HSC Chemistry Guide:
Complete Guide to HSC Titration & Tips
[Draft Edition]
Welcome to your HSC Chemistry Guide for the tips and tricks for different types of titration for the NEW HSC Chemistry Course!
We will be exploring the following titration in order:
Simple Titration (Most common and most likely titration that you will be doing as practical exam)
Tips & Tricks for Simple Titration, Sample Questions on the Day & Answers
Back Titration
Redox Titration
Conductometric Titration
Precipitation Titration
Without further ado, let’s just jump straight in!
Typical HSC Chemistry Simple Titration
You may be asking, what is the mass primary standard do I need to weigh on the day and what volumetric flask should I pick on the day?
Well, there are two common situations that you will encounter.
Perhaps, you are asked on the day that you are to prepare 0.11M of Na2CO3 standardised solution of 250mL. So that you can use it to determine the the unknown concentration of nitric acid via titration. Well, how much sodium carbonate do you need? Well, let’s do the calculation.
c = n/V
n (sodium carbonate) = c x V = 0.11 x 250/1000 = 0.0275
We know that the molar Mass of Sodium Carbonate = ~106 g/mol.
n = m/MM
So, that m = n x MM = 0.0275 x 106 = 2.915 grams.
So, you will be weighing approximately 2.900 grams (or whatever your exact mass is).
Then from there, you recalculate your actual concentration of standardised sodium carbonate which should be close to 0.11M. It does not to be exact as your teacher really wants the concentration of your standardised sodium carbonate to be approximately 0.11M.
Of course, you choosing and using a 250mL volumetric flask for this experiment on the day. You may be using a different volumetric flask (e.g. 150mL) if you are expected to prepare a 150mL volumetric flask.
Alternatively, your teacher on the day may demand you to weigh 2.9 grams of sodium carbonate (or another primary) and obtain a concentrated of 0.11M.
Well, in that case, you will know that you need to choose the 250mL flask. Why?
Well, n (sodium carbonate) = m / MM = 2.9 / 106 = ~0.0273 moles
c = n/V
So, V = n / c = 0.0273 / 0.11 = 0.247 litres = ~ 250mL.
Step 1: Weigh 2.900 grams of anhydrous sodium carbonate in a clean, water-free beaker. [1]
Step 2: Dissolve the sodium carbonate solid with as minimal water as possible, about 50 ml. [2]
Step 3: Transfer the sodium carbonate solution in the beaker into an clean, empty 250ml volumetric flask. [3]
Step 4: Use the wash bottle with distilled water to dissolve and transfer remaining sodium carbonate in the beaker into the volumetric flask. [4]
Step 5: Using a wash bottle, add distilled water to fill up the standard solution’s meniscus up to the graduation mark. Use a dropper to add the final drops of water to fill standard solution’s meniscus to graduation mark. [5]
Step 6: Invert and swirl the volumetric flask five times to ensure even mixing of the now-standardised solution. [6]
Step 7: Rinse the 25ml pipette with distilled water followed by small quantities of standardised solution. [7]
Step 8: Rinse your burette with distilled water followed by small quantities of your titrant (the solution you are adding into the burette, in this case, it is the nitric acid of unknown concentration).
Step 9: Insert your 25ml pipette into your standardised solution to obtain a 25ml aliquot (known volume) of standardised solution (sodium carbonate) where the meniscus touches the 25ml pipette graduation mark.
Step 10: Transfer your 25ml of sodium carbonate solution into a clean, conical flask. [8]
Step 11: Attach your burette to your clamp on retort stand.
Step 12: Transfer approximately 70ml of HNO3 of unknown concentration as your titrant into your burette and record the initial titrant volume that you added into burette (if you added exactly 70ml, then your initiate titrant will be 70ml). Ensure that there is acid and no air bubbles in the space below your burette’s tap. [9]
Step 13: Attach your burette to the clamps connected the retort stand.
Step 14: Add three drops of the appropriate indicator into your conical flask and lower your burette’s tip so that it is just inside the conical flask. [10]
Step 15: Open the stopcock or tap to allow the titrant to react with the solution in the conical flask while swirling the conical flask with your other hand. [11]
Step 16: Reduce the flow of the titrant into the conical flask when you begin to see a colour change. [12]
Step 17: Stop adding titrant when there is permanent colour change and calculate your titre (volume of titrant used)
Step 18: Discard your salt solution in the conical flask to an empty beaker and wash your conical flask thoroughly with distilled water (tap water is fine though). [13]
Step 19: Repeat steps 9 – 15 to obtain three sets titre results that are + or – 0.01ml difference apart and calculate the average titre.
Step 20: Calculate the unknown concentration of your acid or base. In our example, it will be the unknown concentration of HNO3.
Notes To Above Titration Procedure
[1] An anhydrous substance is one that does not absorb water from the air, increasing the mass of your weighed primary standard. Therefore, some of your primary’s mass will be due to water and eventually resulting in your calculated concentration of your unknown acid/base to be higher than in reality.
[2] You will be transferring all this solution into a volumetric flask. If you fill up the beaker with too much water, you may exceed the graduation mark in the volumetric flask which requires you to remake your standard solution.
[3] Different size of volumetric flask allows you to prepare different volumes of your stand solution. You will be given one volumetric flask on the day of your titration. Feel free to rinse the volumetric flask with distilled water to remove any particles that may be present. You can have distilled water remaining in the volumetric flask as you will be adding water after adding your standard solution to fill up to the graduation mark.
[4] You may use a stirring rod to aid the dissolution process and the transfer of standard solution into the volumetric flask. Remember to however transfer any rinsed solution on the stirring rod to the volumetric flask. This is because the stirring rod may contain some of your standard which may be loss if not transferred.
[5] Make sure that your eye level is on the same level as the graduation mark when comparing the meniscus and the graduation mark. If not, you may be adding more or less distilled water than you need. This will affect the concentration of your standardised solution and therefore affect your results of the unknown concentration of acid or base that you want to determine.
[6] I would mix it five times just to give myself certainty that the standardised solution is mixed thoroughly enough. You do not want your aliquot in each titration to have different concentrations affecting your average titre.
[7] Clean your pipette and burette with distilled water to remove any tiny solids (such as dust) that may be present. Afterward rinsing with water, you need to rinsing the pipette with small amounts of the solution that you are pipetting and, for the burette, rinse it with your titrant. This is to ensure that the water you rinsed the instrument is not present in your instrument so it not dilute your solutions. If you don’t rinse you instrument, there may be water present and that will lower the actual volume of solution that you used in titration as you added few mL of water instead. When rinsing, do not use large amounts as you will run out of your solution! You will have limited amount of standard and titrant on the day. You need to do the titration at least 3 times to get three sets of similar results!
[8] Clean your conical flask with distilled water to remove any particles that may be present. You can leave your conical flask wet, this is because your moles of aliquot solution is known. It does not change with volume because the acid and base reaction is dependent on moles of H+ and OH-. This is why it is important to ensure you rinse your pipette and burette with the solution you are adding into them after rinsing with water. Pour out excess water though lol, don’t your aliquot with like large volumes of water in the conical flask since you will be adding your titrant in the conical flask too. It makes swirling easier too with less volume of total solution in the conical flask which helps with obtaining accurate results.
Since you know both the volume of the known solution and unknown solution that will be used at the end of titration, it does not matter whether your put your known solution in the burette or conical flask.
[9] Sometimes there may be dust trapped in the tip that hinder your titrant from leaving the burette. Sometimes, there may be air bubbles trapped in the tip. Either way, release the tap and leave some of the titrant flow out of the burette into a separate beaker to release the air bubbles. Sometimes, it may be hard for you to remove trapped dust or solids in the tip. In that case, let your lab instructor know and go grab a different burette on the day.
[10] If your burette is too high up (from your conical flask) there may be splashes so lowering it down so that it’s just inside the conical flask is good distance.
[11] Make sure you are actively swirling the conical flask throughout the titration process to ensure even mixing or neutralisation of the acid and base.
[12] Once you see a new colour appearing (usually a dot in the centre of your solution), start to reduce the rate at which you are adding into the conical flask by adjusting the stopcock or tap of the burette. While doing this, continue swirling because sometimes you may thing it is a permanent colour change but then after some swirling the new colour fades away. This is because the temporary colour change is due to uneven mixing of the acid and bases and there are still free H+ or OH- ions (depending if your solution in conical flask is an acid or base) in solution.
[13] After you have discarded the salt solution into a separate beaker. Flush your conical flask with water all the way to the top and then pour it out into the sink and do this for at least three times. This will ensure that there will not be any remaining acids or base solutions in your conical flask that will affect your subsequent titration’s results. Yes, if your titration involves a strong acid and a strong base, you can just discard the salt solution into the sink as pH = 7. If not (e.g. strong acid with weak base), transfer it to a separate beaker and dispose it in a waste container that your teacher will have in place for you at the end of your titration session.
Calculating the unknown concentration of HNO3 from our example
Na2CO3(aq) + 2HNO3(aq) -> 2NaNO3(aq) + CO2(aq) + H2O(l)
Suppose we weighted 2.9 grams of Na2CO3 and dissolved it in 250mL volumetric flask. That means that the concentration of Na2CO3 is equal to c = n/V = 2.9 / (250/1000) = 0.1094 .. moles per litre (M)
Suppose that we used a 25mL pipette on the day. This would mean that the volume of aliquot (sodium carbonate) is 25.00mL.
Now suppose that we calculated the average titre of HNO3 (volume of HNO3) is 22.38mL.
Remember that you calculate your average titre by averaging your three titre that is +/- 0.01mL apart.
The mole ratio of Na2CO3: HNO3 is 1:2.
This means that we need twice as much moles of HNO3 to neutralise all of the moles of Na2CO3 added into the conical flask.
We worked at that our concentration of sodium carbonate is 0.1094 M. We used 25.00mL pipette to transfer 25.00mL of sodium carbonate.
n(Na2CO3) in conical flask = c x V = 0.1094M x (25/1000) = 0.00274 moles.
Since we need twice as HNO3 to neutralise 0.00274 moles of sodium carbonate, we therefore need 0.00547 moles of HNO3.
We know that the average titre of HNO3 is 22.38mL.
[HNO3] = n/V = 0.00547 / (22.38/1000mL) = 0.24M.
When should you stop titrating?
You should stop titrating, or adding in the titrant, when you observe a permanent colour change.
This is called the end point.
The end point is the pH range in which the indicator that you selected changes its colour. As we have touched on in Week 5’s notes, different indicators changes colour over a different pH range which means that different indicators have different end points.
The equivalence point is the point at which the total number of moles of H+ ions is EQUAL to the total number of moles of OH– ions being reacted to form pure water.
This means that if you have an acid in your conical flask, the equivalence point is the point where ALL of the hydrogen ions in the conical flask has been neutralised by OH– ions that you adde from the burette.
The vice versa is true if you have a base in your conical flask. That is, all of your OH– ions in the conical flask has been neutralised by the H+ ions that you added from your burette.
What indicator should you use?
You need to select the right indicator by closely matching the pH range in which your indicator changes over with the equivalence point of your acid-base titration.
For example: Suppose your acid-base titration have an equivalence point when the solution has a pH = 7.
The best choice of indicator for such titration would be an indicator that changes colour at pH = 7 (i.e. a titration that has an ‘endpoint’ at pH = 7).
For example, suppose we have an indicator that is blue in solution if the pH is 0 – 6. The indicator also turns yellow when the pH is 7 – 14.
This means that when the indicator changes from blue to yellow, you will know that the equivalence point has been reached.
However, at school, you will mostly likely NOT get such an accurate indicator that has an endpoint exactly equal to the equivalence point. That is, you are unlikely to be given an indicator that has such a sharp ‘endpoint’, i.e. an indicator that changes over colour over a very narrow pH range as we have seen above.
So what do you do?
Well, in this case, the alternative choice of indicator is that the indicator should change colour over a pH range that at least includes the equivalence point.
For example, again, if our equivalence point has a pH = 7, then an alternative suitable indicator could be one that changes colour over the pH range of 6 – 8.
Perhaps, the indicator may be blue when the pH of the solution is 0 – 6 and changes its colour to red when the pH of the solution is 8 – 14 (or any other colours xD).
In this case, if the colour turns permanently red, we know that our equivalence point (where pH = 7) has been reached.
We have already explored why it is OKAY to select an indicator that changes colour over the pH 6 – 8 when the equivalence point is at pH = 7 in Learning Objective #2 in Week 7 Notes.
If you are not unsure of the reason, you can re-visit Learning Objective #2 in Week 7 Notes.
Titration Tips before AND on Practical Day!
Titration Tip #1 – Perform all the steps that we have explored at the beginning of this guide.
Titration Tip #2 – FLUSH the conical flask with tap water, shake and pour water out. Repeat this three times before repeating your titration to get new titre value.
I CANNOT stress enough how much of difference this makes in the reliability of your results!
As Donald Trump might say, “It’s HUGE!” Sorry, I had to do it!
This means that instead of doing 5 titration to get five sets of titre values because they are outside the +/- 0.01mL range, you just need to do the minimal (i.e. 3) to get your average titre.
This saves you time to do the calculation and less stress on the day!
It only takes you like 30 seconds to flush the conical flask three times with water that’s nothing compared to having to do one or more extra titration!
Titration Tip #3 – Choose the right indicator! You may NEVER see the endpoint if you chose the indicator!
Review the sections in Week 7 Notes labelled as :
‘Selecting the right indicator to use in your titration experiment’ section
‘List of common indicators used in HSC Indicators’ section
Pause, think and choose your indicator wisely on the day!
Titration Tip #4 – Practice the titration techniques in class! It is highly likely that your teacher will allow the class to do a practice titration before your practical titration exam.
Make good use of that class time!
Lastly, Do the titration. Believe in yourself. Conquer the task.
Jump to Week 7 Notes for Indicator Information (If you need to revise!)
Areas of Titration Questions with Sample Answers
Typical Simple Titration Questions
1) Evaluate the suitability of Substance “X” as a Primary Standard. Explain how using hygroscopic NaOH affect the concentration of the unknown?
2) Assess ways to improve the accuracy of the titration
3) Sources of error in titration
4) Reasons why there is deviation with other classmates’ results in terms of titre or calculated concentration of the unknown substance (i.e. unknown acid or base).
5) Constructing a titration curve for your titration (or some other titration e.g. weak acid + strong base). This curve will typically be a generic shape and not to scale. However, you are given extra information on the pH at different volume of titrant added on the day, then you can plot a more accurate curve.
6) There will not be calculation questions involving the pH at equivalence point here, as strong acid and strong base titrations will always have an equivalence point equal to pH = 7. Also no ICE table is required as the conjugate base and acid of strong acid and strong base respectively are too weak to interact and react with water molecules to shift the pH of solution away from 7. We have formally discussed why this is the case in Week 7 Notes. So if you have forgotten, feel free to revise for that.
Harder Titration Questions:
NOTE: If you have learnt about Ka and pKa, your teacher could ask for harder titration questions.
For example:
Determining the location of the maximum buffer region.
Determining pKa, Ka, pKb and Kb using titration curve.
If your titration involves a weak acid or weak base and the question requires you to calculate the pH of the resulting solution at endpoint, you will be expected to use ICE table.
We have discussed the pH of the solution at the equivalence point for various acid/base reactions in Week 7 Notes. However, below is a recap:
Equivalence Point for titration involving strong acid and strong base occurs when pH of solution = 7.
Equivalence Point for titration involving strong acid and weak base occurs when pH of solution less than 7.
Equivalence Point for titration involving strong base and weak acid occurs when pH of solution more than 7.
Equivalence Point for titration involving weak base and weak acid could occur at various pH. That is, equal, greater or less than 7.
If the Ka of weak acid > Kb of weak base, then equivalence point occurs at pH < 7.
If the Ka of the weak acid = Kb of weak base then equivalence point occurs at pH = 7.
If the Ka of weak acid < Kb of weak base then the equivalence point occurs at pH > 7.
Other harder titration question include the following:
1) Why is it difficult to perform weak acid – weak base titrations?
Sample Answer: There is no rapid change in pH at the equivalence point or for small volumes of titrant added after equivalence point. Due to this, it is hard to find an indicator that has such as a sharp endpoint (changes colour over a small pH range) that matches the equivalence point exactly for weak base – weak acid titrations. If the indicator does not have a sharp endpoint is used, a lot of excess acid/base woul be added after the equivalence point before the end-point of the indicator will be reach in weak acid – weak base titration. Therefore, the concentration of titrant would appear to be lowered than in reality.
2) Explain why the conical flask can be left wet after cleaning?
Sample Answer: The reaction occurring in the conical flask is dependent on the moles of H+ and OH- ions. As the addition of water (i.e. dilution or changes in volume) does not affect the number of moles of a species (e.g. H+ and OH- ions), the conical flask can be left wet.
3) Justify why it is important to rinse pipette with aliquot solution after water before transferring to conical flask. How will the calculation be affected if the pipette is just rinsed with water?
Sample Answer: By rinsing the pipette with appropriate aliquot solution, it ensures that the water that was used rinsed the pipette is not present in the instrument that would otherwise dilute the aliquot.This is because after rinsing with water, there may be water present and that will lower the actual volume of solution that is used in titration as the actual volume of solution added is less than actual. Therefore, LESS volume of titrant is required (lower titre reading) to neutralise the aliquot solution. Therefore, the concentration of the titrant will be higher than in reality and the concentration of the aliquot will be lower than in reality.
4) Justify why it is important to rinse burette with titrant solution after water before performing the titration. How will the calculation be affected if the pipette is just rinsed with water?
Sample Answer: By rinsing the burette with the appropriate titrant solution, it ensures that the water that was used to rinse the burette is not present in the burette that would otherwise dilute the titrant. This is because after rinsing with water, there may be water present and that will lower the actual volume of solution that is used in titration as the actual volume of solution (titre) added is less than actual. Therefore, MORE volume titrant is required (higher titre reading) to neutralise the aliquot solution. Therefore, the concentration of the titrant will be lower than in reality and the concentration of the aliquot will be higher than in reality.
Back-Titration
So what if we have a very weak acid or base and we wish to determine the concentration?
Well, we can use back-titration rather than regular titration to speed up the process as an alternative for the simple titration that we explored earlier.
Let’s say we want to determine the concentration of a weak acid. We can use back-titration as an alternative to simple titration.
Essentially, you titrate a standardised strong base against a weak acid of unknown concentration. What happens is that you will have excess strong base in the conical flask.
You will then titrate the excess strong base against a standardised strong acid.
From there, you can determine the moles of excess strong base that is neutralised by the strong acid.
Here, you can subtract the moles of excess strong base from the moles present in the original moles of strong base present in the conical flask. This difference will yield the moles of strong base used to neutralise the weak acid.
By comparing the mole ratio between the strong base and weak acid, we can determine the moles of weak acid. If we divide that by the volume of weak acid used in the titration, we can determine the concentration of weak acid.
Redox Titration
Redox Titration is the same as titration whereby it is used to measure the concentration of a solution. However, the titration reaction does NOT involve a neutralisation reaction.
Rather, it is redox (reduction and oxidation) reaction.
So, you need to construct the reduction half equation and the oxidation half equation such as by using the standard reduction potential table.
From that, you can construct the overall redox reaction equation using the reduction & oxidation half equations.
All the other procedures such as washing of equipment, preparation of standardised solution that we mentioned in the simple HSC titration at the beginning of this guide applies to redox titration.
An example of a redox titration question can be found on Question 36 in the Sample HSC Chemistry Paper which you can download below.
Conductometric Titration
It is not likely that you will be performing conductometric titration as a practical exam as there is not much practical skills for students to be assessed on.
Conductometric Titration essentially uses a conductometric probe that is submerged into the reacting solution. The student can then observe the change in conductivity of the solution on a computer screen because the probe is attached to a data logger which is attached to computer.
The calculation in conductometric titration will be same as a simple HSC Chemistry titration that we mentioned at the beginning of this guide.
Conductivity titration can be performed for the following:
Strong Acid – Strong Base
Strong Acid – Weak Base
Weak Acid – Strong Base
Weak Acid – Weak Base
To see the change in conductivity of as titration progresses to reach equivalence point and beyond, please look at Week 7 Notes and go to ‘Conductometric Titration’ section.
We have also covered the explanation for the reason to the change in conductivity as titration progresses for each of the four categories of conductometric titration.
View Conductivity Titration Graphs
Precipitation Titration
We have explored Precipitation Titration in Week 14 Notes, which covers the first Inquiry Question of Module 8.
You can read through the procedure of precipitation titration there.
All the other procedures such as washing of equipment and preparation of standardised solution that we mention at the beginning of this guide applies here.