Aspirin(conducto) Production

EXPERIMENT NO.5 Abstract Determine the amount of salicylic acid in acetyl salicylic acid, trade named as Aspirin. For this we would know the actual amount of salicylic used for each tablet of Aspirin. The following are the reagents & requirements for the experiment; standard solution of NaOH & HCl, solution of Aspirin, burette, pipette, conductometer & volumetric flasks. Here we used conductometric titration between NaOH & Aspirin solution, where NaOH solution is standardized by HCl of 0.1M.

Theory Conductometry [1] is the technique by which quantitative analysis has been made for analysis of different species using the electrical conductance between the two electrodes. This basically based on ability of conductances of different solutions by their respective ions in the solvents. The volume of titration which is required to react completely with the analyzed substance in the titrand is called Equivalence Point. The measure of ability of a solution to conduct electricity is called its conductance[2]. Water is a very poor conductor of electricity. The presence of ionic species (electrolytes) in water increases the conductance considerably. Solutions of electrolytes, like metallic conductors, obey Ohm's law. Thus, the current I passing through a solution of a particular electrolyte is proportional to the applied potential difference V as given by the equation:

I=

V R

Here R is the resistance offered by the solution in ohms (W)

R=

V I

R=

ρ × A

Or The resistance is directly proportional to the length  , of liquid through which the current passes, and inversely proportional to its area of cross-section A. It therefore follows that:

Here r is the constant of proportionality and is called the specific resistivity[2]. It is a constant for an aqueous solution of a given electrolyte of fixed concentration at a particular temperature. In SI units, resistivity has the units of ohm meter (W m). If area of cross-section is increased, more ions will be able to flow and resistance will decrease. The conductance L is defined as the reciprocal of the resistance expressed in units of W –1 or Siemens (S).

L=

1 R

Thus, the conductance of a homogenous body of uniform cross-section is proportional to the cross-sectional area A and inversely proportional to the length  , given by:

L=

A A =κ × ρ × 

Here the reciprocal of resistivity r is replaced by another constantk, which is called the specific conductivity with units W –1m –1 or S m –1 which can be expressed as:

κ=

 1 k = = RA ρ R

  ∴ k =  A 

Here k is called the cell constant. [2] The conductivity of a solution of an electrolyte increases as the temperature is raised. The conductivity of a particular electrolyte at a fixed temperature depends upon: a) the number of ions present in unit volume of solution b) the speed at which the ions oscillate in step with the applied alternating voltage c) ionizability of the electrolyte [2] Conductivity does not solely depend upon the number of ions present in unit volume of the solution. Concentrated solutions of very soluble electrolytes show a maximum conductivity value which then decreases on further concentration of the solution. The reason lies in the fact that ionic distances becomes less as concentration increases. This results in a greater attraction between positive and negative ions, and with a greater reduction in ionic speeds. The phenomenon is called "ionic interference". [2] As the concentration of weak electrolyte is increased, there is an increase in the number of molecules, but a small proportion of them are ionized. By increasing the concentration of a weak electrolyte, the conductivity reaches a maximum value followed by a steady decrease. This is the same effect as was observed for strong electrolyte but in this case, ionic interference is not responsible as a weak electrolyte produces very few ions and few ions can not interfere to a considerable extent.[2] The addition of an electrolyte to a solution of another electrolyte under conditions producing no appreciable change in volume will affect the conductance of the solution according to whether or not ionic reactions occurs, the conductometric conductance may either increase or decrease; thus in addition of base to a strong acid, the conductance decreases owing to the replacement of hydrogen of high conductivity by another cation of lower conductivity. This is the principle underlying conductometric titrations; i.e. the substitution of ions of one conductivity by ions of another conductivity.[1] Aspirin, or acetylsalicylic acid is a salicylate drug often used as an analgesic (to relieve minor aches and pains), antipyretic (to reduce fever), and as an anti-inflammatory. It also has an anti-clotting effect and is used in long-term, low doses to prevent heart attacks and blood clot formation in people at high risk for developing blood clots. High doses of aspirin may also be given immediately after an acute heart attack; these doses may inhibit the synthesis of prothrombin and therefore produce a second and different anticoagulant effect, although this is not well understood.

Aspirin is one of the most frequently used drugs in the treatment of mild to moderate pain, including that of migraines, and fever. It is often combined with other analgesics, even though this has never been shown to be more effective or less toxic than aspirin alone. Aspirin has been used in addition to other non-steroidal anti-inflammatory drugs and opioid analgesics in the treatment of pain associated with cancer. Aspirin can cause blood loss. It can also cause heartburn.

Results Weight of Aspirin

·

=

0.5293 gm

Standardization of NaOH Against 0.1M HCl Serial No. 1 2 3

Initial Volume of NaOH (ml) 0 4.5 8.9

Final Volume of NaOH (ml) 4.5 8.9 13.3

Difference (ml) 4.5 4.4 4.4

Titration Reaction

NaOH + HCl ® NaCl + H2O Concentration of NaOH Concentration of HCl Volume of HCl solution Volume of NaOH solution Concentration of NaOH

= =

M1 = V2 =

= 0.1 M V1 = 5 ml = 4.4 ml (from table) M2 =?

Formula M 1V1 M 2V = n1 n2 Here

n1 = n2 = 1 M2 =

·

Þ

M2 =

M 1V1 n2 V2 n1

5 × 0.1 × 1 = 0.1136M 1 × 4.4

Conductometric Titration Between NaOH And Aspirin Cell constant =

0.01 cm-1

→Addition of 3 ml of NaOH into Aspirin Solution Volume of NaOH added (ml) 0 3 6 9 12 15 18 21 24 27 30 33 36

Conductance (µS)

Specific Conductance (µS/cm)

4.59 15.20 20.90 54.80 30.80 35.30 41.20 48.10 52.23 57.19 60.09 63.34 59.30

0.0459 0.1520 0.2090 0.2480 0.3080 0.3530 0.4120 0.4810 0.5223 0.5719 0.6009 0.6334 0.5930

End Point Volume of NaOH consumed concluded by graph

=

2.5 ml

→Addition of 2 ml of NaOH into Aspirin Solution Volume of NaOH added (ml) 0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 End Point

Conductance (µS)

Specific Conductance (µS/cm)

4.81 6.20 18.20 22.24 24.31 29.90 34.60 36.50 38.29 39.67 40.23 43.80 43.80 43.80 43.80 43.80

0.0481 0.0620 0.1820 0.2224 0.2431 0.2990 0.3460 0.3650 0.3829 0.3967 0.4023 0.4380 0.4380 0.4380 0.4380 0.4380

Volume of NaOH consumed concluded by graph

=

2.0 ml

→Addition of 1 ml of NaOH into Aspirin Solution Volume of NaOH added (ml) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26

Conductance (µS)

Specific Conductance (µS/cm)

4.97 5.09 20.35 22.12 24.65 27.79 29.30 32.33 35.60 37.99 40.10 40.55 41.77

0.0497 0.0509 0.2035 0.2212 0.2465 0.2779 0.2930 0.3233 0.3560 0.3799 0.4010 0.4055

45.36 48.29 50.67

0.4177 0.4536 0.4829 0.5067

54.72 59.80 61.33 65.44 59.50 64.95 68.60 67.66 67.66 67.95 67.95

0.5472 0.5980 0.6133 0.6544 0.5950 0.6495 0.6860 0.6766 0.6766 0.6795 0.6795

End Point Volume of NaOH consumed concluded by graph

=

1.5 ml

→Addition of 0.5 ml of NaOH into Aspirin Solution Volume of NaOH added (ml) 0 0.5

Conductance (µS) 4.91 5.10

Specific Conductance (µS/cm) 0.0491 0.0510

1.0 5.44 1.5 5.38 2.0 15.36 2.5 16.10 3.0 17.50 3.5 18.70 4.0 21.90 4.5 25.70 5.0 26.80 5.5 27.80 6.0 28.65 6.5 29.95 7.0 31.45 7.5 32.02 8.0 33.59 8.5 34.72 9.0 35.51 9.5 36.44 10.0 38.20 10.5 39.65 11.0 40.09 11.5 41.22 12.0 42.25 12.5 43.54 13.0 44.70 13.5 45.70 14.0 45.49 14.5 46.29 15.0 47.01 15.5 47.53 16.0 48.73 16.5 49.21 17.0 50.01 17.5 51.22 18.0 51.65 18.5 52.10 19.0 53.72 19.5 53.72 20.0 53.98 End Point The end point from the graph = V4 = 1.5 ml

·

Average Volume required of NaOH

0.0544 0.0539 0.1536 0.1610 0.1750 0.1870 0.2190 0.2570 0.2680 0.2780 0.2865 0.2995 0.3145 0.3202 0.3359 0.3472 0.3551 0.3644 0.3820 0.3965 0.4009 0.4122 0.4225 0.4354 0.4470 0.4570 0.4549 0.4629 0.4701 0.4753 0.4873 0.4921 0.5001 0.5122 0.5165 0.5210 0.5372 0.5372 0.5398

V =

·

V1 + V2 + V3 + V4 2.5 + 2.0 + 1.5 + 1.5 = = 1.875ml 4 4

Concentration of NaOH while titrating with Aspirin Solution Concentration of NaOH Volume of NaOH Concentration of Aspirin Volume of Aspirin

= = = =

M1 V1 M2 V2

= 0.1136 M = 1.875 ml =? = 30 ml

Titration Reaction

H

O

N

O O C C

C

3H

a O

OO C

+ NaOH ¾¾¾®

C C

3H

+ H2O

Sodiumoacetoxy benzoic acid

Aspirin Formula M 1V1 M 2V = n1 n2 Here

n1 = n2 = 1 M2 =

·

Þ

M2 =

M 1V1n2 V2 n1

0.1136 × 1.875 × 1 = 7.11 × 10 −3 M 1 × 30

Amount of Aspirin (Acetyl Salicylic Acid) in the Given Tablet Concentration of aspirin is found to be 7.11×10-3 M. Therefore, M × Mol.Wt. × V 7.11 × 10 −3 × 164.1604 × 250 = 1000 1000 Amount= = 0.29138 gm The amount of Aspirin is found to be 291.38 mg.

·

Percentage of Aspirin 0.29138 ×100 = 55.05% % weight = 0.5293

·

Percentage Error

The standard value of the amount of aspirin estimated on the label for a single tablet is 300mg. 300 − 291.38 × 100 = 2.873% 300 % error = Hence, The amount of ascorbic acid is found to be 291.38mg in a single given tablet with an error of 2.873%.

Discussion

[1]

The variation in the conductivity of a solution as the concentration of the solution is varied is best represented by plotting conductivity values against the dilution of the solution. The dilution of a solution is the reciprocal of its concentration, and is generally defined as the number of liters of solution containing 1 gm. equivalent of solute. Some substances, such as mineral acids, salt solutions, have much higher specific conductivities than others, such as acetic acid & alcohols. The difference is shown by plotting the results for each type of substance on conductivity-dilution graphs. Those substances with high specific conductivities are known as strong electrolytes; they include most mineral acids, alkalis and most salts. Substances with comparatively small specific conductivities are known as weak electrolytes; they include most organic bases and acids. As a solution gets more and more dilute there are fewer and fewer solute particles in a given volume of the solution. Because it is the solute that causes electrical conductivity (pure water, and other solvents are very poor conductors), on this account, be expected that conductivity would fall with increasing dilution.

a) For Weak Electrolyte For a weak electrolyte, the degree of ionization increases with increasing dilution. The maximum in the conductivity-dilution curves is due to the composite effect of decrease in the total number of solute particles and increase in the degree of ionization. As dilution increases, there are fewer molecules, but a higher proportion of them split up into ions.

b)

For Strong Electrolyte

For a strong electrolyte the other factor is the decrease in ionic interference as the dilution increases. A strong electrolyte is almost fully ionized at all dilutions. At low dilution, the ions will interfere with each other so that their freedom of movement and their speed is restricted. As a solution is diluted, the ionic interference decreases and the ionic speed increases. For a strong electrolyte, increasing dilution gives fewer ions, in a fixed volume of solution, but the ions that are present interfere less with each other, which gives higher ionic speeds. As mentioned earlier the purity of water is critical for the accuracy of our results. Ordinary distilled water is not preferred. Specially pure water, known as conductance water or conductivity water should be used for the preparation of solutions. The purest water referred to as "ultra-pure" has a specific conductance of 0.05 mS / cm at 18°C. Air

must be rigidly excluded when it is being employed. The water we utilized was not pure. It is most probable that it contained impurities and stray electrolytes, which could cause variations in conductance by two ways: a. Contribute the effects of their conductance directly b. Cause ionic interference [2]

Conclusion We conclude that the conductometric method is the best way to determine the equivalence point, if there is a change in conductance is present in titrand solution. Conductomertric titrations are useful for acid-base, precipitation, and complexation titrations.

. 1 R.D. 2

BRAUN, "Introduction to Chemical Analysis" Chemical Process Principles

→Addition of 3 ml of NaOH into Aspirin Solution V olume v/s S p. C onductance

Specific Conductance (us/cm)

0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 0

5

10

15

20

25

30

V olum e of Na OH a dde d (m l)

→Addition of 2 ml of NaOH into Aspirin Solution

35

40

Volume v/s S p.C onductance 0.5 Specific Conductance (us/cm)

0.45 0.4 0.35 0.3 0.25 0.2 0.15 0.1 0.05 0 0

5

10

15

20

25

30

35

Volum e of Na OH a dde d (m l)

→Addition of 1 ml of NaOH into Aspirin Solution Volume v/s Sp.Conductance

Specific Conductance (us/cm)

0.8 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 0

5

10

15

20

Volum e of NaOH adde d (m l)

25

30

→Addition of 0.5 ml of NaOH into Aspirin Solution V o lu m e v /s S p .C o n d u c ta n c e

Specific Conductance (us/cm)

0.6 0.5 0.4 0.3 0.2 0.1 0 0

5

10

15

V o lu m e o f N a O H a d d e d (m l)

20

25