ka determination report

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ABSTRACTThe aim of this experiment is to determine the acid ionization constant or Ka value of the weak acid by using the titration method with sodium hydroxide solution and by measuring the pH value for the weak acid. For the titration method, sodium hydroxide as a base solution was titrated into the solution of weak acid drop by drop. The titration graph thus plotted against the pH value of the solution and the volume of NaOH required to neutralize the weak acid solution in each titration is determined. The half equivalence point and the equivalence point are also determined from the curve. During this experiment too, the average molarity of the unknown acid solution used in the titration also must be calculated. For the second experiment in which the measuring method is applied, the procedure begins with the measurement step to determine the pH value for weak acid. The measurement was done for 3 and times and the data was recorded after the pH values was no longer change against time. After that, further calculations are conducted to determine the value for Ka. According to the result obtained, the Ka values for both methods are close to the Ka value of the acetic acid. Thus, the weak acid use is considered as the acetic acid due to the nearest Ka value determined. However, the titration method shows that its Ka value is more nearest compared to the pH measurement method. The accuracy based on the Ka value also shows the higher percentage for titration method which is equal to 91.80 % while measurement method only represented 70.0 % accuracy. As a conclusion, the titration method is more accurate to determine the Ka value of the weak acid thus at the same time is very effective to determine the type of acid used according to the Ka value compared to the pH measurement method.

INTRODUCTION

The Acid Ionizations Constant, Ka, is the equilibrium constant for the reaction in which a weak acid is in equilibrium with its conjugate base in aqueous solution. Notice that in the equilibrium expression below the concentration of water is not included. This is because water is vastly in excess and the amount changes negligibly on equilibrium being established. Ka can be thought of as a modified equilibrium constant. For example,HA(aq) + H2O(l) = H3O(aq) + A(aq)

Ka = [H3O(aq)][A(aq)] / [HA(aq)]

Therefore, the larger the value of Ka, the stronger is the acid. The value is sometimes expressed as the logarithm of its reciprocal, called pKa. Therefore,pKa = -log Ka

The smaller the values of pKa, the acid become stronger. Ka is a better measure of the strength of an acid than pH because adding more water to the acid solution will not change the value of the equilibrium constant Ka, but it will change the H+ ion concentration on which pH depends. The current investigation set out to test the application of the above described method for determining the Ka of an unknown acid species. The current experiment was designed to employ the use of a pH meter incorporated into the titration of an aqueous solution of a measured amount of unknown acid HA. The data were graphed and analyzed to determine the equivalence point and, thereby, the half way volume and Ka of the unknown acid as described above.

OBJECTIVETo determine the acid ionizations constant, Ka, of weak acid by titration with sodium hydroxide (NaOH) and measuring the pH of the weak acid.

THEORY

All weak acids in solution exist in equilibrium. The equilibrium is between the molecular form of the acid and ionized form of the acid.HA(aq) + H2O(l) = H3O(aq) + A(aq)

For all weak acid this equilibrium lies predominantly on the left. Most of the acid is in molecular form. This result in all weak acids having an equilibrium constant that is less than 1. The largest equilibrium constant is approximately 102. The smallest is about 1013. The equilibrium expression for all monoprotic weak acid is:Ka = [H3O(aq)][A(aq)] / [HA(aq)]Where Ka is the acid ionization constant.Ka can be determined in a variety of ways. The most common is by measuring the pH of a solution of the weak acid. This methods work fine monoprotic acids, but for polyprotic acids, the result is a combination of the various Ka for each acidid protons. Titration is a method that works well for either monoprotic or polyprotic acids. Measuring the pH at various point in the titration and plotting the pH versus volume of the base added gives an indication of the Ka for the acid.Sample calculation:HAH3OA

Initial0.1 M00

Change-3.0x103 M+3.0x103 M+3.0x103 M

Final(0.1 - 3.0x103) M3.0x103 M3.0x103 M

Ka = [H3O(aq)][A(aq)] / [HA(aq)] = (3.0x103)( 3.0x103) / (0.1 - 3.0x103) = 9.3 x 105This is a fairly weak acid

APPARATUS

250 ml beaker Burette Pipette Magnetic barREAGENT: Sodium hydroxide 40 ml Unknown acid Distilled water

EXPERIMENTAL PROCEDURE

A. Determination of Ka value of a weak acid by titration with NaOH

1. First, approximately 40mL of unknown acid solution was obtained. Data was recorded.2. 10.0mL of the unknown acid solution was added to dry 250mL beaker.3. The solution was titrated with NaOH4. The reaction mixture was poured into lab sink, and was flushed with continuous water.5. Next, step 2-4 was repeated to perform 2titration of the weak acid with 0.1M NaOH solution6. The titration curve was draw using excels.7. From the titration curve, the volume of NaOH required to neutralizing the weak acid solution in each titration was determined. Data was recorded.8. The Ka value for titration 1 and 2 was calculated.9. The average Ka value for the weak acid was calculated.10. Using the Ka value calculated in step 10; the unknown weak acid from table 1 was identified. 11. Lastly, the molarity of the unknown acid solution for each titration was calculated.

B. Determination of Ka from the initial concentration and pH of a weak acid solution1. First, 20mL of the unknown weak acid solution (same unknown used in Part A) was added to a dry beaker.2. The pH electrode was removed from the pH 7 buffer solution. The electrode was rinsed with distilled water and the tip of the probe was dried.3. The electrode was inserted into the beaker containing the acid solution from step 1. The solution was stirred. After the pH reading was stabilized, the pH of the solution was recorded.4. The weak acid solution was decanted into waste container.5. Step 1-4 was repeated to perform second pH measurement6. The pH probe was rinsed with distilled water7. Using initial concentration (molarity calculated in step12, part A) the pH of the unknown acid solution, the Ka for each trial was calculated.8. Using the Ka value calculated in step 8, the weak acid was identified from table 1.

RESULTS

Volume of NaOH (mL)pH Titration 1pH Titration 2 pH Titration 3

03.723.753.69

13.903.933.93

24.294.204.16

34.324.324.30

44.434.484.42

54.634.634.55

64.804.774.70

74.954.944.75

85.155.145.02

95.505.435.21

106.096.296.06

1110.7710.7310.74

1211.1911.1111.13

1311.3311.2811.32

1411.4511.3911.42

1511.5311.5511.49

Titration 1

Titration 2

Titration 3

SAMPLE CALCULATIONS

Determination of the Ka value of weak acid by titration with NaOHMethod 1Titration 1Titration 2Titration 3

Equivalence point (mL)10.1910.0210.20

Half equivalence point (mL)5.105.015.10

pH at half equivalence point4.654.854.65

Titration 1Equivalence point = 10.19mLHalf equivalence point = 10.19/2 = 5.10mLThe pH vaue corresponding to 5.10mL is 4.65pKa = pH = 4.65 at half equivalence point.pKa = pHKa = 10-pKa=10-4.65= 2.24 x 10-5

Titration 2Equivalence point = 10.02mLHalf equivalence point = 10.02/2 = 5.01mLThe pH vaue corresponding to 5.01mL is 4.85pKa = pH = 4.85 at half equivalence point.pKa = pHKa = 10-pKa=10-4.85=1.41 x 10-5

Titration 3Equivalence point = 10.20mLHalf equivalence point = 10.20/2 = 5.10mLThe pH value corresponding to 5.10mL is 4.65pKa = pH = 4.65 at half equivalence point.pKa = pHKa = 10-pKa=10-4.65= 2.24 x 10-5

The average Ka value from titration 1, 2, and 3.

= The calculated Ka of the unknown acid is closest to that of acetic acid, .Titration 1No of moles of NaOH :

= CH3COOH(aq)+ NaOH(aq) NaCH3COO(aq)+ H2O(l)From the equation above 1 mole of CH3COOH was neutralized by 1 mole of NaOH.

= CH3COOH.Molarity of CH3COOH :

= 0.1 M

TITRATION 2No of moles of NaOH :


= CH3COOH.Molarity of CH3COOH :

= 0.1 MTITRATION 3No of moles of NaOH :


= CH3COOH.

Molarity of CH3COOH :

= 0.1 M

Average molarity value from titration 1, 2 and 3.

= 0.1 MDetermination of the Ka from the initial concentration and pH of weak acid solution Method 2Titration 1Titration 2Titration 3

Initial pH3.723.753.69

Titration 1The pH of 0.1 M weak acid solution is 3.72 at 25.pH = 3.72[H3O+] = ==1.91

HA + H2O H3O+ + A-Initial:0.10M00Change:-1.91 +1.91M+1.91MEquilibrium:(0.10-1.91)M 1.91M 1.91M

=3.72 Titration 2The pH of 0.1 M weak acid solution is 3.75 at 25.pH = 3.75[H3O+] = ==1.78


= 3.17

Titration 3The pH of 0.1 M weak acid solution is 3.69 at 25.pH = 3.69[H3O+] = ==2.04


= 4.17Average Ka value :

= 1.26 x The average Ka value of the unknown acid is nearest to that acetic acid of 1.8

The accuracy of the first method :

= 91.80%The accuracy of the second method :

= 70.00%

DISCUSSION

Through this experiment, the acid ionizations constant (Ka) of a weak acid is determined by the two methods. Firstly, is by the titration with sodium hydroxide solution and next is by measuring the pH value of that weak acid. For the first experiment by using titration method, Titration 1 shows the Ka value equal to 2.24 x 10-5, Titration 2 shows Ka value equal to 1.41 x 10-5 and Titration 3 shows the Ka value to be in 2.24 x 10-5. From that result, only Titration 1 and Titration 3 shows the same Ka value but Titration 2 shows the significant difference. From the observation, this difference value of Ka is due to that the initial pH value for the acid is not same to each other for all three titration at the beginning of the experiment. As a result, the equivalence point for these three titrations will be differing to each other thus affected the calculation to determined Ka value. Due to that thing, to obtain only one accurate Ka value to represent the titration method, it should be taken in average, which is equal to1.96 x 10-5. From the Table 1, this value is closest to the Ka value of acetic acid.For the second experiment in which the method is by measuring the pH value of the weak acid, the result for Ka values show the significant difference compared to the titration method. Although the initial pH value for the weak acid use is the same with the previous titration experiment, the Ka values determined is not same. According to the result obtained, the Ka value for initial pH of weak acid which respect to 3.72 is 3.72 x 10-5, pH value equal to 3.75 is respect to 3.17 x 10-7 and pH value equal to 3.69 is respect to 4.17 x 10-7 of Ka value. The average value also must be considered in this experiment since all the Ka values is not constant. For this second experiment, the average Ka value is equal to 1.26 x 10-5, thus closest also to the acetic acid.Because of that, the weak acid use can be considered as the acetic acid since both titration and pH calculation method shows the nearest value to this acid according to the Table 1. However, the titration method is more accurate and effective since its shows the nearest Ka value to the acetic acid. Besides, the accuracy calculated for the titration method base on the Ka value of acetic acid shows the larger percentage which is equal to 91.80% compared to the accuracy for the pH calculation method which is equal to 70.0 %. Specifically, the difference of the accuracy between these two methods is due to that in the pH calculation method, the molarity of weak acid is assumed to be fully decreases during the equilibrium point when ionization takes places and the molarity of conjugate ion will fully increasing at another hand.

This can be seen from the equation use in the pH calculation method:

HA + H2O H3O+ + A-Initial:0.10M00Change:-1.91 +1.91M + 1.91MEquilibrium:(0.10-1.91)M 1.91M 1.91M

Thus, the assumption for the fully changes of that molarity thus will affect the further calculation to determine the Ka value since in this method; Ka values are dependent to the changes of acid molarity. Oppositely for the titration method, the Ka values are determine according to the pH value of the half equivalence point occurs at the graph plotted. At this point, concentration of [HA] is equal to the concentration of conjugate base. Thats mean at this point, the titration is reached exactly base required to completely neutralize the acid. Thus, titration method is more practice compared to the pH calculation method which is more theoretically. Plus, during theoretically calculation (second method), it assume perfect condition without any loss. Thats why there is a difference for the Ka value obtained from both methods although the same condition is applied.

CONCLUSION

The experiment was mainly about finding the Ka value to obtain what acid is the unknown acid. By using titration with NaOH, the Ka value calculated for the unknown acid (titration 1) is 2.24 x , for titration 2 is 1.41 x and for titration 3 is 2.24 x . The average Ka value from this method is 1.96 x . The closest value of Ka of acid in table 1 is acetic acid that is 1.810. The second method is by using determination of Ka from the initial concentration and pH of a weak acid solution. The Ka value calculated for the unknown acid (titration 1) is 3.72 x , for titration 2 is 3.17 x and for the unknown acid for titration 3 is 4.17 x . The average Ka value from this method is 1.26 x. The closest value of Ka of acid in table 1 is acetic acid that the Ka is 1.810. From the experiment we could conclude that by using determination of Ka from titration with NaOH is more accurate than using method determination of Ka from the initial concentration and pH of a weak acid solution . This method is more accurate as it minimize the error occurs in the experiment.

RECOMMENDATIONS 1. The experiment should include the indicator as we could see the changes of the titration.2. Avoid parallax error while taking the reading of the burrette when do the titration. The eyes must parallel to the the level meniscus of the burrette.3. The air bubble in the burrette must be removed because to avoid the actual volume delivered will be less than the volume reported.4. Pipetting process has to be accurate in order to avoid excess addition of the titrating substances.5. The beaker, pipette and the flask should be washed properly with distilled water in order to avoid contaminate that will be effect the ph of solutions.

REFERENCES

Robert H.Perry, Don W.Green, Perrys Chemical Engineers Handbook, McGraw Hill,1998.

Engineering Chemistry lab. Dalwani, M., N. E. Benes, et al. (2011) Effect of pH on the performance of polyamide/polyacrylonitrile based thin film composite membranes. Journal of Membrane Science, 372(1-2): 228-238.

Nilsson, M., G. Tragardh, et al. (2008). The influence of pH, salt and temperature on nanofiltration performance. Journal of Membrane Science, 312(1-2): 97-106.

Su, M., D.-X. Wang, et al. (2006) Rejection of ions by NF membranes for binary electrolyte solutions of NaCl, NaNO3, CaCl2 and Ca(NO3)2. Desalination, 191(1- 3): 303-308.

Engineering Chemistry lab

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ka determination report

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