Effect of Reactant Concentration on the Rate of a Reaction

✅ Paper Type: Free Essay	✅ Subject: Chemistry
✅ Wordcount: 5191 words	✅ Published: 18th May 2020

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Open Ended Investigation – Galvanic Cells

Plan, perform and report on an investigation to analyse how the rate of a reaction of a galvanic cell can be affected by concentration of reactants.

Title: The effect of reactant concentration on the rate of a reaction

Aim: To investigate how the rate of a reaction of a galvanic cell can be affected by concentration of one of the reactants (copper sulfate).

Research/theory:

In spontaneous redox reactions, electrons are transferred from one object to another, releasing energy. The reaction must be split into oxidation and reduction half-reactions to measure this. If the reactions are separated a wire can contain the flow of electrons, which can be recorded using any apparatus which can measure electrical current i.e. voltmeter or multimeter.

The half-cells are constructed by placing a metal in a solution containing its ions. For the first half cell, place the copper solid in a solution ofCuSO₄ (source of Cu²⁺ ions) and in the other, place the magnesium ribbon in a solution of MgSO₄ (source of Mg²⁺ ions). The anodic side is where oxidation occurs and the cathode side is where reduction occurs. The more active metal, in this case magnesium will therefore be oxidised. A salt bridge allows for the flow of ions to neutralize the charge build-up in solution, since a charged solution will not give a reading on the multimeter.

The potential difference between the two electrodes is called the cell potential (E_cell). It depends on concentrations and temperature. If the concentrations are all 1M and the temperature is 25°C , the cell potential is referred to as the standard cell potential, E°_cell. Calculated by:

These are the standard potentials in the half-reactions:

When the concentration of a reactant in any chemical reaction increases, the rate of reaction increases, since there are more particles present for collision. The same applies in a galvanic cell – when the concentration of a solution increases there are more anions and cations present, hence increasing the rate of reaction. The standard potentials are thereby altered by concentration, so the Nernst equation can be used to predict the new cell potential 

In the following experiment, the equation would look like this:

Hypothesis:

If the concentration of one of the reactants, copper sulfate is diluted/decreased, the rate of reaction (V) of the galvanic cell will decrease, because there will be less available cations and anions in the solution present for a chemical reaction, and this trend was also predicted using the Nernst equation.

Materials:

Material	Size	Quantity	Disposal
Safety apron	Standard	1	Return to hook
Safety goggles	Standard	1	Return to shelf
Blu Tak	Small piece	2	Return to packaging
Sand paper	App. 2x5cm	1	Return to box
Magnesium electrode	0.2cm x 7cm	1	Wash and return to box
Magnesium solution (1.0 mol.L^-1)	50mL	1	Pour into container of appropriate molarity concentration
Copper electrode	2cm x 6cm	1	Wash and return to box
Copper solution (1.0 mol.L^-1)	150mL	1	Pour into container of appropriate molarity concentration in the equipment trolley
Water	100mL	–	Pour down sink
Beaker	50-100mL	6	Wash and return to cupboard
Measuring cylinder	50mL	1	Wash and return to cupboard
Digital multimeter	Four significant figures	1	Return to plastic bag and place in equipment trolley
Alligator clips	Standard – as long as they fit in multimeter	2	Return to plastic bag and place in equipment trolley
Labels	Small	2	Throw away
Funnel	Small	1	Wash and return to draw
Volumetric pipette	25.00mL	1	Wash and return to equipment trolley
Graduated pipette	10mL	1	Wash and return to equipment trolley
Plastic pipette	2mL	1	Wash and return to draw
Salt bridge	App. 15cm	20	Throw away
Potassium nitrate	Enough to dampen salt bridge	1	Place jar back in equipment trolley

Risk assessment:

Risk	Precaution	Management
Acid can spill into eyes and cause burns, or accidentally be consumed.	Wear safety goggles, and handle with care	Rinse eyes in the eye wash station and notify the supervising teacher immediately. They will contact the Poisons Information Centre if chemical are consumed
Cuts from broken glass equipment	Do not use glass equipment which is wet or if your hands are wet to avoid slippery surfaces	Notify teacher of broken equipment immediately to have it swept away into a disposal bucket Rinse the wound, and with teacher aid, cover cut with a bandage or tape depending on the severity
Touching one lead of multimeter, while the other is in an electrical outlet (if the multimeter is accidentally set to read current). If multimeter is correctly set to read voltage it will have a high resistance and not shock you	Conduct measurements involving the multimeter on a dry surface – keep away from water at all times, especially while connected to electrodes. Mainly, try make sure that it is always set to read voltage, and do not play with the dial.	Notify supervising staff member. They will turn off the multimeter and respond with medical equipment if required, based on amplitude of current.

Method:

Before beginning the experiment, it is important to control and/or measure certain variables

– Using the IPhone ‘Sensors’ app, record the atmospheric pressure

– Use a digital thermometer to measure room temperature

Put on safety apron and goggles
Gather all equipment to workbench
Label beakers accordingly;
1. Magnesium (1.0mol.L^-1)
2. Copper (1.0mol.L^-1)
3. Copper (0.8mol.L^-1)
4. Copper (0.6mol.L^-1)
5. Copper (0.4mol.L^-1)
6. Copper (0.2mol.L^-1)
Fill a volumetric pipette with magnesium solution to the red line until the bottom of the concave meniscusof the water has completely passed the line, and transfer into beaker (a). Repeat again so that the beaker has 50mL of magnesium solution. Do the exact same thing so that beaker (b) is filled with 50mL of the copper solution, BUT remember to use a plastic pipette with distilled water to rinse out all of the magnesium before using copper solution to avoid contamination
Use the graduated pipette to extract copper solution and transfer into appropriate beaker, creating diluted concentrations of copper sulfate (independent variable)
1. Do this 4 times for beaker (c)
2. 3 times for beaker (d)
3. 2 times for beaker (e)
4. Once for beaker (f).
Rinse pipette using distilled water and a plastic pipette, then use the graduated pipette to extract water and transfer into appropriate beaker
1. Once for beaker (c)
2. Twice for beaker (d)
3. Three times for beaker (e)
4. Four times for beaker (f)

Before continuing, double check the following. Layout beakers from left to right in order (a) – (f)

a) 50mL of magnesium sulfate only (1.0mol.L^-1)

b) 50mL of copper sulfate only (1.0mol.L^-1)

c) 40mL of copper sulfate + 10mL of water (0.8mol.L^-1)

d) 30mL of copper sulfate + 20mL of water (0.6mol.L^-1)

e) 20mL of copper sulfate + 30mL of water (0.4mol.L^-1)

f) 10mL of copper sulfate + 40mL of water (0.2mol.L^-1)

Sand both sides of the metal electrodes with sand paper. Place the magnesium electrode in the magnesium solution [beaker (a)], and the copper electrode in beaker (b) – hold electrodes to the side of the beaker with blu-tak
Dunk a salt bridge in the jar of potassium nitrate and place one end in each beaker (not touching the glass walls or the electrode)
Set up the multimeter by connect alligator clips to the sanded electrodes, and connecting the red wire to the copper, and the black wire to the magnesium.
Turn the multimeter on and set to 2000mV setting – divide the 4 digit reading by 1000, and record results in VOLTS. Do this 4 times for each different concentration of copper solution, replacing the salt bridge each time. Make sure to rinse the copper electrode between each change of concentration, and join it to the beaker using same method as described in step 7

Diagram of method (galvanic cell)

Results table + graph:

Table comparing the concentration of copper sulfate to the rate of reaction in a galvanic cell made of copper and magnesium half-cells:

	Rate of reaction (Volts)
Concentration of copper (mol.L^-1)	Trial 1	Trial 2	Trial 3	Trial 4	Average
1.0	1.644	1.641	1.641	1.640	1.642
0.80	1.639	1.641	1.640	1.640	1.640
0.60	1.634	1.640	1.638	1.554	1.637
0.40	1.631	1.637	1.635	1.634	1.634
0.20	1.631	1.627	1.611	1.629	1.629

Discussion of method:

ACCURACY:

Measurements using pipettes and propagating errors:

Step 4 of the method described the process of transferring solutions from a jar to its respective beaker. The 25mL volumetric pipettes used are very accurate, however 50mL of solution needed to be transferred, so the step was repeated twice. This could be made more accurate by using a 50mL pipette which would reduce the inaccuracies in each transfer by half, since there would be half as many transfers of liquid. Essentially, propagating errors were caused due to the increased number of measurements having to be taken. These propagating inaccuracies are the concave meniscus of the water in the pipette – meaning the bottom of the curve had to completely pass the line to achieve the desired volume of solution (seen in diagram to the right). The human parallax error while filling up the pipette was reduced slightly, since the person conducting the experiment squatted to eye level with the pipette. Steps 5 and 6 contain the same inaccuracies, but the propagating errors were heightened since a 10mL graduated pipette was used, hence 5 transfers of water or solution were required to fill each beaker with 50mL of solution. Therefore, the transfer of solution or water from a jar to the desired beaker was very accurate due to the app. 0.5% tolerance of a graduated pipette and four significant figure measurements given by the volumetric pipette, however some slight inaccuracies of the downward concavity of water, and human parallax error while measuring reduced the accuracy. Overall, the accuracy of these steps could be improved (by reducing propagating errors) only with improved equipment. I.e. a 50mL volumetric pipette as well as a 10mL, 20mL, 30mL and 40mL pipette for the beakers with diluted solution.

Other apparatus and precision:

The multimeter was the only other apparatus used for measurements, but since it is digital, involved no human inaccuracies. It was set to the 2000mV setting, giving readings to four significant figures. When the readings were divided by 1000, the result was data measured in Volts, correct to four significant figures/3 decimal places. The accuracy of the final results however were slightly limited by the concentration of the solutions measured only to two significant figures.

Comparison to published data (Nernst equation substitutions):

In the research section, the Nernst equation was used to make theoretical predictions for the rate of reaction (V) of each different concentration. Since the concentrations were only known to two significant figures, the final predicted rate of reactions were only accurately known to two significant figures showing no trend. These predicted figures although, were much larger than data gathered, still showed the same trend. The gap of voltage between gathered results and theory was very similar in all trials, and was actually caused by two main factors;

- The temperature was not at the standard 25^oC, which the constant in the numerator of the Nernst equation (0.0592) actually takes into consideration with its predictions

- The source of copper sulfate was contaminated with an iron solution

- The surface areas of the half-cells were different – magnesium had less surface area

These factors will be discussed in validity, on how they could be controlled, in order to achieve data closer to that of the predictions.

VALIDITY:

Variable	Controlled?	Improvement(s)
Controlled with no improvements needed
Multimeter and wires	Yes – stayed the same	–
Beakers	Yes – stayed the same	–
Magnesium sulfate	Exact same quantity of same source of solution used in each trial – remained unchanged in beaker	–
Magnesium electrode	The same electrode was used throughout, and simply remained in beaker	–
Atmospheric pressure	It is impossible to control, but was measured and remained constant throughout experiment at 1022.495hPa	None required, since research did not show correlation between rate of reaction and atmospheric pressure
Slight improvements able to be made – neglect able errors
Volumetric, graduated and plastic pipettes	Same ones were used throughout	After rinsing between different substance, let pipette dry completely
Salt bridge dunked in Potassium nitrate	Salt bridges were different lengths in each trial, and the amount of potassium nitrate absorbed was not measured, and likely different in each trial	Cut each salt bridge to the same length using scissors, and weigh the masses after dunking in potassium nitrate to ensure the same amount of solution is absorbed
Copper electrode	The same electrode was used throughout, but was rinsed between each concentration change	Dry the electrode before placing in new solution, since some eater drops may be carried in, slightly diluting the solution more than intended.
Major factors which decreased level of validity (reasons for large gap in voltage between data gathered, and predictions – mentioned in ACCURACY)
Variable	Why it impacted the results – how to improve
Temperature	The temperature was measured and remained constant at 18.5^oC, but no measure was taken to ensure the control of this. This temperature is however 6.5^oC below the standard potential energy conditions, meaning that the rate of reaction (V) will be reduced. This is because the predictions of the rate of reaction found from the standard energy potentials and the Nernst equation assume room temperature of 25^oC. When temperature is higher, the ions in the electrolyte solutions move more freely, causing a higher voltage, this contributed to the lower voltage readings than predicted. To improve this, conduct the experiment in a heatd water bath, measured to standard room temperature
Surface area of electrode in solution	Surface area of electrode in respective solution was consistent in each trial, however each half-reaction had a different contact surface area  Magnesium had less surface area in solution than copper. This meant that the copper ions flowing into the magnesium solution couldn’t all be absorbed due to the lower surface area. An improvement is to measure the surface areas using callipers, and submerge the same distance in the beaker.
Copper sulfate solutions (1 – 0.2mol.L-1)	The same source was used throughout, but was visibly contaminated with iron solution – brown/orange substance found in beakers. To fix this problem, use a new, cleaner source of solution, or filter the iron out of the copper prior to measuring and diluting.

The ACCURACY section mentioned that results were different to predictions, and these three variables were described in ways they weren’t controlled – reasons for not meeting expected results, and suggested improvements were given to improve validity for next time.

This table shows many controlled variables and some neglectable variables, the independent variable [CuSO₄] was quantifiable and changed at a consistent interval and the dependant variable was also quantifiable. This kept a fair level of validity of the actual method, however the major uncontrolled variables made the experiment in-valid in proving the hypothesis since there were too many significant factors which could have changed what the experiment was set out to measure.

Discussion of results/trends:

This experiment had very consistent results across each concentration of copper sulfate (independent variable). The margin of errors were:

- 1.0M – 0.004V,

- 0.80M – 0.002

- 0.60M – 0.006

- 0.40M – 0.006

- 0.20M – 0.004

Each concentration had four consistent results, with the maximum accepted error margin being set at 0.010V. Any results outside this margin were deemed outliers. Only trial 4 for 0.60M and trial 3 for 0.20M were not within the given margin of error, however these results were excluded from the average, still letting the experiment be reliable since there were still three consistent results.

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The average results presented the trend that as the concentration of copper sulfate was diluted/decreased, the rate of reaction (Volts) decreased i.e. the more concentrated the independent variable, the higher the voltage. The graph displayed this trend and presented an almost perfectly linear relationship between the voltage of the cathode half-cell (copper) and the rate of reaction (V) of the galvanic cell. The R²value of the line is 0.9734, giving the results trend line a value of R=0.9866.

This means that the data extracted from the experiment, was extremely reliable in supporting the hypothesis, since each change to the concentration of the included results are consistent within a small margin of error (MAX. 0.006 difference), with only two outliers, excluded from final results. The results are also very reliable, reading down the table (or in the graph), presenting an almost perfectly linear relationship between the dilution to concentration and a decrease in rate of reaction (V).

Conclusion:

Therefore, the higher the concentration of one of the reactant solutions (copper sulfate) in a galvanic cell, the higher the rate of reaction (Volts), justifying the hypothesis made.

Bibliography (A-Z by Author(s) surname)