Best method for determining silver ions in solution
Chemistry Coursework
Aim
I am going to investigate which is the best method for determining silver ions in solution. I am interested to see whatever a electrode potential method or a chemical method (e.g. titration) is more appropriate, especially at low concentration.
Background Information
Redox If we consider the following reaction:
2Na(s) + Cl2(g) --> 2Na-Cl+(s)
Salt (Sodium Chloride) is formed when Sodium is added to Chlorine gas. Sodium Chloride is an ionic compound, which is made up of Na+ and Cl- ions. The difference between Na(s) and Na+(s) is that an electron is removed from the Na atom and this a positive charge formed as the number of protons, which are positively charge is not equal to the negatively charged electron. But NaCl(s) is not charged overall, so where is the electron gone? If we look at the equation more closely, we can split the reaction into two parts, and this is called half equations:
2Na(s) --> 2Na+(s) + 2e-
Cl2(g) + 2e- --> 2Cl-(s)
The first equation shows that each sodium atom lost one electron and formed a Sodium ion and an electron. While each chlorine atom gained an electron and form Chloride ions. These two equation is balance and overall there is no electron lose as once a sodium has given up its electron, a chlorine atom will form an ion with this electron. We called this type of reaction Redox - Reduction and Oxidation. The sodium in this example is being oxidized as the atom loses an electron while chlorine is being reduced as it gains one electron.
Electrochemical cell and electrode potential
In the last example, we saw sodium giving up electron and chlorine accepting electron. But sodium could well be accepting electron while chlorine giving up electron. So there must be something deciding the direction of the electron flows. We can investigate this using half cells. If we combine two half cell together, we can form a electrochemical cells. The following diagram shows a simple one:
The voltmeter will register an EMF across the electrode which means that there is a current flowing which means that there is electron flowing in the circuit. If we use a perfect voltmeter (infinite resistance which means that the solutions are not driving a current), the reading on the voltmeter will be the Electrode potential, with the symbol Ecell and units Volts.
Each of this half cell has its own electrode potential. If we look at the copper half cell in the above example, we can write an equilibrium equation:
Cu2+(aq) + 2e- ? 2Cu(s)
The electrode potential measures this equilibrium position between Copper metal and its ion. If the equilibrium lies to the right, the ions has a greater tendency to accept electron and therefore become positive. If we put two of the half cell together, the more positive terminal of the cell will be the one which is more willing to accept electron. The electrode potential also measures how spontaneously this reaction can run.
Standard Electrode Potential.
In order to compare this equilibrium, we need a standard way to measure this constant. The way we do this is to select a common half cell, and measure other against it. The common half cell we used is hydrogen half-cell, which contains hydrogen gas at 1atm at 298K, with a acid solution contains 1 mol dm-3 of H+ ions. The electrode we use is platinum which does not react with both hydrogen gas and ions. We define the potential of this half cell as 0.00V and the standard electrode potential of a half cell is defined as the potential difference between it and the standard hydrogen half cell. The half cell use to find the standard electrode potential is done by dipping metal into 1 mol dm-3 of metal salt solution at 298K. It is given the symbol E??
The Nernst Equation
The relationship between the E(cell), E???temperature and concentration is given by the Nernst Equation, which derived from the free energy of reaction and the 2nd law of thermodynamics:
E(cell) = E? - (RT / nF) ln Q
where:
Q = The reaction quotient
( [product] / [reactant] )
E(cell) = The actual electrode potential measured
E? = The standard electrode potential
T = Temperature of reaction in Kelvin
n = Number of moles of electron transferred
in reaction
F = Faraday Constant (96500 C mol-1)
R = Molar Gas Constant (8.31 Jmol-1K-1)
The Nernst Equation can predict the electrode potential between half-cells under the influence of different temperature and concentration of the solution. The equation shows that it we increase the temperature of the system, the E(cell) will decrease. The different concentration of the reaction also changes the value of Q. Therefore, the concentration of the half cell solution will affect the E(cell) value of the cell. We can work out the concentration of the ions in the cell involve by look at the E(cell) value of the system.
Chemical Analysis : Precipitation Titration
We can also use a chemical analysis to analyse concentration of an ion in solution. We use a titration to determine the concentration of the solution. It involve a reaction in which known concentration of a solution is reacted with an unknown concentration of another solution. Then, an indicator will react with the another solution to give a visual sign, usually a colour change or precipitation. When all the ions in the solution are being used up in the primary reaction between the known and unknown concentration solution, any excess ions added ...
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Chemical Analysis : Precipitation Titration
We can also use a chemical analysis to analyse concentration of an ion in solution. We use a titration to determine the concentration of the solution. It involve a reaction in which known concentration of a solution is reacted with an unknown concentration of another solution. Then, an indicator will react with the another solution to give a visual sign, usually a colour change or precipitation. When all the ions in the solution are being used up in the primary reaction between the known and unknown concentration solution, any excess ions added to the solution will start to react with an indicator. We can use this visual sign to tell us this is the end point of the reaction. By looking at how much of the solution is used, we will able to work out the concentration of the solution and as will know how many moles of ions are used in the reaction.
There are many ways to determine silver ions in the solution using a chemical titration method. I am going to use a precipitation titration between silver nitrate and sodium chloride. The product of the titration reaction, silver chloride, is only sparingly soluble in water. There a is precipitation reaction is happening. The indicator, which is Potassium Chromate, will then form an orange colour precipitate with sliver. The appearance of the orange precipitation is determine by the solubility product of a solid.
Solubility Product
No substance is completely insoluble. There is always a small amount of substance dissolve even they form precipitate in water. The equilibrium position is measured by solubility product, which is defined by:
AB(s) --> xA+(aq) + yB-(aq)
Ksp = [A]x[B]y
We can use this value to predict whether the substance is going to dissolve or not.
Planning the investigation
Brief Outline:
I am going to investigate the way of using electrode potential to measure the concentration and use a titration method. Using Nernst Equation, I am able to work out expected E(cell) value when varying the concentration of [Ag+] ions with standard copper sulphate solution (at 1 mol). Then, by measuring the E(cell) using different concentration of silver ions to very dilute concentration and standard copper sulphate solution, I am able to compare two sets of value and work out the error of this method. Then, by doing a titration, I can compare whether the electrode potential method or the titration is more accurate and compare the value of the concentration.
Apparatus Requirements:
High Resistance voltmeter.
Crocodile clips and leads.
Light bulb.
Breakers and Conical Flask
Filter paper.
Thermometer
Filter Paper strips
Sand paper
Standard Flasks for making up solution
Water bath
Pipette and Burette
Stands
Chemical Requirements:
Copper Sulphate and Sodium Chloride crystals for making up solutions.
0.1M of Silver Nitrate solution.
Copper and Silver metal foil as electrodes.
Potassium Nitrate Solid for making salt bridge
Distilled Water
Potassium Chromate (VI) as indicator
Hazards / Risk Assessment.
For Copper Sulphate (solid) used in this experiment is irritating and can be harmful if ingested. They also cause burns if exposed to skin for a long period. Therefore, care has to be taken to prevent any ingestion or spillage of Copper sulphate. If the chemical got into the eyes, wash with lots of water and seek medical advice in serious case. For spill on skin, it has to be wash in order to prevent burns. If the chemical is spilled, wash with a lot of water and throw it into the drainage.
Silver Nitrate is corrosive to skin and eyes. It also cause blackening of skin. Potassium Chromate is irritant. If spill onto skin or eyes, then wash with lots of water.
I will also wear Laboratory coat and eye protection spectacles in case anything spillage onto the eyes and skin. I will also no be eating in the laboratory in order to prevent any accidental ingestion of chemical.
Experiment Planning Details
First. I will give my experiment preliminary trial to see whether it works or not. First, I am going to make up solutions of Copper Sulphate.
Making solutions for the experiment:
. I will take a volumetric flask (250ml) and according to their molar mass, I will measure out the following amount of solid as in table 1 (except the salt bridge solution, where saturated solution is required.).
. I will carefully put the solid into a breaker and add small amount of distilled water into the breaker. I will dissolve all the solid into this small amount of solution.
. Carefully transfer all the solution into the flask and wash the breaker thoroughly with distilled water and all the liquid should go into the flask.
. Fill the flask up to the mark with distill water and put the plastic top back on the flask.
. Swirl and turn the flask for a minute.
Chemical Used
Molar Mass (g)
Mass required in 1dm3 solution of Concentration 1 mol dm-3 (g)
Mass required in 0.25dm3 solution of concentration 1 mol dm-3 (g)
NaCl
58.5
58.5
4.625
CuSO4.5H2O
60 + 90 = 250
250
62.5
KNO3 (Saturated)
01.1
01.1
25.3
As I am going to investigate the different concentration of the Ag+ ions in solution. I will use the following instructions which enable me to produce a range of solutions from 0.1M to very dilute:
. Using the standard 0.1M AgNO3 solution and a pipette, I will pipette 25ml of them into the standard flask (250ml) and then fill the rest up to the line with distilled water.
. For 0.01M, I will pipette 1/10th of the 1M solution (25ml) and then fill the rest up to the line with distilled water.
. Repeat the above steps, but take the solution from step 2 until all the required solution is made.
By doing this, I am dividing the concentration of the solutions by 1/10th every time. Thus, I will obtain a series of solution of AgNO3(aq) with different concentration. I will also need a series of Sodium Chloride solution with different concentration for titration as well. I will use the above steps.
Working out the E(cell) from Nernst Equation:
Using a spreadsheet on computer (next page), I am able to work out a set of prediction on what E(cell) value is going to be like using the Nernst Equation at 298K (room temperature) using the following half cells, which all of the E(cell) going to be measured against:
2Ag+ | 2Ag || Cu2+ (1M)| Cu
The prediction assume that the concentration of both half cell solution doesn't changes much after I switch on the apparatus as I only used up a very small amount of metal ions and metal.
Finding out the E(cell) by experiment
Once I have worked out the predicted value of the E(cell) , I am going to find out the value experimentally using the following method:
. Using 25ml breakers, I fill up one with Copper (II) Sulphate solution at 1 mol dm-3 and another one fill with Sliver nitrate solution with the concentration which I am going to test. I am going to take note of the temperature of the solution when the experiment is going. I will try to keep the experiment running in constant temperature.
2. I will cut a piece of Copper and Sliver metal strip and remove the oxides from its surface using sand paper.
. Salt bridge is put across the breakers so that ions can pass through the salt bridge. Then, I will connect up the circuit as follows:
2. Once the circuit is connected, I will write down the E(cell) across the cell, which is displayed on the voltmeter. The voltmeter should be high resistance, as this means that no current is drawing from the cell, as stated in the research section.
3. I will repeat this experiment three times to make sure the results are correct.
4. Then, I will repeat the above steps for all the concentration of the
sliver solution I have made up.
Chemical method - Titration
After I have done the above, I will compare the results with a titration method. This titration is using precipitation reaction between chloride and nitrate, with chromate (IV) as indicator which produce red precipitate if excess silver is present in the solution. The reaction between Silver Nitrate and Sodium Chloride is:
NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
The excess Sliver ion will form an red precipitate with chromate as a result.
Method:
. Fill the burette with Silver Nitrate at the concentration which I am going to test (0.1M, for example)
2. Run small amount of the Silver nitrate out to remove the air from the tip. Then record the start position.
3. Pipette 10ml of sodium chloride solution with corresponding concentration into a small conical flask and put about 5 drops of Potassium Chromate into the solution.
4. Slowly add Sliver Nitrate into the solution. I will stop when red colour appear, which doesn't fade away by shaking.
5. Record the reading on the burette.
6. Repeat this three times and get an average results.
I will repeat the above with different concentration of Sliver Nitrate. The concentration of the sodium chloride ion in the flask will be the same as the concentration of the sliver ion in the burette. This is because the a large amount of silver nitrate will be needed if it is not reduced. I expected that 10ml of Sliver Nitrate will needed to reach the end point for each titration. This is because they are same concentration in the flask and burette and from the chemical equation, 1 mole of Silver ions will react 1 mole of Chloride ions to form 1 mole of Sliver Chloride.
Experimental details and results
Preliminary experiments:
I have first tried out the electrochemical cell experiment. I found that the reading of the voltmeter most of the time when I am dealing with high concentration of sliver sulphate are the same and little variation. But at the low concentration, the voltmeter reading varying a lot. This is due to the fact that our voltmeter, although is a digital one and has a very high resistance, is not prefect. Therefore, they still draw a very small current from the solution. Therefore, I repeat all the experiment 4 times and get an average.
For the titration reaction. I found that the indicator doesn't respond (Very little colour change) to sliver nitrate if it is dilute to 0.0001M. As a result, I will need to leave the experiment with concentration below 0.001M.
Results:
Electrochemical cell experment:
Concentration Ag+ (M)
E(cell) (V)
Experiment
2
3
4
5
Average
0.100000
0.40
0.39
0.41
0.41
0.40
0.010000
0.36
0.37
0.35
0.37
0.36
0.001000
0.30
0.32
0.31
0.31
0.31
0.000100
0.24
0.24
0.21
0.24
0.23
0.000010
0.28
0.29
0.15
0.20
0.17
0.17
0.000001
0.11
0.14
0.14
0.13
The graph on the page 11 shows the resulting E(cell) against natural log of the sliver ion concentration (ln[Ag+]). It shows that there is a direct proportional relationship between the concentration of the cell. On page 12 there is another graph showing the prediction value and my experimental results. The graph suggested that the experiments results are fairly accurate compare the prediction by Nernst Equation.
Notice that the E(cell) value from the Nernst Equation we calculated doesn't suggest what the rate of reaction is like. Therefore, it is very difficult to say how quickly the ions is using up. I am not interested in the rate of the reaction in this investigation therefore I make sure that the drop of concentration in both solution is minimal by only connect the circuit for only short time.
The error shown in the E(cell) value may be introduce by the resistance of the wires, voltmeter and the present of salt bridge although these factors should be minimal. The condition of the solution is also important as the reaction is a equilibrium position, especially the temperature of the solution. Therefore, I am going to extent my investigation to the effect of temperature on the E(cell) value.
Titration Experiment
The manipulation of the results from above:
Concentration of AgNO3 being tested / NaCl in the reaction chamber (mol dm-3)
Volume of Ag+ ion used (cm3)
Amount of NaCl in the flask (moles)
Suggested concentration of AgNO3
(mol dm-3)
0.1
1.73
0.001
0.0850
0.01
5.02
0.0001
0.0066
0.001
7.37
0.00001
0.0005
These results seems to be very far from the correct concentration. This is due to the fact that it is quite difficult to decide where is the end point for all the concentration below 0.1M. Sometimes the colour of the solution is too weak and thus the end point in very difficult to determine.
There is also an error introduced in the results. Potassium Chromate does not form a red precipitation instantly when there are excess silver
ion present in the solution. And extra silver ion is needed to create the colour change to be detected by the eyes, especially at low concentration. This is shown in the results as more and more of the solution is required when the solution is more diluted. To determine the amount need to correct, we can use an indicator blank determination. It is done by measuring the volume of silver nitrate is needed to turn a same volume of indicator into the desire visual end point. It is found that it's not significant for 1M solution. For 0.1M solution, it will introduce a 0.04% error to the results while 0.01M solution, it will introduce a 0.4% of error. For 0.001M, its about 2%.
The Effect of temperature on both analysis
After the initial experiment, I am interested in what factors will affect these method of analysis and what their effect are.
For the electrochemical cell experiment, I repeat the experiment I stated at the planning stage, but I also put the breaker contains both solution into a water bath. Then, I will check the temperature inside the breaker and write it down. I made sure that both solution is at the same temperature when I connect the circuit up. I will then connect the circuit and see the value of E(cell) at different temperature.
Concentration of Ag ions (mole)
Temperature
33oC
43oC
50oC
60oC
0.1
0.40
0.40
0.39
0.39
0.01
0.35
0.34
0.34
0.33
0.001
0.28
0.29
0.28
0.26
0.0001
0.22
0.22
0.26
0.20
0.00001
0.19
0.18
0.21
0.20
0.000001
0.17
0.18
0.19
0.19
On the next page it shows a graph which combines all the result into one graph. According to the Nernst Equation, the lines should be parallel with same gradient for each line. The graph suggested that there are some significant change in the E(cell) value when the temperature change. But the change was not as predicted - the higher temperature should decrease the E(cell) value. This may due to the fact that the internal resistance changes inside the solution which affect the current flowing in the system.
The following is the results of doing the titration of silver nitrate at different temperature.
This set of experiment is done using 0.1M of the silver nitrate titrated against 10ml of 0.1M of sodium chloride solution. The reaction vessel (conical flask) is left in the water bath and get to the temperature which to be investigated. The results shows a more accurate than the original experiment. (We expected that there should be 10cm3 of the sliver nitrate used in the titration). I expected that the results should be more error introduced - the solubility product of the sliver nitrate increases with temperature thus the red precipitation is not appearing unless with more sliver nitrate added to the solution.
Conclusion
There are different situation where each of this method whould be used in the determination of the silver. The E(cell) method give you an good approximation of the silver ions concentration. But it depends on the concentration and purity of the solution, as impurity does affect electropotential. The E(cell) experiment can be used in all temperature as the E(cell) value can be calculate using Nernst Equation.
Chemical method, like titration, is used when there are only limited amount of test solution available. The indicator is very important in the success of the determination as we saw that the indicator can introduce an blank error into the experiment.
The temperature of which the determination happening also affects the result. Generally, all the reaction should take place in room temperature as we have found out that at higher temperature, there is a different value obtained.
Bibliography
I have consulted the following books throughout this investigation:
Rendle, Vokins and Davis, 1991. Experimental Chemistry, Second Edition. London: Arnold.
Ottewill and Walsh, March 1996. 'Electrochemical Cell' and 'How to use electrochemical cells' . Chemistry Review Vol:4 Num:4.
Atkins, 1990. Physical Chemistry, Fourth Edition. Oxford: OUP.
Fine and Beall. Chemistry for Engineers and Scientist.. USA: Saunders College Publishing.
Salter Advance Chemistry Course 1994. 'Redox', 'Redox reactions and electrode potentials' in 'Chemical Ideas'. Oxford : Heinemann Educational
Jeffery, Bassett, Mendham, Denney: Vogel's
