# 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.

We Will Write a Custom Essay Specifically
For You For Only \$13.90/page!

order now

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.

7. Rinse the burette 3 times with 5 ml of the 0.1M H2SO4 solution to be titrated. Discard the rinses. With the stopcock open and over a beaker, begin filling the burette with 0.1M H2SO4 solution. Close the stopcock once the tip is draining and the volume is above the stopcock. This will prevent bubbles from forming in the stopcock during the filling process. Finish filling the burette to within a few ml of middle mark i.e near the 25ml mark for 50ml burette. Record this volume as the initial volume in the titration. (Note that this is not the volume of liquid in the burette.)

SETTING UP THE TITRATION (STEPS 5 AND 6 OF THE METHOD)

Clamp the buret into the buret clamp when you are not using them. (See picture below left.) When you clamp the burets in, make sure they are veritcal. The picture below right shows a buret being clamped in crooked at first, but then corrected to be clamped in vertically.

The clearly labelled buret clamped vertically into the buret clamp.

Clamping the buret into the buret clamp. Notice that the buret goes in crooked at first, but is then corrected so that it is vertical.

Before filling and using the burets, you should thoroughly clean the burets so that liquid will not cling to the sides as the burets are drained during the titration experiment (since this would introduce error into your volume measurements). The teflon stopcock of the buret should first be removed from the end and the inside of the buret barrel should be scrubbed with detergent and a buret brush. The buret should then be rinsed thoroughly with tap water to remove all traces of detergent and then rinsed thoroughly with distilled water. (These steps are illustrated below).

A dirty buret. Notice the liquid clinging to the sides. This will introduce error into your volume measurements. Liquid should drain cleanly from the sides if the inside of the buret is clean.

The teflon stopcock should pull out easily from the end of the buret. Remove the stopcock before cleaning the buret.

Scrub out the barrel of the buret with detergent and a buret brush.

Rinse thoroughly with tap water, then distilled water.

Before filling the burets they should be rinsed out once more, but now with the solution you plan on filling the buret with. First put the teflon stopcock back in with the valve shut, then pour about 10 mL of dilute acid or base solution into the buret, tilt the buret on its side, and rotate the buret so that the solution rinses the sides. Once the sides are thoroughly rinsed, open the stopcock valve to let all of the solution drain through the stopcock tip. (See pictures below.)

Put the teflon stopcock back in with the valve closed and pour about 10 mL of the solution you want to fill the buret with, into the buret.

Tilt the buret on its side and wash down the walls by rotating the buret, then right the buret and let all of the solution drain out through the stopcock tip.

Now fill the buret with the dilute acid or base solution until the liquid level is close to the top of the measuring range. Open the stopcock valve and let a few mL of solution run through until all air bubbles are purged from the tip. Close the stopcock valve and read the initial volume on the buret

Open the stopcock and let a few mL of solution flow through until all air bubbles are purged from the tip.

Then read the initial volume on the buret (do only for the base buret in the first titration). A black line on a white piece of paper held behind the buret will darken the meniscus and make it a little easier to read the volume. I would read this volume as 3.42 mL, do you agree with that?

Diagram 9

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 ï¿½.0006 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.

10.

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.

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)

9.0 M NaOH

=

0.014 L

=

14 mL

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).

(2 mol NaOH)

MNaOH =

[{(0.1500 M H2SO4)(0.03000 L)}

(1 mol H2SO4)]

0.02000 L

=

0.45 M NaOH

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

mol base =

mol acid x

(1 mol base)

(1 mol acid)

MolarMass Acid =

grams acid

moles acid

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.

MolarMass Acid =

1.100 g acid

[{(0.2100 M NaOH)(0.02610 L)}(1 mol acid/ 1mol NaOH)]

=

200.7 g

mol

RESULTS AND CALCULATIONS:

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

Final volume (methyl orange)/cm3:

Initial volume/cm3:

Titre (methyl orange)/cm3:

Final vol (phenolphthalein)/cm3:

Initial volume/cm3:

Titre (phenolphthalein)/cm3:

Final volume (sulphuric acid)

Initial volume / cm3:

Table 3

Molar mass of Na2CO3, M

106 gmol-1

Mass of bottle and contents before transfer, M1

g

Mass of bottle and contents after transfer, M2

g

Mass of Na2CO3, m = (M1 – M2)

g

Amount of Na2CO3 , n = m/M

mol

Volume of solution, V

0.25 dm3

Concentration of Na2CO3 , c = n/V

moldm-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:

Equipment

Measurement to the nearest:

Balance (1 decimal place)

0.08 g

Balance (2 decimal place)

0.008 g

Balance (3 decimal place)

0.0008 g

Measuring Cylinder (25 cm3)

0.5 cm3

0.04 cm3

0.08 cm3

0.2 cm3

1. 1.00 g on a 2 decimal place balance

2. 10.00 g on a 2 decimal place balance

3. 1.00 g on a 3 decimal place balance

4. 10 cm3 in a 25 cm3 measuring cylinder

5. 25 cm3 in a 25 cm3 measuring cylinder

7. 25 cm3 in a 50 cm3 burette (Grade B)

1. 8%

2. 0.8%

3. 0.08%

4. 5%

5. 2%

6. 0.16%

7. 0.32%

8. 0.08%

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 ï¿½0.5 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 ï¿½1 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 ï¿½0.05 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.

* 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 ï¿½ Na2SO4 + CO2 + H2O

Ratio in moles 1 : 1 ï¿½ 1 : 1 : 1

Moles = mass = volume ï¿½ concentration.

MR

H2SO4

Na2CO3

moles

0.0025

0.0025

volume

25

1000 dm3

25

1000 dm3

concentration

0.1M

0.1M

Table 8

Since moles = mass = volume ï¿½ concentration

MR

= mass = 25 ï¿½ 0.1M

106 1000 dm3

mass = 0.025 ï¿½ 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).

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.