level to low potential energy level. Molar concentration of solution is a main factor that affecting cell potential which is inversely proportional with Ecell.

REFERENCES Galvanic Cells, Rice University, retrieved from https://opentextbc.ca/chemistry/chapter/17-2galvanic-cells/ Acap, (2016), Galvanic Cell retrieved from https://www.pdfcoke.com/doc/314274311/GalvanicCell-docx Dependence of Cell Potential on Concentration retrieved from https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_(Zumda hl_and_Decoste)/11%3A_Electrochemistry/11.4%3A_Dependence_of_Cell_Potential_on _Concentration