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Titration with a primary standard.

Extracts from this document...

Introduction

TITRATION WITH A PRIMARY STANDARD 1. Wear all the protective attire required and make sure that all the equipment is collected and is thoroughly cleaned with distilled water leaving it dirt free, so that the chemicals being used will not be contaminated during the experiment. 2. Determine the mass of the weighing boat and then weigh out 0.265g g of dry Na2CO3. Record actual mass used to nearest 0.01 g. Transfer the solid to a 250 ml Erlenmeyer flask, rinse the weigh paper (or weighing bottle) into the flask with a small amount of distilled water and dissolve the solid with 25 ml of distilled water. Using a stirring rod, stir the solid and water in the beaker to dissolve, adding more water is necessary. 3. Transfer carefully the solution to the 250cm3 volumetric flask (pouring the solution down the funnel to avoid spillage). Rinse the beaker three times to make sure all the solution goes into the volumetric flask, each time pouring the solution down the stirring rod to rinse it. 4. Carefully make up the solution to about 1cm of the mark on the neck of the flask using distilled water.Insert the stopper and shake to mix the contents. 5. Using a dropping pipette, add enough distilled water to bring the bottom of the meniscus on the mark. Now, mix it thoroughly, by turning the volumetric flask upside down twice, to ensure complete mixing 6. Clean and check the flow rate of the burette . Setting and Determination of the flow rate: Put a 25 ml graduated cylinder underneath the end of the tubing. Turn on the air pump, and collect a certain volume (e.g., 20 ml) of the titrant in the cylinder. Measure the required time (t). Calculate the flow rate (F) as follows: F= V(ml) / t (s) A flow rate of about 1-3 ml/min (0.0166-0.05 ml/sec) is appropriate Do not change the settings once you have measured the flow rate. ...read more.

Middle

Stop once the minimum is reached. Place the white tile under the flask to better observe any colour changes. 8. Start to add the NaOH solution drop by drop, swirling after each drop. 9. When a hint of pink appears, swirl the flask well. The colour may disappear. If it doesn't, the endpoint is reached. Record this volume. If the colour fades, proceed to step 10. 10. Add 1/2 drop to the flask. This is achieved by opening the stopcock until only part of a drop is hanging from the tip of the burette. Touch the tip to the inside of the flask. Rinse the drop into the flask with some distilled water and swirl well. Repeat if the solution is colourless. Once the endpoint is reached record your final volume. NOTE: If the colour of the solution is a strong bright pink, you have passed the endpoint and your final volume is not accurate. 11. Repeat steps 7-10 for the other two acid samples. Be sure your initial volume of NaOH is always over 30 ml before you start. (It need not be near 0 ml. as this wastes the solution) Flush the contents of the beaker down the drain with lots of water. Rinse the burette thoroughly with tap water (3 times) then 3 times with distilled water. Verify that the final rinse is neutral by placing a drop from the burette on red litmus paper. If it is still basic, the red litmus will turn blue. Making a solution from solid base. mol NaOH = MNaOHVNaOH g NaOH = mol NaOH x MolMass NaOH The mass of NaOH needed to make 500.0 mL of 0.2000 M NaOH g NaOH = (0.2000M NaOH x 0.5000 L) x 40.00 g NaOH/mol NaOH = 4.000 g NaOH Making a solution from a more concentrated base solution. Vinitial = (MfinalVfinal) Minitial The volume of 9.0 M NaOH needed to make 50.00 mL of 2.5 M NaOH VNaOH = (2.5 M NaOH x 0.05000 L) ...read more.

Conclusion

Solution Steps: 1. Write a balanced equation: 2NaOH + H2SO4 Na2SO4 + 2H2O 2. Determine the number of moles of the standard NaOH solution used: 3. Use the mole ratio from the balanced equation to convert moles of NaOH to moles of H2SO4: 4. .Use the volume of acid solution used to determine the molarity of the acid solution: 5. Notice that the 1dm3/1000cm3 and the 1000cm3/1dm3 will offset each other. One may shorten the problem by skipping these conversions EXPERIMENTAL ERROR: In order to calculate your experimental error for each of your reactions, use the equation below. The theoretical value is the heat of reaction, for your acid, per mole of water formed, shown in Table 1. The experimental value is the actual heat of reaction you determined in each of your experiments (per mole of water formed). To obtain %Error for a measurement, you need to know the Theoretical and Experimental values for your measurement, or set of calculations. Use this formula to obtain the %Error, which is always a positve (absolute) value. FURTHER STUDY: 1. Repeat procedure with the concentrations of H2SO4 and NaOH reduced 10 and 100 times. Study the effect of concentration on the pH change. 2. Plot the first derivative (dpH/dt) against time of all the previous experiments and locate the end point in each case. The end point in the first derivative curves are defined as the point at the maximum value. The first derivative curves can be obtained by graphing programs such as MicroCal Origin. SOURCES CONSULTED: * Rendle, Vokins and Davis, 1991. Experimental Chemistry, Second Edition. London: Arnold. * Ottewill and Walsh, March 1996. 'Electrochemical Cell' and 'How to use electrochemical cells' . Chemistry Review Vol:4 Num:4. * Atkins, 1990. Physical Chemistry, Fourth Edition. Oxford: OUP. * Fine and Beall. Chemistry for Engineers and Scientist.. USA: Saunders College Publishing. * Salter Advance Chemistry Course 1994. 'Redox', 'Redox reactions and electrode potentials' in 'Chemical Ideas'. Oxford : Heinemann Educational Jeffery, Bassett, Mendham, Denney: Vogel's Textbook of Quantitative Chemical Analysis, 5th e ...read more.

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