Determining the concentration of acid in a given solution

Mr (Na2CO3·10H2O)  = (23+23+12+48+20+160) = 286

250cm3 of 0.1 mol dm-3 contains (0.1 x 250/1000) moles

I will need (0.1 x 250 x 286/1000) = 7.15g of Na2CO3·10H2O in 250cm3

The error on a 2dp balance would be ± 0.005g. I would need to tare the balance before I weigh my solid, so this error would be doubled for every weighing because the balance needs to estimate 0.00 and the actual weight of the solid.

If I measured out 7.15g the percentage error for each reading is (0.005 x 2/ 7.15) x 100= 0.139%

This is quite a low percentage error so I could measure 7.15g of sodium carbonate and mix it directly into 250cm3 of water to make the solution.

Alternatively, to reduce the percentage error of weighing the solid, I could make up 250cm3 of a solution with the concentration of 1.0 mol dm-3 and then take 25cm3 of the stronger concentration and dilute it down in a conical flask to make 250cm3.

This would mean weighing out 0.25 of a mole of sodium carbonate to dilute into 250cm3 of water. (250/1000 x 286) = 71.5g

The percentage error of weighing 71.5g on a 2dp balance is (0.005 x 2/ 71.5) x100 =0.014%

This way of measuring the concentration needed may work out to have a higher percentage error as I would need to use a volumetric flask twice and also use a pipette. I would make up the 1 mol dm-3 solution in a 250cm3 volumetric flask which has the percentage error of 0.08%. As I would then take 25cm3 out using a 25cm3 pipette, this would mean further errors. The percentage error of the pipette is 0.24%. I would add that 25cm3 to another volumetric flask and fill up with water to get the concentration I needed. This would mean the overall precision percentage error for weighing out a higher concentration to reduce errors on the balance would be (0.014 + 0.24 + (0.08 x 2)) = 0.414%

The second method has a much higher overall percentage error than measuring out the correct mass of solid and adding it straight into the volumetric flask. The first way I explained would reduce procedural uncertainties as well as I will only have to judge the line on volumetric flask once instead of judging it twice and judging the line on the pipette.

I will therefore measure 7.15g of hydrated sodium carbonate and dilute it straight into a 250cm3 volumetric flask.

I could make up my solution by weighing the sodium carbonate in the volumetric flask as this would mean I wouldn’t have to move the solid, but it would be hard to get the exact amount as I couldn’t really get to the solid once it was in, to adjust the amount inside, as the neck of the flask is so narrow.

I could measure the solid into a small beaker on the balance, and then dilute it and pour the liquid into the volumetric flask. I would have to rinse the beaker with distilled water quite a few times and pour the liquid into the volumetric flask to make sure I have every last bit of the sodium carbonate inside the volumetric flask. I think this is a more accurate and suitable way of measuring out the solid as I can ensure that I have the correct amount of sodium carbonate in my solution to make the solution 0.1 mol dm-3.

Indicators 5 6

The concentrations I will be using of sodium carbonate and sulfuric acid are weak. Sulfuric acid is a strong acid, yet sodium carbonate is a weak alkali. I will use methyl orange as my indicator as it is suitable for a strong acid and a weak base and will show the endpoint on the acidic side between pH 3 and 4.5. This is important as the base is being added to neutralise the acid solution, so it will take place on the acidic side of the endpoint in the solution. The colour range of methyl orange is smaller than methyl red, so the endpoint will be clearer and easier to judge. Methyl orange will change the solution from red to yellow, which I think will be quite easily visible for me. It is also a readily available indictor in my school.  

Apparatus Needed

• Solid hydrated sodium carbonate
• Sulfuric acid solution
• Methyl red indicator
• 50cm3 burette
• 25cm3 glass pipette
• Pipette filler
• 2dp Balance
• 2 x 100ml glass beaker
• 100ml plastic beaker for waste
• Metal spatula
• Distilled water bottle
• Glass rod
• 250cm3 volumetric flask
• Conical flask
• Dropping pipette
• Clamp stand
• White tile
• Methyl red indicator

Risk Assessment 3

Procedure

Making up the sodium carbonate solution

• Turn on a 2dp balance and ensure it is on a stable and flat surface
• Wipe the balance to make sure there is no excess of previous weighing on the plate which could affect my reading
• Rinse a 100ml glass beaker with distilled water and dry to ensure it is clean
• Place the beaker on the balance
• Tare the balance
• Measure out 7.15g exactly of sodium carbonate into the beaker using a spatula
• Read the balance to ensure the weight is correct and write down the value next to the balance on my paper
• Add distilled water to the beaker and mix using a glass rod until the solid is dissolved.
• Hold the glass rod down the neck of the volumetric flask and pour the solution down the glass rod into the volumetric flask
• Add more distilled water to the beaker, rinse around, and pour down the glass rod again into the conical flask.
• Repeat the rinsing of the beaker 4 times to ensure that all of the solid is now in the volumetric flask.
• Be careful that the level of liquid isn’t near the 250cm3 line on the flask and use more distilled water to pour down the glass rod and ensure there is no sodium carbonate left on that.
• Pour distilled water around the inside edge of the volumetric flask to make sure there is no sodium carbonate left on the edges of the volumetric flask.
• Fill the volumetric flask up a bit more, but be careful not to go too close to the line.
• Get a small dropping pipette and rinse with distilled water to make sure it isn’t contaminated
• Place waste solutions in a plastic beaker
• Get another small glass beaker and rinse that with distilled water as well
• Add distilled water to the beaker and use the pipette in the beaker so that it doesn’t contaminate the distilled water bottle.
• Add drops of the distilled water from the pipette until the level of liquid is very close to the line, then bend down to have your eyes in line and add the last drops to ensure that the liquid doesn’t exceed the line.
• Place the lid on the volumetric flask and gently tip it upside down lots of times to ensure that the solution has an even concentration you throughout

If the volumetric flask wasn’t tipped then the concentration of the sodium carbonate would be higher at the bottom of the flask. When the solution is taken out of the volumetric flask, the pipette usually reaches to the bottom of volumetric flask. This means that the first titres would be using a stronger concentration of acid, as the majority of the sodium carbonate would be at the bottom. The concentration of the pipettes taken would then be weaker, meaning that the amount of solution needed to react with the acid would steadily increase and provide inaccurate and unreliable results.

Titrating

• Get the burette in a clamp and rinse it by pouring sodium carbonate solution down and letting it out the tap into a plastic waste beaker
• Make sure there are no bubbles in the spout by pouring more solution through and tapping the tip to remove bubbles of air
• Place the burette where measurements can be read clearly
• Pour the sodium hydroxide solution in the top of the burette
• Check the burette is straight from the front and side
• Read the measurement on the burette to the nearest 0.05cm3
• Record the value on a piece of paper as the initial burette reading
• Use conical flask and rinse it with distilled water
• Use the 25cm3 glass pipette and pipette filler and suck up sulfuric acid solution, swirl it round and pour excess into the sink or plastic beaker. Do this twice to ensure the pipette is not contaminated with anything but sulfuric acid.
• Pipette 25cm3 of the sulfuric acid solution and squeeze into the conical flask
• Touch the last drop from the pipette on the surface of the liquid
• Rinse down the edges using a distilled water bottle to ensure that all the acid particles are down in the base of the solution
• Place the burette and stand back on the bench
• Place the conical flask on the stand underneath the burette, with a white tile underneath so that the colour change is more easily visible
• Place 2 drops of methyl red into the acid solution in the conical flask and rinse down the sides again to ensure all of the indicator is in the solution.
• Adjust the burette so that the tip is inside the main body of the conical flask
• Wrap your left had around the tap of the burette and hold the tap between your fingers
• Hold the conical flask with your right hand
• Turn the tap on so that the sodium carbonate solution pours into the conical flask
• Mix up the solution in the flask by moving the conical flask in circular motions keeping the tip of the burette well inside the conical flask
• For the rough titration slow down the flow of sodium carbonate solution when about 25cm3 has passed through the taps so that the end point can be reached more accurately with a slow flow of solution
• Ensure the burette is vertical and read the reading again to the nearest 0.05cm3
• Record the value on the paper as the final burette reading
• Work out how much sodium carbonate solution was needed to reach the rough end point
• Fill up the burette again so that there is plenty of solution for the next titre
• Read the value to the nearest 0.05cm3 and record it on the paper
• Get a new conical flask and rinse it to ensure it isn’t contaminated or pour out the waste solution and rinse the conical flask with distilled water three or four times to ensure that there are no remainders from the previous titre
• Pipette 25cm3 of sulfuric acid solution into the conical flask, touching the tip onto the surface of the liquid again and rinse down with distilled water
• Add 2 drops of methyl red and rinse down the sides again with distilled water
• Repeat the titration again, flowing the solution out fast until you are 2-3cm3 away from the rough titre so that the endpoint can be reached much more accurately with just a few drops of sodium carbonate solution changing the colour
• Repeat the titration steps by topping up burette, ensuring it is straight, recording measurements, rinsing the conical flask, pipette 25cm3 into the conical flask and doing the reaction again, until three titres have been achieved that are less than 0.1cm3 apart 

Results

Mass of solid sodium carbonate used: 7.19g

I stopped collecting data when I finished ‘Titration 4’ as I had 3 results within 0.1cm3.

Analysis 8

The average titre is: (( 24.55 + 24.60 + 24.60 ) / 3 ) = 24.58cm3

I will round this up to 24.6cm3 as this is accurate to 3 significant figure. The fourth figure I recorded from my burette isn’t a real significant figure as it could only be a 0 or a 5.

My rough titrations were used to see the approximate volume of acid needed to neutralise the sodium carbonate. I ran the sulfuric acid through quickly and made sure that the end point was defiantly reached by creating a darker pink solution.

My second titration I did, ‘Titration 1’ I used to see how far from the endpoint I was. I ran out 25cm3 of the sulfuric acid as I knew the endpoint was around there. The colour was slightly darker than the colour of the endpoint so I knew that slightly less acid was needed to reach the endpoint.

For ‘Titration 2’ I ran out 22cm3 of sulfuric acid straight as I knew the endpoint was slightly under 25cm3 and didn’t want to miss the endpoint by adding too much. I then slowed down the tap so that drops were entering the conical flask and any colour change was easier to detect against the white tile. I added the acid drop by drop and eventually reached a very pale peachy, pink colour from the yellow, so I knew the endpoint was reached. I kept this conical flask so that further titrations could be compared against this endpoint colour.

For the last two titres, I let 23cm3 of acid flow straight into the conical flask as I knew that the endpoint was around 24.5cm3. I slowed down the flow of acid and then added it drop at a time to ensure that the colour change could be detected instantly as each drop affects the colour change and could mean missing the endpoint.

I achieved three measurements of my titres within 0.1cm3 which means they are reliable results.

The ideal concentration of sodium carbonate that I was going to make up was

0.1 mol dm-3. I worked out that I would need 7.15g of sodium carbonate in a 250cm3 volumetric flask using the calculation: (0.1 x 250 x (286/1000)) = 7.15g

When I was carrying out my experiment I weighed out exactly 7.19g of sodium carbonate. I can reverse the calculation I used to work out the exact concentration of sodium carbonate I used. This would be (7.19 / 250 / (286 / 1000)) = 0.10056 mol dm-3.

I will round this to 0.101 mol dm-3 as it is accurate to 3 significant figures as all the other data I used to work this value out was also accurate to 3 significant figures.

H2SO4 (aq)    +     Na2CO3 (aq)    →             Na2SO4 (aq)    +    CO2 (g)   +    H2O (l)

Moles  =  Concentration (mol dm-3)  x  Volume (dm3)

No. of moles of sodium carbonate used: 0.101 x (25.0cm3 / 1000) = 0.002525 moles

I will round this to 3 significant figures as the other values I used are accurate to 3 significant figures. This means that 0.00253 moles of Na2CO3 reacted.

The sodium carbonate and sulfuric acid react in a 1:1 ratio

This means that 0.00253 moles of sulfuric acid reacted with the sodium carbonate, so there is 0.00253 moles of this concentration of sulfuric acid in 24.60cm3.

Concentration (mol dm-3)  =  Moles  /  Volume (dm3)

Concentration of sulfuric acid: 0.00253 / (24.6/1000) = 0.10285 mol dm-3 

I will round the concentration of sulfuric acid down to 0.103 mol dm-3 (3 sig figs)

Evaluation 7, 8

Procedural Uncertainties

There are lots of possible errors that could have occurred whilst I was doing the experiment.

I had to ensure that all the solution measured in my burette was sulfuric acid, as the volume measured is very important in calculating my titre. I therefore rinsed the burette with the sulfuric acid to ensure it wasn’t contaminated with anything else.

I rinsed the volumetric flask out with distilled water to ensure that it didn’t have any sodium carbonate in it already or any other contaminate. This means that I know the exact mass of sodium carbonate in the volumetric flask, so therefore know the exact concentration of the solution. If there was more sodium carbonate in the flask before I added my known mass, then the concentration would be higher than I calculated. This would mean that more acid would be needed to neutralise it, and I would think that the acid had a lower concentration than it actually did.

I rinsed the pipette out with the sodium carbonate solution before using it so that it didn’t have any water or any other liquids inside. If there was water inside then this would dilute the sodium carbonate that I transferred in my pipette. It would make the concentration of the sodium carbonate in the 25cm3 lower, so less acid would be needed to neutralise it. I would therefore presume that the sulfuric acid was of a higher concentration that it actually is as the volume needed to react would be less.

I rinsed the conical flasks thoroughly with distilled water between each titration to ensure that there was no excess sodium carbonate or sulfuric acid in the conical flask which could react with any new sodium carbonate. If I went over the endpoint and I tipped out the solution, but there was excess sulfuric acid inside the volumetric flask then the sodium carbonate added would react with the sulfuric acid already in the conical flask. This would mean that less acid from the burette would be needed and I would therefore think that the sulfuric acid in the burette was of a higher concentration that it actually was.

When I added the sodium carbonate into the conical flask I also rinsed down the edges with distilled water to ensure that all of the sodium carbonate was in the bottom of the flask. It didn’t matter how much distilled water was in the flask, but it was important that all of the sodium carbonate molecules were in the solution at the bottom. I could calculate the moles so sodium carbonate in the conical flask and calculate the moles of sulfuric acid from that value. If not as much sulfuric acid was needed to neutralise the alkali as not all of the alkali reacted, then I would presume that the acid was of a higher concentration than it actually was as the volume needed was lower.

I inverted the volumetric flask when I made up the sodium carbonate and between each titre. This is to ensure that the concentration of the sodium carbonate was even throughout the volumetric flask. If I didn’t invert the flask then the first titres would be using the solution from the bottom of the flask which would have a higher concentration of sodium carbonate as it was added to the flask first. The concentration of the alkali taken would decrease for each titration making the volume of sulfuric acid needed to react decrease each titre. This would provide me with unreliable and inconsistent results.

When I was filling the volumetric flask, the pipette and the burette I had to ensure that the meniscus was resting on the calibration line to ensure an accurate result. If I had filled the volumetric flask up with the meniscus slightly above the calibration line, the concentration of the sodium carbonate would be slightly lower. This would mean that less acid would be needed to neutralise the alkali and I would think that the acid was of a higher concentration than it actually was so the titres would not be accurate.

If I read the meniscus slightly below the calibration line in the volumetric flask then the concentration of sodium carbonate used would be slightly higher that I would think. This would mean that more acid would be needed to neutralise the sodium carbonate, which would make me think that the concentration of the sulfuric acid was lower than it actually was meaning my calculations would not be accurate.

The endpoint of the reaction was quite hard to determine as the colour change was not very obvious. The solution turned from pale yellow to a pale peachy, pink colour. I didn’t want to add too much indicator as the indicator can make the solution ore acidic. If I added too much indicator to try to make the endpoint clearer, then less acid would be needed to neutralise the sodium carbonate and I would think that the concentration of acid used was higher than it actually was.

To make judging the endpoint easier, I kept the volumetric flask of a solution which I thought was the colour of the endpoint of the reaction. I used this to compare my other solutions too. This made my results reliable, as I stopped titrating when the same colour was reached, but not necessarily completely accurate.

I continually swirled the conical flask whilst I was titrating to ensure that the sulfuric acid was being mixed around the whole solution and would react evenly throughout. If I didn’t swirl the solution then the area in the middle of the conical flask where the sulfuric acid was falling would turn pink first as the OH- particles and H+ particles in that part of the solution would have reacted. I could mistake this as the endpoint if the colour wasn’t mixed throughout the solution until all of the particles had reacted and a colour change appeared throughout. If I mistook the small area of colour change as the endpoint then I would think that the concentration of sulfuric acid was much higher than it actually was as only a small amount was needed to create the colour change.

Precision Uncertainties 7, 8

The percentage errors calculated above have a total percentage uncertainty of:

(0.20 + 0.24 + 0.08) = 0.52%

To work out the range of accuracy of my results I will calculate 0.52% of the concentration, 0.103 mol dm-3: ( 0.103 / 100 ) x 0.52 = 0.000536 mol dm-3

The total precision uncertainty for the concentration of sulfuric acid calculated in my analysis would therefore be +/- 0.000536 mol dm-3.

If the concentration was inaccurate by + 0.000536 mol dm-3 this would results in me calculating a concentration of (0.000536 + 0.10285) = + 0.10338 mol dm-3.

If this was rounded down to 3 significant figures it would also give me a result of 0.103 mol dm-3, so it wouldn’t have any effect on my overall results, but if my measurements were accurate to 4 significant figures then this error would mean my concentration would be 0.1035 mol dm-3.

If the concentration was inaccurate by -0.000536 mol dm-3 it would mean my overall concentration calculated would be (0.10285– 0.000536) = 0.10234 mol dm-3 

If this was accurate to 3 significant figures then my concentration could possibly be 0.102 mol dm-3. This proves that the precision uncertainties of the equipment I used are great enough to have a possible impact on my results and mean that they are not completely accurate.

References

1  Revised Nuffield Advanced Book of Data (1984) Longman Group Limited, pg 93

2  Salters Advanced Chemistry (2000) Chemical Ideas, Heinmann, pg 12

3 CLEAPSS (2008) Science Publications HazCards 2007 Edition

4 Mr Churchill (February 2008) Thinking about practical work, ME 18

5 Vogel (1989) Vogel’s Textbook of Quantitative Chemical Analysis 5th Edition, Longman Scientific and Technical, pg 262

6 Salters Chemistry (2008) AS Experimental and Investigative Skill, Information on the use of indicators in acid-alkali titrations

7 Wycombe High School, Chemistry Department (February 2008) Appendix 2: Precision and accuracy in measurements

8 Wycombe High School, Chemistry Department (Feburary 2008) Measurement, uncertainty, and working with significant digits