Sulphur Dioxide
Sulphur Dioxide is a very polluting gas and therefore is converted to sulphuric acid where possible. This can then be sold as a useful by-product. Sulphur dioxide is emitted from car exhausts, burning coal and from ore production, although recently steps have been taken to make sure that the sulphur dioxide produced is converted to sulphuric acid. But sulphuric acid is also produced naturally from volcanoes. Sulphur dioxide is a very soluble gas and combines readily with water and oxygen in the atmosphere to form sulphuric acid. This then falls to earth as acid rain. Acid rain is a very harmful environmental problem. Acid rain can kill vegetation, for example the Black Forest in Germany has suffered from the effects of acid rain. Acid rain kills aquatic life and turns lakes acidic, it can ruin statues or limestone buildings and dissolves metals out of the soil leading to poisonous metals being released into water sources.
Acid rain is the reason why sulphur dioxide should be converted into sulphuric acid after extracting metals.
Acids and their reactions
Acids and Alkalis are classified as strong or weak depending on how far they can form ions in water. Strong acids or alkalis are completely ions in solution but weak acids or alkalis are only partially ions in solution.
Acids are aqueous solutions that contain hydrogen (H+ ions). They turn litmus paper red and remove CO2 from carbonates.
Acids react with carbonates to form salt, carbon dioxide and water.
Acid + Carbonate Salt + CO2 + H2O
e.g.- 2HCl(aq) + CaCO3(s) CaCO2(aq) + CO2(g) +H2O(aq)
In the practical that is being performed the reaction involves an acid and a carbonate will be used, therefore the products that will be produced are a salt, carbon dioxide and water.
Sulphuric Acid
The acid that is being used in this experiment is sulphuric acid. Sulphuric acid has a formula of H2SO4 and is a combination of H+ and HSO4- ions. It is a very strong acid. A molecule of Sulphuric Acid, H2SO4, consists of two atoms of hydrogen, one atom of sulphur and four atoms of oxygen
Sulphuric acid is a colourless viscous corrosive oily liquid, its properties are as follows:
- Melting Point : 10.3C
- Boiling Point : 338C
- Formula weight 98.08
- Specific gravity or density 1.94
- Flash point none
Sulphuric acid is the strong acid produced by dissolving sulphur trioxide in water.
Sulphur Trioxide + Water Sulphuric Acid
SO3 + H2O H2SO4
Indicators
Indicators are used in titration to show the point between an acid and alkali. Different indicators are used for different strengths or weaknesses of acids or alkalis.
For example for a:
For a weak acid and a weak alkali pH meter, conductivity meter or a temperature probe has to be used.
In this experiment a strong acid, sulphuric acid, and a weak alkali, sodium carbonate will be used. Therefore methyl orange is the appropriate indicator.
All acid-base indicators are weak acids this means that just a small proportion of their molecules ionise in an aqueous solution. This can be illustrated in the following equation:
HIn <=> H+ + In– (In – indicator)
HIn and In– have different colours, for example with methyl orange HIn is red and In– is yellow.
As indicators are weak acids mainly HIn will form. With methyl orange this is coloured red, with Phenolphthalein this is colourless. Adding an acid or an alkali changes the amount of ions that are in the solution. This then changes the colour of the solution. For example adding an alkali to methyl orange changes its colour. This is because the alkali reacts with the ions in the solution. The HIn then ionises to produce more H+ and In- ions, eventually there are more In- ions and the indicator will appear yellow. In methyl orange this change will occur when there are 10 times more In- ions than there is HIn. With methyl orange a H+ ion has attached itself to a Nitrogen atom
Reactions with indicators are reversible so the indicator can start at either colour.
All indicators change colour over a particular pH range and this has to be considered when we are choosing an indicator for a particular titration. The pH ranges of the above indicators are as follows:
Strong acid and weak alkali titrations such as sulphuric acid and sodium carbonate are likely to have an endpoint in the region of pH 3- 5. This corresponds to the pH range that methyl orange indicator changes colour over, therefore methyl orange indicator is a good choice for this combination.
Phenolphthalein can only be used for a combination of a weak acid and a strong alkali as this titration is likely to have an endpoint within the range of 8-10.
Calculations
A mole is a certain number of particles, generally atoms or molecules. The number is 6.023 x 10 23, this is called Avogadros constant. In practice atoms or molecules are not counted, therefore it is more important to know the mass of a mole of something. The mass of a mole of atoms is the same as its relative atomic mass in grams
Therefore,
Moles = Mass
RAM
For a molecule,
Moles = Mass
RMM (relative molecular mass)
Thos formula will be used in calculating the mass of sodium carbonate needed.
No. of moles = volume x concentration
This formula will be used in calculating the concentration of the sulphuric acid.
Other techniques
Techniques such as colorimetry have recently been developed to find out the concentration of a solution. This technique is very effective but cannot be carried out with the equipment provided.
Apparatus
- Sulphuric Acid solution
- Anhydrous Sodium Carbonate Solution
- Weighing Bottle
- Balances
-
2 100cm3 beakers
-
250cm3 conical flask
-
250cm3 volumetric flask
- Burette
-
50cm3 pipette an filler
- 2 small dropping pipettes
- 2 Small funnel
- Distilled water
- Methyl Orange indicator
- White Tile
- Clamp
- Glass rod
(Safety Apparatus)
Method
Titration involves using a solution with a known concentration to determine the concentration of another solution using molar calculations.
The burettes and pipettes must be both clean and dry or rinsed with the solution that they are going to contain, sulphuric acid. As there is not a fixed number of moles in the burette this method is an effective way of avoiding dilution. A small amount of solution will be added to the burette and pipettes, this solution must then be swirled around for a few moments and then panned out. The apparatus is then wet with the solution.
After the burette has been cleaned, it will then be clamped. The acid should be run through the burette to fill up the jet and the tap can be used to remove any air bubbles in the jet. . A funnel will then be placed at the top of the burette and the valve set so that no solution can flow through. The burette and funnel should be set up so that the funnel is below eye level for safety and practical reasons. Remove the funnel after use as otherwise more solution may drip into the burette from the funnel and an accurate volume would not be recorded. The conical flask should be placed underneath the burette so that the solution would run down into it if the valve were open. The conical flask should be placed on a white tile. This makes it clearer to see the colour change of the indicator as the colour of the laboratory bench may interfere with the colour of the indicator and the exact end-point would be difficult to determine.
A rough titre needs to be performed first to obtain an approximate value of the volume needed to reach end-point. The burette must be filled with a 100cm3 to the 0cm3 mark so that the meniscus is just touching the graduation mark. If the meniscus is slightly below the graduation mark then the value should be recorded (for example 0.2cm3) but if the meniscus is above the graduation mark then some of the solution must be removed as an accurate value for the volume cannot be obtained.
The standard solution (method described later) of sodium carbonate should then be measured and put in the conical flask Before putting the solution in the flask the flask must be rinsed out with water to make sure that it is clean and then dried. As the conical flask has a certain number of moles measured into it rinsing it with the solution would increase the number of moles in the solution. Therefore distilled water must be used. The flask must be empty before the solution is poured into it, otherwise the amount of moles in the solution might be effected.
Then using the pipette and filler a volume of exactly 20cm3 should be measured into the conical flask. This volume should be measured with a pipette and pipette filler so that the volume is accurate. Add two drops of the methyl orange indicator using a small dropping pipette into the solution in the conical flask. This ensures that the same amount of indicator is used for each solution, this then gives a constant environment and ensures that all titration end-points are based on the same amount of indicator. A large amount of indicator would change colour quicker as it has a greater concentration and therefore more contact with the solution than a small amount. Then by turning the tap run the solution into the conical flask until the solution turns orange.
Then record the reading on the burette. Swirl the mixture in the conical flask around as the solution runs from the burette into the conical flask.
This will ensure that the two solutions mix together completely.
This rough reading will give an approximate value at which the endpoint occurs. In the following titrations run the burette into the flask until it is approximately 2-3cm3 less than the rough titre. Then run it through slowly until you reach the endpoint. After each titre remove the solution in the conical flask and if necessary refill the burette. At the beginning and end of every titration record the value on the burette. Always make sure that there is enough solution in the burette to carry out the experiment.
The experiment will be repeated until 3 titres that agree to within 0.1cm3 of each other are recorded. The average of these 3 results will then be taken for use in later calculations. This will provide an accurate result and will rule out any anomalous results that might have affected the experiment.
Making the Standard Solution
First a solution of known concentration will be prepared so that the concentration of the unknown solution can be found. This will involve dissolving the solid anhydrous sodium carbonate in an amount of distilled water. The mass of the solid that is dissolved has to be correctly calculated to create a specific concentration of solution.
The solution used to find out the concentration of the sulphuric acid must be in the same range as the concentration of the sulphuric acid as otherwise the solution would be neutralised too quickly or too slowly. Therefore the sodium carbonate solution must have a concentration between 0.5 and 0.15moldm-3. A solution of concentration 0.1moldm-3 will be used as this is the average of the range of values given for the concentration of the sulphuric acid.
The stochiometric equation of the reaction is:
Na2C)3 + H2SO4 Na2SO4 +CO2 +H2O
1 mole + 1 mole 1 mole + 1mole + 1mole
The stochiometric equation is found out as follow:
The ratio of moles of sodium carbonate to sulphuric acid is 1:1, as one mole of each substance combines together to give the products shown.
It can be seen from the formula that it takes one mole of sodium carbonate to neutralise 1 mole sulphuric acid. The relative molecular mass of sodium carbonate is 106 and the relative molecular mass of sulphuric acid is 98 grams, therefore it would take 106 grams of sodium carbonate to neutralise 98 grams of sulphuric acid if their molar concentration was equal. As the exact concentration of sulphuric acid is not known an accurate prediction cannot be made. But the concentrations of the solutions are similar, therefore it should take roughly an equal amount of each substance in order that neutralisation occurs. This helps to determine the volume of the sample of sodium carbonate to be used. A 25cm3 volume of sodium carbonate would in theory produce a reading on the burette of a roughly similar amount. This would be a good reading as the percentage error is not that high and the titration would not take too long. A volume less than 25cm3 would increase the percentage error and would run through the burette too quickly. A volume more than 25cm3 would be difficult to control the swirling of and would take too long. The percentage error however would be reduced but the mixing of the two solutions would be harder. Therefore a solution of volume 2hcm3 is the volume that will be used.
To find the mass of sodium carbonate needed to create a solution of 0.1molar the below equations have to be used.
Mass of substance = Molar needed x RMM
Therefore:
Mass = 0.1 x 106
= 10.6
Therefore the mass of sodium carbonate needed to create a solution of 0.1moldm-3 is 10.6g.
But 1dm3 is equal to 1000cm3 and this volume is not necessarily needed. The amount of titrations performed is determined by when three concordant results are recorded plus the rough titre. Therefore more than 4 samples of sodium carbonate must be available, this number should be doubled to take into account possible errors. There should then be some solution available in case three concordant results are not found with the first two readings. Therefore a solution with the capacity to remove 10 samples should be made,
10 x 25cm3 – 250cm3
Therefore,
Mass needed = 10.6 x 250/1000
= 2.65g
This mass will be measured out accurately to the nearest 0.01 g on a top-pan balance in a weighing bottle as this will provide an accurate value for the mass. Time should not be wasted in obtaining an exact measurement, instead record the mass that can easily be obtained .The value for the mass will be taken three times or until the balance produces three consecutive accurate readings to ensure that the correct amount of substance is obtained. Then if the mass is not exactly 2.65g perform the following calculation to ensure that the accurate number of moles is recorded and used in subsequent calculation.
0.100 mol dm-3 x recorded balance = ~0.09 –1.01mol dm-3
2.65
The substance will then be dissolved in 100g of distilled water in a beaker. The solution will be stirred constantly with a glass rod until a clear and colourless solution appears. The solution should the be poured into a 250am3 volumetric flask using a small funnel, then add distilled water until it is approximately 1cm3 below the graduation line, then carefully using a pipette add distilled waster until the meniscus is level with the graduation mark. Then put a bung in the top and invert several times to ensure that the solution is completely amalgamated. Then label the flask with the name and concentration of the solution.
The amount of moles in a sample needs to be calculated so that the concentration of the sulphuric acid can be found out at a later point.
The number of moles used in each sample can be calculated as follows:
Moles = Mass
RMM
The mass of each sample is not 2.65 as this is the mass used for the 250cm3 of solution, instead look at the amount of sodium carbonate in 1000cm3 and then divide to reach the mass for 25 m3,
10.6 x 25/1000 = 0.265g
Therefore the number of moles used in each sample will be,
Mass = 0.265 = 0.0025moles
RMM 106
As the ratio of mole in the reaction between the sodium carbonate and the sulphuric acid is 1:1, it should take 0.0025 moles of sulphuric acid to neutralise the sodium carbonate.
The final concentration of the sulphuric acid can then be calculated as follows:
Titre = X (volume of sulphuric acid needed in cm3)
25cm3 of Na2CO3 of 0.01M (0.01M means 0.01 mol in one 1dm3 (1000cm3))
Xcm3 = (X/1000)dm3
Moles of NA2CO3 = 0.0025
1:1 mole ratio, therefore,
0.0025 moles of Sulphuric acid in Xcm3 of solution, therefore,
concentration = moles
volume
concentration = 0.0025
(X/1000)
Safety Precautions
Some of the substances that will be used during this experiment might be hazardous to health. It is important to discover what problems the substances may cause and how to deal with spillages and their disposal.
Sulphuric Acid
Sulphuric acid causes severe burns, solutions with a concentration high than 1.5M are defined as corrosive but solutions between 0.5M and 1.5M should be labelled as an irritant. Therefore it is safe to assume that sulphuric acid with a concentration less than 0.5M offers no serious health problems but still should be handled with care. It should not be fumed and is dangerous with the following substances:
- Water
- Hydrochloric acid
- Chloride
- Chlorates
- Manganates
- Reactive Metals e.g. – potassium
None of these substances are being used directly with the sulphuric acid in this experiment.
Protection:
- Safety Goggles
- Overall
- Gloves (if spilt)
Disposal - Dilute cautiously with water and add anhydrous sodium carbonate.
If spilt wear eye protection and gloves, cover with mineral absorbent, add anhydrous sodium carbonate and leave to react. Then add water and rinse area.
The main problems associated with sulphuric acid involve a high concentration of acid. As the concentration is between 0.05 and 0.15 no serious precautions need to be taken but care should be taken at all times.
Sodium Salts
Sodium Carbonate is an irritant, it is irritating to the eyes, skin and respiratory system.
Protection
If spilt scoop up and clean area.
Dyes and Indicators
Dyes and indicators offer a minimal hazard but avoid contact with skin as skin discolouration will occur. It may be flammable.
Disposal - Dilute with 10l of water and dispose.
General Safety
Care must be taken at all times. Safety Goggles and overalls must always be worn.
If any substance is swallowed medical attention should be sought as soon as possible. If any substance splashes in the eye the eye should be rinsed for 10 minutes and medical attention should be sought. If spilt on skin wash area with excess water and remove contaminated clothing., If a large area is effected or blistering occurs seek medical attention.
Bibliography
- Chemistry Review Volume 12, number 2, November 2002
- Cleaps Hazard Cards
- Making Standard Solution – Margaret Ferguson
- Chemical Ideas – Salters Advanced Chemistry
- Chemical Storylines – Salters Advanced Chemistry
- GCSE CGP Chemsitry Revision Guide
- http://ecommerce.hach.com/stores/hach/pdfs/wah/eng_pdf/AcidityMethylOrange_BT_Other_SHB_Eng_WAH.pdf