Titration with a primary standard.

To pipet 25.00 ml (note the significant figures) of sodium carbonate, you will use a 25 ml TD pipet and a rubber bulb. If the pipet does not fill above the mark, remove the rubber bulb squeeze it and reattach to continuing the sucking up of the liquid. During this process you finger should be held tightly on the top of the pipet to prevent loss of the liquid already in the pipet. When the pipet is filled above the mark, remove the bulb and put your thumb tightly over the top of the pipet to prevent loss of the liquid in the pipet. 

Release the pressure of this finger slightly and let the contents of the pipet drain into a "slop beaker" until the bottom of the meniscus is just above the mark. Reapply pressure to the top to stop the outflow of the liquid. Touch the tip to a clean side of the slop bucket to remove any partial drop from the tip. Place the tip of the pipet in a clean Erlenmeyer flask and release your finger from the top.

Let the pipet drain by gravity, 10-20 seconds after the pipet is empty, touch it to the side of the Erlenmeyer flask to draw off the last drop. In the process of delivering the exact volume of liquid to the flask using this kind of pipet, some pipets are designed to have the last amount of liquid blown out or even washed out, but the pipet you are using is a to deliver (TD) pipet, and is designed to empty by gravity.. 

Before starting the titration makes sure that the white tile that your flask containing your 25cm3 sodium carbonate solution, will lye on during the titration is clean so that you can clearly see when the sodium carbonate solution has changed its colour.

8. Begin titrating the sodium carbonate solution with the sulphuric acid solution by adding small increments of the titrant from the burette. Swirl the reaction mixture between additions. Use smaller increments of H2SO4 at each time so as not to surpass the precise endpoint. The acid-base indicator, methyl orange, will eventually turn from red, which indicates an excess of Na2CO3, to yellow, which indicates an excess of H2SO4. The endpoint will be when the solution remains orange, which indicates that the stoichiometric quantity of H2SO4 has been added.
9. Once the solution is green, bring it to a gentle boil on a hot plate to expel all dissolved carbon dioxide. Remove it from the hot plate to a wire gauze, and rinse the inside surfaces with a small amount of distilled water. If the solution stays green, the titration is complete. Record this volume as the final volume of titrant. If it turns blue, allow it to cool, and complete the titration by adding a few drops more of titrant. If the solution is yellow, I would have overshot the endpoint and must repeat the titration.

Step 8 is necessary because one of the products of the titration is carbon dioxide. The carbon dioxide will react with water to form carbonic acid which can cause the reaction mixture to become acidic prior to the addition of enough H2SO4 to neutralize the Na2CO3. If this happens, the indicator will change color early, and one will have a "premature endpoint".

• Flush the neutralized solution down the drain with lots of water. Repeat the titration until you have 3 titrations whose molarities are within ±.0006 of the average. At the end of each titration, record the value that is shown on the burette.  When starting a new titration make sure that the volumetric flask that contained the sodium carbonate solution is fully rinsed out so that it does not mix with the new batch of sodium carbonate solution and therefore affect the result.

11. 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 blue litmus paper. If it is still acidic, the blue litmus will turn pink.

For an even more accurate result, one could carry out another titration using a secondary standard. For this, I would use NaOH.

TITRATION WITH A SECONDARY STANDARD

METHOD

NOTE:  Before using glassware, inspect it for cleanliness.  If it is cloudy, (or if water 'beads' in the buret) rinse it with a small volume of rinsing acid.  Once clear (or if already clear), rinse the glassware with tap water.  Then rinse with a small volume of distilled water.

1. Prepare 0.10 M H2SO4 stock, using a clean dry beaker, a 20 ml transfer pipette, and a 250 ml volumetric flask. Rinse the pipette with a small volume of the stock acid before transferring the stock solution to the flask. Immediately after transferring, rinse the pipette with tap water, then distilled water. Carefully pour some of the solution to a clean, dry beaker.
2. Rinse the transfer pipette with a small volume of the prepare 0.1 M H2SO4 solution.  Transfer 20 ml of the solution into a rinsed Erlenmeyer flask.  Add two drops of phenolphthalein solution.  Repeat this step with two more flasks.
3. Using a clean dry beaker, obtain about 70 ml of the unknown NaOH solution.
4. Transfer a small volume of the NaOH solution to the buret (already rinsed with water and distilled water) and rinse the buret wall with the solution.  Drain out through open stopcock.
5. Set up the buret onto the stand.  Using a funnel, add the NaOH solution to the buret, filling it under the 0.00 ml mark, near the centre point.
6. Allow some of the solution to drain from the burette until the solution level is slightly below the 0.00 ml mark.  Make sure any air bubbles are expelled from the beneath the stopcock.  Record the initial volume.
7. Slowly allow the NaOH solution to drain into the flask (prepared in step 2), swirling the flask.  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.

The mass of NaOH needed to make 500.0 mL of 0.2000 M NaOH

Making a solution from a more concentrated base solution.

The volume of 9.0 M NaOH needed to make 50.00 mL of 2.5 M NaOH

The concentration of NaOH, if 30.00 mL of 0.1500 M H2SO4 neutralizes 20.00 mL of NaOH solution according to 2NaOH(aq) + H2SO4(aq) ----> Na2SO4(aq) + 2H2O(l).

The molar mass of a solid acid determined from the mass of acid used and the volume of known concentration base solution required to neutralize that sample of acid. In this example titration, the acid and base react in this titration in a 1:1 mole ratio (monoprotic acid).

mol base = MbaseVbase

The molar mass of solid acid if 1.100 g of acid is neutralized by 26.10 mL of 0.2100 M NaOH. The acid and NaOH react in a 1:1 mole ratio.

RESULTS AND CALCULATIONS: 

I will record the data from the experiment in the tables below:

Table 3

Table 4

Calculate the concentration of the stock H2SO4  solution for each titration and report an average molarity using the following equation:

 

M1V1 = M2V2

If the first titre (phenolphthalein), x cm3, represents the hydroxide plus half the carbonate, and the difference between the phenolphthalein and methyl orange titres, y cm3, represents half the carbonate, then the hydroxide is equivalent to (x – y) cm3 of acid, and the carbonate to 2y cm3 of acid.

This will help in finding the concentration of each of the compounds in mol dm-3.

Percentage  Error of some equipment.

Uncertainty measurements and percentage errors of some of the  equipment I will use, are given below:

Table 5

Preliminary investigations:

I will use laboratory trials, to find out if I need to modify any part of my plan.

It is well worth spending some time trying things out on a test tube scale. This will give me an idea of the way the chemicals behave and how quickly they react, and the type of gases evolved. This type of preliminary work will help me to decide on the key variables and make a better plan.
I will then be confident enough to decide on what measurements I will be making and how many I will make.
 
Calculations – quantities 
Uses theory to justify procedures and scale of working
 
A good starting point is to use the equation for the reaction and to do some preliminary calculations. Suppose you want to collect a gas and can measure up to 100 cm3 – then you can calculate roughly how much of the reagents you need to produce about the right amount of gas.
 
Measuring instruments 

I have decided on the apparatus, which I feel will make accurate measurements which are critical to my results.

Every time I take measurements in an experiment I am limited by the accuracy of the measuring instruments. It is possible to calculate the limits of accuracy of different measuring instruments

Errors which are not quantifiable also occur in experiments. I cannot calculate their exact effect, but will consider them as possible causes of anomalies

Measuring Volumes:
One can measure the volume of a liquid with a measuring cylinder, pipette or burette.
 
Measuring cylinders 
Measuring cylinders are least accurate. This may not matter if another measurement is even less accurate, or if I  want excess of a liquid. If the cylinder has graduations every 1 cm3, then when I measure 10 cm3 I can be sure I have more than 9.5 cm3 but less than 10.5 cm3. In this case the error is ±0.5 cm3 in 10 cm3, and the percentage error is 0.5/10 x 100 = 5%.
If I had measured 50 cm3 with the same measuring cylinder, the error would have been 0.5/50 x 100 = 1% so the bigger the reading the smaller the percentage error.

If I use a big measuring cylinder the graduations may be every 2 cm3. If I measured 50 cm3 with one of these, I could be sure that I had more than 49 cm3 but less than 51 cm3, so the  error would be ±1 cm3 in my reading.
 
Pipettes 
Pipettes are more accurate than measuring cylinders. Most school pipettes are made to an accuracy of one drop when they are used correctly.The volume of one drop = 0.05 cm3. A 10 cm3 pipette has an error of ±0.05 cm3 . In 10 cm3, the percentage error is 0.05/10 x 100 = 0.5%.
 
Burettes 
Burettes are also more accurate than measuring cylinders. They have graduations every 0.1 cm3, so when I  take a reading it should not be more than 0.05 cm3 too high or too low.
 
When you use a burette you, however, take a reading at the start and the end , so you have two errors of 0.05 cm3 i.e. total error = 0.10 cm3. If you are using your burette to do a titration there may be another error of one or two drops which is due to your judgement of when the indicator changes colour. This means that in a titration (as opposed to just using a burette to measure a volume) you may have an error of 0.2 cm3.
 
Measuring Mass: 
 
Balances 
Very accurate balances read to 0.001 g.This means a reading of 1.000 g is more than 0.9995 g but less than 1.0005 g.The percentage error in a reading of 1.000 g is 0.0005/1.000 x 100 = 0.05%.
 
Balances may also read to 2 decimal places i.e. they have readings every 0.01 g.This means a reading of 1.00 g is more than 0.995 g but less than 1.005 g.The percentage error in a reading of 1.00 g is 0.005/1.00 x 100 = 0.5%.
 
Calculating % error:

Calculating the percentage uncertainty (often called percentage error) ..

To calculate the error in using different measuring instruments,

• Look at the graduations on the measuring scale.
• Decide what reading you can be certain about.
• Usually the error is half the size of the graduation.
   
  If your judgement is involved, for example when doing a titration or measuring a time, then you can only estimate the error.
   
  If there are several errors in measurement the final answer I reach, will be affected by all of the errors. As a rough guide I can assume that the total error is the sum of each of the individual percentage errors.
   
  To improve the accuracy of my results, I need to reduce the size of the biggest error. For example, if you have a small error in measuring mass, but a big error in reading temperature, then measure temperatures more accurately

Calculate the volume of NaOH added.  (Final Volume - Initial Volume) for each trial.

Calculate the NaOH concentration from the results of each trial.  Average these results

PROCESSING THE DATA

Part I Titration with a Primary Standard:

For each titration, calculate the molarity of the sulphuric acid solution using the volume of H2SO4 needed to achieve the endpoint of the titration, the mass of the sodium carbonate and a balanced equation for the reaction, as shown in the example below. Determine the average and verify that you have three results that are within .0006 of your average. One must not forget to propagate the error through the calculations.

                         H2SO4   +   Na2CO3   →   Na2SO4   +  CO2  +  H2O

Ratio in moles       1          :          1      →          1         :      1     :    1

                            Moles = mass    = volume × concentration.

                                           MR

Table 8

Since moles = mass    =  volume          × concentration

                        MR

                    = mass     =   25                ×  0.1M

                        106         1000   dm3 

                       mass     =  0.025 ×  106

                                   =   2.65g

Therefore, the mass of Na2CO3 required is 2.65g.

Part II Titration with a Secondary Standard:

Calculate the molarity of the sulphuric acid, using the volume of NaOH needed to reach the endpoint, the molarity of the standard NaOH solution, the volume of H2SO4 pipetted and a balanced equation for the reaction. One must not forget to propagate the error through the calculations.

Compare the molarity determined using a primary standard and the molarity determined using a secondary standard. State which molarity you think is the most reliable. Use this molarity as your accepted value to calculate the %error.

Titration Problem Example: 

If 20 cm3 of a 0.3 M solution of NaOH is required to neutralize 30.0 cm3 of a sulfuric acid solution, what is the molarity of the acid solution?

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).

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