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IQ3 How are solutions of acids and bases analysed?
IQ3 How are solutions of acids and bases analysed?
IQ3 SYLLABUS STATEMENTS
3.1 conduct practical investigations to analyse the concentration of an unknown acid or base by titration
3.2 investigate titration curves and conductivity graphs to analyse data to indicate characteristic reaction profiles, for example:
– strong acid/strong base
– strong acid/weak base
– weak acid/strong base
3.3 model neutralisation of strong and weak acids and bases using a variety of media
3.4 a) explore the use of Keq for different types of chemical reactions, including but not limited to: dissociation of acids and bases
b) calculate and apply the dissociation constant (Ka) and pKa (pKa = -log10 (Ka)) to determine the difference between strong and weak acids
3.5 explore acid/base analysis techniques that are applied:
a) in industries
b) by Aboriginal and Torres Strait Islander Peoples
c) using digital probes and instruments
3.6 conduct a chemical analysis of a common household substance for its acidity or basicity (ACSCH080), for example:
– soft drink
– wine
– juice
– medicine
3.7 describe the importance of buffers in natural systems
3.8 conduct a practical investigation to prepare a buffer and demonstrate its properties
3.1 conduct practical investigations to analyse the concentration of an unknown acid or base by titration
Volumetric analysis is an important technique in chemistry. It is used to measure the volume of a solution of known concentration needed to react with a solution of unknown concentration. The technique used is called titration.
•A titration is used to determine the concentration of acid or base in a solution.
•It is a process in which a solution of known concentration (titrant) is dispensed from a burette to a known quantity of an unknown solution (analyte) until the stoichiometric equivalence point (number of moles of acid = number of moles of alkali) is reached.
•Indicators, such as phenolphthalein or methyl red, are usually added to the analyte to show when the equivalence point has been reached.
•The progress of a titration reaction is monitored by a pH or titration curve.
An important component of a titration is the solution of known concentration. This is called the standard solution. It can be determined by titration against a primary standard, or it can be prepared by adding a known weight of a reagent in a definite volume of solution to create a primary standard solution.
Prior to preparing a standard solution, a primary standard must be selected. Not all substances are suitable to be used as standards. NaOH, for example, is hygroscopic (absorbs water from the surrounding air) and thus its weight may change. However, these may be later standardised by reacting them with other known solutions.
A primary standard has important characteristics. It is a solution prepared from a solid that:
is water-soluble
has a high purity
has a known formula
is stable in air (low reactivity)
has low hygroscopicity (won't absorb water from air, which would alter weight)
is non-toxic
is cheap and readily available
One reagent often used as a primary standard in acid-base titration is sodium carbonate, for standardisation of aqueous acids: hydrochloric, sulfuric acid and nitric acid solutions (but not acetic acid)
Determine an unknown concentration by titrating it against a known concentration.
https://www.wikihow.com/Perform-a-Titration
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. (We use the conical flask since often in industry you will have smaller quantities of this than your standard solutions.)
A. The Titration
1. Clean general glassware: Rinse out all glassware you will be using, including any beakers or flasks, with deionised water. Allow to dry completely.
2. Clean the burette: To get an accurate reading from your titration, your burette must be completely clean.
Close the stopcock and fill the burette with deionised water. Swirl the water around a few times before opening the stopcock and allowing it to drain.
Repeat the rinsing process at least 3 times with water.
Perform the same rinsing process at least 2 times with your titrant (standard solution - see below for preparation).
3. Prepare the burette:
Make sure the stopcock is closed. Use a funnel to fill the burette with an excess amount of titrant. Clamp the burette carefully to a retort stand.
Place a beaker under the stopcock. Open the stopcock slightly, let the excess titrant flow out until it reaches the zero mark. [1]
4. Prepare the analyte in the conical flask:
Rinse a 25ml pipette with deionised water, followed by small quantities of analyte. [2]
Insert a pipette into the analyte (unknown concentration) solution to obtain a 25ml aliquot.
Transfer the 25ml of analyte into a clean conical flask. [3]
Place the conical flask containing the analyte under the burette, which is positioned so that the nozzle is just inside the flask. [4]
5. Monitor the titration:
Choose the appropriate indicator for the anticipated endpoint [5] and add 3-4 drops to the analyte AND/OR
Insert a pH meter into the analyte.
6. Titrating: first titration (approximating results)
Open the stopcock with one hand while carefully swirling the conical flask with the other. [6]
Allow 5mL aliquots to run through until you near the titration endpoint (at this stage you will see flashes of colour) [7].
Slow the flow to dropwise with constant stirring until you record a permanent colour change. [8]
7. Titrating: three titrations (accurate results)
Repeat the titration steps from 3 to 5.
Run the titrant through, slowing to 1 mL aliquots and then dropwise when you near the endpoint.
Notes
[1] Sometimes there may be air bubbles trapped in the tip. Releasing the tap to let some of the titrant flow out of the burette into a separate beaker will release the air bubbles. This also allows you to check the stopcock is working properly, and to get a feel for the speed of the flow.
[2] You can leave the conical flask wet because the moles of aliquot solution are known: this does not change with volume. Don’t use large volumes in the conical flask since you will be adding your titrant too. A smaller volume of total solution in the conical flask also makes swirling easier, which helps with obtaining accurate results.
[3] Do not use large amounts as you will run out of your solution: you need to do the titration at least 3 times to get three sets of similar results.
[4] If your burette is too high up from your conical flask there may be splashes; lowering it to just inside the conical flask is a good distance.
[5] The end point is the pH range in which the indicator that you selected changes its colour. Different indicators change 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. This means that if you have an acid in your conical flask, the equivalence point is the point where ALL of the hydrogen ions has been neutralised and vice versa if you have a base in your conical flask.
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 and it 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, you will mostly likely NOT get such an accurate indicator, one that has such a sharp ‘endpoint’ or such a narrow pH range over which it changes colour. The alternative choice of indicator is that it 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 a 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). In this case, if the colour turns permanently red, we know that our equivalence point (where pH = 7) has been reached. This alternative method of ‘selecting an indicator with a pH range includes the equivalence point’ is good enough and will not significantly affect titration results (for our purposes).
[6] Make sure you are actively swirling the conical flask throughout the titration process to ensure even mixing or neutralisation of the acid and base.
[7] 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.
[8] If you notice a colour change, close the stopcock and swirl for a few seconds. If the colour disappears, open the stopcock slightly and continue to add the titrant drop by drop until you get a permanent change.
[9] Reducing errors:
Different people have different perception and sensitivity to colour changes. Therefore, it is possible to misjudge when the endpoint of the indicator has been reached during titration. Place a piece of white paper underneath your conical flask to allow easier judgement of colour change.
Misreading the volume of titrant used – parallax error. The eye level must be at the same level as the meniscus to accurately determine the volume of the standardised solution that is pipetted and amount of titrant used in burette. Place a piece of white paper behind the burette or volumetric flask to allow easier comparison of the graduation mark and meniscus.
Not transferring all solid/liquid when preparing standardised solution – it may happen that some of the primary standard was still left in the beaker or stirring rod. Thoroughly wash and transfer your wash solution into the volumetric flask multiple times to ensure all of the primary standard is transferred to the volumetric flask.
B. Selecting the standard solution:
A primary standard is a substance that has fixed chemical composition, high purity, chemical stability and solubility so that it can be used to prepare a standardised solution with an accurately known concentration.
High molecular mass to minimise percentage error when measuring mass.
High solubility in deionised water so that the standard completely dissolves (calculated known concentration is accurate).
High purity so that the calculated known concentration of the standard is accurate.
Chemically stable so that it does not undergo chemical reactions with deionised water or moisture in the air.
Anhydrous - DOES NOT absorb water from the air and so it does not form bonds with water. A hydrous standard will result in lower concentration of dissociated ions when dissolved in the water in the volumetric flask.
Making the standard solution: NaOH as an example
1. Weigh calculated amount of anhydrous sodium hydroxide pellets in a clean, water-free beaker. [1]
2. Dissolve the sodium hydroxide with minimal deionised water, about 50 ml. [2]
3. Transfer the sodium hydroxide solution in the beaker into a clean, empty 250ml volumetric flask.
4. Wash the beaker with deionised water to dissolve any remaining sodium hydroxide and transfer into the volumetric flask.
5. Add deionised water to fill the standard NaOH solution’s meniscus up to the graduation mark. Use a dropper to add the final drops. [3]
6. Invert and swirl the volumetric flask five times to ensure even mixing of the now-standardised NaOH solution. [4]
Notes
[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] If you fill the beaker with too much water, you may have too much solution for the volumetric flask, which will require you to remake your standard solution.
[3] 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.
[4] Inadequate mixing can mean that your aliquot in each titration would have different concentrations.
Practical 3.1.1 Measuring the acid concentration of vinegar
PA 6.6 Determination of acid concentration of vinegar PP88-90
Introduction
Volumetric analysis is essentially the reaction of a known quantity of solution with the sample to be analysed. Volumetric analysis employs precision glassware (pipette and burette) to measure out solutions with great accuracy. (Pearson Education Australia, 2019)
Vinegar is a common household item containing acetic acid, as well as other chemicals. This experiment is designed to determine the molar concentration of acetic acid in a sample of vinegar by titrating it with a standard solution of NaOH.
CH3COOH(aq) + NaOH(aq) -> CH3COONa(aq) + H2O(l)
Adding the sodium hydroxide, a basic solution, to the acetic acid, an acidic solution, a neutralisation reaction occurs. An indicator known as phenolphthalein, is added to the vinegar. This indicator turns the solution to a dark pink when excess NaOH makes the solution more basic. When the solution turns light pink, it has successfully been neutralised.
The amount of NaOH used to standardise the vinegar can then be used to determine the amount of acetic acid in the vinegar as they are at a 1:1 ratio as seen in the above equation. The moles of NaOH used to neutralise the acid equal the number of the moles of acetic acid present in the vinegar.
Preparing the standard solution
Warning: Wear safety glasses, gloves and lab coats! NaOH is caustic. Avoid skin contact. Clean up spills immediately.
1 Prepare 250 mL of 0.10 mol L–1 NaOH:
This solution must be prepared carefully.
Weigh out the sodium hydroxide pellets into a beaker and record the exact amount.
Add approximately 50 mL of deionised water to dissolve the NaOH pellets, then carefully transfer the solution to a 250 mL volumetric flask.
Rinse the stirring rod and beaker with deionised water, using a wash bottle, and transfer the rinse water into the flask. Repeat this two more times.
Add deionised water to the flask until the solution volume is exactly 250 mL.
Stopper the volumetric flask and invert and shake 5 times.
At this stage, you would usually standardise by carrying out 3 titrations against oxalic acid, for example, to ensure you have made up your solution accurately. Oxalic acid meets the standard criteria to a high degree.
2 When the solution has been thoroughly mixed, carefully transfer it into a bottle. Label the bottle with a description of the contents, the date and your name. Include the hazard warning: Caustic. Allow the solution to cool with the bottle lightly capped. When cool, tightly cap the bottle.
Titrating
1. Prepare and set up as in steps 1-3 in the titration procedure in the white box above.
2. Close the stopcock of the burette. Fill the burette with the standardised NaOH solution using a small funnel. Open the stopcock and drain a few millilitres of solution into a small beaker until the tip of the burette is free of air bubbles. Refill the burette and adjust the meniscus to 0.00 mL.
3. Prepare a diluted stock solution of vinegar: Using a clean, dry pipette, transfer exactly 25.0 mL into a 250 mL volumetric flask. Add distilled water to the flask until the solution volume is exactly 250 mL. Stopper the volumetric flask and shake.
4. Using a different pipette, transfer exactly 25mL of the diluted vinegar solution to a conical flask.
Now choose one set of Steps 5-8 from the two choices below
Using a pH meter
5. Immerse a clamped pH meter into the flask. Position the tip of the burette close to the mouth of the flask.
6. Take a pH reading at 0.0 mL.
7. Add 1.0 mL of NaOH from the burette, then take a pH reading.
8. Continue adding NaOH in 1.0 mL increments until 50.0 mL has been added and the pH has changed dramatically.
Using indicator
5. Add three drops of phenolphthalein indicator to the conical flask. Position the tip of the burette close to the mouth of the flask. Open the stopcock and adjust the flow to a rapid drop rate.
6. Swirl the flask constantly to mix the NaOH and vinegar. View the colour of the reaction mixture against a white background. Occasionally stop the titration and wash down the sides of the flask with a wash bottle.
7. Continue the addition of NaOH until the phenolphthalein indicator just becomes a faint pink colour. This critical stage of the titration is called the endpoint. Record the exact volume of NaOH delivered from the burette. This volume is called the titre.
8. Rinse out the flask three times and shake dry or use another clean, dry flask if available. Then perform several more titrations. Use a slower burette drop rate when within 1 mL of the expected endpoint. Record all titres. Remember: Stop the titration just at the point where the faint pink colour appears.
Note: A good analyst can achieve successive titres with a variance of ± 0.1 mL. The first titre is usually rejected because of over-titration beyond the endpoint, i.e. adding too much NaOH.
Adapted from (Pearson Education Australia, 2019)
VIEW PPT
Titrations Technique PPT
12C6 3.1a Titrations Technique PPT.pptVIEW Videos:
Performing a titration:
EAR#18 https://www.youtube.com/watch?v=nno4lmw6y10&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS&index=18 [6.56 mins]
Setting up and performing https://www.youtube.com/watch?v=sFpFCPTDv2w
Titration techniques: https://www.youtube.com/watch?v=Lf7H83ryrII&index=59&list=PLB755F0C836855EAB 20.37 [20.37 mins]
REVIEW OF TITRATION CALCULATIONS
VIEW PPT
Titrations Technique PPT
12C6 3.1a2 Titration Calculations 2 PPT.pptVIEW Videos:
Titration calculations: https://www.youtube.com/watch?v=gOqzUKIfqlE&index=31&list=PLeFSFSJ9WqSCKs3ELNcsclaXnKU0qvg7t [7.09 mins]
Acid and base titrations: https://www.youtube.com/watch?v=UwL5-lFG_RE&index=60&list=PLB755F0C836855EAB [14.23 mins]
Acid-Base Titration Prof Dave https://www.youtube.com/watch?v=dLNsPqDGzms [2.39 mins]
Titration Calculations https://www.youtube.com/watch?v=2z4mlE6MK0U [2.33 mins]
HSC Q4: https://www.youtube.com/watch?v=-V-8JKcMAF4&index=64&list=PLB755F0C836855EAB 8.40
TASK 3.1.1
Online simulation involving calculations http://employees.oneonta.edu/viningwj/sims/titrations_t.html
TASK 3.1.2
1. Define the following terms:
a) Volumetric analysis
b) Primary Standard
c) Standard Solution
d) Suitable Indicator
e) Titration
f) Equivalence Point
g) Endpoint
2. If you were to titrate a solution of acetic acid against sodium hydroxide, what equation would you write? What type of reaction would this be? Which indicator would you choose for this reaction?
3. Why is NaOH not chosen as a primary standard?
4. What are the criteria for primary standards to be used in acid-base titrations?
A good primary standard meets the following criteria:
high level of purity
low reactivity ( high stability, especially to water and oxygen)
high equivalent weight (to reduce error from measuring the mass)
not likely to absorb moisture from the air (hygroscopic) to reduce changes in mass in humid versus dry environments
non-toxic
inexpensive and readily available
In practice, few chemicals used as primary standards meet all of these criteria, although it's critical that a standard is of high purity. Also, a compound which may be a good primary standard for one purpose may not be the best choice for another analysis.
For example, sodium hydroxide (NaOH) tends to absorb moisture and carbon dioxide from the atmosphere, thus changing its concentration. A 1-gram sample of NaOH may not actually contain 1 gram of NaOH because additional water and carbon dioxide may have diluted the solution. To check the concentration of NaOH, a chemist must titrate the NaOH against a primary standard (in this case a solution of potassium hydrogen phthalate (KHP). KHP does not absorb water or carbon dioxide, and it can provide visual confirmation that a 1 gram solution of NaOH really contains 1 gram. NaOH, once its concentration has been validated through the use of a primary standard, is often used as a secondary standard.
There are many examples of primary standards; some of the most common include:
Sodium chloride (NaCl) is used as a primary standard for silver nitrate (AgNO3) reactions.
Sodium carbonate for standardisation of aqueous acids: hydrochloric, sulfuric acid and nitric acid solutions (but not acetic acid)
5. (2003, Q14)
In a titration of a strong base with a strong acid, the following procedure was used:
1. A burette was rinsed with water and then filled with the standard acid.
2. A pipette was rinsed with some base solution.
3. A conical flask was rinsed with some base solution.
4. A pipette was used to transfer a measured volume of base solution into the conical flask.
5. Indicator was added to the base sample and it was titrated to the endpoint with the acid.
Which statement is correct?
(A) The calculated base concentration will be correct.
(B) The calculated base concentration will be too low.
(C) The calculated base concentration will be too high.
(D) No definite conclusion can be reached about the base concentration.
ANS: C: If you add water to the burette, you are diluting the standard acid. The [H+] will actually be lower than what you think it is. Because it is lower, you will use less OH- than you should. Because you are using less OH-, it will seem to be stronger than it is ie the calculated [OH-] will be too high. If you rinse the conical flask with the base solution, you are adding moles OH- to it than you think you have. You will use more acid to neutralise it, the [OH-] will again appear to be higher than it is.
6. When preparing our standard solution what is the correct procedure for each of the following pieces of apparatus?
Pipette
Burette
Conical flask
7. Why is each of the procedures in Q7 important?
8. (2008, Q28, 6 marks)
A standard solution was prepared by dissolving 1.314 g of sodium carbonate in water. The solution was made up to a final volume of 250.0 mL.
(a) Calculate the concentration of the sodium carbonate solution. (2 marks)
This solution was used to determine the concentration of a solution of hydrochloric acid. Four 25.00 mL samples of the acid were titrated with the sodium carbonate solution. The average titration volume required to reach the end point was 23.45 mL.
(b) Write a balanced equation for the titration reaction.
(c) Calculate the concentration of the hydrochloric acid solution.
ANS:
(a) Better responses provided the correct formula for sodium carbonate and calculated the number of moles used. These candidates correctly calculated the concentration of the solution.
(b) Better responses correctly wrote a balanced chemical equation for the neutralisation reaction.
(c) Better responses showed clear working with the correct calculation of the number of moles of sodium carbonate used in the titration. Better responses correctly applied a 2:1 mole ratio of hydrochloric acid to sodium hydroxide and showed the calculation of hydrochloric acid calculation. The best responses did not round off at any step and gave answers correctly using four significant figures. Weaker responses could only show one correct step of the calculation. Weaker responses did not convert volume of solutions into litres, or used the acid volume to calculate the number of moles of sodium hydroxide.
9. (2004, Q16, 5 marks)
a) Outline the procedure you would use to prepare a standard solution of sodium hydrogen carbonate from solid sodium hydrogen carbonate. (3 marks)
b) Calculate the mass of solid sodium hydrogen carbonate required to make 250 mL of 0.12 mol L−1 solution. (2 marks)
ANS
(a) Although most candidates identified that the solid needed to be weighed, few completely outlined that it was transferred and dissolved in a volumetric flask and then filled it up to the calibration line. Candidates need to sequence their steps more appropriately. The terminology associated with equipment must be known.
(b) Most candidates correctly calculated the number of moles of solid, but fewer candidates were able to calculate the mass and express it to two significant figures. Appropriate setting out of working is crucial. Rounding-off needs to be done at the end of the calculation
6. (2002, Q8)
In a titration, an acid of known concentration is placed in a burette and reacted with a base that has been pipetted into a conical flask. What should each piece of glassware be rinsed with immediately before the titration?
Burette Pipette Conical flask
(A) Acid Base Water
(B) Water Water Water
(C) Acid Base Base
(D) Water Water Base
REVIEW 3.1
WS 6.7 Volumetric Analysis PP 70-71
REVIEW 3.1.2
PAST HSC QS (HSC EXAMS IN EARLIER YEARS HAD 15 MC QS, CURRENTLY THERE ARE 20.)
2016 29A
A solution of hydrochloric acid was standardised by titration against a sodium carbonate solution using the following procedure. • All glassware was rinsed correctly to remove possible contaminants. • Hydrochloric acid was placed in the burette. • 25.0 mL of sodium carbonate solution was pipetted into the conical flask. The titration was performed and the hydrochloric acid was found to be 0.200 mol L–1.
(a) Identify the substance used to rinse the conical flask and justify your answer. 2
(b) Seashells contain a mixture of carbonate compounds. The standardised hydrochloric acid was used to determine the percentage by mass of carbonate in a seashell using the following procedure.
• A 0.145 g sample of the seashell was placed in a conical flask.
• 50.0 mL of the standardised hydrochloric acid was added to the conical flask.
• At the completion of the reaction, the mixture in the conical flask was titrated with 0.250 mol L–1 sodium hydroxide.
The volume of sodium hydroxide used in the titration was 29.5 mL. Calculate the percentage by mass of carbonate in the sample of the seashell.
ANS:
29A Criteria Marks • Provides correct reason and substance used 2 • Provides correct substance 1
Sample answer: Water should be used to rinse the conical flask as this will not change the number of moles of Na2CO3 placed in it, whereas rinsing with the solution will, and so will result in an inaccurate calculation.
29B Criteria Marks • Provides correct answer with relevant working 4 • Provides a substantially correct answer with relevant working 3 • Provides some relevant steps 2 • Provides a relevant step 1
2015 Q2 D
2013 Q19 A
2013 Q28 See below
2013 Q28A
Criteria Marks • Identifies the mistake made AND • Provides a proposed change that would improve the validity of the result 2 • Identifies the mistake made 1
Sample answer: Instead of blowing through the pipette, student should have touched the end of pipette to surface of flask to draw out the liquid.
2013 Q28B
Criteria Marks • Relates correctly the decreased number of moles caused by both steps to a decrease in volume of HCl to produce an overall increase in calculated concentration 3 • Identifies both mistakes and relates them to a reduction of moles OR • Relates reduced moles of one step to increased HCl concentration as calculated 2 • Identifies the reduced number of moles caused by either step OR • Identifies both mistakes 1
Sample answer: In step 2, rinsing the pipette with water would decrease the number of moles of Na2CO3 it contains. This would be compounded by not filling to the mark made in step 3, resulting in the pipette having fewer moles of Na2CO3. Hence, a lower volume of the HCl would be added from the burette, but the student would think that there were more moles of HCl in this volume. So, they will calculate the concentration of the acid solution to be higher than the actual value.
2011 Q26 (6 marks)
A manufacturer makes lemon cordial by mixing flavouring, sugar syrup and citric acid. The concentration of the citric acid is determined by titration with NaOH. The sodium hydroxide solution is prepared by dissolving 4.000 g of NaOH pellets in water to give 1.000 L of solution. This solution is standardised by titrating 25.00 mL with a 0.1011 mol L–1 standardised solution of HCl. The average titration volume is found to be 24.10 mL. To analyse the lemon cordial 50.00 mL of the cordial is diluted to 500.0 mL. Then 25.00 mL of the diluted solution is titrated with the NaOH solution to the phenolphthalein endpoint. The following data were collected during one of the analysis runs of the lemon cordial. Titration #1 volume 26.55 mL Titration #2 volume 27.25 mL Titration #3 volume 27.30 mL Titration #4 volume 27.20 mL
(a) Why is the calculated concentration of the standardised NaOH solution different 2 from the concentration calculated using the mass given, assuming no human error occurred? 2 marks
(b) Determine the concentration of citric acid in the lemon cordial. 4 marks
ANS:
Q26A
Criteria Marks • Provides a valid reason • Correctly calculates both NaOH concentrations OR • Indicates actual value is less than theoretical value 2 • Provides a valid reason OR • Correctly calculates TWO NaOH concentrations 1
Q26B
Criteria Marks • Correctly calculates concentration of original solution of citric acid 4 • Calculates concentration of original solution of citric acid using incorrect NaOH concentration or incorrect mole ratio or incorrect volume of HCl OR • Correctly calculates concentration of diluted solution of citric acid 3 • Calculates concentration of citric acid solution with TWO errors 2 • Provides ONE correct step in the calculation 1
2010 Q28 (see diagram below)
ANS:
Criteria Marks • Demonstrates a thorough understanding of all THREE steps by describing features of steps with appropriate reference to techniques and equipment used • Determines concentration of HCl correctly with correct equation 7–8 • Demonstrates a sound understanding of all THREE steps by outlining features of steps with appropriate reference to some techniques and equipment used • Provides ONE correct calculation and correct equation 5–6 • Demonstrates a limited understanding of all THREE steps. Refers to some techniques and equipment used • Calculates moles or mass of Na2CO3 or correct equation 3–4 • Demonstrates a basic understanding of some steps or refers to some techniques or equipment used OR • Calculates formula mass of Na2CO3 or titrated moles or mass of Na2CO3 with some outline of procedure OR • Calculates correct concentration of HCI OR • Provides correct equation and refers to some techniques or equipment used 2
3.2 investigate titration curves and conductivity graphs to analyse data to indicate characteristic reaction profiles
INTRODUCTION
Acid - base reactions are usually colourless and this makes it difficult to locate the equivalence point.
The equivalence point is the point at which #moles H+ exactly equals #moles OH-
As a result we need some chemical substance or electronic device to help us determine the equivalence point. This may be an indicator, a pH probe (data logger) or pH meter, or a conductivity meter.
The pH probe measures changes in the pH as the reaction progresses. If we start with an acid (either weak or strong) the pH will be low. It will rise as the base is added. How quickly it will rise will depend on the nature of the acid and base.
The conductivity meter measures the concentration of ions in solution. A strong acid will fully ionise, while a weak acid only partially ionises. As the base is added, OH- ions react with H+ ions and so they are converted into molecular water particles. This reduces the ion concentration of the strong acid and shifts the equilibrium of the weak acid to release more ions (due to Le Chatelier's Principle). Once neutralisation occurs, additional ions of the base are added un-neutralised, and so the conductivity will rise again.
Select one of the following websites for titration simulations: http://www.oup.com/uk/orc/bin/9780199277896/01student/activities/acidBaseLayoutf.swf
NOTES:
Acid titrated ‘with’ a base means that you are slowly adding a base into the conical flask (containing acid) and therefore the base is your titrant/analyte (solution in burette).
Base titrated ‘with’ an acid means that you are slowly adding an acid into the conical flask (containing base) and therefore the acid is your titrant/analyte (solution in burette).
The curves below are just generic titration curves. The volume of acid/base added from the burette into the conical flask will vary depending on the concentration and volume of your unknown solution, as well as your standardised solution.
a) Strong acid/Strong base
a) Strong acid/Strong base
In these reactions, both species are completely dissociated, so the net ionic equation will be:
H+(aq) + OH-(aq) → H2O(l)
NOTE: The mole ratio here is 1:1. This ratio will change if we use a diprotic acid (eg H2SO4) or base (eg Mg(OH)2) or a triprotic acid (eg H3PO4). It is essential to write out an equation for the reaction so that the mole ratio can be established before starting the calculations.
At the equivalence point for all strong acid and strong base titration, the pH of the solution is always equal to 7.
Suppose we have a strong acid/strong base titration such as:
HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)
The conjugate acid of a strong base formed as product in titration is very a weak acid and is unable to react with water to act as a proton donor (i.e. unable to donate a H+ to water). For instance, the conjugate acid in our example would be the Na+ ion. Since the sodium ion is the conjugate acid of the strong base, NaOH, it is unable to react with water as an acid, i.e. sodium ion cannot donate H+ to water.
Similarly, the conjugate base of a strong acid as a product in titration is a very weak base and is unable to react with water to act as a proton acceptor (i.e. unable to accept a H+ from water). The conjugate base in our example is the Cl– ion. Since the chloride ion is the very weak conjugate base of the strong acid, HCl, it is unable to react with water as a base, i.e. chloride ion cannot accept H+ from water.
This means that only water has an effect on the pH of the solution at equivalence point. Since pure water is neutral (pH=7) and the conjugate acid/base components of the salt product (eg NaCl) DO NOT react with water to affect the pH, the pH of the solution at equivalence is equal to 7.
If you keep adding excess acid/base after the equivalence point, the pH will change as shown in the titration curve diagram as the strong acid/base will react with water and affect the pH.
View Video:
Acid Base Titration Curves - Strong acid/Strong base: https://www.youtube.com/watch?v=uO8DWyRJMHo [0.00-3.11 mins]
Titration Curve
The titration curve of a strong acid/strong base will have an equivalence point of pH=7, and will look like the one above.
Equivalence (Stoichiometric) Point – is the point in the titration when enough titrant (substance in the burette) has been added to react exactly with the substance being titrated (solution in the flask); defined by stoichiometry where moles H+ = moles OH-.
The equivalence point of a titration is not necessarily the same as the end point (where the indicator changes colour). When we choose an indicator for a titration, we want the indicator end point (where the color changes) and the titration equivalence point to be as close as possible. Careful selection of the indicator will ensure that the difference between them, and therefore the error, is negligible.
The dramatic change in pH near the equivalence point makes it easier to choose an indicator because the colour changes will be sharp, usually occurring with the addition of a single drop of titrant.
Conductivity Graph:
For an electric current to be conducted by a solution, ions must be present to carry the current. Conductivity relies on ions in solution.
VIEW Video:
Conductivity Titrations - Strong acid/Strong base: https://www.youtube.com/watch?v=J7EWciv_4iY [0.00-6.54 mins]
The conductivity graph below shows strong base added to the strong acid. The conductivity starts high with the strong acid having high dissociation, drops as the H+ ions join with OH- ions to form covalent (non-ionic, therefore non-conducting) water molecules. After the equivalence point is reached, excess OH- ions are being added, and the conductivity will increase again.
Extra for experts: Although we are also adding spectator ions (eg Na+), they do not conduct electricity strongly as the H+ and OH- ions do (for reasons of size and electronegativity vs mobility)
b) Strong acid/Weak base
b) Strong acid/Weak base
Titration Curve:
Since the base is a weak base, much of it will remain in molecular form. The H+ from the acid reacts directly with the un-ionised base. The reaction goes to completion. One example is:
NH3 + H+ → NH4+
View Video:
Acid Base Titration Curves - Strong acid/Weak base: https://www.youtube.com/watch?v=uO8DWyRJMHo [3.11-4.53 mins]
Conductivity Graph:
VIEW Video:
Conductivity Titrations - Strong acid/Weak base: https://www.youtube.com/watch?v=J7EWciv_4iY [7.05-9.46 mins]
Look at the conductivity graph on the right.
What species is in the flask to start?
What part of the graph shows you that the acid is a strong acid?
Why does the conductivity drop?
Why does the conductivity not increase again?
c) Weak acid/Strong base
c) Weak acid/Strong base
Weak acids only partially ionise, and so much of the acid remains in molecular form. Consider acetic acid (CH3COOH). The OH- ions from the base react directly with the molecular acid. The equilibrium constant is very large, hence the reaction proceeds to completion. The equation is:
CH3COOH + OH- → CH3COO- + H2O
NOTE: This is a 1:1 ratio, but not all weak acids are monoprotic (eg H2CO3), so write out an equation before you start to solve the problem.
Titration Curve:
View Video:
Acid Base Titration Curves - Weak acid/Strong base: https://www.youtube.com/watch?v=uO8DWyRJMHo [4.54-8.01mins]
Conductivity Graph:
VIEW:
Conductometric Titration - Weak acid/Strong base https://www.youtube.com/watch?v=iQFQVuicW74 [9.46- 12.49 mins]
Look at the conductivity graph on the right.
What species is in the flask to start?
What part of the graph shows you that the acid is a weak acid?
Why does the conductivity rise?
Why does the conductivity finish up above the starting level?
View PPT
12C6 3.2a Titration Curves PPT.pptTASK 3.2.1
Complete WS below (Geometric method is
12C6 Titration Curves WS.pdf
VIEW Videos::
ABR#19 Titration Curves https://www.youtube.com/watch?v=VJp4zOzDgVg&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS&index=19 [9.10]
http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html
Understanding an acid base titration curve https://www.youtube.com/watch?v=YPM-bbvasC4
TASK 3.2.2
Work through the PPT
12C6 3.1a2 Titrations Calculations PPT.pptTASK 3.2.3
1. (2009, Q3)
Which of the following groups contains ONLY acidic substances?
(A) Antacid tablets, baking soda, laundry detergents
(B) Blood, oven cleaner, seawater
(C) Milk, tea, drain cleaner
(D) Vinegar, wine, aspirin
ANS: D Vinegar is acetc acid, during the course of winemaking and in finished wines, acetic acid can play a significant role, aspirin is acetylsalicylic acid
2. (2006, Q9)
Which statement best describes the equivalence point in a titration between a strong acid and a strong base?
(A) The point at which the first sign of a colour change occurs
(B) The point at which equal moles of acid and base have been added together
(C) The point at which equal moles of H+ ions and OH– ions have been added together
(D) The point at which the rate of the forward reaction equals the rate of the reverse reaction
3. (2005, Q10)
A titration was conducted by adding NaOH from a teflon-coated burette to HCl in a conical flask. The pH in the flask was recorded during the titration and Curve A was produced. The table shows appropriate indicators used to identify the equivalence point in titrations. For NaOH and HCl the appropriate indicator is bromothymol blue.
Indicator Acidic Colour Range of Colour Change Basic Colour
Methyl orange Red 3.1 – 4.4 Yellow
Methyl red Red 4.4 – 6.2 Yellow
Bromothymol blue Yellow 6.0 – 7.6 Blue
Cresolphthalein Colourless 8.1 – 9.7 Red
Alizarin yellow Yellow 10.1 – 12.0 Red
A second titration was conducted by adding NaOH to a different acid. The pH in the flask was recorded during the titration and Curve B was produced. What is the appropriate indicator for Curve B using the table?
(A) Methyl orange
(B) Methyl red
(C) Cresolphthalein
(D) Alizarin yellow
4. (2009, Q21, 6 marks)
The graph shows changes in pH for the titrations of equal volumes of solutions of two monoprotic acids, Acid 1 and Acid 2.
a) Explain the differences between Acid 1 and Acid 2 in terms of their relative strengths and concentrations. (3 marks)
b) Name the salt produced by the reaction of an acid of the same type as Acid 2 with KOH(aq). (1 mark)
c) Calculate the concentration of hydrogen ions when 20 mL of KOH(aq) has been added to Acid 1. (1 mark)
d) Why would phenolphthalein be a suitable indicator for both titrations? (1 mark)
5. The question below is taken from the website: https://www.chegg.com/homework-help/questions-and-answers/graphs-figure-p1610-show-conductivity-solution-function-volume-titrant-added-graphs-best-r-q12966141
The graphs below show the conductivity of a solution as a function of the volume of titrant added.
Which of the graphs best represents the titration of
a) a strong acid with a strong base? Explain your answer.
b) a weak acid with a strong base? Explain your answer.
6. (1998, Q8)
20 mL of KOH solution was titrated with 0.20 mol.L-1 H2SO4 solution in a conductivity cell. The data obtained were plotted to give the graph shown below.
The concentration of the KOH solution was
A) 0.30 mol.L-1 B) 0.15 mol.L-1 C) 0.12 mol.L-1 D) 0.075 mol.L-1
REVIEW 3.2
WS 6.6 Acid-base curves P68
PAST HSC QS (QS WITH NO MARKS INDICATED ARE MULTIPLE CHOICE. HSC EXAMS IN EARLIER YEARS HAD 15 MC QS, CURRENTLY THERE ARE 20.)
2017 Q1; Q13; Q24C 3
2016 Q29A 2; 29B 4
2015 Q14; Q26C 4
2014 Q30A 1; 30B 4
2013 Q28B 3
2012 Q4; Q19; Q30A 2; Q30B1 4
2011 Q15; Q18; Q26A 2; 26B 4
2010 Q28 8
2009 Q14; Q21A 3; Q21B 1; 22A 2; 22B 1; 22C 2
2008 Q28A 2; 28B 1; 28C 3
2007 Q21C 3
2006 Q24B 1; 24C 3
3.3 model neutralisation of strong and weak acids and bases using a variety of media
All other things being equal, a neutralisation reaction involves the formation of water molecules from the reaction between hydrogen (or hydronium) ions with hydroxide ions.
H+(aq) + OH-(aq) ⇌ H2O(l)
OR H3O +(aq) + OH-(aq) ⇌ 2H2O(l)
However, the reality of neutralisation reactions can change depending upon the type of reaction.
For a strong acid + strong base, the neutralisation reaction above is a good approximation and the equivalence point should have a pH of 7. [The centre point of a titration is the equivalence point and this describes the neutralisation point where all hydrogen ions are neutralised by hydroxide ions (Silove, 2018)].
But despite the fact that we describe both the reactions of a strong base and weak acid, as well as a strong acid and weak base as neutralisation reactions, these latter two do not have an equivalence point at a pH of 7.
To understand different types of neutralisation reactions, we use models.
TASK 3.3.1 Investigate simulations:
1. http://college.cengage.com/chemistry/discipline/mml3/#gen_16_7
2. https://phet.colorado.edu/en/simulation/acid-base-solutions (click on image below left)
3. https://phet.colorado.edu/sims/html/ph-scale/latest/ph-scale_en.html (click on image below right)
VIEW Videos:
EAR #20 https://www.youtube.com/watch?v=79TI9NWoteA&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS&index=20 [7.25 mins]
Explaining a pH > 7 for a neutralisation reaction between a strong base and a weak acid. Visichem www.visichem.thelearningfederation.edu.au/topic12.html You will need to click on this link.
12C6 2.2a Neutralisation Model MOV.mov12C6 2.4a AcidBase Model WS for MOV.pdfTASK 3.3.2
1. Discuss one model of neutralisation you have used which helped you understand neutralisation reactions.
2. Explain the meaning of the following statement: “When a weak acid dissolves in water an equilibrium is established between the acid molecule and its constituent ions”.
3. Explain the following observations:
a) The titration of a strong acid against a strong base produces an equivalence point around a pH of 7.
b) The titration of a weak acid against a strong base produces an equivalence point around a pH of 9-11.
c) The titration of a strong acid against a weak base produces an equivalence point around a pH of 3-5.
3.4 a) explore the use of Keq for different types of chemical reactions, including but not limited to:
dissociation of acids and bases
VIEW Videos:
EAR# 21 https://www.youtube.com/watch?v=uxcYQJmzSd8&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS&index=21 [8.46 mins]
Ka and acid strength various videos at this link https://www.khanacademy.org/science/chemistry/acids-and-bases-topic/copy-of-acid-base-equilibria/v/ka-and-acid-strength
3.4 b) calculate and apply the dissociation constant (Ka) and pKa (pKa = -log10 (Ka)) to determine the difference between strong and weak acids
We know that many acids are only partially ionised in water. The stronger the acid, the more readily it is ionised. The weaker the acid, the less it is ionised.
An equilibrium exists between the acid molecule and its constituent ions in a weak acid.
HA(aq) ⇌ H+(aq) + A-(aq)
The extent to which acids are ionised can be determined using a special equilibrium constant known as the acid ionisation constant, Ka.
For weak monoprotic acids which undergo partial ionisation:
[H+] [A-]
Ka = ------------
[HA]
The smaller the value of Ka the weaker the acid, the lower the numerator (ie ionic products) and the higher the value of the denominator (ie molecular acid reactants ).
Both acid strength and pH can be situated on a log scale. Strong acids ionise completely and weak acids only partially ionise. Hence we can use another indicator of acid strength, pKa. As with pH, pKa is calculated as follows:
pKa = -log10(Ka)
As the acid gets weaker, pKa increases.
ATAR Notes Key Point
A high Ka will indicate a strong acid.
A low Ka will indicate a weak acid.
A high pKa will indicate a weak acid.
A low pKa will indicate a strong acid. (Silove, 2018, p. 42)
For comparison:
the Ka for sulfuric acid is 1.2 x 10-2, whilst its pKa is 1.92. Sulfuric acid is a strong acid.
the Ka for carbonic acid is 4.8 x 10-11, whilst its pKa is 10.32. Carbonic acid is a weak acid.
Type 1: Calculating pKa
The pKa is calculated using the expression:
pKa = - log (Ka)
where "Ka" is the equilibrium constant for the ionization of the acid.
Example: What is the pKa of acetic acid, if Ka for acetic acid is 1.78 x 10-5?
pKa = - log (1.78 x 10-5) = - ( - 4.75) = 4.75
Type 2: Calculating Ka from the pKa
The Ka for an acid is calculated from the pKa by performing the reverse of the mathematical operation used to find pKa.
Ka = 10-pKa or Ka = antilog ( - pKa)
Example: Calculate the value of the ionization constant for the ammonium ion, Ka, if the pKa is 9.74.
9.74 = - log (Ka)
-9.74 = log (Ka)
Ka = 10-9.74 = 1.82 x 10-10
On a calculator calculate 10-9.74, or "inverse" log ( - 9.74).
Type 3: Calculating pKb
The pKb is calculated using the expression:
pKb = - log (Kb)
where Kb is the equilibrium constant for the ionization of a base.
Example: What is the pKb for methyl amine, if the value of Kb for methyl amine is 4.4 x 10-4?
pKb = - log (4.4 x 10-4) = - ( - 3.36) = 3.36
Type 4: Calculating Kb from pKb
The Kb for an acid is calculated from the pKb by performing the reverse of the mathematical operation used to find pKb.
Kb = 10-pKb or Kb = antilog ( - pKb)
Example: Calculate the value of the ionisation constant, Kb, for aniline if the pKb is 9.38.
9.38 = - log (Kb)
-9.38 = log (Kb)
Kb = 10-9.38 = 4.17 x 10-10
On a calculator calculate 10-9.38, or "inverse" log ( - 9.38).
12C6 3.4b Ka Data Sheet.docx
TASK 3.4.1
Complete WS 6.4
VIEW Videos:
ABR#21 https://www.youtube.com/watch?v=uxcYQJmzSd8 [8.41 mins]
TASK 3.4.2
Complete online
http://web.mst.edu/~gbert/WeakAB/AB.html
Example:
A 0.30 M solution of this base is 25.0% ionised.
BOH ---- B+ + OH-
Calculate the ionisation constant for this base.
Example solution:
If a 0.30 M solution of this base is 25.0% ionised, the concentration of ions will be:
[OH-] = [B+] = (0.30) x (25/100) = 0.075 M
The concentration of undissociated BOH will be:
[BOH] = (0.30) - 0.075 = 0.225
Kb = [OH-] [B+] / [BOH]
K = (0.075)(0.075)/(0.225) = 0.025
PRACTICAL 3.4.1
PA 6.7 Determination of two acid dissociation constants PP91-2
TASK 3.4.3
1. Why is it incorrect to identify a strong acid solution, such as hydrochloric acid in water, as an equilibrium?
2. (1998, Q21, 4 marks)
The hydrogen ion concentration of 1.00 mol.L-1 ethanoic acid and chloroethanoic acid at 25oC are given in the table.
Acid Formula [H+] (mol.L-1)
Ethanoic acid CH3COOH 4.18 x 10-3
Chloroethanoic acid ClCH2COOH 3.74 x 10-2
a) What is the pH of 1.00 mol.L-1 chloroethanoic acid? (1 mark)
b) What is the Ka of chloroethanoic acid? (1 mark)
c) Compare the strength of ethanoic acid with that of chloroethanoic acid and explain your answer in terms of Ka. (2 marks)
3. (1998, Q29b, 2 marks)
Phosphoric acid is a polyprotic acid. It dissociates in water in three steps.
Step 1 H3PO4 + H2O ⇌ H3O+ + H2PO4- Ka1
Step 2 H2PO4- + H2O ⇌ H3O+ + HPO42- Ka2
Step 1 HPO42 + H2O ⇌ H3O+ + PO43- Ka3
Which step would have the highest Ka value? Explain your answer.
REVIEW 3.4.1
WS 6.4 Indicators and Ka P66
Review 8.5
REVIEW 3.4.2
Complete Qs below
12C6 3.4 Acid Dissociation Qs.docx3.5 explore acid/base analysis techniques that are applied:
in industries
by Aboriginal and Torres Strait Islander Peoples
using digital probes and instruments
VIEW Videos:
ABR#22 https://www.youtube.com/watch?v=ViiCduBKsE8&t=3s [5.24 mins]
TASK 3.5.1
The techniques associated with acid/base reactions, volumetric analysis and the nature of salt production have wide applications. In this section you will become an expert on one of these applications and teach it to the other groups. Applications of acid/base techniques fall into the categories of:
a) Industrial applications
Select an application from one specific industry eg. food, mining, pharmaceuticals, petrochemicals or viticulture.
EXAMPLE: Acidity of wine contributes to taste, balances sugars and alcohol and supports fermentation, enhancing the wine’s colour and flavour. (Davis, Disney, & Smith, 2018) https://www.awri.com.au/industry_support/winemaking_resources/frequently_asked_questions/acidity_and_ph/#title9
b) Aboriginal and Torres Strait Island cultural applications
Kuku Yalanji people in the Daintree region in Northern Queensland have used the medicinal use of the soap tree (Alphitonia excelsa). Its lather can be used as an antiseptic and cleanser. Subsequent analysis has shown the presence of acidic saponins, while the wood, bark and leaves contain methyl salicylate. (Davis, Disney, & Smith, 2018)
A few interesting articles and websites below may help you contextualise your knowledge and focus your additional research.
Aboriginal healing practices and Australian bush medicine by Philip Clarke, published in the Journal of the Anthropological Society of SA V 33 in 2008. http://www.friendsofglenthorne.org.au/wp-content/uploads/Clarke-Vol-33-2008.pdf
Chemistry in Australia May 2016 issue https://chemaust.raci.org.au/sites/default/files/pdf/2016/CiA_May%202016_final_0.pdf
Indigenous Engagement with Science: Towards deeper understandings
Indigenous Futures https://www.csiro.au/en/Showcase/Indigenous-Futures
Indigenous medicine – a fusion of ritual and remedy https://www.csiro.au/en/Showcase/Indigenous-Futures
c) Technological applications (using digital probes and instruments)
EXAMPLES:
Blood gas sensor technology https://acutecaretesting.org/en/articles/understanding-the-principles-behind-blood-gas-sensor-technology
NSW Environment Protection Authority www.epa.nsw.gov.au
Titration techniques in the Food industry https://www.newfoodmagazine.com/article/1485/titration-techniques-in-the-food-industry/
pH adjustment for wastewater treatment http://www.phadjustment.com/Technical_Articles.html
3.6 conduct a chemical analysis of a common household substance for its acidity or basicity, for example:
– soft drink
– wine
– juice
– medicine
VIEW Videos:
ABR#23 https://www.youtube.com/watch?v=9wcA21LVhVk [6.48]
PRACTICAL 3.6.1 See 3.1.1
PA 6.6 Determination of acid concentration of vinegar PP88-90
PRACTICAL 3.6.2
Introduction
In the last experiment you used the technique of titration to determine the concentration of acetic acid in vinegar. However vinegar is not the only acidic or basic substance which can be found in the home. Many substances contain acids or bases in solution and hence may be analysed using titration. You need to bring together your knowledge about titration techniques, strong and weak solutions and the appropriate selection of an indicator, particularly if your test solution is already coloured.
You will need to write your own method, select appropriate reagents and other materials, and consider aspects of safety, accuracy, reliability and validity. You will also need to show your calculations and evaluate your final result.
You can choose to analyse a soft drink, some wine, some juice or a medicine. You will need to research which main acid(s) or base(s) you are seeking to neutralise in your titration.
Practical 3.1.1 Measuring the acid concentration of vinegar
PA 6.6 Determination of acid concentration of vinegar PP88-90
Introduction
Volumetric analysis is essentially the reaction of a known quantity of solution with the sample to be analysed. Volumetric analysis employs precision glassware (pipette and burette) to measure out solutions with great accuracy. (Pearson Education Australia, 2019)
Vinegar is a common household item containing acetic acid, as well as other chemicals. This experiment is designed to determine the molar concentration of acetic acid in a sample of vinegar by titrating it with a standard solution of NaOH.
CH3COOH(aq) + NaOH(aq) -> CH3COONa(aq) + H2O(l)
Adding the sodium hydroxide, a basic solution, to the acetic acid, an acidic solution, a neutralisation reaction occurs. An indicator known as phenolphthalein, is added to the vinegar. This indicator turns the solution to a dark pink when excess NaOH makes the solution more basic. When the solution turns light pink, it has successfully been neutralised.
The amount of NaOH used to standardise the vinegar can then be used to determine the amount of acetic acid in the vinegar as they are at a 1:1 ratio as seen in the above equation. The moles of NaOH used to neutralise the acid equal the number of the moles of acetic acid present in the vinegar.
Preparing the standard solution
Warning: Wear safety glasses, gloves and lab coats! NaOH is caustic. Avoid skin contact. Clean up spills immediately.
1 Prepare 250 mL of 0.10 mol L–1 NaOH:
This solution must be prepared carefully.
Weigh out the sodium hydroxide pellets into a beaker and record the exact amount.
Add approximately 50 mL of deionised water to dissolve the NaOH pellets, then carefully transfer the solution to a 250 mL volumetric flask.
Rinse the stirring rod and beaker with deionised water, using a wash bottle, and transfer the rinse water into the flask. Repeat this two more times.
Add deionised water to the flask until the solution volume is exactly 250 mL.
Stopper the volumetric flask and invert and shake 5 times.
At this stage, you would usually standardise by carrying out 3 titrations against oxalic acid, for example, to ensure you have made up your solution accurately. Oxalic acid meets the standard criteria to a high degree.
2 When the solution has been thoroughly mixed, carefully transfer it into a bottle. Label the bottle with a description of the contents, the date and your name. Include the hazard warning: Caustic. Allow the solution to cool with the bottle lightly capped. When cool, tightly cap the bottle.
Titrating
1. Prepare and set up as in steps 1-3 in the titration procedure in the white box above.
2. Close the stopcock of the burette. Fill the burette with the standardised NaOH solution using a small funnel. Open the stopcock and drain a few millilitres of solution into a small beaker until the tip of the burette is free of air bubbles. Refill the burette and adjust the meniscus to 0.00 mL.
3. Prepare a diluted stock solution of vinegar: Using a clean, dry pipette, transfer exactly 25.0 mL into a 250 mL volumetric flask. Add distilled water to the flask until the solution volume is exactly 250 mL. Stopper the volumetric flask and shake.
4. Using a different pipette, transfer exactly 25mL of the diluted vinegar solution to a conical flask.
Now choose one set of Steps 5-8 from the two choices below
Using a pH meter
5. Immerse a clamped pH meter into the flask. Position the tip of the burette close to the mouth of the flask.
6. Take a pH reading at 0.0 mL.
7. Add 1.0 mL of NaOH from the burette, then take a pH reading.
8. Continue adding NaOH in 1.0 mL increments until 50.0 mL has been added and the pH has changed dramatically.
Using indicator
5. Add three drops of phenolphthalein indicator to the conical flask. Position the tip of the burette close to the mouth of the flask. Open the stopcock and adjust the flow to a rapid drop rate.
6. Swirl the flask constantly to mix the NaOH and vinegar. View the colour of the reaction mixture against a white background. Occasionally stop the titration and wash down the sides of the flask with a wash bottle.
7. Continue the addition of NaOH until the phenolphthalein indicator just becomes a faint pink colour. This critical stage of the titration is called the endpoint. Record the exact volume of NaOH delivered from the burette. This volume is called the titre.
8. Rinse out the flask three times and shake dry or use another clean, dry flask if available. Then perform several more titrations. Use a slower burette drop rate when within 1 mL of the expected endpoint. Record all titres. Remember: Stop the titration just at the point where the faint pink colour appears.
Note: A good analyst can achieve successive titres with a variance of ± 0.1 mL. The first titre is usually rejected because of over-titration beyond the endpoint, i.e. adding too much NaOH.
Adapted from (Pearson Education Australia, 2019)
TASK 3.6.1
1. (2010, Q7) Equal volumes of four 0.1 mol L−1 acids were titrated with the same sodium hydroxide solution.
Which one requires the greatest volume of base to change the colour of the indicator?
(A) Citric acid
(B) Acetic acid
(C) Sulfuric acid
(D) Hydrochloric acid
2. (2003, Q23, 4 marks)
25.0 mL of 0.12 mol L−1 standard barium hydroxide solution was titrated with nitric
acid. The results are recorded in the table.
Titration Volume of nitric acid used (mL)
1 20.4
2 18.1
3 18.2
4 18.1
(a) Write a balanced chemical equation for the reaction of barium hydroxide with
nitric acid. (1 mark)
(b) Calculate the concentration of the nitric acid. (3 marks)
3. (2010, Q28, 8 marks)
The flowchart shown outlines the sequence of steps used to determine the concentration of an unknown hydrochloric acid solution.
Describe steps A, B and C including correct techniques, equipment and appropriate calculations. Determine the concentration of the hydrochloric acid.
4. (2009, Q22, 7 marks)
The nitrogen content of bread was determined using the following procedure:
A sample of bread weighing 2.80 g was analysed.
The nitrogen in the sample was converted into ammonia.
The ammonia was collected in 50.0 mL of 0.125 mol L−1 hydrochloric acid. All of the ammonia was neutralised, leaving an excess of hydrochloric acid.
The excess hydrochloric acid was titrated with 23.30 mL of 0.116 mol L−1 sodium hydroxide solution.
(a) Write balanced equations for the TWO reactions involving hydrochloric acid. (2 marks)
(b) Calculate the moles of excess hydrochloric acid. (1 mark)
(c) Calculate the moles of ammonia. (2 marks)
(d) Calculate the percentage by mass of nitrogen in the bread. (2 marks)
5. (2005, Q24, 5 marks)
An antacid tablet is known to contain calcium carbonate (CaCO3). To determine the mass of calcium carbonate in the tablet, the following procedure was used.
The tablet was crushed and then placed in a beaker.
A pipette was used to add 25.0 mL of 0.600 mol L–1 hydrochloric acid to the crushed tablet in the beaker.
Once the reaction between the calcium carbonate and hydrochloric acid had stopped, phenolphthalein indicator was added to the reaction mixture.
A teflon-coated burette was then used to add 0.100 mol L–1 sodium hydroxide to the beaker to neutralise the excess hydrochloric acid.
The phenolphthalein changed from colourless to pink after 14.2 mL of the sodium hydroxide solution had been added.
a) Write a balanced chemical equation for the reaction that occurred between the calcium carbonate in the tablet and the hydrochloric acid. (1 mark)
b) How many moles of hydrochloric acid were added to the tablet? (1 mark)
c) Calculate the mass of calcium carbonate in the original antacid tablet. (3 marks)
6. (2009, Q14)
Citric acid, the predominant acid in lemon juice, is a triprotic acid. A student titrated 25.0 mL samples of lemon juice with 0.550 mol L–1 NaOH. The mean titration volume was 29.50 mL. The molar mass of citric acid is 192.12 g mol–1.
What was the concentration of citric acid in the lemon juice?
(A) 1.04 g L–1
(B) 41.6 g L–1
(C) 125 g L–1
(D) 374 g L–1
7. (2007, Q21, 5 marks) Look at the table below.
(a) State what colour the red cabbage indicator would be in a 0.005 mol L–1 solution of H2SO4? Show your working.
(b) Using the red cabbage indicator, what colour would the solution be if 10 mL of 0.005 mol L–1 H2SO4 was diluted to 100 mL?
(c) What volume of 0.005 mol L–1 KOH is required to neutralise 15 mL of the diluted solution of H2SO4?
3.7 describe the importance of buffers in natural systems
“A buffer is a chemical added to water to minimise a change in the pH if acids or bases are added.” (Selinger & Barrow, 2017, p. 318)
“A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. It has the property that the pH of the solution changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. Many life forms thrive only in a relatively small pH range...” (Hegarty, 2018, p. 56)
A buffer controls the pH of a solution.
A buffer solution is usually a mixture of:
a weak acid and its conjugate base (eg HF and NaF)
or
a weak base and its conjugate acid (eg NH3 and NH4Cl)
If acid is added to the buffer, the excess hydrogen ions are removed by the base. If base is added to the buffer, excess hydroxide ions are removed by the H+ ions.
The net effect is that the pH of the solution containing the buffer remains relatively stable within the working pH range of the buffer solution (they have a capacity which dictates how much acid/base can be neutralised before the pH changes, and the amount by which it will change.)
Human blood contains a buffer of carbonic acid (H2CO3) and bicarbonate anion (HCO3-) in order to maintain blood pH between 7.35 and 7.45: a value higher than 7.8 or lower than 6.8 can lead to death. In this buffer, hydronium and bicarbonate anion are in equilibrium with carbonic acid. The carbonic acid in the first equilibrium can decompose into CO2 gas and water, resulting in a second equilibrium system between carbonic acid and water. Because CO2 is an important component of the blood buffer, its regulation in the body, as well as that of O2 , is extremely important, especially when the human body is subjected to strenuous conditions.
In the body, there is another equilibrium between hydronium and oxygen which involves the binding ability of hemoglobin. An increase in hydronium causes this equilibrium to shift towards the oxygen side, thus releasing oxygen from hemoglobin molecules into the surrounding tissues/cells. This system continues during exercise, providing continuous oxygen to working tissues.
The blood buffer is:
H3O+ + HCO−3 ⇌ H2CO3 + H2O
With the simultaneous equilibrium:
H2CO3 ⇌ H2O + CO2
Swimming Pools are another situation where buffers help maintain the pH (Selinger & Barrow, 2017, pp. 316-317). Explain how the equilibrium shown below could act as a buffer in a pool.
OCl−(aq) + H2O(l) ⇌ HOCl(aq) + OH−(aq)
Rising carbon dioxide levels in the atmosphere may be having an impact on the oceans, again due to buffers. As is the case with all buffers, there is a point beyond which continual addition of acid can no longer be balanced by the buffer, and the pH will start to drop.
When CO2 dissolves in the ocean it combines with water to form hydrogen (H+) and bicarbonate (HCO3-) ions:
CO2(g) + H2O(l) ⇌ H+(aq) + HCO3-(aq)
Some of the hydrogen ions combine with carbonate (CO32-(aq)) ions to form additional bicarbonate ions resulting in a decrease in the carbonate ions and an increase in the bicarbonate ions:
H+(aq) + CO32-(aq) ⇌ HCO3-(aq)
The net effect of these changes when carbon dioxide dissolves in seawater is for the concentrations of H+, CO2, and HCO3- to increase, and the concentration of CO32- to decrease.
The problem is that CO32 is a major component of corals and the shells of many marine organisms. Lowering the CO32 concentration can therefore have a negative impact on the calcification of exoskeletons for corals, marine molluscs, foraminifera, etc.
More information can be found here: http://ocean-acidification.net/ or here https://climate.nasa.gov/news/13/climate-change-seeps-into-the-sea/
TASK 3.7.1
Try this simulation
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf
VIEW
Buffers PPT
12C6 3.7a Buffers PPT.pptxVIEW Videos:
ABR# 25 https://www.youtube.com/watch?v=4RRdVJmT2VA [8.03 mins]
Buffer system in the blood https://www.youtube.com/watch?v=2mTUj33Np74 [6.32 mins]
Natural buffer systems: https://www.youtube.com/watch?v=tdQHXUXkjn8&list=PLB755F0C836855EAB&index=56 [11.46 MIBS]
TASK 3.7.1
1. (2008, Q26, 4 marks)
Explain how a buffer works with reference to a specific example in a natural system.
3. (2011, Q25, 3 marks)
Explain the role of the conjugate acid/base pair, H2PO4-/HPO42-, in maintaining the pH of living cells. Include chemical equations in your answer.
3.8 conduct a practical investigation to prepare a buffer and demonstrate its properties
VIEW Videos:
ABR#24 https://www.youtube.com/watch?v=9iOHqJalSQc [7.53 mins]
VIEW PPT
Buffers
12C6 3.7b Buffers PPT.ppt12C6 3.7e Buffer Calculations.docPRACTICAL
PA 6.8 Buffers PP 93-4
How valid is it?
How reliable is it?
How accurate is it? What are the sources of error in the experimental design? How could it be improved?
TASK 3.8.1
Complete WS 6.8 Buffers P72
TASK 3.8.2
1. (2005, Q9) Which of the following pairs would form a buffer solution?
(A) HCl(aq) / Cl−(aq)
(B) H2PO4−(aq) / PO43−(aq)
(C) H2SO4(aq) / HSO4−(aq)
(D) CH3COOH(aq) / CH3COO−(aq)
2. (2012, Q8)
Which of the following acid / base pairs could act as a buffer?
a) H3O+ / H2O b) H2O / OH–
c) HNO3 / NO3- d) H2PO4- / HPO42-
3. (2003, Q15) Which of the following graphs shows how pH will vary when dilute HCl is added to 100 mL of dilute natural buffer solution with an initial pH of 7.0?
A
3. (2003, Q15) Which of the graphs shows how pH will vary when dilute HCl is added to 100 mL of dilute natural buffer solution with an initial pH of 7.0?
A
4. Dilute HCl was added to 100mL of dilute natural buffer solution with an initial pH of 7.0 and the change in pH was plotted against the volume of the HCl added. Explain the shape of the pink line on the graph. (Hegarty, 2018, p. 57)
REVIEW 3.8.1
WS 6.8 Buffers P72
REVIEW 3.8.2
PAST HSC QS (QS WITH NO MARKS INDICATED ARE MULTIPLE CHOICE. HSC EXAMS IN EARLIER YEARS HAD 15 MC QS, CURRENTLY THERE ARE 20.)
2017 Q14
2015 Q24B 3
2012 Q8
2011 Q25 3
2008 Q26 4
2005 Q9
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