Equilibrium of the Iron Thiocyanate Reaction

Equilibrium of the Iron Thiocyanate Reaction

Many chemical reactions are found to proceed to an equilibrium at which a mixture of both reactant and products is present. The extent to which reactants are converted to products is expressed by the equilibrium constant, K.

This experiment will probe the equilibrium of Fe(III) ions reacting with the thiocyanate ion, SCN(. The reaction has two possible products given below, in lab this week you will determine which of these two reactions actually occurs.

Fe3+ + SCN( ( FeSCN2+Rxn 1

Fe3+ + 2SCN( ( Fe(SCN)2+Rxn 2

The nice aspect of this equilibrium is that the reactants are colorless, but the product is a deep red color. The concentration of the products at equilibrium will be measure by observing the absorption of blue light.

By measuring the concentration of the product produced from five different reactant concentrations, you will determine which equilibrium above occurs for the iron thiocyanate reaction.

Measuring Concentration by Optical Absorption:

Along with studying a system at chemical equilibrium, this laboratory will introduce the concept of absorption spectroscopy. The red product of the reaction is absorbing blue light, allowing the compound to transmit red room light to your eye. There is a simple relationship between the amount of light a sample will absorb and the concentration of the absorbing molecule.

In this relationship A is the absorbance (read by a spectrophotometer), c is the solution concentration (mol/L), l is the path length of the sample (cm) and ( is the extinction coefficient of the molecule (in this case it has units of L / mol cm). ( indicates how easily the molecule absorbs light at a specific wavelength. This relationship between solution absorbance and analyte concentration is known as Beers law.

The important aspect of Beers law in this instance is that Absorbance is linearly related to the concentration of the absorbing species. This allows the creation of a calibration curve to be constructed between absorbance and known concentrations of the product. In Figure 1 the iron thiocyanate ion absorbance at 470 nm is plotted as a function of concentration. This plot allows the rapid determination of solution concentration from solution absorbance. As you perform these absorbance measurements, consider how long it took to make a concentration measurement using freezing point depression. (The advantage of freezing point depression is you do not need to know the identity of the compound.)

Figure 1. An absorbance vs. concentration plot for the product of the iron thiocyanate equilibrium.

Experimental Procedure:

Preparing the Five Equilibrium Mixtures:

From the front of the room obtain a 5 mL pipet, a graduated pipet, a red pipet bulb and two spectrophotometer tubes. Place the spectrophotometer tubes in your test tube rack.

In your smallest beakers obtain from the reagent bottles 30 mL of 0.002 M Fe(NO3)3 and 20 mL of 0.002 M KSCN.

The five solutions you will prepare will have equal amounts of iron ions and differing quantities of thiocyanate ions. Each solution will have the same 10 ml volume, made by adding varying amounts of H2O.

In your test tube rack add five clean test tubes. Label your test tubes 1 ( 5 by placing your rack on a sheet of paper and writing the number on the paper. Fill the test tubes with the reagents indicated in Table 1. It is best to fill all tubes first with the amounts of thiocyante using the graduated pipet. Place the pipet bulb gently over the end of the pipet. The bulb should be on so loosely that it would fall off if you held the pipet upside down. Do not jam the bulb onto the pipet. Allow your partner to hold your reagent beaker. To deliver 1.0 mL of SCN(, draw into the pipet 3 or 4 mL of solution. Remove the bulb and allow the solution to drain back into the reagent beaker. Place your index finger over the top of the pipet to control the flow. Slow the flow down with your finger, stopping it at the 1 mL mark (or 9 mL mark depending on the labeling of your pipet). Expect to draw up and drain the solution several times before you get the volume correctly. Transfer the SCN( solution to the correct test tube. Repeat the addition for the other four test tubes. (Be sure that all members of the group have the pleasure of pipetting.)

After the SCN( solutions have been added, use the 5.0 mL pipet to measure the Fe3+ solutions. The drawing up in this 5 mL pipet is more difficult because there is only a small volume between the 5 mL mark and the top. Have the bulb on loosely, so that it is easily removed as the liquid nears the mark. Try hard not to fill the red bulb with Fe(NO3)3. If the solution gets on your hands wash them immediately. Expect to perform several extractions before the correct volume is reached.

Rinse the graduated pipet with water and add the necessary amounts of water to each test tube. This should take the volume of each test tube to 10 mL. You will need a little beaker of DI water from which you will draw the water into the pipet.

mL of 0.002M SCN(mL of 0.002M Fe3+mL H2O

Tube 11.05.04.0

Tube 22.05.03.0

Tube 33.05.02.0

Tube 44.05.01.0

Tube 55.05.00.0

Obtain from the front of the room a piece of parafilm so that the test tubes can be shaken. Cover the top of each test tube with the parafilm and press down with your thumb. Invert the tube several times to completely mix the solution. Start with the most dilute SCN( solution so that as the next solution is mixed the parafilm does not accidentally contaminate a less dilute solution.

Any reagent remaining in your beakers should be placed in the waste-beaker in the hood.

Calculate the initial concentrations of Fe3+ and SCN( in each test tube 1 - 5. This is just a dilution calculation.Measurement of the Product Concentration:

Take your solutions to a Vernier colorimeter to measure the absorbance of the FeSCN ion-complex.

Since you have the absorbance vs. concentration curve at a wavelength of 470 nm, select the 470 nm wavelength with the arrow keys on the touch pad of the colorimeter The colorimeter needs to be blanked. This means setting colorimeter to read a water sample as having zero absorbance. Pour a small amount of the blanking solution (DI water) into a colorimeter cuvet, about full with DI water. Wipe the outside of the spectrometer tube with a Chemwipe or paper-towel. Insert the cuvet into the colorimeter so that the clear sides of the cuvet are along the axis of the colorimeter. Close the colorimeter lid and press the Cal button on the colorimeter. Wait for the lights to stop flashing. The LabQuest unit should read 0.000 Abs. Discard DI water down the sink.

After the spectrometer has been calibrated, measure the absorbance for each FeSCN solution. Start with the most dilute solution and end with the most concentrated solution. Obtain a waster beaker for your table. Pour into the cuvet solution #1 (the most dilute). Pour this solution into a waster beaker at your table. Now you can add solution #1 to the cuvet and be sure that the concentration in the cuvet is the same as the test tube you prepared the solution in. Place the solution in the spectrometer and record the absorbance.

After recording the absorbance in a sample empty the cuvet into the waste beaker. Then make an absorbance measurement on the next solution, by rinsing the cuvet once with the solution you intend to measure.

Determine the concentration of the FeSCN ion-complex in each reaction tube by converting all calculated absorbance to concentration using the linear equation in the Figure 1 calibration curve.

All of the test tube solutions can be placed in the waste beaker. The colorimeter cuvet should be rinsed thoroughly with DI water and returned to the front of the room.

Ask your instructor if you are to prepare for next weeks lab on Acid-Base Titration, by prepare approximately 500 mL of 0.1 M NaOH solution from a stock NaOH solution of 6 M. This requires you to do a dilution calculation. The 500 mL 0.1M NaOH solution can be stored in a corked flask until next weeks lab.

Determining the Equilibrium Constant for the Two Proposed Reactions:

In this experiment data collection is rather quick and data analysis is rather lengthy. As you did in the pre-lab assignment, determine K for each proposed reaction using all five initial concentrations. Recall that you have measured the [FeSCN]eq concentration and you will have to use the reaction stoichiometry to calculate [Fe3+]eq and [SCN(]eq from the initial concentrations of [Fe3+]0 and [SCN(]0. It is probably a good idea to write the ICE tables for each reaction and create a data table that contains the components needed to calculate K (ex. [Fe3+]eq, [SCN(]eq and [Fe(SCN)??+]eq) for each test tube.

From the reproducibility of the equilibrium constant, which equilibrium reaction

Fe3+ + SCN( ( FeSCN2+Rxn 1

Fe3+ + 2SCN( ( Fe(SCN)2+Rxn 2

occurs in the reaction of Fe3+ with SCN(?

Report the equilibrium constant, K, as an average of the five experimental measurements. The uncertainty in your measurement can be given as ( the standard deviation in your measurement. If you are unfamiliar with standard deviation calculations, ask the TA to show you how this can quickly be calculated in Excel.

In your abstract briefly describe the system studied and the experimental method used to determine K. Also report the equilibrium constant and its uncertainty.

Turn in your lab sheets before you leave lab.

Place your 0.1 M NaOH solution at the front of the room. It will be stored until next week. EMBED Excel.Chart.8 \s

_1203044491.unknown

_1297343455.xlsChart1

0.6683745

0.442303

0.32926725

0.2162315

0.10319575

0.0353743

0.78141025

[Fe(SCN)?]?+ mol/L

Absorbance

[Fe(SCN)?]?+ calibration curve at 470 nm

Abs = 4521.4[conc]

Sheet1

concabs

1.12E-040.514

9.20E-050.3858

4.50E-050.1864

2.40E-050.1088

1.50E-046.68E-01

1.00E-044.42E-01

7.50E-053.29E-01

5.00E-052.16E-01

2.50E-051.03E-01

1.00E-053.54E-02

1.75E-047.81E-01

Sheet1

0

0

0

0

Sheet2

0

0

0

0

0

0

0

[Fe(SCN)?]?+ mol/L

Absorbance

[Fe(SCN)?]?+ calibration curve at 460 nm

Abs = 4521.4[conc]

Sheet3