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GCSE & A level Chemistry Calculations: Volumetric analysis: acid-alkali titrations

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ACID - ALKALI TITRATIONS - methods and calculations

Analytical methods, procedures and calculations

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study examples carefully12. Introducing Volumetric Analysis - titration calculations e.g. acid-alkali titrations AND introduction to how to do an acid-alkali titration via an antacid indigestion tablet investigation.

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Sub-index for this titration page

1. Introduction - reminders of what molarity is

2. Examples of how to do acid - alkali titration calculations

3. The basic procedure for carrying out an acid - alkali titration

4. How do you choose the right indicator for an acid - alkali titration?

5. Antacid indigestion tablet investigation (an introductory exercise).

6. Why we need to repeat titrations for increased accuracy of results

7. (Titration 1) How to titrate a weak base-alkali like ammonia with standardised hydrochloric acid using methyl orange indicator

8. (Titration 2) Titrating a weak/strong acid with a standard sodium hydroxide solution using phenolphthalein indicator.

9. (Titration 3) Titrating a strong base-alkali (e.g, sodium hydroxide) with a standardised hydrochloric acid using phenolphthalein indicator.

10. (Titration 4) Titrating calcium hydroxide solution with standard hydrochloric acid using phenolphthalein indicator to find out the solubility of calcium hydroxide in water.

and on separate pages

Make a standard solution of known concentration (mass/volume) or molarity (mol/dm3)

On-line Quantitative Chemistry Calculations

See also Advanced level GCE-AS-A2 acid-alkali titration calculation questions

How to make up a standard solution (mass/volume, molarity mol/dm3) on separate page

GCSE/IGCSE self-assessment Quizzes on titrations:

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Keywords: Quantitative chemistry calculations How to you do acid-alkali titration calculations? What is the procedure for doing acid-alkali titrations? How do you do a titration? What apparatus do you need to do a titration? Help for problem solving in doing volumetric titration calculations. Practice revision questions on titrations, using experiment data. This page describes and explains, with fully worked out examples, how to do simple titration calculations involving acids and alkalis. These methods of calculation involve a knowledge of the mole concept e.g. the interconversion of mass-moles-formula mass (mol = mass/Mr) and know how to calculate and use molarity (molarity = mol/volume in dm3). Overall the page gives a description and explanation of simple example of volumetric analysis preceded by how to do volumetric titrations. Online practice exam chemistry CALCULATIONS and solved problems for KS4 Science GCSE/IGCSE CHEMISTRY and basic starter chemical calculations for A level AS/A2/IB courses. These revision notes and practice questions on how to do acid-alkali titration calculations and worked examples should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.


See also GCSE/IGCSE Acid & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids

1. Introduction - reminders of what molarity is
  • This page is all about a quantitative chemical procedure called a titration.
    • We use a titration technique to find out exactly, the concentration of a specified solute in solution.
    • e.g. you can find out precisely how much of an acid solution (this volume is called the titre) is needed to neutralise an accurately measured out volume of an alkali, or the other way round.
    • e.g. if you know the concentration of the acid, together with this, and the two volumes of the solutions required for neutralisation, you can calculate the unknown concentration of the other alkaline solution.
    • How to do titrations to find the concentration of an acid or alkali from the relative volumes used and the concentration of one of the two reactants are described in the method and apparatus sections of further down the page.
  • Before studying this section you should work your way through section 11. Molarity, volumes and solution concentrations
  • You should be able to carry out calculations involving neutralisation reactions in aqueous solution given the balanced equation or from your own practical results.
  • The examples in section 7. moles and mass. and section 11. concentration will help you follow the calculations below.
    • Note again: 1dm3 = 1 litre = 1000ml = 1000 cm3, so dividing cm3/1000 gives dm3.
    • and other useful formulae or relationships are:
      • molarity (mol/dm3) = mol / volume (dm3 = cm3/1000),
      • moles = molarity (mol/dm3) x volume (dm3 = cm3/1000),
      • volume = moles / molarity
      • 1 mole = formula mass in grams.
    • In most volumetric calculations of this type, you first calculate the known moles of one reactant from a volume and molarity.
    • Then, from the equation, you relate this to the number of moles of the other reactant, and then with the volume of the unknown concentration, you work out its molarity.
  • I haven't quoted separately the atomic masses used in the titration calculations, but they quoted as part of the calculation and its pretty obvious what is what!

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2. Examples of how to do acid - alkali titration calculations

  • Titration calculation Example 12.1
    • Given the equation: NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)
      • This is read as 1 mol  +  1 mol  reactants  ===> 1 mol  +  1 mol  products
    • 25.0 cm3 of a sodium hydroxide solution was pipetted into a conical flask and titrated with a standard solution of 0.200 mol dm-3 (0.2M) hydrochloric acid (mol dm-3 means mol/dm3).
    • Using phenolphthalein indicator for the titration it was found that 15.0 cm3 of the acid was required to neutralise the alkali. In the appendix this is titration procedure 3.
      • Calculate the molarity of the sodium hydroxide and its concentration in g/dm3.
        • moles = molarity x volume (in dm3 = cm3/1000)
        • moles HCl = 0.200 x (15.0/1000) = 0.003 mol
        • moles HCl = moles NaOH (1 : 1 in equation)
        • so there is 0.003 mol NaOH in 25.0 cm3
        • scaling up to 1000 cm3 (1 dm3), there are ...
        • 0.003 x (1000/25.0) = 0.12 mol NaOH in 1 dm3
        • molarity of NaOH is 0.120 mol dm-3  (or 0.12M)
        • since mass = moles x formula mass
        • and Mr(NaOH) = 23 + 16 + 1 = 40
        • concentration in g/dm3 = molarity x formula mass
        • concentration in g/dm3 is 0.12 x 40 = 4.80 g/dm3 

     

  • Titration calculation Example 12.2
    • Given the equation: 2KOH(aq) + H2SO4(aq) ==> K2SO4 + 2H2O(l)
      • This is read as 2 mol  +  1 mol  reactants  ===> 1 mol  +  2 mol  products
    • 20.0 cm3 of a sulphuric acid solution was titrated with a standardised solution of 0.0500 mol/dm3 (0.05M) potassium hydroxide.
    • Using phenolphthalein indicator for the titration, the acid required 36.0 cm3 of the alkali KOH for neutralisation what was the concentration of the acid?  In the appendix this is titration procedure 2.
      • moles = molarity x volume (in dm3 = cm3/1000)
      • mol KOH = 0.0500 x (36.0/1000) = 0.0018 mol
      • mol H2SO4 = mol KOH / 2 (because of 2 : 1 ratio in equation above)
      • mol H2SO4 = 0.0018/2 = 0.0009 (in 20.0 cm3)
      • scaling up to 1000 cm3 of solution = 0.0009 x (1000/20.0) = 0.0450 mol
      • mol H2SO4 in 1 dm3 = 0.0450
      • so molarity of H2SO4 = 0.0450 mol dm-3 (0.045M)
      • since mass = moles x formula mass
      • and Mr(H2SO4) = 2 + 32 + (4x16) = 98
      • concentration in g/dm3 is 0.045 x 98 = 4.41 g/dm3 
    • -
  • Titration Calculation Example 12.3
    • Given the equation: NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)
      • This is read as 1 mol  +  1 mol  reactants  ===> 1 mol  +  1 mol  products
    • 25.00 cm3 portions of a dilute hydrochloric acid solution were titrated with a standard solution of sodium hydroxide of concentration 0.250 mol/dm3 (mol dm-3). In the appendix this is titration procedure 2.
    • Using phenolphthalein indicator for the titration, it was found that the average titration was 18.50 cm3 of sodium hydroxide, calculate (i) the molarity of the hydrochloric acid and (ii) its concentration in g/dm3.
      • moles NaOH = molarity NaOH x volume of NaOH (in dm3 = cm3/1000)
      • moles NaOH = 0.250 x (18.5 / 1000) = 0.004625
      • In the equation 1 mole of HCl reacts with 1 mole of HCl
      • therefore in the titration reaction: moles HCl = moles NaOH
      • therefore there were 0.004625 moles HCl in 25.00 cm3.
      • molarity = moles / volume in dm3     (1 dm3 = 1000 cm3)
      • (i) molarity HCl = 0.004625 / (25.00/1000) = 0.004625/0.025 = 0.185 mol/dm3
      • (ii) concentration = molarity x formula mass
        • formula mass HCl = 1 + 35.5 = 36.5
        • = 0.185 x 36.5 = 6.75 g/dm3
    • -
  • Titration Calculation Example 12.4
    • Given the equation: 2NaOH(aq) + H2SO4(aq) ==> Na2SO4 + 2H2O(l)
      • This is read as 2 mol  +  1 mol  reactants  ===> 1 mol  +  2 mol  products
    • 25.00 cm3 portions of a sodium hydroxide were titrated with a standardised solution of 0.75 mol/dm3 sulphuric acid solution using phenolphthalein indicator. In the appendix this is titration procedure 1.
    • If the average titration was 17.70 cm3 of sulfuric acid, what is the molar concentration of the sodium hydroxide?
      • moles H2SO4 in titration = molarity H2SO4 x volume in dm3
      • moles H2SO4  = 0.75 x (17.70/1000) = 0.013275 mol
      • From the balanced equation, for every mole of H2SO4, two moles of NaOH react
      • Therefore moles NaOH = 2 x moles H2SO4
      • moles NaOH = 0.013275 x 2 = 0.02655 mol
      • molarity of NaOH = moles NaOH / volume in dm3
      • molarity NaOH = 0.02655 / (25.00/1000) = 1.062 mol/dm3
    • -
  • Titration Calculation Example 12.5
    • Given the equation: NH3(aq) + HCl(aq) ==> NH4Cl(aq)
      • This is read as 1 mol  +  1 mol  reactants  ===> 1 mol  products
    • 5.00 cm3 portions of household ammonia were titrated with a standard hydrochloric acid solution of 1.00 mol/dm3. In the appendix this is titration procedure 2.
    • If the average titration, using methyl orange indicator, was 22.5 cm3 of hydrochloric acid, calculate ...
    • (i) the molarity of the ammonia solution, and its concentration in (ii) g/dm3 and (iii) g/cm3.
      • (i) mol HCl in titration = molarity HCl x volume of HCl in dm3
      • mol HCl = 1.00 x (22.50/1000) = 0.0225 mol HCl
      • From the equation, 1 mole of NH3 reacts with 1 mole of HCl
      • Therefore mol HCl  = mol NH3 = 0.0225
      • molarity of NH3 = mol NH3 / volume NH3in dm3
      • molarity of NH3 = 0.0225 / (5.00/1000) = 4.50 mol/dm3
      • (ii) concentration = formula mass x molarity
      • formula mass NH3 = 14 + (3x1) = 17
      • concentration of NH3 = 17 x 4.50 = 76.5 g/dm3
      • (iii) since there are 1000 cm3 in 1 dm3
      • concentration of NH3 = 76.5/1000 = 0.0765 g/cm3
    • -
  • Titration Calculation Example 12.6
    • Given the equation: CH3COOH(aq) + NaOH(aq) ==> CH3COONa(aq) + H2O(l)
      • This is read as 1 mol  +  1 mol  reactants  ===> 1 mol  +  1 mol  products
    • 25.00 cm3 portions of vinegar (ethanoic acid, CH3COOH, 'acetic acid' solution) from a local supermarket were pipetted into a conical flask and titrated with a standardised solution of sodium hydroxide, of concentration 0.2000 mol/dm3 using phenolphthalein indicator. In APPENDIX 2 this is titration is described in procedure 2.
    • (i) If the average titration value was 14.70 cm3 of the sodium hydroxide solution, what is the molarity of the ethanoic acid in the vinegar?
      • moles NaOH in titration = molarity x volume in dm3
      • mole NaOH = 0.200 x (14.70/1000) = 0.00294 moles
      • from the equation moles NaOH = moles CH3COOH = 0.00294
      • molarity CH3COOH = mol CH3COOH/volume CH3COOH in dm3
      • molarity CH3COOH = 0.00294 / (25.00/1000) =  0.1176 mol/dm3
    • (ii) What is the concentration of the ethanoic acid in g/dm3?
      • mass = moles x formula mass
      • formula mass of CH3COOH = 12 + 3 + 12 + 16 + 16 + 1 = 60
      • mass CH3COOH = 0.1176 x 60 = 7.056g (in 1 dm3, 1000 cm3)
      • concentration = 7.06 g/dm3 (2dp, 3sf)
        • for other mass/volume values ...
        • eg if you were asked for the concentration in g/100cm3,
          • divide the g/dm3 by 10 giving 0.706g/100cm3
        • or if you were asked for the concentration in g/cm3,
          • divide the g/dm3 by 1000 giving 0.00706 g/100cm3
      • You can also work out the vinegar concentration from the original moles and volume of the titration eg
      • mass = moles x formula mass
      • mass = 0.00294 x 60 = 0.1764g in the 25 cm3 sample titrated.
      • Since 25cm3 = 25/1000 = 0.025 dm3
      • concentration = 0.1764 / 0.025 = 7.06g/dm3
    • (iii) If a bottle of vinegar from the super-market contained 250 cm3 of liquid, how many grams of ethanoic acid are in the solution?
      • concentration CH3COOH = formula mass CH3COOH x molarity of CH3COOH
      • formula mass CH3COOH = 12 + (3x1) + 12 +(2x32) + 1 = 60
      • concentration CH3COOH = 60 x 0.1176 = 7.056 g/dm3
      • Since 1 dm3 = 1000 cm3, there must be proportionately
      • 7.056 x (250/1000) = 1.76 g of CH3COOH in the bottle of vinegar (3 sf)
    • -
  • Titration Calculation Example 12.7 A titration to determine the solubility of calcium hydroxide in water.
    • Approximately 250 cm3 of water was shaken with solid calcium hydroxide (slaked lime) until no more appeared to dissolve and the solution filtered to remove any excess undissolved solid. You now have a saturated solution known as limewater.
    • 25.00 cm3 of the filtered calcium hydroxide solution (limewater) was pipetted into a conical flask and titrated with standard hydrochloric solution (0.100 mol/dm3) using phenolphthalein indicator until the pink colour just disappeared (the end-point). The experiment was repeated several times giving an average HCl titration of 10.0 cm3.
    • Calculate the concentration of the calcium hydroxide in mol/dm3 and g/dm3.
    • The equation for the titration is: Ca(OH)2 + 2HCl ==> CaCl2 + 2H2O
      • This is read as 1 mol  +  2 mol  reactants  ===> 1 mol  +  2 mol  products
    • moles = molarity x volume (dm3)
    • mol HCl in titration = 0.10 x 10.0/1000 = 0.0010
    • Therefore from the equation: mol Ca(OH)2 = mol HCl/2 = 0.0005
    • Molarity of Ca(OH)2 = moles/volume = 0.0005 / (25.00/1000) = 0.0005 / 0.025 = 0.02 mol/dm3
    • concentration g/dm3 = molarity x formula mass
    • Atomic masses: Ca = 40, O = 16, H = 1, Mr[CaOH)2] = 74
    • Concentration of Ca(OH)2 = 0.02 x 74 = 1.48 g/dm3
    • -

NOTE: If you have a question based on nitric acid (HNO3), they are the same as for hydrochloric acid (HCl) because both acids have one acidic proton (H), so balancing equations involves the same numbers.

Self-assessment quizzes and questions on acid - alkali titrations

Type in answer click me for QUIZ!Honly  or multiple choice click me for QUIZ!Honly

Advanced level GCE-AS-A2 acid-alkali titration calculation questions

More questions involving molarity in

section 11. introducing molarity   and   section 14.3 on dilution


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3. The basic procedure for carrying out an acid - alkali titration

volumetric apparatus for a titration volumetric apparatus

A variety of apparatus you might come across, particularly the pipette (1) for measuring accurately volumes of solutions to be analysed by titration with a standard solution in a burette (3) with the chemicals mixing in a conical flask (4).

  • The diagram above shows the typical apparatus (1)-(6) used in manipulating liquids and on the left a brief three stage description of titrating an acid with an alkali:
    • (i) An accurate volume of acid is pipetted into the conical flasks using a suction bulb and pipette for health and safety reasons. Universal indicator is then added, which turns red in the acid.
    • The alkali, of known accurate concentration, is put in the burette and you can conveniently level off the reading to zero (the meniscus on the liquid surface should rest on the zero -- graduation mark).
      • Note other possibilities are:
        • (ii) An accurate volume of alkali is measured into a flask and titrated with an acid solution of known concentration.
        • (iii) A small amount of accurately weighed solid acid is dissolved in water and titrated with alkali.
        • (iv) A small amount of accurately weighed solid alkali is dissolved in water and titrated with acid.
          • This can procedure can be used to compare the effectiveness of ant-acid indigestion tablets, which are designed to neutralise excess acid in the stomach. (more details further down the page)
          • A known and equal mass of each brand of indigestion tablet is crushed and mixed with some water eg 20 cm3 (fair test points).
          • Make sure the mixture is gently swirled to completely dissolve the crushed tablet powder.
          • The burette is filled with a standard solution of hydrochloric acid and zeroed to the top calibration mark of 0.00 cm3.
          • Universal indicator is added to the flask and the indigestion powder should turn it blue - alkaline.
          • The acid is carefully and slowly added until the indicator turns green - neutral at the end-point of the titration.
          • You then read the volume of acid required to neutralise the ant-acid powder.
          • The bigger the volume of acid required for neutralisation, the more effective the indigestion powder per mass of powder.
          • Repeat the procedure with another brand of indigestion powder using the same standard acid solution (fair test).
    • For (i) The alkali is then carefully added by running it out of the burette in small quantities into the acid solution, controlling the flow with the tap, until the indicator seems to be going yellow-pale green.
      • The conical flask should be carefully swirled after each addition of alkali to ensure all the alkali reacts.
    • Near the end of the titration, the alkali should added drop-wise until the universal indicator goes green.
      • This is called the end-point of the titration and the green means that all the acid has been neutralised.
      • The volume of alkali needed to titrate-neutralise the acid is read off from burette scale, again reading the volume value on the underside of the meniscus.
      • The calculation can then be done to work out the concentration of the alkali.

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4. How do you choose the right indicator for an acid - alkali titration?

  • Universal indicator, and most other acid-base indicators, work for strong acid and alkali titrations, but universal indicator is a somewhat crude indicator for other acid-alkali titrations because it gives such a range of colours for different pH's. Examples of more accurate and 'specialised' indicators are:
    • titrating a strong alkali with a strong acid (or vice versa):
      • e.g. for sodium hydroxide (NaOH) - hydrochloric/sulphuric acid (HCl/H2SO4) titrations, use ...
      • phenolphthalein indicator (pink in alkali, colourless in acid-neutral solutions), the end-point is the pink <==> colourless change.
      • Litmus works too, the end point is the red <==> purple/blue colour change AND with universal indicator you see green at the endpoint, but these two indicators are NOT as accurate as phenolphthalein.
        • They are ok as simple demonstration to illustrate the principles of a titration, BUT, they should not be used for quantitative work.
    • titrating a weak alkali with a strong acid:
      • e.g. for titrating ammonia (NH3) with hydrochloric/sulfuric acid (HCl/H2SO4), use ...
      • methyl orange indicator (red in acid, yellowish-orange in neutral-acid), the end-point is an 'orange' colour, not easy to see without experience.
      • screened methyl orange indicator is a slightly different dye-indicator mixture that is reckoned to be easier to see than methyl orange, the end-point is a sort of 'greyish orange', but still not easy to do accurately.
    • titrating a weak acid with a strong alkali:
      • e.g. for titrating ethanoic acid (CH3COOH) with sodium hydroxide (NaOH), use ...
      • phenolphthalein indicator (pink in alkali, colourless in acid-neutral solutions, pink in alkali), the end-point is the first permanent pink. when the solution just becomes alkaline.
      • methyl red indicator (red in acid, yellow in neutral-alkaline), the end-point is 'orange'.
    • titrating a weak acid with a weak alkali (or vice versa):
      • These are NOT practical titrations because the pH changes at the end-point are not great enough to give a sharp colour change with any indicator.

Indicator

colour in acid pH<7 colour in neutral pH=7 colour in alkali pH >7

litmus

red 'purple' blue
phenolphthalein colourless colourless >9 pink
methyl orange <3.5 red, orange about pH 5, > 6 yellow yellow yellow
methyl red <5 red, orange, >6 yellow yellow yellow
bromothymol blue <6 yellow green >8 blue

* Used in titrations - see section 10, litmus is NOT a good indicator for use in accurate titrations.

titration indicator colours

Summary: Three common indicators for titrations and their colours:

litmus: red in acid, blue in alkali, doesn't give that sharp a colour change except with moderately concentrated solutions, the neutralisation end-point is a fairly sharp change red<=>blue

phenolphthalein: colourless in acid, pink in alkali, the neutralisation end-point is the sharp change colourless<=>pink

methyl orange: red in acid, yellow in alkali, the neutralisation end-point is an orange


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EXAMPLES OF TITRATIONS

5. An Antacid Indigestion Tablet Investigation

This first titration description also acts as an introduction as how to do a TITRATION with a burette and conical flask etc. AND it doesn't involve complex titration calculations, other titrations are described in Appendix 2. which do involve calculations.

There is also a paragraph on errors and reasons for repeating a titration several times.

Introduction

There are many brands of antacid indigestion medications on the market and through the following experimental investigation you can check out their value for money. There are various ways you can approach the investigation e.g. you can compare tablet with tablet in terms of the recommended dose as to which neutralises the most acid. You can compare the cost of each tablet with the amount of acid it neutralises.

Experimentally, the investigation involves titrating an indigestion tablet (weighed) with standardised hydrochloric acid (in the burette, to represent stomach acid) using methyl orange indicator to observe when the tablet has been completely neutralised, that point is known as the end-point of the titration.

The apparatus, chemicals and indicator colours are illustrated in the diagram on the left and the procedure described in five stages below.

 

Preparing the sample and titration procedure

Initially the burette is clamped carefully in a vertical position and filled with standard hydrochloric acid of known concentration. A burette is a long glass tube (open at the top), and accurately calibrated for volume in cm3 an 1/10th cm3 intervals, with a tap and tip at the lower end.

For safety reasons wear safety glasses, initially work below eye level, using a funnel, the acid is carefully added from the stock bottle down in to the burette, avoid spillage, until the level is above the 0.00 cm3 scale mark.

This burette filling is done below eye level in case of spillage down into your eyes - goggles/safety glasses, which you should be wearing, do not protect from liquid falling onto your forehead! 

The acid is run through to expel any air bubbles in the tip or tap until the reading below the meniscus is 0.00 cm3 (the reading on the above diagram is 7.00 cm3, which could represent a small titration value).

All burette readings should be made at eye level and taking the value exactly below the meniscus on the burette scale.

The burette is usually calibrated to a maximum 50.00 cm3 (only 10.00 cm3 in diagram - I couldn't fit rest of scale on!). Now we are ready to take the antacid indigestion tablet!

The weighed tablet is crushed up and dissolved in e.g. 25 cm3 or 50 cm3 of pure water (for other tablets keep to the same volume of water as part of the fair test).

Add a few drops of methyl orange indicator to the tablet solution and it should turn yellow for an alkali. Carefully place the conical flask under the tip of the burette so drops don't go astray!

The titration: You carefully add small portions of the hydrochloric acid, swirling after each addition and checking the colour of the indicator (although not shown in the diagram, its good to stand the flask on white tile to see the colour changes better).

At the start of the titration the methyl orange indicator is yellow. As you add the acid you get 'splurges' of reddish-orange colour until the mixture is swirled in the conical flask and the yellow is temporarily restored.

The swirling of the flask contents is important, it ensures all the added hydrochloric acid reacts with the antacid indigestion tablet solution i.e. everything gets well mixed up and reacted.

Try to add dropwise (to avoid overshooting) when you seem to be near the orange colour at the end-point as the yellow indicator colour begins to fade.

The indicator colour at the end-point is orange and indicates all the dissolved tablet has been neutralised.

At the end-point you take the titration reading by carefully reading the burette scale under the meniscus (think of the underside of the meniscus as sitting on your reading - see diagram on the right illustrating a reading of 24.5 cm3).

To get the titration value you subtract the1st reading from the 2nd.

The first reading might be zero (0.00 cm3) BUT you can do subsequent titrations without refilling the burette every time, again you just subtract the 1st reading from the 2nd.

You continue to use the burette like this until in needs refilling for further titrations.

If you 'overshoot' the titration with excess acid, the methyl orange indicator turns red and the result is invalid.

The first titration you do is likely to be the most inaccurate until you 'get your eye in'. This is called the rough titration and you can do the next titration quite rapidly until near the end-point and then proceed slowly and accurately to the end-point itself. 

The titration should be repeated several times with the same brand of and the average (mean) titration value calculated to use in any subsequent calculations.

This makes the experimental results more valid and reliable, as will any subsequent calculations and conclusions based on the data recorded.

The procedure should then be repeated with different brands of antacid indigestion tablets and the results compared.

You should keep to the same volumes of water and the same concentration of hydrochloric acid throughout the whole class/individual investigation.

Using a whole class you could amass quite a bit of data by dividing the work up amongst the pupils.

Data and analysis of the results

There are various ways in which you can interpret the results, so here are a few ideas.

(a) Initially you can compare the volume of acid needed to neutralise an individual tablet, which is simply X cm3 of HCl neutralised per tablet. This gives a straightforward comparison, the bigger the titration the more stomach acid would be neutralised.

(b) If you have weighed the tablet, which I recommend you do, you can compare the effectiveness of the antacid tablets in terms of acid per mass of tablet i.e. X cm3 of HCl neutralised per gram tablet (cm3/g). So this measure the effectiveness of the tablet based on mass ('weight').

(c) If you know the cost of the packet of indigestion tablets, you can work out the cost of an individual tablet. Then you can calculate the 'cost effectiveness' of the medication by dividing the titration value by the cost per tablet e.g. X cm3 of HCl neutralised per cost of tablet (cm3/p)


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6. Extra note - Why repeat titrations several times?
 
  • To be reasonably sure of your final result, you cannot base it on one titration, no matter how perfectly you think you have done it.
  • You need repeated AND consistent burette readings.
  • There can be many source of errors, some random, and, worst of all, some systematic error to do with your laboratory equipment or your titration technique.

    Errors may arise from ...

    dirty apparatus, grease in the burette can cause traces of liquid to remain, so giving a false larger than actual titration reading,

    sometimes a result doesn't fit in at all with the other results (referred to as an 'outlier'), may be due to carelessness in observing the end-point or poor pipetting etc. - human error,

    apparatus may be faulty,

  • If your results are close together, this is a good sign - consistent readings imply you have been accurate in your titrations.

    Your first titration is very likely to be a bit inaccurate, and is referred to as the initial rough titration.

    This gives you an approximate value to aim for much more accurately in subsequent titrations.

  • After repeating the titrations accurately e.g. three times, you then average your results and use the average in your calculations.

    You titration, at best, is likely to be measured to the nearest 0.05 cm3, but you can use the average to two decimal places because statistically, that is the most probable value.

    e.g. suppose you got burette reading titration values of 21.45, 21.40, 21.45 (which would be impressive!),

    the average is 21.43 cm3, and that goes into the calculation of ....

    However, if a titration reading seems way out, 'posh word', anomalous, it should be ignored and not included in you calculation e.g. if a titration value of 21.75 was obtained, its not close enough to the other three values to be considered of value.

 

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APPENDIX 2 - More examples of HOW TO DO TITRATIONS - apparatus and procedures

7. Titration 1. The titration of a weak base-alkali with a strong acid

e.g titrating ammonia (pipetted) with standardised hydrochloric acid (of known concentration in the burette) using methyl orange indicator

The apparatus, chemicals and indicator colours are illustrated in the diagram on the left.

Initially the burette is clamped carefully in position and filled with standard hydrochloric acid (e.g. 0.10 to 1.0 mol/dm3, but accurately known, preferably to 4 sig. figs.).

You wear safety glasses and work below eye level when filling the burette safely using a funnel until a bit above the 0.00 cm3 mark.

The acid is run through until the reading below the meniscus is 0.00 cm3 (the reading in the diagram is 7.00 cm3, which could represent a small titration value). The burette is usually calibrated on the scale to 50.00 cm3 in 0.10 cm3 increments (only 10.00 cm3 in diagram - couldn't fit rest of scale on!)

The ammonia solution is accurately measured out into the conical flask with e.g. a 25 cm3 pipette and suction bulb (see diagram further down).

You suck the ammonia solution up to a few cm above the calibration mark and then let it run down until the meniscus rests on the calibration mark (top of right diagram).

Add a few drops of methyl orange indicator to the ammonia solution and it should turn yellow for an alkali. Carefully place the conical flask under the tip of the burette so drops don't go astray!

The titration: You carefully add small portions of the acid, swirling after each addition and checking the colour of the indicator (not shown in the diagram, but its good to stand the flask on white tile).

At the start of the titration the methyl orange indicator is yellow.

As you add the acid you get 'splurges' of reddish-orange colour until the mixture is swirled in the conical flask.

The swirling of the flask contents is important, it ensures all the added hydrochloric acid reacts with the ammonia solution.

Try to add dropwise when you seem to be near the orange colour at the endpoint.

The end-point is an orange colour indicating when all the ammonia is neutralised by which ever acid you are using.

You taking the reading of the volume of the hydrochloric acid on the scale line the meniscus lies on (see lower part of diagram above on the right).

If you 'overshoot' the titration with excess acid, the methyl orange indicator turns red and the result is invalid.

The titration should be repeated several times with other 25 cm3 portions and the average (mean) titration value calculated to use in any subsequent calculations.

How to calculate the concentration of the weak alkali is explained in the top half of the page.

The theory of which indicator to use is explained on the Changes in pH in a neutralisation, choice and use of indicators page.

There is also a paragraph on errors and reasons for repeating a titration several times.

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8. Titration 2. The titration of a weak/strong a strong acid with sodium hydroxide solution

e.g titrating ethanoic acid or hydrochloric acid (pipetted) with standardised sodium hydroxide solution (of known concentration in burette) using phenolphthalein indicator

The apparatus, chemicals and indicator colours are illustrated in the diagram on the left.

Initially the burette is clamped carefully in position and using a funnel, it is filled with standard sodium hydroxide solution (e.g. 0.10 to 1.0 mol/dm3, but accurately known, preferably to 4 sig. figs.).

Wearing goggles, pour in  using a funnel and work below eye level.

The sodium hydroxide solution is run through until the reading below the meniscus is 0.00 cm3 (the reading in the diagram is 7.00 cm3, which could represent a titration value). The burette is usually calibrated on the scale to 50.00 cm3 in 0.10 cm3 increments (only 10.00 cm3 in diagram - couldn't fit rest of scale on!)

The acid solution is accurately measured out into the conical flask with e.g. a 25 cm3 pipette and suction bulb (see diagram further down).

Add a few drops of phenolphthalein indicator to the acid solution and it should turn colourless for an acid. Carefully place the conical flask under the tip of the burette so drops don't go astray!

The titration: You carefully add small portions of the sodium hydroxide, swirling after each addition and checking the colour of the indicator (not shown in the diagram, but its good to stand the flask on white tile).

At the start of the titration the phenolphthalein indicator is colourless.

As you add the alkali you get 'splurges' of pink colour until the mixture is swirled in the conical flask.

The swirling of the flask contents is important, it ensures all the added sodium hydroxide reacts with the acid in the flask.

Try to add dropwise when you seem to be near the faint pink colour of the endpoint.

The end-point is the first faint, but permanent pink colour, that is when all the acid is neutralised by the sodium hydroxide.

You taking the reading of the volume of the sodium hydroxide on the scale line the meniscus lies on (see lower part of diagram above on the right).

If you 'overshoot' the titration with excess alkali, the phenolphthalein indicator becomes an even deeper pinkish-red and the result is invalid.

The titration should be repeated several times with other 25 cm3 portions and the average (mean) titration value calculated to use in any subsequent calculations.

How to calculate the concentration of the weak acid is explained in the top half of the page.

The theory of which indicator to use is explained on the Changes in pH in a neutralisation, choice and use of indicators page.

There is also a paragraph on errors and reasons for repeating a titration several times.

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9. Titration 3. The titration of a strong base-alkali with hydrochloric solution

e.g titrating sodium hydroxide solution (pipetted) with standardised hydrochloric solution (of known concentration in burette) using phenolphthalein indicator

The apparatus, chemicals and indicator colours are illustrated in the diagram on the left.

Initially the burette is clamped carefully in position and filled with standard hydrochloric acid (e.g. 0.10 to 1.0 mol/dm3, but accurately known, preferably to 4 sig. figs.).

Wear safety glasses and initially work below eye level - so nothing can fall on your face!

Using a funnel, the hydrochloric acid is run through until the reading below the meniscus is 0.00 cm3 (the reading in the diagram is 7.00 cm3, which could represent a titration value). The burette is usually calibrated on the scale to 50.00 cm3 in 0.10 cm3 increments (only 10.00 cm3 in diagram - couldn't fit rest of scale on!)

The alkali solution is accurately measured out into the conical flask with e.g. a 25 cm3 pipette and suction bulb (see diagram further down).

Add a few drops of phenolphthalein indicator to the alkali solution and it should turn deep pink for an alkali. Carefully place the conical flask under the tip of the burette so drops don't go astray!

The titration: You carefully add small portions of the acid, swirling after each addition and checking the colour of the indicator (not shown in the diagram, but its good to stand the flask on white tile).

At the start of the titration the phenolphthalein indicator is a deep reddish-pink in alkaline solution.

As you add the acid you get 'splurges' of colourless solution until the mixture is swirled in the conical flask.

The swirling of the flask contents is important, it ensures all the added hydrochloric acid reacts with the sodium hydroxide.

Try to add dropwise when you seem to be near the faintest pinkness left in the solution near the endpoint.

The end-point is when the last trace of pink colour first disappears from the solution.

You taking the reading of the volume of the hydrochloric acid on the scale line the meniscus lies on (see lower part of diagram above on the right).

 If you 'overshoot' the titration with excess acid, it still stays colourless and the result is invalid.

The titration should be repeated several times with other 25 cm3 portions and the average (mean) titration value calculated to use in any subsequent calculations.

How to calculate the concentration of the alkali is explained in the top half of the page.

The theory of which indicator to use is explained on the Changes in pH in a neutralisation, choice and use of indicators page.

There is also a paragraph on errors and reasons for repeating a titration several times.

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10. Titration 4. Determination of the concentration of calcium hydroxide in limewater

This almost the same experiment as APPENDIX 2. Example 3. above

A practical exercise for Advanced A level students

Calcium hydroxide is slightly soluble in water and the resulting solution is commonly known as 'limewater.

It is a relatively strong alkali and readily titrated with standardised hydrochloric acid using phenolphthalein indicator.

titrating calcium hydroxide solution with hydrochloric acid using phenolphthalein indicatorRequirements

  • A bottle of water shaken well for quite a few minutes with excess solid calcium hydroxide and left to stand so that the excess solid settles out.

    • It should be left to stand for too long, otherwise the solution absorbs the acidic gas carbon dioxide.

  • 50 cm3 burette (with 0.1 cm3 graduations on the scale)

  • 25.0 cm3 pipette (you can use a 10 cm3 pipette, but smaller and less accurate titration)

  • 250 cm3 conical flask and white tile

  • Squeezy wash bottle of deionised/distilled water.

  • Standardised 0.1 or 0.05 mol dm-3 (M for short) hydrochloric acid.

  • Phenolphthalein acid-base indicator solution (pink >pH 9, colourless <pH 9

Method-Procedure

  • how to use a calibrated pipetteWear safety glasses, initially work below eye level and use a funnel to fill the burette.

  • Fill the burette with the standard acid solution of known concentration and level off so that the bottom of the meniscus rests on the 0.00 calibration mark.

  • Using a safe suction bulb, suck the calcium hydroxide solution until its a few cm above the calibration mark, then run it down until the meniscus rests on the calibration mark (diagram on right).

  • Pipette 25.0 cm3 of the clear liquid above the excess solid in the bottle of limewater into the conical flask.

  • Add a few drops of phenolphthalein indicator solution (it should turn pink - alkaline)

  • Place the conical flask on the white tile under the burette tip.

  • how to use a burette how to read the scale on a burette line under meniscusCarefully titrate the limewater until ONE drop of the hydrochloric acid removes the last of the pink colour (the end-point is pink to colourless).

  • Take the reading of the burette scale - the line on which the meniscus rests (right diagram).

  • Repeat the procedure at least twice to obtain good concordance of titration results.

  • Your results should be over a maximum of 0.2 cm3 and +/- 0.1 cm3 range would be very good.

  • In doing subsequent titrations you do not have to keep on

Typical Results and Calculations ('borrowed' Q19 from one of my volumetric titration calculation pages)

The solubility of calcium hydroxide in water can be measured reasonably accurately to 3sf by titrating the saturated solution with standard hydrochloric acid.

In the calculation below assume the molarity of the standardised hydrochloric acid is 0.1005 mol dm-3.

At 25oC, a few grams of solid calcium hydroxide was shaken with about 400 cm3 of deionised water, and then filtered.

  1st 2nd 3rd 4th 5th -
2nd burette reading/cm3           average
1st burette reading/cm3           titration
titration value/cm3            
titration value used in average           -

50.0 cm3 samples of the 'limewater' gave an average titration of 15.22 cm3 of 0.1005 mol dm-3 hydrochloric acid using phenolphthalein indicator.

(a) The equation for calcium hydroxide reacting with hydrochloric acid.

Ca(OH)2(aq) + 2HCl(aq) ==> CaCl2(aq) + 2H2O(l)

(b) The reacting mole ratio of Ca(OH)2 : HCl and hence calculate the moles of them involved in the titration.

from equation, mole ratio Ca(OH)2 : HCl is 1:2

and since moles solute = molarity x volume in dm3   (dm3 = cm3/1000)

mol HCl used in titration = 0.1005 x 15.22/1000 = 0.001530 mol HCl

therefore mol Ca(OH)2 = 0.001558/2 = 0.000765 mol Ca(OH)2

(c) Calculate the molarity of the solution in terms of mol Ca(OH)2 dm-3.

Scaling up from mol Ca(OH)2 in 50 cm3 to 1 dm3 (1000 cm3)

molarity Ca(OH)2 = 0.00765 x 1000/50 = 0.1558 = 0.153 mol dm-3 (to 3sf)

(d) What is the approximate solubility of calcium hydroxide in g Ca(OH)2 per 100g water?

Mr[Ca(OH)2] =74, mass = moles x formula mass

so solubility in g dm-3 = 0.153 x 74 = 1.13 g dm-3 (3 sf)

(If needed in another question, this is equal to 1.13/1000 = 1.13 x 10-3 g/cm3 water)

Since density of water is ~1.0 g cm-3, so the solubility per 100g is 1/10th of the solubility per dm3,

solubility of Ca(OH)2 is about 0.113 g/100 g H2O

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Self-assessment Quiz on titrations

type in titration answer click me for QUIZ! QUIZ  or multiple choice titration click me for QUIZ! QUIZ

See also Advanced level GCE-AS-A2 acid-alkali titration calculation questions


See also GCSE/IGCSE Acid & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids


OTHER CALCULATION PAGES

  1. What is relative atomic mass?, relative isotopic mass and calculating relative atomic mass

  2. Calculating relative formula/molecular mass of a compound or element molecule

  3. Law of Conservation of Mass and simple reacting mass calculations

  4. Composition by percentage mass of elements in a compound

  5. Empirical formula and formula mass of a compound from reacting masses (easy start, not using moles)

  6. Reacting mass ratio calculations of reactants and products from equations (NOT using moles) and brief mention of actual percent % yield and theoretical yield, atom economy and formula mass determination

  7. Introducing moles: The connection between moles, mass and formula mass - the basis of reacting mole ratio calculations (relating reacting masses and formula mass)

  8. Using moles to calculate empirical formula and deduce molecular formula of a compound/molecule (starting with reacting masses or % composition)

  9. Moles and the molar volume of a gas, Avogadro's Law

  10. Reacting gas volume ratios, Avogadro's Law and Gay-Lussac's Law (ratio of gaseous reactants-products)

  11. Molarity, volumes and solution concentrations (and diagrams of apparatus)

  12. How to do acid-alkali titrations and calculations, diagrams of apparatus, details of procedures (this page)

  13. Electrolysis products calculations (negative cathode and positive anode products)

  14. Other calculations e.g. % purity, % percentage & theoretical yield, dilution of solutions (and diagrams of apparatus), water of crystallisation, quantity of reactants required, atom economy

  15. 14.1 % purity of a product 14.2a % reaction yield 14.2b atom economy 14.3 dilution of solutions

  16. 14.4 water of crystallisation calculation  14.5 how much of a reactant is needed? limiting reactant

  17. Energy transfers in physical/chemical changes, exothermic/endothermic reactions

  18. Gas calculations involving PVT relationships, Boyle's and Charles Laws

  19. Radioactivity & half-life calculations including dating materials

  20. Some Advanced A Level Practical Exercises and Calculations

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