Chemical Equilibrium: Finding a Constant, Kc

 Abstract

The equilibrium constant  is present in every single reaction. When iron(III) and thiocyanate are added together equilibrium is established. Therefore, the value of  can be determined by changing the volume of KSCN.  The final equilibrium constant  was determined to be and the percent error was 28%.

Introduction

In this experiment the main idea is to determine the equilibrium constant  for the following reaction.

When the reactions iron (III) and thiocyanate react and produce thiocyanoiron(III).  The reactions produce an equilibrium constant  Therefore, the concentration of each ion must be known in order to find the equilibrium constant. Four trials were conducted with different volumes of KSCN and , to find the concentrations of each reactant and product in the equation above. Finding the concentrations would lead to finding the equilibrium constant by the expression below from the Law of Mass action, since all of the concentration are known.

The order of each product and reaction in the above relationship is one, due to the equilibrium established in the reaction.  A standard concentration was prepared with a very high concentration  and a very low concentration of . Due to this high difference of concentration the reaction uses up all the concentration of [and forces the reaction to occur. This behavior is described by LeChatelier’s principle, which states that due to this disturbance in the difference of concentration the reaction returns to equilibrium. From the balanced equation above since all the  is used the initial  is equal to the standard

In order to find the concentration of  at equilibrium Beer’s Law was used. Beer’s Law establishes a relationship between the concentration and the absorbance of the reaction.  Shown in figure 1 as the concentration of  increases the absorbance of the blue light in the calorimeter would also increase. Since the solution changes color to light red the absorbance of blue light can be measured with the colorimeter. Therefore, the following relationship is established.

Where setting over the absorbance in standard equal to  over the absorbance at equilibrium can yield  by rearranging the equation above since the rest is known yields.

 

By finding  at equilibrium  initial can be found. Therefore the following relationship was used.

Since the total volume is 10mL and the volumes are given in Table 1  initial concentrations were found. By knowing the initial concentrations of , the following relationship was used to determine at equilibrium due to the 1:1 ratio in the reaction.

Therefore, the equilibrium constant can be determined. Furthermore, the accepted value for the equilibrium constant is

Data

Table 1: Required volume

Test Tube number  (mL) KSCN (mL)  (mL)
1 5 2 3
2 5 3 2
3 5 4 1
4 5 5 0
5 9 1 0

 

Table 2: Concentrations for initial reactants

Trial  (M)   (M)
1 0.001 0.0004
2 0.001 0.0006
3 0.001 0.0008
4 0.001 0.001
5 0.0002 0.0002

 

 

 

 

 

 

Results

Table 3: Calculated

Trial (M)   (M) Absorbance
1 0.00005649 0.000944 0.000344 174.30 0.198
2 0.00007874 0.000921 0.000521 163.98 0.276
3 0.00010556 0.000894 0.000694 169.95 0.37
4 0.00013181 0.000868 0.000868 174.87 0.462

 

Average

Discussion

The average equilibrium constant was calculated to be.  The accepted value for  is 133, which caused a 28 percent error using the equation below.

The experiment required a lot of perscion measuring each solution in the test tube and drying the water particles after washing the serological pipet.  The water particles in the pipet can cause a slight error when mixing with the solution.  Four test tubes were prepared with different volumes as shown in table 1. By changing the volume of KSCN while maintaining an entire volume of 10mL. Adding water leads to different initial concentration forshown in table 2. However,  concentration stays the same since all four trials require a volume of 5mL. The standard  was important in order to find the equilibrium concentrations for  in each trial.  Table 3 shows all four trials with the calculated equilibrium constants. Shown in table 3 trials 1, 3 and 4 are nearly close to each other, by averaging the three trials the average is 131. This indicates that the is constant in the reaction. However, trial 2 has a low  the main source of this low equilibrium constant is small amounts of water particles in the serological pipet and in the test-tube, since they were washed with water.  Moreover, some of the factors maybe from not mixing the solution in the test-tube thoroughly.

 

Sample calculation:

Chemical Equilibrium: Finding a Constant, Kc

 Abstract

The equilibrium constant  is present in every single reaction. When iron(III) and thiocyanate are added together equilibrium is established. Therefore, the value of  can be determined by changing the volume of KSCN.  The final equilibrium constant  was determined to be and the percent error was 28%.

Introduction

In this experiment the main idea is to determine the equilibrium constant  for the following reaction.

When the reactions iron (III) and thiocyanate react and produce thiocyanoiron(III).  The reactions produce an equilibrium constant  Therefore, the concentration of each ion must be known in order to find the equilibrium constant. Four trials were conducted with different volumes of KSCN and , to find the concentrations of each reactant and product in the equation above. Finding the concentrations would lead to finding the equilibrium constant by the expression below from the Law of Mass action, since all of the concentration are known.

The order of each product and reaction in the above relationship is one, due to the equilibrium established in the reaction.  A standard concentration was prepared with a very high concentration  and a very low concentration of . Due to this high difference of concentration the reaction uses up all the concentration of [and forces the reaction to occur. This behavior is described by LeChatelier’s principle, which states that due to this disturbance in the difference of concentration the reaction returns to equilibrium. From the balanced equation above since all the  is used the initial  is equal to the standard

In order to find the concentration of  at equilibrium Beer’s Law was used. Beer’s Law establishes a relationship between the concentration and the absorbance of the reaction.  Shown in figure 1 as the concentration of  increases the absorbance of the blue light in the calorimeter would also increase. Since the solution changes color to light red the absorbance of blue light can be measured with the colorimeter. Therefore, the following relationship is established.

Where setting over the absorbance in standard equal to  over the absorbance at equilibrium can yield  by rearranging the equation above since the rest is known yields.

 

By finding  at equilibrium  initial can be found. Therefore the following relationship was used.

Since the total volume is 10mL and the volumes are given in Table 1  initial concentrations were found. By knowing the initial concentrations of , the following relationship was used to determine at equilibrium due to the 1:1 ratio in the reaction.

Therefore, the equilibrium constant can be determined. Furthermore, the accepted value for the equilibrium constant is

Data

Table 1: Required volume

Test Tube number  (mL) KSCN (mL)  (mL)
1 5 2 3
2 5 3 2
3 5 4 1
4 5 5 0
5 9 1 0

 

Table 2: Concentrations for initial reactants

Trial  (M)   (M)
1 0.001 0.0004
2 0.001 0.0006
3 0.001 0.0008
4 0.001 0.001
5 0.0002 0.0002

 

 

 

 

 

 

Results

Table 3: Calculated

Trial (M)   (M) Absorbance
1 0.00005649 0.000944 0.000344 174.30 0.198
2 0.00007874 0.000921 0.000521 163.98 0.276
3 0.00010556 0.000894 0.000694 169.95 0.37
4 0.00013181 0.000868 0.000868 174.87 0.462

 

Average

Discussion

The average equilibrium constant was calculated to be.  The accepted value for  is 133, which caused a 28 percent error using the equation below.

The experiment required a lot of perscion measuring each solution in the test tube and drying the water particles after washing the serological pipet.  The water particles in the pipet can cause a slight error when mixing with the solution.  Four test tubes were prepared with different volumes as shown in table 1. By changing the volume of KSCN while maintaining an entire volume of 10mL. Adding water leads to different initial concentration forshown in table 2. However,  concentration stays the same since all four trials require a volume of 5mL. The standard  was important in order to find the equilibrium concentrations for  in each trial.  Table 3 shows all four trials with the calculated equilibrium constants. Shown in table 3 trials 1, 3 and 4 are nearly close to each other, by averaging the three trials the average is 131. This indicates that the is constant in the reaction. However, trial 2 has a low  the main source of this low equilibrium constant is small amounts of water particles in the serological pipet and in the test-tube, since they were washed with water.  Moreover, some of the factors maybe from not mixing the solution in the test-tube thoroughly.

 

Sample calculation:

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