Titrating Sodium hydroxide with an unknown molarity, against hydrochloric acid to find its' molarity.
Titrating Sodium hydroxide with an unknown molarity, against hydrochloric acid to find its' molarity.
P2aP2aAim: - To find out the concentration of sulphuric (VI) acid by performing a titration with sodium carbonate solution.
Introduction
Sodium carbonate is an alkali, meaning that it has basic properties and has a pH that is greater than pH7, which is neutral. All pHs that are lower than pH7 are considered to be acidic. The pH scale runs from pH1, to pH14.
A base is a solution/substance, which receives hydrogen ions (H+, also referred to as a proton) willingly, whereas acids are proton donators, this means that they give hydrogen ions off when in solution. The more basic something is, the more hydrogen ions it can receive and the more acidic something is, the more hydrogen ions it gives out, pH is a measure of this.
In this particular titration, we will be neutralising the acid (H2SO4) with an alkali (Na2CO3), which is a carbonate. When an acid and a carbonate are reacted, the resulting products are a salt, carbon dioxide and water. These equations show what reactants and products are involved in this experiment: excellent shows you are thinking coherently about the problem evidence for P11a
H2SO4 (?q) + Na2CO3 (?q) --> Na2SO4 (?q) + H2O (?) + CO3 (g)
mole + 1 mole --> 1 mole + 1 mole + 1 mole
Hydrochloric acid (?q) + Sodium hydroxide (?q) --> Sodium chloride (?q) + Water (l)
The acid (H2SO4, in this case) will release protons when the base (Na2CO3) is added to it, which will receive them. Once the solution has been neutralised, it will no longer be able to receive any more protons. We can test the pH of solutions by using various different types of indicators. Good chemical knowledge P11a
Methyl orange is the indicator that will be used within this experiment. When methyl orange is present in acidic solutions, it is red; in basic solutions it is yellow. When the concentrations of two solutions are equal, methyl orange is orange; a mixture of the two colours. When methyl orange just changes from orange to red, this is when we know that we have neutralised the solutions, this is the end point. We know the end point of the titration, as the solution within the conical flask will just turn a light pink. The end point of the titration is the visible colour change that can be seen. The calculated end point is known as the stoichiometric point, this is the actual volume of acid, which is used in order to neutralise the alkali. This should be 25cm3 of the acid as we are using this volume of alkali within the conical flask.
In order to find out what the concentration of the sulphuric acid is, I need to have a standard solution. A standard solution is a solution, good P11a that has been made, the concentration of which is known. As I know the concentration of the standard solution, I can work out what the concentration of the sulphuric acid with the unknown concentration is.
A primary standard needs to meet certain specification. They are: good P11a
* It must be available in a highly pure state
* It must be stable in air at ordinary temperatures
* It must be easily soluble in water
* It must have a high molar mass
* It must be in solution when used in volumetric analysis , must undergo complete and rapid reaction
The concentration of the sulphuric acid is said to be between 0.15moldm-3 and 0.05moldm-3. In order to make up the standard solution, I need to estimate the concentration of the sulphuric acid that we are using. To do this I need to take the mean of the concentrations that we are told the acid could be.
0.05moldm-3 + 0.15moldm-3 = 0.10moldm-3
2
good P11 a and the calculation part is P11b
Therefore, I need make up a standard solution of the sodium carbonate solution with a concentration of 0.10moldm-3. I will need to make up this solution with the appropriate concentration of sodium carbonate solution, so that approximately 25cm3 of the base reacts will react with 25cm3 of the sulphuric acid.
To do this I need to calculate the concentration of sulphuric acid, I know that the concentration of the acid is roughly 0.10moldm-3. We must assume that it is this number as we do not know otherwise. P11b
From the equation for this reaction above, I know that 1 mole of sodium carbonate will react with 1 mole of sulphuric acid. There is an equation that links moles, volume and concentration.
moles = volume x concentration
= 0.025dm3 x 0.10moldm3
= 0.0025 moles of sulphuric acid
Therefore 0.025dm3 of sodium carbonate solution is equivalent to 0.0025 moles of sulphuric acid. We know that the sodium carbonate solution needs to be 0.10 moldm-3 in order to neutralise the sulphuric acid.
Therefore to neutralise 25cm3 of 0.10moldm-3 sulphuric acid I need to use 0.0025 moles of sodium carbonate solution, I can now work out the mass of anhydrous sodium carbonate that I will need to use in order to make up the standard solution. There is an equation that links together moles, the mass and the molar mass of the substance.
Moles = mass
molar mass
0.0025 = mass
(2x23+12+3x16)
0.0025 = mass
106g
Mass = 0.0025x106
= 0.265g
Therefore 0.265g of solid anhydrous sodium carbonate would be needed to be diluted with distilled water to produce 0.025dm3 of sodium carbonate solution.
However, we need to make up 0.25dm3 of sodium carbonate solution; this means that in order to find the correct mass of anhydrous sodium carbonate that needs to be used, so I therefore need to multiply the above answer by ten.
Mass = 0.265 x 10
= 2.65g
As the top pan balance that I have access to only reads to 0.01g, I will need to round the mass up so that it will be possible to actually weigh out this amount. This will mean that my results will not be as precise as they should be. However, as this is the most precise weighing equipment there is, I must use it. Good point !
The reaction between the anhydrous sodium carbonate and the distilled water will be an exothermic one. This is as energy is released as the anhydrous solid is being hydrated by the water molecules; most of this energy is given out in the form of heat energy, which is why it is an exothermic reaction. The energy that is used within this system -24.6 ?H/kJmol-1 ?, as the number of a minus figure, we can tell that the energy is being given out and not taken in.
? Chemical Data Book by JG Stark and HG Wallace. Second ...
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The reaction between the anhydrous sodium carbonate and the distilled water will be an exothermic one. This is as energy is released as the anhydrous solid is being hydrated by the water molecules; most of this energy is given out in the form of heat energy, which is why it is an exothermic reaction. The energy that is used within this system -24.6 ?H/kJmol-1 ?, as the number of a minus figure, we can tell that the energy is being given out and not taken in.
? Chemical Data Book by JG Stark and HG Wallace. Second edition in SI, reprinted in 2003, published by John Murray.
Therefore, the sodium carbonate solution will need to be left to cool to room temperature before it is used within the titration. This is as rate of reaction is affected by temperature, the higher the temperature, the faster the rate of reaction. If the solution was to be used straight away, then the first titrations would need less acid in order to neutralise the alkali. This would mean that the results I obtain would be incorrect.
Equipment
* 2.65g solid, anhydrous sodium carbonate
* 50cm3 Sulphuric acid
* 250 cm ³ conical flask
* Grade B, 50 cm³ Burette excellent detail
* Grade B, 25 cm ³ pipette
* Pipette filler
* Dropping pipette
* Small filter funnel
* 100 cm³ beaker
* Wash bottle and distilled water
* Eye protection
* Boss, clamp and stand
* White tile
* White card
* Methyl Orange indicator
* Spatula
* Weighing bottle
* Glass rod
* Top pan balance excellent _2b
In this experiment we will be using 0.10M sulphuric acid and 0.10M sodium carbonate solution, which are dangerous chemicals.
Sodium carbonate, irritant, wear eye protection/protective gloves when weighing solid1
Sulphuric acid, very corrosive, wear eye protection2
As, when using any corrosive of dangerous chemicals, lab coats should be worn in order to protect clothing and skin. Care should be taken when handling these chemicals and any spillages should be dealt with immediately.
Care should also be taken to avoid inhaling any fumes that may be released by the reacting chemicals as they may cause damage to the respiratory system.
Eye protection must be worn as contact of these chemicals with the eyes could cause severe damage.
Accuracy and precision of equipment used:
* 25cm³ (grade B) graduated pipette: this reads correct to the nearest 0.5cm³
* Top pan balance: this reads to the nearest 0.01g
* 50cm3 (grade B) burette: this reads correct to the nearest 0.5cm3
Method
To ensure that all equipment is clean, first wash them all out with distilled water, and then wash out those that will contain chemicals in them, with the chemicals that will be put in them.
. Clamp the burette to the stand, making sure that the burette is parallel to the stand as if it is not, this will affect where the meniscus is and so what readings are obtained from the burette. Ensure that the tap on the burette is turned to the off position.
2. Using a weighing bottle, weigh out 1.33g of anhydrous sodium carbonate (making sure to zero the balances before placing the Na2CO3 in the weighing bottle to ensure that the correct mass is achieved).
3. Place the anhydrous sodium carbonate into the 100cm3 beaker, washing out the weighing bottle with distilled water and transferring it into the beaker. Do this at least twice (to ensure that all of the solid is used so that I will have the correct molarity of sodium carbonate solution).
4. Using the glass rod and a little more distilled water ensuring that all of the Na2CO3 is dissolved.
5. Place the sodium carbonate solution into the 250cm3 graduated flask, washing out the beaker twice with distilled water and transferring the contents into the flask (ensuring that the molarity of the solution will be correct and therefore the results will be precise).
6. Fill the flask with distilled water until the meniscus of the solution almost reaches the line to indicate that there is 250cm3 of solution within the flask.
7. Ensuring that the meniscus of the solution is at eye level, use the dropping pipette to fill the flask with distilled water so that the bottom of the meniscus of the solution touches the line (ensures that there will be the accurate amount of solution within the flask with the correct molarity, therefore ensuring precision of results obtained).
8. Ensuring that there is a lid on the graduated flask, shake the bottle. This mixes the solution to guarantee that it is a homogenous standard solution (same concentration throughout). This will help to make the results more precise. Then leave the solution for a while before use to allow for cooling.
9. Using the filter funnel and a clean, dry 100 cm³ beaker, pour some sulphuric acid into the burette until it is nearly 0 on the scale. Remove the filter funnel after use (ensures
from Hazcards #61
2 from Hazcards #98
that no extra liquid from the funnel is able to drip into the burette, so the accurate volume of acid is in the burette).
0. Run a little of the solution out to ensure that the jet is full.
1. Using a dropping pipette; pipette some hydrochloric acid into the burette until it is around 0. Using the piece of white card, make sure that the bottom of the meniscus of the solution is touching the line at 0. This is to ensure that the correct volume of H2S04 is used and so the results are precise.
2. Place the white tile onto the bottom of the stand (so it is easier to see the colour change in the solution later).
3. Fill the 25 cm ³ pipette, using the pipette filler, with sodium carbonate solution, then pour into the conical flask and place onto the white tile.
4. Place a couple of drops of methyl orange indicator into the sodium carbonate solution (not much more, otherwise it will be too hard to see the slight colour change later, which aids in ensuring that the results are as precise as possible) and swirl the contents of the flask
5. Record the volume reading on the burette before starting the titration.
6. Add the sulphuric acid in small quantities, swirling the flask after each addition. Once you begin to near the end point, add the acid, drip by drip. The yellow colour of the solution will turn orange, then pink once the sodium carbonate solution has been neutralized.
7. Record the final burette reading and calculate the volume of solution that you have run out into the flask.
8. Repeat the titration process several more times, until you have three volumes which agree within 0.10cm³.
Sources of reference used
* Titrimetric analysis for A & S levels (SI Edition), J G Stark, published by John Murray
* Chemical Ideas (Second edition), Salters Advanced Chemistry, published by Heinemann
* Hazcards publisher?
* Chemistry Review, December 2001.
P10/11 excellent start
Analysis
Table to show the starting points, end points and titre for the titrations perfomed.
Start (cm3)
End (cm3)
Titre (cm3)
Rough
50
26.70
23.30
50
26.75
23.25
2
50
26.75
23.25
3
50
27.70
23.30
To ensure that my results were as accurate and as precise as possible, I needed to have 3 titrations that were within 0.10cm3 of one another. I managed to obtain these results within the first three titrations that I did, however, I decided to perform another titration just to make sure that these results were accurate. They were as they were all within 0.05cm3 of one another.
I need to find the mean (average) titre in order to find out the average amount of sulphuric acid that was needed to neutralise the sodium carbonate solution. To do this I simply need to find the sum of the titres used and then divide this number by the number of titrations performed.
Mean titre = 23.30 + 23.25 + 23.25 + 23.30
4
= 93.1
4
= 23.275 cm3 H2SO4
The concentration of sodium carbonate solution that was used in the titration can be worked out by using two equations:
. Number of moles = mass (g) ÷ Relative Molecular Mass (gmol-1)
2. Number of moles = volume of solution (dm3) x concentration of solution (moldm-3)
I know that I used 2.65g of anhydrous sodium carbonate in order to make up the standard solution of sodium carbonate. From this knowledge I can work out the number of moles of sodium carbonate that is in 2.65g of anhydrous sodium carbonate. To do this, I can use the first equation.
Number of moles = 2.65g ÷ 106gmol-1
= 0.025 moles of Na2CO3
However, the mass that was used was 2.65g; this was because I made up a solution that would be enough for ten titrations that were using 25cm3 of sodium carbonate solution. Therefore I need to divide the figure I obtained above by ten in order to find the correct number of moles of sodium carbonate that was used.
Number of moles = 0.025 ÷ 10
= 0.0025 moles of Na2CO3
To calculate the concentration of the sodium carbonate solution that was used I need to use the second equation. However, it needs to be rearranged before I can use it:
Concentration of solution (moldm-3) = number of moles ÷ volume of solution (dm3)
Concentration of solution = 0.0025 ÷ 0.025
= 0.10 moldm-3
So, the concentration of the sodium carbonate solution that I made up to use in the titration was 0.10moldm-3.
H2SO4 (?q) + Na2CO3 (?q) --> Na2SO4 (?q) + H2O (?) + CO3 (g)
mole + 1 mole --> 1 mole + 1 mole + 1 mole
This equation shows what happened within the titration and from it I can see that 1 mole of sodium carbonate will react with 1 mole of sulphuric acid, therefore within the titration the number of moles that are in the sodium carbonate solution will be neutralised by the same number of moles in the sulphuric acid that is being used. Given that the molarity of the solutions is the same.
For each titration I withdrew 25cm3 of sodium carbonate solution from the graduated flask, which contained my standard solution of sodium carbonate. This means that each titration should therefore have used 25cm3 of the solution.
The number of moles that were used in the titrations can be worked out by using the following method:
1 of mass/g = 2.65g x 1 = 0.0025 moles
0 RMM/gmol-1 106gmol-1 10
From this equation I can see that the stoichiometric point of 0.0025 moles of sulphuric acid that is added to sodium carbonate should also be 0.0025 moles.
I have already calculated the concentration of the sodium carbonate solution to be 0.10moldm-3 as well as having found out the number of moles that were used (0.0025moles) within the titration. Now it is possible to work out the concentration of the sulphuric acid that was used within the titration as this was the unknown concentration. This can be done by simply using the following equation:
Concentration of solution (moldm-3) = number of moles
volume of solution (dm3)
I need to take the average titre that was used within the titrations that I performed as the volume of solution so that I can calculate the concentration of the sulphuric acid that was used. As the current volume of sulphuric acid that was used is in cm3, I need to convert this into dm3 by dividing the figure by 1000.
Concentration of solution = 0.0025
(23.275 ÷ 1000)
= 0.0025
0.023275
= 0.107 moldm-3
This shows that the concentration of the sulphuric acid used is roughly 0.10moldm-3. To check whether this was the correct molarity of the sulphuric acid, I first asked the other students within my class what molarity they calculated the acid to be. More or less the same molarity was calculated, there were differences of about 0.002moldm-3. However, I counted these types of figures as negligible.
Later, the chemistry technician, who made up the sulphuric acid, was asked what the molarity of the sulphuric acid that was used within the titration was. It was 0.10moldm-3.
Evaluation
To evaluate this experiment where I titrated a standard solution of sodium carbonate against sulphuric acid of an unknown concentration, I need to look at the uncertainties and limitations that could have occurred within this experiment.
There were no anomalies within my results, as can be seen in the table of results in the analysis section.
Accuracy
Accurate - the result is close to a reference value or average of the data is close to a reference value1.
Within this titration, I ensured that my results would be as accurate as possible by:
* Starting off with a rough titration, where I could practice all the practical procedures within the titration, ascertain the rough end point and stoichiometric end point.
* Having at least 3 results which were within 0.1cm3 of each other.
Precision
Precise - the data points are close together (there can be a random error) 1.
There were some procedures that I carried out in order to ensure that this experiment and the results that were obtained were as precise as possible:
* Ensuring that the temperature was constant throughout the titrations; as temperature affects the rates of reactions. The higher the temperature, the faster the rate of reaction, so if the temperature increased towards the end of the experiment, the end points would have been reached faster than before.
* I used distilled water to ensure that there were no other compounds or elements that could cause impurities within the sodium carbonate solution and so would disrupt the concentration of the standard solution that would be made.
* I teared the top pan balance after placing the weighing bottle on to it and before weighing out the 2.65g of anhydrous sodium carbonate into it. This is to ensure that the correct mass of solid is weighed out and used.
* The mass of the solid was weighed to two decimal places, which is the highest degree of accuracy we can get with the top pan balances that are provided for me to use.
The nature of Science, published by the Royal society for Chemistry, section 1.
* I ensured that the solid was fully dissolved by firstly adding small amounts of Na2CO3 into the distilled water and stirring the solution with the blunt end of a glass rod.
* I took care to shake the volumetric flask with the sodium carbonate solution in, to ensure that the molarity of the solution was what it should be. This is as the solute is denser than the water molecules and so may sink to the lower regions of the solution. This would mean that the molarity of the solution would not be the same throughout the graduated flask.
* All of the equipment used was washed with distilled water and the chemical that would be placed in it or dried before the titration took place. This was to ensure that the chemicals used would not be affected by any chemicals that may have been in the equipment beforehand.
* I ensured that the jet of the burette that was used was full; this was done by running some of the acid out of the burette and then filling the burette up to 0 again. This is as the liquid in the jet of the burette is part of the 50cm3 that is within the burette, if the jet was empty, the burette would not actually contain 50cm3 of acid. This would mean that the reading of acid that I obtained at the end would be incorrect.
* I ensured that the bottom of the meniscus of the acid in the burette just touched the 0 line, also in the volumetric flask and pipette; I ensured that the bottom of the meniscus of solution just touched the line. This was to ensure that the correct volume of liquids in each case was available; this would mean that the concentration of the sodium carbonate solution was correct and the volume of acid in the burette would be correct too.
* After filling the burette with acid, I removed the funnel. This is as there may still be some acid in the funnel and due to gravity, a drop or two may fall into the burette if the funnel was left in it. This would mean that the final reading of acid obtained at the end would be affected and so incorrect.
* A white tile was placed underneath the conical flask, this was to make be able to see the colour change that indicated the end point of the titration easier and more clear. Without the white tile, due to the colour of the wooden stand, it would be very hard to see the subtle colour change from orange to pink.
* Only two drops of Methyl Orange indicator was used in each of the titrations. This was because if more was added, the colour change from orange to pink would be harder to see and would occur later, due to the deeper orange colour that would occur if more than two drops were put into the conical flask.
* The volume of acid that was read off at the end of each titration was read to three significant figures, as this is the largest amount of significant figures that is able to be read off of the burette.
* I had to judge what I felt was the end point of the titration (a light pink) and with each titration I had to make sure that the end point was the same or as near to the same colour as possible.
* A new conical flask that had been washed out with distilled water and then dried was used for each titration to avoid contamination.
Uncertainties
In titrations there are uncertainties, which can occur. They are usually related to the precision of the equipment that is being used.
The percentage uncertainty (sometimes called error) can be calculated using the following equation:
Percentage uncertainty (error) = error x 100
reading
"Generally the limit of precision is taken to be half a division on either side of the smallest units on the scale you are using."1
The above statement is how you can work out the 'error' that is to be used in the equation for percentage uncertainty. The error is the difference in all the other possible readings, divided by two.
A grade B, 50cm3 burette can be read to 0.01cm3. Therefore the end volume reading can be:
0.01 ÷ 2
= ± 0.005cm3
Percentage uncertainty of burette= 0.005 x 100
23.275
= 0.021%
So there is a 0.02% uncertainty with my average volume of acid that was used within the
Salters Advanced Chemistry 2000. Activities and assessment pack, second edition, publisher - Heinemann. Page 2.
titrations. This is a very small figure though and should not have affected my results much, as the figure is so small it is safe to say that this piece of equipment is very precise.
The top pan balance can be read to 0.01g, therefore the end mass reading can be:
0.01 ÷ 2
= ± 0.005g
Percentage uncertainty of balance = 0.005 x 100
2.65
= 0.189%
The error of a grade B, 250cm3 volumetric flask is 0.2cm3 1.
Percentage uncertainty of flask = 0.2 x 100
250
= 0.080%
The error of a grade B, 25cm3 pipette, when used correctly (allowed to drain and retain the last drop), is 0.06cm3 1.
Percentage uncertainty of pipette = 0.06 x 100
25
= 0.240%
All of these percentage uncertainties on their own are not very great and so it looks as if they would not affect my end reading much and so would not affect the calculated concentration of the acid much. However, once all of these percentage uncertainties have been combined, it is possible that they will have a great affect on the end reading that I obtained. These percentage uncertainties will affect my end results by changing the concentration of the sulphuric acid used that I calculated.
Table to show percentage uncertainties of equipment used and total percentage uncertainty.
Equipment
Percentage uncertainty (%)
Grade B, 50cm3, burette
0.021
Top pan balance
0.189
Grade B, 250cm3 volumetric flask
0.080
Grade B, 25cm3 pipette
0.240
Total percentage uncertainty
0.530
The total percentage uncertainty of all of the equipment used is 0.53%, this means that the concentration of the sulphuric acid I calculated could be 0.53% higher or lower than that of the actual concentration.
To work out the possible concentrations of sulphuric acid, due to the percentage errors, I can use the following equation:
Conc. of acid ± sum of percentage uncertainties as a percentage of the conc.
0.11 ± 0.53%
= 0.11 ± (0.11 x 0.53 )
100
= 0.11 ± 0.000583
0.11 + 0.000583 or 0.11 - 0.000583
= 0.110583 = 0.109417
= 0.11 moldm-3 = 0.11moldm-3
As you can see, even with the percentage uncertainties added or taken away from the concentration of the acid that I calculated, the concentration still rounds up or down to 0.11moldm-3.
This means that the percentage uncertainties are very low and so the accuracy and precision of the equipment used is very high. It also means that the molarity of the sulphuric acid that I calculated it to be is very precise.
However, I still need to look at the reliability of my results, as they can be precise without being accurate of reliable.
Reliability
Reliability - this is assessed through comparison of an individual results with a reference or class mean1.
By looking at the concentrations of sulphuric acid that the rest of my class calculated, I can see whether or not the concentration of acid I calculated is reliable. I can also check with the chemistry technician, who made up the sulphuric acid which was used within the titrations, to see whether the concentration that I calculated is correct and so reliable.
The nature of Science, published by the Royal society for Chemistry, section 1.
Table to show the concentrations of sulphuric acid that the class calculated.
Number
Concentration (moldm-3)
0.10
2
0.11
3
0.11
4
0.11
5
0.11
6
0.11
Total
0.65
Average Concentration (moldm-3)
0.65 ÷ 6
= 0.108
= 0.11
The table shows the concentrations that were calculated by the rest of my chemistry class as well as the average concentration. As you can see the average concentration is 0.11moldm-3, which is the molarity that I calculated the sulphuric acid to be.
Also, the chemistry technician made up the sulphuric acid to be 0.10moldm-3, as the concentration that I calculated it to be is only 0.01moldm-3 out, I think that it is safe to say that the concentration of sulphuric acid I calculated is reliable.
Improvements
The experiment could be improved by changing some procedures and equipment, if this equipment were available, then the experiment would become more accurate and precise. As well as this the results obtained would too be more accurate and reliable.
* Using a colorimeter to assess the end point of each titration. A colorimeter works by calculating the percentage of light either absorbed or allowed (transmitted) through a solution. As the colour of the indicator changes, so will the percentage of light transmitted or absorbed. In each of the titrations we could take a certain percentage as our end point, this way the readings for the end points should be more reliable.
* Using grade A burettes, volumetric flasks and pipettes. These have a lower percentage uncertainty to them as they are more precise and so more accurate. These would lower the overall percentage uncertainties within the experiment, making it more precise.
The experiment could be further extended by titrating the sulphuric acid against other alkali solutions with a molarity of 0.1moldm-3 and by using different indicators, which may give different end points to the titrations.
Hanna Cheung
