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Titration experiment - write up

To analyse a solution of dilute sulphuric acid and calculate the concentration of acid it

contains

 

Aim:

 

To determine the precise concentration of sulphuric acid solution, using a solution of accurately known strength, called a ‘standard solution’.

 

Introduction:

 

A titration is a laboratory technique used in chemical analysis. A solution containing an unknown quantity of a known substance is placed in a conical flask, and a solution of known strength is then added from a burette. 1 The addition of this solution continues until an end point of the titration is reached. This is signified by a colour change in an indicator (or some other visible effect) and at this point the titration is stopped. The equivalence point of the titration occurs when the two solutions have reacted exactly. An indicator is a substance that changes colour when the reaction is complete. In an acid – base titration, the indicator is one colour at one pH and a different colour at another pH. The indicator I will be using is methyl orange.

 

I will be using sodium carbonate, as it is suitable for use as a primary standard for titrations of strong acids. The strong acid I will be using is sulphuric acid, which is a strong mineral acid. It is soluble in water at all concentrations. It has many applications, and is one of the top products of the chemical industry.2 A problem is that it produces acid rain. Acid rain is rain or any other form of precipitation that is unusually acidic. It has harmful effects on plants, aquatic animals and buildings. Acid rain is mostly caused by human emissions of sulfur and nitrogen compounds which react in the atmosphere to produce acids.3

 

 An acid is a substance which contains hydrogen and when dissolved in water furnishes hydrogen ions.

e.g.:  H2SO4 = 2H+ + SO4 2−

 

A base is a substance which will react with hydrogen ions to give a salt and water only. The base I will be using is sodium carbonate. The alkalis are substances which when dissolved in water furnish hydroxyl ions.

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e.g.:  NaOH  = Na+ + OH −

 

According to the Bronsted-Lowry theory the definition of an acid is a proton donor and a base is a proton acceptor.

 

 

 

 

Usefulness of titrations:

Titrations are important to understand how acids and bases react with each other, ways to understand detection of experimental error, and to experimentally calculate or figure out what a chemical or substance consists of or the properties it can contain. Titrations are a key tool in industry laboratories and for educational purposes. In addition a titration is not complicated, and is a fast and easy procedure, giving accurate results which are simple to reproduce. It is applied to biodiesel, used in the petrochemical and food industry.

 

Plan:

 

The amount of sodium carbonate required to make up 250 cm³ of the solution:

I plan to make up 250cm³ of sodium carbonate solution (Na2CO3); I have chosen to do so because the pipette I will be using has a volume of 25cm³, which means, with these volumes I can repeat my titration ten times, without the need of making up a new solution, this will save time.

 

The concentration of sulphuric acid (H2SO4) is between 0.05 and 0.15 mol dm -³, hence it is appropriate to choose the mid- point between these concentrations, which is 0.1 mol dm -³.

 

The balanced equation for the reaction:

1) Na2CO3 (aq)      +        H2SO4 (aq)      Na 2SO4 (aq)       +    H2O (aq)       +     CO2 (g)

2) Base                +     acid               salt                 +    water         +    gas

 

There is a 1:1 ratio of the compounds sodium carbonate and sulphuric acid, meaning one mole of Na2CO3 neutralises one mole of H2SO4.

 

Before making up my known 250cm³ of Na2CO3 solution, I will first need to find the amount of Na2CO3 needed to be dissolved in 250 cm³of distilled water for making a one molar solution. To do this I will find the molar mass of Na2CO3 then divide the molar mass (Mr) by 1000 to convert the units from cm³ to mol dm-³, I will then divide by 4 (because 1000/250 = 4) to get the exact amount of anhydrous sodium carbonate.

 

The Mr of Na2CO3:

Na =   23  x   2 = 46

C   =   2    x   1 = 12

O   =   6    x   3 = 48

 

46 + 12 + 48 = 106 grams of Na2CO3 = 1 mol of Na2CO3.

 

Hence, 106 grams dissolved in distilled water is the mass for a 1 litre solution which would give a concentration of 1 mol dm -³. As stated earlier I will use a 250 cm solution, therefore 1000cm would be an unwanted volume, as it is far too much. To decrease it to 250 cm I must divide my Mr by 4 and then by 10 to work out the mass needed for a 0.1 molar solution.

 

106/4 = 26.5 g

 

26.5 g/10 = 2.65 g

 

1000/4 = 250 cm of distilled water

 

Consequently, I will use 2.65g of Na2CO3, to give a 0.1 molar solution.

 

The indicator:

Acid-base indicators are generally weak acids, with a dissociation which can be represented as:

HIn (aq)         H+ (aq) + In- (aq)

 

The weak acid HIn and/or its conjugate base In-  is coloured. A change in pH causes a shift (in the equilibrium in the above equation) which causes a change in colour.

 

A good indicator should show a dramatic colour change in order to make it easy to detect the end point of the titration. I will be using the recommended indicator methyl orange. In acid its colour is orange. In alkali its colour is pink.

 

 

Risk Assessment:

All substances must be handled with much care and precaution. During the experiment safety goggles will be worn at all times to protect the eyes, and protective clothing i.e. a lab coat will also be worn to protect the skin and clothing from any spillages.

 

Name of chemical Hazard Safety Precautions

Sodium carbonate (anhydrous)

 

 

 

An irritant which can irritate the lungs if inhaled. Can cause irritation to the skin and eyes.

Eye contact: Immediately flush the eye with water for 10 minutes.

 

 

 

 

 

 

 

 

Sulphuric acid

 

 

 

 

 

 

 

 

 

 

 

Methyl orange

Harmful if swallowed.

 

 

 

 

 

An irritant which can cause severe burns and tissue damage due to dehydration and the heat liberated by the reaction with water. Damage to gastrointestinal tract.

Methyl orange is thought to be toxic if swallowed or inhaled. Avoid skin contact. The solution may be made up of ethanol or an ethanol/water mixture meaning it could be poisonous

If irritation persists, call for medical help. Skin contact: Wash off with water. If swallowed: Call for medical help if the amount swallowed is large.

Eye contact: Immediately flush the eye with water for 10 minutes. If irritation persists, call for medical help. Skin contact: Wash off with water. If swallowed: have 2 -3 glasses of water and seek medical help.

and flammable.

Eye contact: Immediately flush the eye with water. If irritation persists, seek medical attention.

Skin contact: Wash off with plenty of water. If swallowed: Drink plenty of water and call for medical help.

 

 

Apparatus:

      sodium carbonate (2.65g)

      sulphuric acid (between 0.05 and 0.15 mol dm-3)

      A bottle of methyl orange - indicator solution

      A burette (50 cm3)

      A pipette (25 cm3)

      Volumetric flask (250cm3)

      Conical flask

      2 bottles of distilled water (300cm)

      Electronic balance (to 2 decimal places)

      2 Beakers (100cm3 and 250cm3)

      Glass stirring rod

      Pipette filler

      Spatula

      Weighing bottle

      Funnel

      White tile

      Goggles

      Lab coat

 

 

 

 

 

 

 

 

 

 

 

Figure 1.1

Diagram to show the set up apparatus:

       

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Reasons for certain apparatus being used:

Pipette Pipettes are more accurate than measuring cylinders. They are easy and safe to transfer a specific volume of solution accurately. It is a precise piece of equipment designed to deliver a certain volume of liquid. The pipette is filled above the mark by sucking solution into it and this liquid is allowed to drain away. This process of filling and allowing the liquid to run away will be repeated to ensure that the pipette contains nothing but the solution. The pipette has a percentage error of 0.06. 

Burette Burettes are also more accurate than measuring cylinders. They have graduations every 0.1 cm3, so when you take a reading it should not be more than 0.05 cm3 too high or too low. The

volume of solution can be easily and accurately read. A burette is an accurate piece of equipment and usually has a capacity of 50 c.c. The burette is first washed out with the solution it is to contain, the washings being allowed to run through the jet so as to wash this part also. A second washing is desirable to ensure complete elimination of impurities. The burette has a percentage error of 0.05. 

Volumetric flask Volumetric flasks are specially designed containers for very accurate mixing and diluting.  Like volumetric pipettes, volumetric flasks are designed to measure one volume only.  The long neck makes it easy to determine when the final volume has been reached. It has a specific graduated meniscus line. But it is safer to make up near the mark, then homogenize, then make up to the mark using a small pipet and doing it drop by drop at eye level to the meniscus. It is safe for mixing/inverting the solution because of the stopper. 

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%.4

 

 

 

 

 

 

 

 

 

 

 

 

 

Method

 

Preparation of a standard solution:

      Wear safety goggles, a lab coat and gloves for heath and safety.

      In any titration, the concentration of one of the solutions must be accurately known.

      Firstly, calculate the mass of the solid needed to make a solution of the required concentration.

      Place the weighing bottle on the electronic weighing balance, then press the ‘tare’ button so that the balance discounts for the weight of the weighing bottle and reads 0.00g.

      Carefully add 2.65g of anhydrous sodium carbonate into the weighing bottle using a spatula.

Record the mass of anhydrous Sodium carbonate.

 

 Making up a solution of sodium carbonate, of known concentration:

1)      Before beginning, make sure that all the glassware has been properly cleaned and are dry, so that the solution doesn’t get contaminated.

2)       Put the sodium carbonate powder into a clean 250 cm3 beaker. Rinse out the weighing bottle three times with distilled water and add it into the beaker. This

will ensure all the powder has gone into the beaker, and so the concentration is kept as accurate as possible. 

3)      Dissolve solid sodium carbonate in the 250 cm3 beaker, add around 50 cm3 of distilled water and mix using a stirring rod to ensure it has fully dissolved.

4)     

Figure 1.2A diagram showing meniscus is touching graduation mark. 

Clean the stirring rod using distilled water and wash the water into the beaker.

5)      Pour further distilled water into the baker until it is filled half way (up to 150 cm3)

6)      (the funnel prevents spillages, hence, keeping the concentration as precise as possible)

7)      Wash out the beaker very well, rinse it out 2 or 3 times with distilled water, so that all the solution is transferred using a funnel, from the beaker to the volumetric flask.

8)      Add more distilled water to the volumetric flask until it is one cm below the graduation line. At the point continue to add distilled water but with a dropping pipet drop by drop until the bottom of the meniscus lines up exactly with the graduation mark on the neck of the volumetric flask. Do this at eye level so that it is accurate.

9)      Place the stopper on the flask and invert it by continuously shaking about 10 to 15 times.

 

 

 

 

Preparing the burette:

1)      Before beginning, make sure that all the glassware has been properly cleaned and are all dry, so that the solution doesn’t get contaminated. (burette, pipette, conical flask and volumetric flask)

2)      Set up the boss and clamp stand.

3)      Rinse out the burette with sulphuric acid twice; this will help avoid inaccuracy of results due to contamination. 

4)      Close the burette tap and use a 250 cm3 beaker and funnel to fill the burette with sulphuric acid solution.

5)      Open the tap, and allow the solution to run out space under the tap is full of solution.

6)      Using a white tile let the solution run out of the burette until it reads 0.00, and the bottom of the meniscus is exactly on the line. Make sure to be eye level with the meniscus so that an accurate reading can be made.

7)      Check for air bubbles if there are some present, empty then fill the burette again. If there are no air bubbles then precede the experiment.

8)      Remove the funnel from the top of the burette.

9)      Record the volume of the acid before starting the experiment to the nearest 0.05 cm3

10)  Rinse out the pipette with sodium carbonate solution twice. Twirl the pipette round with your fingers and hold it at an angle to the body, so that the solution rinses it out completely and there are no inaccuracies.

11)  Pipette out sodium carbonate solution using the pipette filler, do this to the graduation mark on the pipette.

12)  Carefully transfer the solution in the pipette in to the conical flask. Try to avoid any spillages as this would give major errors in results, they would be inaccurate and incorrect.

13)  Let the pipette drain freely, once the solution has run out, touch the tip of the pipette to the side of the flask, to deliver exactly 25 cm3 into the flask.

14)  Add 3 drops of methyl orange indicator, (make sure each time the amount of drops are added is the same. It must remain constant to make it a fair test.)

15)  Place the conical flask on the stand so that the tip of the burette slightly enters the neck of the conical flask. Then place the white tile underneath the conical flask.(so that a colour change can be recognised more easily)

16)  Swirl the conical flask to mix the three drops of indicator with the 25 cm3 of sodium carbonate solution.

17)  Firstly, do a ‘rough’ titration. Open the tap and allow the solution in the burette to run freely into the conical flask. Add 1-2 cm3 at a time then swirl and repeat.

18)  Watch out for the endpoint of the methyl orange. The colour change is from slightly orange to pale pink. Once the colour change is observed, swirl the conical flask until the colour change is permanent. Read off the reading on the burette, this is your rough titration reading and will help with the next few accurate titres.

19)  Rinse out the conical flask with tap water, and then with distilled water.  

20)   Repeat steps 11 – 16. This time, once the burette reading is 3 cm3 above the rough titration reading. Add drop by drop, controlling the tap very carefully. Use one hand to swirl, and the other to operate the tap. Repeat the titration until three concordant titres are achieved all within 0.1 of each other.

 

How to ensure greater accuracy of results:

1)      Make sure all glassware is clean and dry.

2)      Weigh the anhydrous sodium carbonate accurately.

3)      Check electronic balance for any dust or spillages.

4)      Add exact amount of distilled water for making solution.

5)      Rinse the pipette out with the solution it will contain.

6)      Rinse the burette out with the solution it will contain.

7)      Check that the burette tap is not too stiff nor too lose, and that it is not leaking.

8)      Do a rough titration to give a rough indication of where the colour change occurs.

9)      Keep the number of indicator drops added constant (e.g. – always add 4 drops each time.)

10)  Use a white tile, so that the colour change is recognized more easily.

11)  After letting the pipette run its containing solution into the conical flask, touch the tip of it on o the side of the conical flask so that the last drop is accounted for.

12)  Remove the funnel after filling up the burette.

13)  After each titration rinse out the conical flask with tap water and then with distilled water.

14)  Continuously swirl the conical flask as the sulphuric acid is being added.

15)  Always make a reading at eye level.

16)  Work in a clean, spacious and hazard free environment.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Results

 

Using my results I will now determine the precise concentration of sulphuric acid solution (H2SO4) that is required in order to neutralize 25 cm3 of 0.1 mol dm-3 of sodium carbonate solution. (Na2CO3)

 

Results table:

 

Titration (cm3) rough 1 2 3 4 5

Final burette reading 27.4 27.8 27.4 28.5 28.5 28.5

Initial burette reading 0.00 0.00 0.00 0.00 0.00 0.00

Titre cm3 27.8 28.2 27.4 28.5 28.5 28.5

 

 

Average table of results:

 

Titration (cm3) 1 2 3

Final burette reading 28.5 28.5 28.5

Initial burette reading 0.00 0.00 0.00

Titre cm3 28.5 28.5 28.5

 

In the above table I have placed the best results which are all 0.1 cm3 of each other.

 

The average titre:

To work out the average titre volume I will take three out of the five readings that are all within 0.1 cm3 of each other. I will add the three readings together, then divide them by the number of readings (3)

 

 

 

 

 

Average titre value = 28.5 cm3    

 

 

 

 

The balanced equation shows that 1 mole of  Na2CO3 reacts with 1 mole of H2SO4, therefore at neutralisation the number of moles of will equal the number of moles of

But in solution the number of moles is equal to concentration multiplied by the volume (dm3.)

N      =    C    x    V

 

Na2CO3 (aq)      +        H2SO4 (aq)      Na 2SO4 (aq)       +    H2O (aq)       +     CO2 (g)

          1           :          1 

 

 

 

 

 

The concentration of the solution, Na2CO3, is known and the volume of the pipette (25cm3) is also known. Therefore, I can work out the number of moles of Na2CO3 in the conical flask

N      =    C    x    V

N      =    0.1  x    25/1000

N      =    0.0025 mol

 

At neutralisation there are the same number of moles of H2SO4.We can re-arrange the formula N = C  x  V, to work out the concentration of the acid, H2SO4.

 

Number of moles = 0.0025 mol    x    volume   =   average

 

Titre value = 28.5 cm3

 

C = N/V

 

0.0025

28.5/1000

 

C  = 0.087 mol dm3, I will round this value up to give me 0.09

Hence, the concentration of the sulphuric acid solution is = 0.09 mol dm -3

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Evaluation

 

During my second titration I obtained an anomalous result, it was 0.4 cm3 in difference, if I compare it, with my rough titration. It could have come about for a number of reasons, for instance I may have added too much sulphuric acid to neutralise 25cm³ of sodium carbonate solution. An experimental error may have occurred in terms of practical skills, as during the experiment the tap became stiff to turn, hence, extra solution may have run into the conical flask neutralising it at an incorrect titre. Leaving the funnel in the burette would also account for inaccurate results as more solution would have dripped into the burette which would affect the readings.

 

 A constant variable which I tried to keep throughout my experiment was the number of drops of methyl orange indicator being added into the conical flask. However, I may have accidentally added an extra drop or the size of the drops I added could have differed. For example, during my second titration I added three big drops, this would have caused the solution to neutralise faster, hence, giving me a lower titre reading, and this error could have caused the anomaly in my experiment.

 

Furthermore analytical skills play a key role. As a chemist I had to observe, make my own judgment and decide when I thought a colour change had occurred. (Using

methyl orange as an indicator was difficult, as the colour change is not very significant.) To make sure my observations were accurate I could have used a pink/peach coloured bit of paper to compare each colour change in my conical flask. During every titration I had to stop adding the acid at the point where I thought the colour change was the same as the last, the colour change had to remain constant as this would make it a fair test. However, my judgment of the indicator changing colour, may have not always been exactly correct.

 

There may have been a spillage of the sodium carbonate powder when transferring it from the weighing boat/bottle into the beaker. The amount needed is 2.65g; if the amount is less then this will result in decreasing the concentration, making it incorrect. Hence, my calculations would be inaccurate. Dropping any powder onto the balance would have caused an incorrect reading, therefore I had to make sure it was very clean, I used a tissue to do so. I could have also used a more accurate scale, perhaps, an analytical balance which measures to 0.0001g as this instrument would have a high degree of precision. To ensure that that the sodium carbonate was completely dissolved I could have used a magnetic stirrer.

 

A volumetric flask is a very accurate piece of equipment in measuring a volume of a liquid, and it must be filled up exactly to the graduation mark to make up 250cm3.To make sure I was accurate I used a small dropping pipette to add the final few drops into the flask making sure the meniscus touched the bottom of the line. I also used a white tile and made sure I was at the correct eye level, so that I wasn’t looking at the meniscus from a wrong angle. Getting this wrong would have caused me to begin my experiment again. Hence, I made sure I was extremely precise and cautious. The volumetric flask has a percentage error of 0.24%. During my experiment I made sure to shake the volumetric flask, so that, I could ensure that the standard solution was mixed thoroughly.

 

In addition if the equipment was not properly cleaned then my solution would have been contaminated resulting in inaccurate results. I cleaned out the burette and pipette with the solutions which were to be added to each so that they had been rinsed out and a contamination error could not occur. I used distilled water to clean my conical flask. After filling my burette I wiped it clean with a tissue so that readings could be made accurately and no spillages into my conical flask occurred as this would have altered the concentration, hence, corrupting my results. The apparatus used have percentage errors these were calculated so that they are accounted for and my results are as correct as possible.

 

When adding the sulphuric acid through the burette tap some of the solution dripped onto the sides on the conical flask, to account for those drops I washed out the conical flask with distilled water every time I saw a drop on the side or the neck of the flask.

Distilled water would not alter my concentration therefore it was effective and a good way of accounting for extra drops. I made sure I repeated the titration until concordant results were achieved.

 

Another evaluative point is that, I worked in a very crowded environment, which was unsafe. It could have caused spillages or major accidents. There were many people working in the classroom, which was hazardous and also dangerous for heath and safety reasons. If I was to repeat the experiment, I would do so in a less dangerous, less crowded environment.

 

Overall my experiment went very well, and I managed to obtain solid results which led me to achieve my aim by using a set of calculations and equations to work out the concentration of the acid. The outcome was that at neutralisation the number of moles of Na2CO3 will equal the number of moles of H2SO4, and an indicator change signaled the acid-base reaction was complete.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Percentage errors

 

Experimental error is always with us; it is in the nature of scientific measurement that uncertainty is associated with every quantitative result. This may be due to inherent limitations in the measuring equipment, or of the measuring techniques, or perhaps the experience and skill of the experimenter. Certain apparatus have different percentage errors, as they are all unlike in volume etc.

 

The formula for working out percentage error is:

 

                        

 

 

 

 

 

25cm3 pipette

The error for 25 cm3 pipette is 0.06

The equation is: % error = 0.06/25 x 100 = 0.24%

 

50cm3 burette

The error for 50 cm3 burette is 0.05

% error  = calculated value – actual value  x 100

calculated value

  

  

The equation is: % error = 0.05/28.5 x 100  =  0.175  = 0.18%

 

250cm3 volumetric flask

The error for 250 cm3 volumetric flask is 0.6

The equation is: % error = 0.6/250 x 100  =  0.24 %

 

 

Balance error

The error for the weighing balance is 0.005

 The equation is: % error = 0.005/2.65g x 100 = 0.18%

 

To find the overall % error in my experiment I will add the four errors of these four pieces of apparatus.

 

25cm3 pipette = 0.24%

50cm3 burette = 0.18%

250cm3 volumetric flask = 0.24%

Weighing balance = 0.18%

 

Overall % error = 0.24 + 0.18 + 0.24 + 0.18 = 0.84

Therefore the overall % error in this experiment is = 0.84%

 

% errors can affect the results especially if they are large. The highest % error = 0.24% for the pipette and the volumetric flask.

 

Statistics

The concentration of the sulphuric acid solution is = 0.09 mol dm -3

 

 

 

 

2) 0.84/100   x   0.09 =  7.56 x 10 -4

   

 

 

The resulting values would be my boundaries.

 

    0.09   +  7.56 x 10 -4  is my maximum

-4 

0.09 – 7.56 x 10-4  is my minimum error. 

3) 0.09      +        7.56 x 10 -

4                                                       - 

 

 

 

 

 

 

 

 

I am confident about my values, the minimum and maximum error boundaries show that I have accounted for all the inaccuracies that could have altered my results. They are not too high nor to low which shows that all the errors during my experiment were not highly significant.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Bibliography

 

Books:

      George Facer - AS Chemistry - Edexcel - Philip Allan updates 2005 - pages 284, 285, and 286.

 

      Ann and Patrick Fullick - Chemistry second edition - Heinemann advanced science – page 273 and 275.

 

Websites:

      http://www.adrian.edu/chemistry/th/Somelinks/Spages/jfakult/jfakult.php

 

      http://ptcl.chem.ox.ac.uk/~hmc/hsci/chemicals/sodium_carbonate.html

 

      http://ptcl.chem.ox.ac.uk/~hmc/hsci/chemicals/methyl_orange.html

 

      http://images.google.co.uk

 

       http://www.bioquest.org/icbl/projectfiles/titration.jpg

 

      http://en.wikipedia.org/wiki/Meniscus

 

Articles:

      Zahoor Al-Haq – Chemistry review – volume 16 – number 4 – ‘How to be a lab success’ – page 13 – Philip Allan updates – April 2007.

 

References:

      1 - Ann and Patrick Fullick - Chemistry second edition - Heinemann advanced science – page 273 and 275.

      2 - http://en.wikipedia.org/wiki/Sulfuric_acid#Industrial_hazards

      3 - http://en.wikipedia.org/wiki/Acid_rain

      4 - http://www.chemistry-react.org/go/Tutorial/Tutorial_4428.html

 

 

      Figure 1.1 - http://images.google.co.uk

 - http://www.bioquest.org/icbl/projectfiles/titration.jpg

 

      Figure 1.2   - http://en.wikipedia.org/wiki/Meniscus

      http://www.dartmouth.edu/~chemlab/techniques/titration.html - photos of method

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gcse chemistry titration fuel combustion organic chemistry alcohols write experiment

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