Sample Lab Report - Grade: A PDF

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Sample Lab Report...

Description

2P – Lee Determination of Sodium Carbonate Mass from Soda Ash: Back Titration I.

Purpose

The goal of this experimentation is to experience the procedure of back titration which is used in order to determine the quantity of a substance in a multiple step reaction. In the case of this lab, the quantity of pure sodium carbonate in crude soda ash can be found using back titration. Ultimately, the use of back titration allows for easier identification of the mass percent of sodium carbonate and allows for the use of aqueous solutions with known molar concentrations. With this, the calculation of the percent of sodium carbonate in soda ash is easier and more accurate. II.

Introduction

The term soda ash refers to the crude form of the compound sodium carbonate; therefore it is produced from mineral trona or sodium-carbonate-bearing brines or can be made through various manufacturing chemical reactions. This inorganic molecule is essential for the assembly of glass, detergents and other chemicals thus causing it to one of the leading chemicals in the US industrial industry. Due to soda ash being the impure version of sodium carbonate, the impurities are inert to acid allowing the titration with a strong acid. The following steps illustrate the reaction:

CO32− + H + ↔ HCO3− + H + ↔ H2CO3 ↔ H2O + CO2

Part 1

Part 2

Part 3

The technique of back titration aids in finding the amount of total carbonate present in the soda ash; while an indicator can be used to test this for the two step reaction the back titration method is easier since it only involves one indicator. This involves the adding of excess acid which reacts with the carbonate, and then the solution is boiled to get rid of the carbon dioxide and lastly cooled. The ultimate step is to titrate the left over acid with standardized NaOH. During the addition of the correct amount of acid, the carbonate ions will change into HCO3 ions. Furthermore the second addition of acid will cause the change of HCO3

−

−

ions to

H2CO3 . In order to ensure that excess acid is present, an extra 10 mL of acid is necessary. Ultimately the important aspects of back titration are that an excess of acid is present and that a known amount of acid is added. III.

Procedure

1. Obtain a sample of a dried sample of unknown soda ash from the oven and allow it to cool before measuring the mass. The drying of the sample allows for a more accurate mass due to the moisture from the environment causing a hydrate to form with the soda ash. 2. Next, measure out two samples of approximately 0.4 g from the unknown soda ash using the weigh boats provided. Add each sample to two different 250 mL Erlenmeyer flasks

with labels of “Sample 1” and “Sample 2”. Make sure to record the mass of the two samples to four decimal places. 3. Then add 50 mL of distilled water and dissolve the soda ash in first flask with the label “Sample 1” by swirling. Afterwards, add three drops of the indicator phenolpthalein. Repeat this for the second flask. 4. Set up the apparatus for the titrations, there should be two burets on one ring stand(refer to Figure 1 for more detail). Using a buret rinsed with HCl , add more standardized HCl solution of .1 M to the buret and titrate both flasks until the pink color in the flask is not present and the solution is clear.

Figure 1: Buret Apparatus 5. Add approximately the same amount of HCl added in Step 4 in order to make the solution clear again to both flasks. Next, add an extra 10 mL of standardized HCl in order to ensure presence of excess acid. 6. After titrating, allow both flasks to boil for about five minutes; once the boiling process is done let the flasks cool to room temperature and then add three drops of indicator chemical once more to both. 7. In another buret, rinse with NaOH and fill the buret with the same .1 M standardized NaOH solution. Now titrate the cooled samples with the NaOH until the endpoint is reached in which a faint constant pink color is present (Refer to Figure 2 for correct pink colors).

Figure 2: From left to right: first shade is clear for how the color should be after HCl, second shade is the faint pink color after back titrating with NaOH, third shade is too dark of a pink.

IV.

Data

Table 1: Mass and Volume for the Back Titration of Soda Ash (Raw Data) Molarity of HCl

0.09935 M

Molarity of NaOH

0.09990 M Trial 1

Trial 2

Sample Mass(±.0001 g)

.4242

.4322

Initial HCl Volume(±.01 mL)

9.94

9.92

Final HCl Volume(±.01 mL)

27.63

26.93

Total HCl Vol. (±.01 mL)

17.69(43.38 with excess)

44.02

Initial NaOH Volume(±.01 9.83 mL)

9.87

Final NaOH Volume(±.01 11.12 mL)

11.11

Total NaOH Vol. (±.01 mL)

1.24

1.29

Calculations A. Equations Used:

● VolumeT itrant (Mo ●Netmm olHCl

= mm

1mm olNa2CO3 ) = m molNa2CO3titrated 2mmolHCl 105.3mg ) = gNa2CO3 ● (mmolNaCO )( 1mmol MassofNa2CO3 ● Sample MassofSod aAsh (100%) = %Na 2CO3 ●

(Netmm olHCl )(

B. Sample 1-

●

43.38m L( . 09935m M ) = 4.310mmolHCl

●

1.29m L( . 09990m M ) = . 1289mmolNaOH

●

4.181Netmm olHCl = 4.310m molHCl − . 1289m molNaOH

●

(4.181mmolHCl )(

●

(2.0905mmolNaCO)(

1mm olNa2CO3 ) = 2.0905m molNa2CO3titrated 2mmolHCl 105.3mg ) = . 2201gNa2CO3 1 mmol

. 2201g (100%) = 51.89%Na 2CO3 . 4242g C. Sample 2-

●

44.02m L( . 09935m M ) = 4.3733mmolHCl

●

1.24m L( . 09990m M ) = . 1244mmolNaOH

●

4.2490Netmm olHCl = 4.3733m molHCl − . 1244m molNaOH

1mm olNa2CO3 ) = 2.1245m molNa2CO3titrated (4.249mmol  HCl )( ● 2mmolHCl 105.3mg (2.0905 mmol  NaCO)( ) = . 2237gNa2CO3 ● 1 mmol

. 2237g (100%) = 51.76%Na 2CO3 . 4242g

Average = 51.83%Na 2CO3

V.

Discussion and Conclusion

The technique of back titration provided numerous quantitative data which aided in finding the overall pure mass percentage of sodium carbonate in soda ash. When comparing the mass percentages of the two samples, a very small difference of .13% is found. This difference is very insignificant therefore leading to precise data. The average of 51.83% illustrates how the majority of the sample is pure sodium carbonate; this could be due to the specific sample provided or due to other experimental error. To start off one possible experimental error could have occurred when finding the mass of the two samples with unknown concentrations of sodium carbonate. Due to the top of the scale missing, the mass kept fluctuating because of the air. This could have lowered or increased the percentage of the sodium carbonate in the sample depending on the actual mass of the sample; but it would only have a small effect on the percentage due to the fluctuations only occurring in the ±.001 region. Another possibility of error could be because of not allowing the titrated solutions with excess acid to boil properly. Even though the flasks were emitting steam, the production of bubbles was limited. The bubbles signify the release of carbon dioxide, therefore the mass of the sodium carbonate found could be less than the actual mass due to the dissociation not being complete. Another observation which could be a possibility of error is the back titration using the NaOH. During the process, only a few milliliters of standardized base were necessary in order to produce a pink hue in the solution. Such could have occurred due to not enough excess acid present therefore leading to the possibility of a higher concentration of sodium carbonate found since the volume of NaOH used was small. While there were numerous possible sources of error, the lab in its entirety

was done with caution in order to prevent contamination or deviation from precision and accuracy....