describe the steps of an acid base titration experiment

• Analytical Chemistry

Acid Base Titration

An acid–base titration is a method of quantitative analysis for determining the concentration of an acid or base by exactly neutralizing it with a standard solution of base or acid having known concentration.

Table of Content

What is acid-base titration, recommended videos.

• Titration Curve

Solved Example

• Frequently Asked Questions – FAQs

An acid-base titration is an experimental technique used to acquire information about a solution containing an acid or base.

Hundreds of compounds both organic and inorganic can be determined by a titration based on their acidic or basic properties. Acid is titrated with a base and base is titrated with an acid. The endpoint is usually detected by adding an indicator.

Learning Objective

To determine the concentration of an acid or base by measuring the volume of titrant (of known concentration) that reacts with it according to a stoichiometric proton-transfer reaction.

An acid-base titration involves strong or weak acids or bases. Specifically, an acid-base titration can be used to figure out the following.

• The concentration of an acid or base
• Whether an unknown acid or base is strong or weak.
• pK a of an unknown acid or pK b of the unknown base.

Let us consider acid-base reaction which is proceeding with a proton acceptor. In water, the proton is usually solvated as H 3 O + . H 2 O is added to the base to lose (OH – ) or gain (H 3 O + ). Acid-base reactions are reversible.

The reactions are shown below.

HA + H 2 O → H 3 O + + A – (acid)

B – + H 2 O → BH + OH – (base)

Here [A – ] is the conjugate base, B-H is conjugate acid. Thus we say

Acid + Base ⇋ Conjugate base + Conjugate acid

It is possible to give an expression for [H + ] in terms of K A , K B and K w for a combination of various types of strong and weak acids or bases.

• Titration – A process where a solution of known strength is added to a certain volume of a treated sample containing an indicator.
• Titrant – A solution of known strength of concentration used in the titration.
• Titrand – The titrand is any solution to which the titrant is added and which contains the ion or species being determined.
• Titration curve – A plot of pH Vs millilitres of titrant showing the manner in which pH changes Vs millilitres of titrant during an acid-base titration.
• Equivalence point – The point at which just an adequate reagent is added to react completely with a substance.
• Buffer solution – A solution that resists changes in pH even when a strong acid or base is added or when it is diluted with water

Types of Acid-Base Titration

The types and examples of strong/weak acids and bases are tabulated below.

Titration Curve & Equivalence Point

In a titration, the equivalence point is the point at which exactly the same number of moles of hydroxide ions have been added as there are moles of hydrogen ions . In a titration, if the base is added from the burette and the acid has been accurately measured into a flask. The shape of each titration curve is typical for the type of acid-base titration.

The pH does not change in a regular manner as the acid is added. Each curve has horizontal sections where a lot of bases can be added without changing the pH much. There is also a very steep portion of each curve except for weak acid and the weak base where a single drop of base changes the pH by several units. There is a large change of pH at the equivalence point even though this is not centred on pH 7. This is relevant to the choice of indicators for each type of titration.

Choice of Indicators

Acid-base indicators are substances which change colour or develop turbidity at a certain pH. They locate equivalence point and also measure pH. They are themselves acids or bases are soluble, stable and show strong colour changes. They are organic in nature.

A resonance of electron isomerism is responsible for colour change. Various indicators have different ionization constants and therefore they show a change in colour at different pH intervals.

Acid-base indicators can be broadly classified into three groups.

• The phthaleins and sulphophthaleins (eg; Phenolphthalein)
• Azo indicators (eg; Methyl orange)
• Triphenylmethane indicators (eg; Malachite green)

The two common indicators used in acid-base titration is Phenolphthalein and methyl orange. In the four types of acid-base titrations, the base is being added to the acid in each case. A graph is shown below where pH against the volume of base added is considered. The pH range over which the two indicators change colour. The indicator must change within the vertical portion of the pH curve.

pH Range Over which the Two Indicators Change Colour

The Choice of indicators based on the type of titration is tabulated below.

Recommended Video

What is an indicator what are acid-base indicators.

A 1.2gm sample of a mixture of (Na 2 CO 3 + NaHCO 3 ) is dissolved and titrated with 0.5N HCl. With phenolphthalein, the endpoint is at 15ml while after further addition of methyl orange a second endpoint is at 22ml. Calculate the percentage composition of the mixture.

15 + 15 = 30ml acid is necessary to neutralize Na 2 CO 3 completely.

Total volume needed = 15 + 22 = 37ml

(37-30) = 7ml acid is needed for neutralizing NaHCO 3

Therefore, Na 2 CO 3 composition (%) is

= [(30 x 0.5 x 0.053)/1.2] x 100 = 66.25%

Frequently Asked Questions – FAQs

Why is acid base titration important.

The purpose of a strong acid-strong base titration is to determine the acid solution concentration by titrating it with a basic solution of known concentration, or vice versa until there is neutralization. The reaction between a strong acid-base and a strong base will, therefore, result in water and salt.

What are the applications of acid-base titration?

An acid-base titration is used to determine the unknown acid or base concentration by neutralizing it with an acid or a known concentration basis. The unknown concentration can be calculated using the stoichiometry of the reaction.

Which indicator is used in acid-base titration?

Using a phenolphthalein indicator, a strong acid- strong base titration is performed. Phenolphthalein is selected because it changes colour between 8.3-10 in a pH range. In basic solutions, it will appear pink, and clear in acidic solutions.

What are two acid-base indicators?

Examples of acid-base indicators include red cabbage juice, litmus paper, phenolphthalein and. An acid-base indicator is a weak acid or weak base which dissociates in water to produce the weak acid and its conjugate base, or the weak base and its conjugate acid. The species and their conjugate are of different colours.

Define the Equivalence point.

The point at which just an adequate reagent is added to react completely with a substance.

Put your understanding of this concept to test by answering a few MCQs. Click ‘Start Quiz’ to begin!

Select the correct answer and click on the “Finish” button Check your score and answers at the end of the quiz

Visit BYJU’S for all Chemistry related queries and study materials

Your result is as below

Request OTP on Voice Call

Register with BYJU'S & Download Free PDFs

Register with byju's & watch live videos.

Chemistry Learner

It's all about chemistry.

• Chemical Bonds
• Chemical Reactions
• Materials Chemistry
• Organic Chemistry
• Periodic Trends
• Periodic Table Groups
• How to Read Periodic Table
• Naming Covalent Compounds Worksheets
• Net Ionic Equation Worksheets
• Types of Chemical Reactions Worksheets
• Word Equations Worksheets
• Valence Electrons Worksheets
• Graphing Periodic Trends Worksheets
• Periodic Trends Ionization Energy Worksheets
• Atomic Structure And Isotopes Worksheets

Acid-Base Titration

An acid-base titration is a fundamental process in analytical chemistry to determine the concentration of unknown acids or bases. It is based on the neutralization reaction , where an acid and a base react to form water and a salt. This method allows scientists and chemists to measure the amount of one substance precisely by reacting it with a known quantity of another. The process is crucial in understanding the strength and concentration of various solutions, providing valuable insights into their chemical properties and enabling accurate quantitative analysis in multiple industries, from pharmaceuticals to environmental sciences. [1-4]

When free H⁺ and OH⁻ ions are present in a water-based solution, they unite to form water molecules (H 2 O), a process known as neutralization. When an acid and a base mix, they create a solution in water that contains a dissolved salt composed of the base’s cation and the acid’s anion. For instance, the combination of hydrochloric acid (HCl) and sodium hydroxide (NaOH) results in an aqueous solution of sodium chloride (NaCl), as represented by this equation: [1,2]

HCl + NaOH → NaCl + H 2 O

Acid-base titration involves the controlled addition of a solution with a known concentration (titrant) to a solution of the substance being analyzed (analyte). The goal is to reach the equivalence point, where the moles of the acid and base are stoichiometrically equal, leading to complete neutralization. This neutralization is indicated by a sudden change in the solution’s pH. By carefully monitoring this process and noting the volume of titrant required to reach the equivalence point, scientists can calculate the concentration of the unknown solution.

The success of an acid-base titration hinges on carefully selecting indicators, which are substances that undergo distinct color changes at specific pH levels. Indicators play a crucial role in signaling the approach to the equivalence point, aiding in determining the endpoint of the titration. Indicators function by altering their molecular structure in response to variations in hydrogen ion concentration. It is accompanied by the transition from acidic to basic and vice versa. Common indicators, such as phenolphthalein or methyl orange, change color at different pH ranges, allowing chemists to visualize the transition and precisely pinpoint the endpoint. [1,4]

Acid-base titrations come in various types, each tailored to analyze different combinations of acids and bases. The image below shows the different acid-base titration curves. [1-3]

1. Strong Acid-Strong Base

Strong acid-strong base titrations involve neutralizing a strong acid like hydrochloric acid with a strong base like sodium hydroxide. The reaction between these compounds is rapid and complete, reaching the equivalence point with a pH of 7. This type of titration serves as a foundational model due to its clear endpoint and predictable outcomes, making it a common starting point for students studying titration techniques.

2. Weak Acid-Strong Base

Weak acid-strong base titrations involve a weak acid, such as acetic acid, reacting with a strong base like sodium hydroxide. Unlike strong acid-strong base titrations, these titrations require careful consideration of equilibrium reactions. The equilibrium between the weak acid and its conjugate base influences the pH at the equivalence point, resulting in a slightly basic endpoint. Understanding the dissociation constants and equilibrium reactions becomes crucial for accurately determining the unknown concentration.

3. Strong Acid-Weak Base

In this titration, a strong acid, like hydrochloric acid, undergoes neutralization with a weak base, like ammonia. The strong acid, known for its complete dissociation in water, donates protons readily. At the same time, the weak base accepts these protons at a slower rate due to its partial ionization. Consequently, the equivalence point in this titration does not result in a neutral pH of 7 but rather in an acidic pH due to the excess of the undissociated weak base in the solution.

4. Weak Acid-Weak Base

Weak acid-weak base titrations are less popular than the other three. This type involves the titration of a weak acid, like acetic acid, with a weak base, like ammonia. There is not a steep change in the titration curve. The absence of significant pH shifts during titration makes it challenging to perform such titration, limiting the amount of information that can be derived from such a curve.

• Chem.libretexts.org
• Khanacademy.org
• Opentextbc.ca
• Openstax.org

Trending Topics

© 2024 ( Chemistry Learner )

Your browser is not supported

Sorry but it looks as if your browser is out of date. To get the best experience using our site we recommend that you upgrade or switch browsers.

Find a solution

• Skip to main content
• Skip to navigation

• Back to parent navigation item
• Primary teacher
• Secondary/FE teacher
• Early career or student teacher
• Higher education
• Curriculum support
• Literacy in science teaching
• Periodic table
• Interactive periodic table
• Climate change and sustainability
• Resources shop
• Collections
• Post-lockdown teaching support
• Remote teaching support
• Starters for ten
• Screen experiments
• Assessment for learning
• Microscale chemistry
• Faces of chemistry
• Classic chemistry experiments
• Nuffield practical collection
• Anecdotes for chemistry teachers
• On this day in chemistry
• Global experiments
• PhET interactive simulations
• Chemistry vignettes
• Context and problem based learning
• Journal of the month
• Chemistry and art
• Art analysis
• Pigments and colours
• Ancient art: today's technology
• Psychology and art theory
• Art and archaeology
• Artists as chemists
• The physics of restoration and conservation
• Ancient Egyptian art
• Ancient Greek art
• Ancient Roman art
• Classic chemistry demonstrations
• In search of solutions
• In search of more solutions
• Creative problem-solving in chemistry
• Solar spark
• Chemistry for non-specialists
• Health and safety in higher education
• Analytical chemistry introductions
• Exhibition chemistry
• Introductory maths for higher education
• Commercial skills for chemists
• Kitchen chemistry
• Journals how to guides
• Chemistry in health
• Chemistry in sport
• Chemistry in your cupboard
• Chocolate chemistry
• Adnoddau addysgu cemeg Cymraeg
• The chemistry of fireworks
• Festive chemistry
• Education in Chemistry
• Teach Chemistry
• On-demand online
• Live online
• Selected PD articles
• PD for primary teachers
• PD for secondary teachers
• What we offer
• Chartered Science Teacher (CSciTeach)
• Teacher mentoring
• UK Chemistry Olympiad
• Who can enter?
• How does it work?
• Resources and past papers
• Top of the Bench
• Schools' Analyst
• Regional support
• Education coordinators
• RSC Yusuf Hamied Inspirational Science Programme
• RSC Education News
• Supporting teacher training
• Interest groups

• More navigation items

Titrating sodium hydroxide with hydrochloric acid

In association with Nuffield Foundation

In this experiment students neutralise sodium hydroxide with hydrochloric acid to produce the soluble salt sodium chloride in solution. They then concentrate the solution and allow it to crystallise to produce sodium chloride crystals

You have to decide if this experiment is suitable to use with different classes, and look at the need for preliminary training in using techniques involved in titration (see Teaching notes). What follows here assumes that teachers have judged the class to be capable of doing this experiment using a burette with reasonable expectation of success.

Assuming that the students have been given training, the practical work should, if possible, start with the apparatus ready at each work place in the laboratory. This is to avoid vulnerable and expensive glassware (the burette) being collected from an overcrowded central location.

Source: © Getty Images

Students doing a titration experiment in a school science laboratory.

Time required

Filling the burette, measuring out the alkali into the flask, and titrating it until it is neutralised takes about 20 minutes, with false starts being likely for many groups. In practice it does not matter if the end-point is overshot, even by several cubic centimetres, but the aim is to find the proportions for a roughly neutral solution.

Producing a neutral solution free of indicator, should take no more than 10 minutes.

Evaporating the solution may take the rest of the lesson to the point at which the solution can be left to crystallise for the next lesson. Watching solutions evaporate can be tedious for students, and they may need another task to keep them occupied – eg rinsing and draining the burettes with purified water.

• Eye protection
• Burette, 30 or 50 cm 3 (note 1)
• Conical flask, 100 cm 3
• Beaker, 100 cm 3
• Pipette, 20 or 25 cm 3 , with pipette filter
• Stirring rod
• Small (filter) funnel, about 4 cm diameter
• Burette stand and clamp (note 2)
• White tile (optional; note 3)
• Bunsen burner
• Pipeclay triangle (note 4)
• Evaporating basin, at least 50 cm 3 capacity
• Crystallising dish (note 5)
• Microscope or hand lens suitable for examining crystals in the crystallising dish

Apparatus notes

• If your school still uses burettes with glass stopcocks, consult the CLEAPSS Laboratory Handbook, section 10.10.1, for their care and maintenance. This experiment will not be successful if the burettes used have stiff, blocked or leaky stopcocks. Modern burettes with PTFE stopcocks are much easier to use, require no greasing, and do not get blocked. Burettes with pinchcocks of any type are not recommended; while cheap, they also are prone to leakage, especially in the hands of student beginners.
• Burette stands and clamps are designed to prevent crushing of the burette by over-tightening, which may happen if standard jaw clamps are used.
• The optional white tile is to go under the titration flask, but white paper can be used instead.
• Ceramic gauzes can be used instead of pipeclay triangles, but the evaporation then takes longer.
• The evaporation and crystallisation stages may be incomplete in the lesson time. The crystallisation dishes need to be set aside for crystallisation to take place slowly. However, the dishes should not be allowed to dry out completely, as this spoils the quality of the crystals. With occasional checks, it should be possible to decide when to decant surplus solution from each dish to leave good crystals for the students to inspect in the following.
• Sodium hydroxide solution, 0.4 M (IRRITANT), about 100 cm 3 in a labelled and stoppered bottle
• Dilute hydrochloric acid, 0.4 M, about 100 cm 3 in a labelled and stoppered bottle
• Methyl orange indicator solution (or alternative) in small dropper bottle

Health, safety and technical notes

• Read our standard health and safety guidance .
• Wear eye protection throughout.
• Sodium hydroxide solution, NaOH(aq), (IRRITANT at concentration used) – see CLEAPSS Hazcard  HC091a and CLEAPSS Recipe Book RB085. The concentration of the solution does not need to be made up to a high degree of accuracy, but should be reasonably close to the same concentration as the dilute hydrochloric acid, and less than 0.5 M.
• Dilute hydrochloric acid, HCl(aq) – see CLEAPSS Hazcard  HC047a and CLEAPSS Recipe Book RB043. The concentration of the solution does not need to be made up to a high degree of accuracy, but should be reasonably close to the same concentration as the sodium hydroxide solution, and less than 0.5 M.
• Methyl orange indicator solution (the solid is TOXIC but not the solution) – see CLEAPSS Hazcard  HC032  and CLEAPSS Recipe Book RB000. 

Source: Royal Society of Chemistry

Apparatus for titrating sodium hydroxide with hydrochloric acid to produce sodium chloride.

• Using a small funnel, pour a few cubic centimetres of 0.4 M hydrochloric acid into the burette, with the tap open and a beaker under the open tap. Once the tip of the burette is full of solution, close the tap and add more solution up to the zero mark. (Do not reuse the acid in the beaker – this should be rinsed down the sink.)
• Use a pipette with pipette filler to transfer 25 (or 20) cm 3  of 0.4 M sodium hydroxide solution to the conical flask, and add two drops of methyl orange indicator. Swirl gently to mix. Place the flask on a white tile or piece of clean white paper under the burette tap.
• Add the hydrochloric acid to the sodium hydroxide solution in small volumes, swirling gently after each addition. Continue until the solution just turns from yellow-orange to red and record the reading on the burette at this point. This coloured solution should now be rinsed down the sink.
• Refill the burette to the zero mark. Carefully add the same volume of fresh hydrochloric acid as you used in stage 1, step 3, to another 25 (or 20) cm 3  of sodium hydroxide solution, to produce a neutral solution, but this time without any indicator.
• Pour this solution into an evaporating basin. Reduce the volume of the solution to about half by heating on a pipeclay triangle or ceramic gauze over a low to medium Bunsen burner flame. The solution spits near the end and you get fewer crystals. Do not boil dry. You may need to evaporate the solution in, say, 20 cm 3  portions to avoid overfilling the evaporating basin. Do not attempt to lift the hot basin off the tripod – allow to cool first, and then pour into a crystallising dish.
• Leave the concentrated solution to evaporate further in the crystallising dish. This should produce a white crystalline solid in one or two days.
• Examine the crystals under a microscope.

Looking for an alternative method?

Check out our  practical video on preparing a salt  for a safer method for evaporating the solution, along with technician notes, instructions and a risk assessment activity for learners.

Teaching notes

Titration using a burette, to measure volumes of solution accurately, requires careful and organised methods of working, manipulative skills allied to mental concentration, and attention to detail. All of these are of course desirable traits to be developed in students, but there has to be some degree of basic competence and reliability before using a burette with a class. The experiment is most likely to be suited to 14–16 year old students. This is discussed further below, but what follows here assumes that you have judged the class to be capable of doing this experiment using a burette with reasonable expectation of success.

Students need training in using burettes correctly, including how to clamp them securely and fill them safely. You should consider demonstrating burette technique, and give students the opportunity to practise this. In this experiment a pipette is not necessary, as the aim is to neutralise whatever volume of alkali is used, and that can be measured roughly using a measuring cylinder.

It is not the intention here to do quantitative measurements leading to calculations. The aim is to introduce students to the titration technique only to produce a neutral solution.

Alternative indicators you can use include screened methyl orange (green in alkali, violet in acid) and phenolphthalein (pink in alkali, colourless in acid).

Leaving the concentrated solutions to crystallise slowly should help to produce larger crystals. The solubility of sodium chloride does not change much with temperature, so simply cooling the solution is unlikely to form crystals.

Under the microscope (if possible, a stereomicroscope is best) you can see the cubic nature of the crystals. If crystallisation has occurred in shallow solution, with the crystals only partly submerged, ‘hopper-shaped’ crystals may be seen. In these crystals, each cube face becomes a hollow, stepped pyramid shape.

Student questions

What substances have been formed in this reaction? Write a word equation and a symbol equation.

Why must you use another 25 cm 3  of sodium hydroxide solution, rather than making your crystals from the solution in stage 1?

What shape are the crystals?

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry.

Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

The experiment is also part of the Royal Society of Chemistry’s Continuing Professional Development course:  Chemistry for non-specialists .

© Nuffield Foundation and the Royal Society of Chemistry

• 14-16 years
• 16-18 years
• Practical experiments
• Practical skills and safety
• Acids and bases

Specification

• 1.8.18 demonstrate knowledge and understanding of how pure dry samples of soluble salts can be prepared by: adding excess insoluble substances to acid; adding alkali to acid, or vice versa, in the presence of an indicator; and repeating without indicator…
• 8. Investigate reactions between acids and bases; use indicators and the pH scale
• Mandatory eexperiment 4.2A - A hydrochloric acid/sodium hydroxide titration, and the use of this titration in making the sodium salt.
• 3. Find the concentration of a solution of hydrochloric acid
• 2a Determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration.
• The volumes of acid and alkali solutions that react with each other can be measured by titration using a suitable indicator.
• Students should be able to: describe how to carry out titrations using strong acids and strong alkalis only (sulfuric, hydrochloric and nitric acids only) to find the reacting volumes accurately
• Salt solutions can be crystallised to produce solid salts.
• Students should be able to describe how to make pure, dry samples of named soluble salts from information provided.
• 5.9C Carry out an accurate acid-alkali titration, using burette, pipette and a suitable indicator
• 3.18 Describe how to carry out an acid-alkali titration, using burette, pipette and a suitable indicator, to prepare a pure, dry salt
• C5.4.7 describe and explain the procedure for a titration to give precise, accurate, valid and repeatable results
• 6 Titration of a strong acid and strong alkali to find the concentration of the acid using an appropriate pH indicator
• 7 Production of pure dry sample of an insoluble and soluble salt
• C5.1b describe the technique of titration
• PAG 6 Titration of a strong acid and strong alkali to find the concentration of the acid using an appropriate pH indicator
• C5.3.6 describe and explain the procedure for a titration to give precise, accurate, valid and repeatable results
• C4 Production of pure dry sample of an insoluble and soluble salt
• In an acid-base titration, the concentration of the acid or base is determined by accurately measuring the volumes used in the neutralisation reaction. An indicator can be added to show the end-point of the reaction
• Titration is used to determine, accurately, the volumes of solution required to reach the end-point of a chemical reaction.
• (j) titration as a method to prepare solutions of soluble salts and to determine relative and actual concentrations of solutions of acids/alkalis
• (f) acid-base titrations

Related articles

Titration screen experiment

Four out of five

Give students the opportunity to conduct their own titration experiment on a computer or tablet. This resource also includes a redox titration experiment.

Chromatography challenge | 16–18 years

By Andy Markwick

Explore analytical techniques and their applications with a chromatography investigation and research activity

Demonstrate concentration and density with a transition metal colloid cell

2024-02-19T10:06:00Z By Declan Fleming

Boost 11–14 learners’ understanding of diffusion and transition metal chemistry

6 readers' comments

Only registered users can comment on this article., more experiments.

‘Gold’ coins on a microscale | 14–16 years

By Dorothy Warren and Sandrine Bouchelkia

Practical experiment where learners produce ‘gold’ coins by electroplating a copper coin with zinc, includes follow-up worksheet

Practical potions microscale | 11–14 years

By Kirsty Patterson

Observe chemical changes in this microscale experiment with a spooky twist.

Antibacterial properties of the halogens | 14–18 years

By Kristy Turner

Use this practical to investigate how solutions of the halogens inhibit the growth of bacteria and which is most effective

• Contributors
• Email alerts

Site powered by Webvision Cloud

Personalised lessons and regular feedback to ensure you ace your exams! Book a free consultation today

100+ Video Tutorials, Flashcards and Weekly Seminars

Gain hands-on experience of how physics is used in different fields. Experience life as a uni student and boost your university application with our summer programme!

• Revision notes >
• A-level Chemistry Revision Notes >
• AQA A-Level Chemistry Revision Notes

Acids and Bases - Titrations (A-Level Chemistry)

Titration of acids and bases, what is titration.

A titration is an experiment through which the concentration of an unknown acid or alkali can be calculated.

When an acid and alkali react together, a neutralisation reaction occurs. If we have an unknown concentration of acid , we can do a titration to see how much acid is needed to neutralise a standard solution of alkali (of known concentration), and vice-versa.

Steps of Titration

The steps in the titration are as follows:

1. Fill up a burette with acid. A burette is used to place the acid of unknown concentration. A burette is preferred to a measuring cylinder, as it is more precise .

2. Use a pipette to add alkali to a conical flask. A volumetric pipette is used to deliver a fixed volume of standard solution of alkali into a conical flask. This is also more precise than a measuring cylinder.

3. Add an indicator to the alkali. An indicator is added to the alkali in the conical flask, which causes a colour change to occur.

4. Slowly add the acid into the alkali. The acid from the burette is allowed to run into the alkali, swirling constantly to ensure the reactants mix well.

5. Slow the adding of acid when the indicator changes colour. When the colour from the indicator changes, the flow of the acid is slowed so that it is added drop by drop .

6. Stop adding acid when there is a permanent colour change. When the colour of the indicator produces a permanent colour change. The flow of acid is stopped.

7. Record the titre. The volume of acid delivered, known as the titre , is recorded.

8. Repeat the titration. The titration is repeated until concordant (within 0.1 cm³) titres are obtained.

Calculating Concentration of Acid

The known volume and concentration of the alkali enable the moles to be calculated.

The balanced symbol equation enable the moles ratio of acid:alkali to be determined. Hence the number of moles of acid can be calculated.

The mean titre is the volume of the acid. Using the moles and the volume, the concentration of the unknown acid can be determined.

Titration is a process of measuring the concentration of an unknown solution by adding a known amount of another solution of known concentration until the reaction between the two is complete. Titration can be used to determine the concentration of an acid or a base in a solution. In acid-base titration, an acid or a base of unknown concentration is reacted with a base or an acid of known concentration, respectively. The reaction between the two solutions is typically monitored using a pH indicator, which changes color as the pH of the solution changes. The endpoint of the titration is reached when the reaction is complete, and the pH indicator changes color. At this point, the amount of the known solution that has been added is used to calculate the concentration of the unknown solution. The equation for an acid-base titration is typically: acid + base → salt + water

The purpose of a Titration in A-Level Chemistry is to determine the concentration of an unknown substance. This helps in determining the strength of an acid or a base and understanding how much of it is required to react with another substance.

An Acid in Chemistry is a substance that releases hydrogen ions (H+) in water. These ions are responsible for making the solution acidic.

A Base in Chemistry is a substance that releases hydroxide ions (OH-) in water. These ions are responsible for making the solution basic or alkaline.

The difference between a Strong Acid and a Weak Acid is their concentration of hydrogen ions (H+) in water. Strong Acids have a high concentration of H+ ions and Weak Acids have a low concentration.

The difference between a Strong Base and a Weak Base is their concentration of hydroxide ions (OH-) in water. Strong Bases have a high concentration of OH- ions and Weak Bases have a low concentration.

The steps involved in a Titration Experiment are: Prepare a solution of the known substance (base) Add a few drops of an indicator to the unknown substance (acid) Slowly add the known substance (base) to the unknown substance (acid) until a color change occurs Measure the volume of the known substance (base) used in the reaction Calculate the concentration of the unknown substance (acid)

An Indicator is used in a Titration Experiment to determine when the reaction between the known substance (base) and the unknown substance (acid) is complete. The color change of the indicator indicates the end point of the reaction.

It is important to have an accurate measurement in a Titration Experiment because it directly affects the accuracy of the concentration of the unknown substance (acid). Any error in measurement can result in incorrect results.

Still got a question? Leave a comment

Leave a comment, cancel reply.

Save my name, email, and website in this browser for the next time I comment.

AQA 3.1.1 Atomic structure

Ionisation energies (a-level chemistry), atomic structure – electron arrangement (a-level chemistry), atomic structure – electrons in atoms (a-level chemistry), atomic structure – mass spectrometry (a-level chemistry), atomic structure – element isotopes (a-level chemistry), atomic structure – atomic and mass number (a-level chemistry), atomic structure – subatomic particles (a-level chemistry), aqa 3.1.10 equilibrium constant kp for homogeneous systems, equilibrium constant for homogenous systems – le chatelier’s principle in gas equilibria (a-level chemistry), equilibrium constant for homogenous systems – gas equilibria and kp (a-level chemistry), equilibrium constant for homogeneous system – changing kp (a-level chemistry), equilibrium constant for homogenous systems – gas partial pressures (a-level chemistry), aqa 3.1.11 electrode potentials and electrochemical cells, acids and bases – drawing ph curves (a-level chemistry), acids and bases – acid-base indicators (a-level chemistry), acids and bases – dilutions and ph (a-level chemistry), electrode potentials and electrochemical cells – commercial applications of fuel cells (a-level chemistry), electrode potentials and electrochemical cells – electrochemical cells reactions (a-level chemistry), electrode potentials and electrochemical cells – representing electrochemical cells (a-level chemistry), electrode potentials and electrochemical cells – electrode potentials (a-level chemistry), electrode potentials and electrochemical cells – half cells and full cells (a-level chemistry), aqa 3.1.12 acids and bases, acids and bases – titrations (a-level chemistry), acids and bases – buffer action (a-level chemistry), acids and bases – ph of strong bases (a-level chemistry), acids and bases – ionic product of water (a-level chemistry), acids and bases – more ka calculations (a-level chemistry), acids and bases – the acid dissociation constant, ka (a-level chemistry), acids and bases – the ph scale and strong acids (a-level chemistry), acids and bases – neutralisation reactions (a-level chemistry), acids and bases – acid and base strength (a-level chemistry), acids and bases – the brønsted-lowry acid-base theory (a-level chemistry), aqa 3.1.2 amount of substance, amount of substance – percentage atom economy (a-level chemistry), amount of substance – calculating percentage yields (a-level chemistry), amount of substance – stoichiometric calculations (a-level chemistry), amount of substance – balancing chemical equations (a-level chemistry), amount of substance – empirical and molecular formulae (a-level chemistry), amount of substance – further mole calculations (a-level chemistry), amount of substance- the mole and the avogadro constant (a-level chemistry), amount of substance – measuring relative masses (a-level chemistry), amount of substance – the ideal gas equation (a-level chemistry), aqa 3.1.3 bonding, periodicity – classification (a-level chemistry), bonding – hydrogen bonding in water (a-level chemistry), bonding – forces between molecules (a-level chemistry), bonding – bond polarity (a-level chemistry), bonding – molecular shapes (a-level chemistry), bonding – predicting structures (a-level chemistry), bonding – carbon allotropes (a-level chemistry), bonding – properties of metallic bonding (a-level chemistry), bonding – properties of covalent structures (a-level chemistry), bonding – covalent bonds (a-level chemistry), aqa 3.1.4 energetics, aqa 3.1.5 kinetics, kinetics – the maxwell–boltzmann distribution and catalysts (a-level chemistry), kinetics – the collision theory and reaction rates (a-level chemistry), aqa 3.1.6 chemical equilibria, calculations with equilibrium constants (a-level chemistry), chemical equilibria applied to industry (a-level chemistry), chemical equilibria and le chatelier’s principle (a-level chemistry), aqa 3.1.7 oxidation, reduction and redox, oxidation, reduction and redox equations – balancing redox equations (a-level chemistry), oxidation, reduction and redox equations – redox processes (a-level chemistry), oxidation, reduction and redox equations – oxidation states (a-level chemistry), aqa 3.1.8 thermodynamics, thermodynamic – calculations involving free energy (a-level chemistry), thermodynamic – gibbs free energy (a-level chemistry), thermodynamic – entropy change predictions (a-level chemistry), thermodynamic – total entropy changes (a-level chemistry), thermodynamic – introduction to entropy (a-level chemistry), thermodynamic – calculating enthalpy changes of solution (a-level chemistry), thermodynamic – enthalpy of solution (a-level chemistry), thermodynamic – enthalpy of hydration (a-level chemistry), thermodynamic – calculations involving born-haber cycles (a-level chemistry), thermodynamic – construction of born-haber cycles (a-level chemistry), aqa 3.1.9 rate equations, rate equations – reaction determining steps (a-level chemistry), rate equations – reaction half lives (a-level chemistry), rate equations – uses of clock reactions (a-level chemistry), rate equations – determining orders of reactions graphically (a-level chemistry), rate equations – determining order of reaction experimentally (a-level chemistry), rate equations – temperature changes and the rate constant (a-level chemistry), rate equations – the rate constant (a-level chemistry), rate equations – introduction to orders of reactions (a-level chemistry), rate equations – the rate equation (a-level chemistry), rate equations – measuring rate of reaction (a-level chemistry), aqa 3.2.1 periodicity, periodicity – trends along period 3 (a-level chemistry), aqa 3.2.2 group 2, the alkaline earth metals, uses of group 2 elements and their compounds (a-level chemistry), reactions of group 2 elements (a-level chemistry), group 2, the alkaline earth metals (a-level chemistry), aqa 3.2.3 group 7(17), the halogens, the halogens -halide ions and their reactions (a-level chemistry), the halogens – disproportionation reactions in halogens (a-level chemistry), the halogens – reactions with halogens (a-level chemistry), the halogens – group 7, the halogens (a-level chemistry), aqa 3.2.4 properties of period 3 elements, properties of period 3 elements – properties of period 3 compounds (a-level chemistry), properties of period 3 elements – reactivity of period 3 elements (a-level chemistry), aqa 3.2.5 transition metals, transition metals – autocatalysis of transition metals (a-level chemistry), transition metals – transition metals as homogeneous catalysts (a-level chemistry), transition metals – transition metals as heterogeneous catalysts (a-level chemistry), transition metals – examples of redox reactions in transition metals (a-level chemistry), transition metals – iodine-sodium thiosulfate titrations (a-level chemistry), transition metals – carrying titrations with potassium permanganate (a-level chemistry), transition metals – redox titrations (a-level chemistry), transition metals – redox potentials (a-level chemistry), transition metals – redox reactions revisited (a-level chemistry), transition metals – ligand substitution reactions (a-level chemistry), aqa 3.2.6 reactions of ions in aqueous solution, reactions of ions in aqueous solutions – metal ions in solution (a-level chemistry), aqa 3.3.1 introduction to organic chemistry, introduction to organic chemistry – structural isomers (a-level chemistry), introduction to organic chemistry – e/z isomerism (a-level chemistry), introduction to organic chemistry – reaction mechanisms in organic chemistry (a-level chemistry), introduction to organic chemistry – general formulae (a-level chemistry), introduction to organic chemistry – introduction to functional groups (a-level chemistry), introduction to organic chemistry – naming and representing organic compounds (a-level chemistry), aqa 3.3.10 aromatic chemistry, aromatic chemistry – friedel-crafts acylation and alkylation (a-level chemistry), aromatic chemistry – halogenation reactions in benzene (a-level chemistry), aromatic chemistry – electrophilic substitution reactions in benzene (a-level chemistry), aromatic chemistry – improved benzene model (a-level chemistry), aromatic chemistry – introduction to benzene (a-level chemistry), aqa 3.3.11 amines, amines – nitriles (a-level chemistry), amines – properties and reactivity of amines (a-level chemistry), amines – amine synthesis (a-level chemistry), amines – introduction to amines (a-level chemistry), aqa 3.3.12 polymers, polymer disposal (a-level chemistry), polymer biodegradability (a-level chemistry), condensation polymers (a-level chemistry), polyamide formation (a-level chemistry), aqa 3.3.13 amino acids, amino acids, proteins and dna – dna replication (a-level chemistry), amino acids, proteins and dna – dna (a-level chemistry), amino acids, proteins and dna – enzyme action (a-level chemistry), amino acids, proteins and dna – structure of proteins (a-level chemistry), amino acids, proteins and dna – structure of amino acids (a-level chemistry), aqa 3.3.14 organic synthesis, organic synthesis – considerations in organic synthesis (a-level chemistry), organic synthesis – organic synthesis: aromatic compounds (a-level chemistry), organic synthesis – organic synthesis: aliphatic compounds (a-level chemistry), aqa 3.3.15 nmr, analytical techniques – high resolution ¹h nmr (a-level chemistry), analytical techniques – types of nmr: hydrogen (a-level chemistry), analytical techniques – types of nmr: carbon 13 (a-level chemistry), analytical techniques – nmr samples and standards (a-level chemistry), analytical techniques – nuclear magnetic resonance spectroscopy (a-level chemistry), aqa 3.3.16 chromatography, analytical techniques – different types of chromatography (a-level chemistry), analytical techniques – chromatography (a-level chemistry), aqa 3.3.2 alkanes, alkanes – obtaining alkanes (a-level chemistry), alkanes – alkanes: properties and reactivity (a-level chemistry), aqa 3.3.3 halogenoalkanes, halogenoalkanes – environmental impact of halogenalkanes (a-level chemistry), halogenoalkanes – reactivity of halogenoalkanes (a-level chemistry), halogenoalkanes – introduction to halogenoalkanes (a-level chemistry), aqa 3.3.4 alkenes, alkenes – addition polymerisation in alkenes (a-level chemistry), alkenes – alkene structure and reactivity (a-level chemistry), aqa 3.3.5 alcohols, alcohols – industrial production of alcohols (a-level chemistry), alcohols – alcohol reactivity (a-level chemistry), alcohols – alcohol oxidation (a-level chemistry), alcohols – introduction to alcohols (a-level chemistry), aqa 3.3.6 organic analysis, organic analysis – infrared (ir) spectroscopy (a-level chemistry), organic analysis – identification of functional groups (a-level chemistry), aqa 3.3.7 optical isomerism, optical isomerism (a-level chemistry), aqa 3.3.8 aldehydes and ketones, aldehydes and ketones – reactions to increase carbon chain length (a-level chemistry), aldehydes and ketones – testing for carbonyl compounds (a-level chemistry), aldehydes and ketones – reactivity of carbonyl compunds (a-level chemistry), aldehydes and ketones – carbonyl compounds (a-level chemistry), aqa 3.3.9 carboxylic acids, carboxylic acids and derivatives – structure of amides (a-level chemistry), carboxylic acids and derivatives – acyl groups (a-level chemistry), carboxylic acids and derivatives – properties and reactivity of esters (a-level chemistry), carboxylic acids and derivatives – properties and reactivity of carboxylic acids (a-level chemistry), 21: organic synthesis, 29: intro to organic chemistry, aromatic chemistry – benzene nomenclature (a-level chemistry), cie 1: atomic structure, bonding – ion formation (a-level chemistry), cie 10: group 2, cie 11: group 17, cie 13: intro to as organic chemistry, cie 14: hydrocarbons, cie 15: halogen compounds, cie 16: hydroxy compounds, cie 17: carbonyl compounds, cie 18: carboxylic acids and derivatives, cie 19: nitrogen compounds, cie 2: atoms, molecules and stoichiometry, cie 20: polymerisation, cie 22: analytical techniques, cie 23: chemical energetics, cie 24: electrochemistry, cie 25: equilibria, cie 27: group 2 elements, cie 28: chemistry of transition elements, transition metals – colour in transition metal ions (a-level chemistry), transition metals – optical isomerism in complex ions (a-level chemistry), transition metals – cis-trans isomerism in complex ions (a-level chemistry), transition metals – complex ion shape (a-level chemistry), transition metals – ligands (a-level chemistry), transition metals – introduction to complex ions (a-level chemistry), cie 3: chemical bonding, bonding – properties of ionic bonding (a-level chemistry), cie 30: hydrocarbons, aromatic chemistry – reactivity of substituted benzene (a-level chemistry), cie 31: halogen compounds, cie 32: hydroxy compounds, cie 33: carboxylic acids and derivatives, cie 34: nitrogen compounds, cie 35: polymerisation, cie 36: organic synthesis, cie 37: analytic techniques, analytical techniques – deuterium use in ¹h nmr (a-level chemistry), cie 4: states of matter, cie 6: electrochemistry, cie 7: equilibria, cie 8: reaction kinetics, cie 9: the periodic table, cie: 26: reaction kinetics, catalysts, edexcel topic 1: atomic structure and the periodic table, edexcel topic 10: equlibrium 1, edexcel topic 11: equilibrium 2, edexcel topic 12: acid-base equilibria, edexcel topic 13: energetics 2, edexcel topic 14: redox 2, edexcel topic 15: transition metals, edexcel topic 16: kinetics 2, edexcel topic 17: organic chemistry 2, edexcel topic 18: organic chemistry 3, organic synthesis – practical purification techniques (a-level chemistry), organic synthesis – practical preparation techniques (a-level chemistry), edexcel topic 19: modern analytical techniques 2, edexcel topic 2a & b: bonding and structure, edexcel topic 3: redox 1, edexcel topic 4: inorganic chemistry & the periodic table, the halogens – testing for ions (a-level chemistry), edexcel topic 5: formulae, equations and amounts of substance, edexcel topic 6: organic chemistry 1, edexcel topic 7: modern analytical techniques 1, edexcel topic 8, edexcel topic 9: kinetics 1, ocr 2.1.1 atomic structure and isotopes, ocr 2.1.2 compounds, formulae and equations, ocr 2.1.3 amount of substance, ocr 2.1.4 acids, ocr 2.1.5 redox, ocr 2.2.1 electron structure, ocr 2.2.2 bonding and structure, ocr 3.1.1 periodicity, ocr 3.1.2 group 2, ocr 3.1.3 the halogens, ocr 3.1.4 qualitative analysis, ocr 3.2.2 reaction rates, ocr 3.2.3 chemical equilibrium, ocr 4.1.1 basic concepts of organic chemistry, ocr 4.1.2 alkanes, ocr 4.1.3 alkenes, ocr 4.2.1 alcohols, ocr 4.2.2 haloalkanes, ocr 4.2.3 organic synthesis, ocr 4.2.4 analytical techniques, ocr 5.1.1 rates, equilibrium and ph, ocr 5.1.2 how fast, ocr 5.1.3 acids, bases and buffers, ocr 5.2.1 lattice enthalpy, thermodynamic – enthalpy key terms (a-level chemistry), thermodynamic – lattice enthalpies (a-level chemistry), ocr 5.2.2 enthalpy and entropy, ocr 5.2.3 redox and electrode potentials, ocr 5.3.1 transition elements, precipitation reactions of metal ions in solution (a-level chemistry), ocr 5.3.2 qualitative analysis, ocr 6.1.1 aromatic compounds, ocr 6.1.2 carbonyl compounds, ocr 6.1.3 carboxylic acids and esters, ocr 6.2.1 amines, ocr 6.2.2 amino acids, amides and chirality, ocr 6.2.3 polyesters and polyamides, ocr 6.2.4 carbon–carbon bond formation, ocr 6.2.5 organic synthesis, ocr 6.3.1 chromatography and qualitative analysis, ocr 6.3.2 spectroscopy, related links.

• A-level Chemistry Past Papers

Boost your A-Level Chemistry Performance

Get a 9 in A-Level Chemistry with our Trusted 1-1 Tutors. Enquire now.

100+ Video Tutorials, Flashcards and Weekly Seminars. 100% Money Back Guarantee

Gain hands-on experience of how physics is used in different fields. Boost your university application with our summer programme!

Learn live with other students and gain expert tips and advice to boost your score.

Let's get acquainted ? What is your name?

Nice to meet you, {{name}} what is your preferred e-mail address, nice to meet you, {{name}} what is your preferred phone number, what is your preferred phone number, just to check, what are you interested in, when should we call you.

It would be great to have a 15m chat to discuss a personalised plan and answer any questions

What time works best for you? (UK Time)

Pick a time-slot that works best for you ?

How many hours of 1-1 tutoring are you looking for?

My whatsapp number is..., for our safeguarding policy, please confirm....

Please provide the mobile number of a guardian/parent

Which online course are you interested in?

What is your query, you can apply for a bursary by clicking this link, sure, what is your query, thank you for your response. we will aim to get back to you within 12-24 hours., lock in a 2 hour 1-1 tutoring lesson now.

If you're ready and keen to get started click the button below to book your first 2 hour 1-1 tutoring lesson with us. Connect with a tutor from a university of your choice in minutes. (Use FAST5 to get 5% Off!)

Want to create or adapt books like this? Learn more about how Pressbooks supports open publishing practices.

89 Acid-Base Titrations

Learning objectives.

By the end of this section, you will be able to:

• Interpret titration curves for strong and weak acid-base systems
• Compute sample pH at important stages of a titration
• Explain the function of acid-base indicators

As seen in the chapter on the stoichiometry of chemical reactions, titrations can be used to quantitatively analyze solutions for their acid or base concentrations. In this section, we will explore the underlying chemical equilibria that make acid-base titrimetry a useful analytical technique.

Titration Curves

A titration curve is a plot of some solution property versus the amount of added titrant. For acid-base titrations, solution pH is a useful property to monitor because it varies predictably with the solution composition and, therefore, may be used to monitor the titration’s progress and detect its end point. The following example exercise demonstrates the computation of pH for a titration solution after additions of several specified titrant volumes. The first example involves a strong acid titration that requires only stoichiometric calculations to derive the solution pH. The second example addresses a weak acid titration requiring equilibrium calculations.

Calculating pH for Titration Solutions: Strong Acid/Strong Base

A titration is carried out for 25.00 mL of 0.100 M HCl (strong acid) with 0.100 M of a strong base NaOH (the titration curve is shown in (Figure) ). Calculate the pH at these volumes of added base solution:

(a) 0.00 mL

(b) 12.50 mL

(c) 25.00 mL

(d) 37.50 mL

(a) Titrant volume = 0 mL. The solution pH is due to the acid ionization of HCl. Because this is a strong acid, the ionization is complete and the hydronium ion molarity is 0.100 M . The pH of the solution is then

(b) Titrant volume = 12.50 mL. Since the acid sample and the base titrant are both monoprotic and equally concentrated, this titrant addition involves less than a stoichiometric amount of base, and so it is completely consumed by reaction with the excess acid in the sample. The concentration of acid remaining is computed by subtracting the consumed amount from the intial amount and then dividing by the solution volume:

(c) Titrant volume = 25.00 mL. This titrant addition involves a stoichiometric amount of base (the equivalence point ), and so only products of the neutralization reaction are in solution (water and NaCl). Neither the cation nor the anion of this salt undergo acid-base ionization; the only process generating hydronium ions is the autoprotolysis of water. The solution is neutral, having a pH = 7.00.

(d) Titrant volume = 37.50 mL. This involves the addition of titrant in excess of the equivalence point. The solution pH is then calculated using the concentration of hydroxide ion:

pH = 14 − pOH = 14 + log([OH − ]) = 14 + log(0.0200) = 12.30

Check Your Learning

Calculate the pH for the strong acid/strong base titration between 50.0 mL of 0.100 M HNO 3 ( aq ) and 0.200 M NaOH (titrant) at the listed volumes of added base: 0.00 mL, 15.0 mL, 25.0 mL, and 40.0 mL.

Titration of a Weak Acid with a Strong Base

Consider the titration of 25.00 mL of 0.100 M CH 3 CO 2 H with 0.100 M NaOH. The reaction can be represented as:

Calculate the pH of the titration solution after the addition of the following volumes of NaOH titrant:

(b) 25.00 mL

(c) 12.50 mL

(a) The initial pH is computed for the acetic acid solution in the usual ICE approach:

(b) The acid and titrant are both monoprotic and the sample and titrant solutions are equally concentrated; thus, this volume of titrant represents the equivalence point. Unlike the strong-acid example above, however, the reaction mixture in this case contains a weak conjugate base (acetate ion). The solution pH is computed considering the base ionization of acetate, which is present at a concentration of

Base ionization of acetate is represented by the equation

Note that the pH at the equivalence point of this titration is significantly greater than 7, as expected when titrating a weak acid with a strong base.

(c) Titrant volume = 12.50 mL. This volume represents one-half of the stoichiometric amount of titrant, and so one-half of the acetic acid has been neutralized to yield an equivalent amount of acetate ion. The concentrations of these conjugate acid-base partners, therefore, are equal. A convenient approach to computing the pH is use of the Henderson-Hasselbalch equation:

(pH = p K a at the half-equivalence point in a titration of a weak acid)

(d) Titrant volume = 37.50 mL. This volume represents a stoichiometric excess of titrant, and a reaction solution containing both the titration product, acetate ion, and the excess strong titrant. In such solutions, the solution pH is determined primarily by the amount of excess strong base:

Calculate the pH for the weak acid/strong base titration between 50.0 mL of 0.100 M HCOOH( aq ) (formic acid) and 0.200 M NaOH (titrant) at the listed volumes of added base: 0.00 mL, 15.0 mL, 25.0 mL, and 30.0 mL.

Performing additional calculations similar to those in the preceding example permits a more full assessment of titration curves. A summary of pH/volume data pairs for the strong and weak acid titrations is provided in (Figure) and plotted as titration curves in (Figure) . A comparison of these two curves illustrates several important concepts that are best addressed by identifying the four stages of a titration:

initial state (added titrant volume = 0 mL): pH is determined by the acid being titrated; because the two acid samples are equally concentrated, the weak acid will exhibit a greater initial pH

pre-equivalence point (0 mL < V < 25 mL): solution pH increases gradually and the acid is consumed by reaction with added titrant; composition includes unreacted acid and the reaction product, its conjugate base

equivalence point ( V = 25 mL): a drastic rise in pH is observed as the solution composition transitions from acidic to either neutral (for the strong acid sample) or basic (for the weak acid sample), with pH determined by ionization of the conjugate base of the acid

postequivalence point ( V > 25 mL): pH is determined by the amount of excess strong base titrant added; since both samples are titrated with the same titrant, both titration curves appear similar at this stage.

Acid-Base Indicators

The equilibrium in a solution of the acid-base indicator methyl orange, a weak acid, can be represented by an equation in which we use HIn as a simple representation for the complex methyl orange molecule:

The anion of methyl orange, In − , is yellow, and the nonionized form, HIn, is red. When we add acid to a solution of methyl orange, the increased hydronium ion concentration shifts the equilibrium toward the nonionized red form, in accordance with Le Châtelier’s principle. If we add base, we shift the equilibrium towards the yellow form. This behavior is completely analogous to the action of buffers.

The perceived color of an indicator solution is determined by the ratio of the concentrations of the two species In − and HIn. If most of the indicator (typically about 60−90% or more) is present as In − , the perceived color of the solution is yellow. If most is present as HIn, then the solution color appears red. The Henderson-Hasselbalch equation is useful for understanding the relationship between the pH of an indicator solution and its composition (thus, perceived color):

In solutions where pH > p K a , the logarithmic term must be positive, indicating an excess of the conjugate base form of the indicator (yellow solution). When pH > p K a , the log term must be negative, indicating an excess of the conjugate acid (red solution). When the solution pH is close to the indicator pKa, appreciable amounts of both conjugate partners are present, and the solution color is that of an additive combination of each (yellow and red, yielding orange). The color change interval (or pH interval ) for an acid-base indicator is defined as the range of pH values over which a change in color is observed, and for most indicators this range is approximately p K a ± 1.

There are many different acid-base indicators that cover a wide range of pH values and can be used to determine the approximate pH of an unknown solution by a process of elimination. Universal indicators and pH paper contain a mixture of indicators and exhibit different colors at different pHs. (Figure) presents several indicators, their colors, and their color-change intervals.

The titration curves shown in (Figure) illustrate the choice of a suitable indicator for specific titrations. In the strong acid titration, use of any of the three indicators should yield reasonably sharp color changes and accurate end point determinations. For this titration, the solution pH reaches the lower limit of the methyl orange color change interval after addition of ~24 mL of titrant, at which point the initially red solution would begin to appear orange. When 25 mL of titrant has been added (the equivalence point), the pH is well above the upper limit and the solution will appear yellow. The titration’s end point may then be estimated as the volume of titrant that yields a distinct orange-to-yellow color change. This color change would be challenging for most human eyes to precisely discern. More-accurate estimates of the titration end point are possible using either litmus or phenolphthalein, both of which exhibit color change intervals that are encompassed by the steep rise in pH that occurs around the 25.00 mL equivalence point.

The weak acid titration curve in (Figure) shows that only one of the three indicators is suitable for end point detection. If methyl orange is used in this titration, the solution will undergo a gradual red-to-orange-to-yellow color change over a relatively large volume interval (0–6 mL), completing the color change well before the equivalence point (25 mL) has been reached. Use of litmus would show a color change that begins after adding 7–8 mL of titrant and ends just before the equivalence point. Phenolphthalein, on the other hand, exhibits a color change interval that nicely brackets the abrupt change in pH occurring at the titration’s equivalence point. A sharp color change from colorless to pink will be observed within a very small volume interval around the equivalence point.

Key Concepts and Summary

The titration curve for an acid-base titration is typically a plot of pH versus volume of added titrant. These curves are useful in selecting appropriate acid-base indicators that will permit accurate determinations of titration end points.

Chemistry End of Chapter Exercises

Instructions

Complete the following exercises by providing answers to each question. There are a total of 4 questions in this section. Please be aware that certain questions may contain subsections.

1. Explain how to choose the appropriate acid-base indicator for the titration of a weak base with a strong acid.

Answer(s): At the equivalence point in the titration of a weak base with a strong acid, the resulting solution is slightly acidic due to the presence of the conjugate acid. Thus, pick an indicator that changes color in the acidic range and brackets the pH at the equivalence point. Methyl orange is a good example.

2. Explain why an acid-base indicator changes color over a range of pH values rather than at a specific pH.

(a) no KOH added

(b) 20 mL of KOH solution added

(c) 39 mL of KOH solution added

(d) 40 mL of KOH solution added

(e) 41 mL of KOH solution added

Answer(s): (a) pH = 2.50; (b) pH = 4.01; (c) pH = 5.60; (d) pH = 8.35; (e) pH = 11.08

• 1 Titration of 25.00 mL of 0.100 M HCl (0.00250 mol of HCI) with 0.100 M NaOH.
• 2 Titration of 25.00 mL of 0.100 M CH 3 CO 2 H (0.00250 mol of CH 3 CO 2 H) with 0.100 M NaOH.

Chemistry 2e 2019 Copyright © 2018 by OSCRiceUniversity is licensed under a Creative Commons Attribution 4.0 International License , except where otherwise noted.

Share This Book

Acid-Base Titrations ( CIE IGCSE Chemistry )

Revision note.

Chemistry Lead

Acid-base titrations

• Titrations are a method of analysing the  concentration of solutions
• They can determine exactly how much alkali is needed to neutralise a quantity of acid – and vice versa
• You may be asked to perform titration calculations to determine the moles present in a given amount or the concentration /  volume required to neutralise an acid or a base
• Titrations can also be used to prepare salts
• 25 cm 3  volumetric pipette
• Pipette filler
• 50 cm 3  burette
• 250 cm 3  conical flask
• Small funnel
• 0.1 mol / dm 3  sodium hydroxide solution
• Sulfuric acid of unknown concentration
• A suitable indicator
• Clamp stand, clamp & white tile

The steps in performing a titration

• Use the pipette and pipette filler and place exactly 25 cm 3   sodium hydroxide solution into the conical flask
• Fill the burette with hydrochloric acid , place an empty beaker underneath the tap. Run a small portion of acid through the burette to remove any air bubbles
• Record the starting point on the burette to the nearest 0.05 cm 3
• Place the conical flask on a white tile so the tip of the burette is inside the flask
• Add a few drops of a suitable indicator to the solution in the conical flask
• Perform a rough titration by taking the burette reading and running in the solution in 1 – 3 cm 3  portions, while swirling the flask vigorously
• The endpoint is when one drop causes a sharp colour change
• Make sure your eye is level with the meniscus
• Repeat the titration with a fresh batch of sodium hydroxide
• As the rough end-point volume is approached, add the solution from the burette one drop at a time until the indicator just changes colour
• Record the volume to the nearest 0.05 cm 3  
• Repeat until you achieve two concordant results (two results that are within 0.1 cm 3   of each other) to increase accuracy

Common errors during a titration include:

• This can lead to some liquid dripping into the burette and cause false / high readings
• The jet space is the part of the burette after the tap
• Not filling this space can lead to false readings
• Readings should be taken from the bottom of the meniscus
• Careful: The scale on the burette has 0.0 cm 3 at the top and 50 cm 3 (typically) at the bottom
• Indicators are used to show the endpoint in a titration
• However, methyl orange and phenolphthalein are very suitable
• Some of the most common indicators with their corresponding colours are shown below:

Common acid-base indicators

You've read 0 of your 0 free revision notes

Get unlimited access.

to absolutely everything:

• Downloadable PDFs
• Unlimited Revision Notes
• Topic Questions
• Past Papers
• Model Answers
• Videos (Maths and Science)

Join the 100,000 + Students that ❤️ Save My Exams

the (exam) results speak for themselves:

Did this page help you?

Author: Stewart

Stewart has been an enthusiastic GCSE, IGCSE, A Level and IB teacher for more than 30 years in the UK as well as overseas, and has also been an examiner for IB and A Level. As a long-standing Head of Science, Stewart brings a wealth of experience to creating Topic Questions and revision materials for Save My Exams. Stewart specialises in Chemistry, but has also taught Physics and Environmental Systems and Societies.

Practical Science

Table of Contents

Titration Acid-Base (Simple)

Titration is a fundamental laboratory technique used to determine the concentration of a substance in a solution. In this article, we will walk you through the steps of performing a simple acid-base titration using sodium hydroxide and hydrochloric acid. We will also discuss important safety precautions to keep in mind, such as wearing appropriate protective gear, handling chemicals with care, and disposing of waste properly. Additionally, we will address environmental concerns related to titration, including the safe disposal of chemical waste and glassware. By following these guidelines, you can perform a titration safely and responsibly while protecting yourself and the environment.

The purpose of this experiment is to determine the concentration of a solution of sodium hydroxide by titration against a standard solution of sodium hydroxide.

Introduction

Hydrochloric acid is a monoprotic acid in that it produces one mole of hydrogen ions per mole of compound, we can simplify the formula to HA. This simple formula is often used to represent an acid.

Sodium hydroxide reacts with hydrochloric acid according to the equation:

NaOH + HCl –> NaCl + H 2 O

To show you when the reaction is complete – the stoichiometric point or equivalence point – you use an indicator called phenolphthalein, which is colourless in acid and pink in alkaline solution. The point at which the addition of one drop (or even less) of alkali changes the solution from colourless to just faintly pink is called the end-point and, in this case, shows that the reaction is just complete.

Requirements

• safety spectacles
• filter funnel, small
• burette, 50 cm 3
• 2 beakers, 100 cm 3
• sodium hydroxide solution, approx. 0.5 mol dm 3  NaOH ( CORROSIVE )
• pipette, 25 cm 3
• pipette filler
• Hydrochloric acid solution, 0.5 mol dm 3  ( IRRITANT )
• conical flask, 250 cm 3
• phenolphthalein indicator solution
• wash-bottle of distilled water
• Using the funnel, rinse the burette with the sodium hydroxide solution and fill it with the same solution. Do not forget to rinse and fill the tip. Record the initial burette reading in the ‘rough’ column of the Results Table.
• Using a pipette filler, rinse the pipette with some of the Hydrochloric acid solution and carefully transfer 25.0 cm 3  of the solution to a clean 250 cm 3  conical flask.
• Add 2-3 drops of the phenolphthalein indicator solution.
• Run sodium hydroxide solution from the burette swirling, until the solution just turns pink.
• Refill the burette with the sodium hydroxide solution, and again record the initial burette reading to the nearest 0.05 cm 3  (one drop).
• Using the pipette, transfer 25.0 cm 3  of the hydrochloric acid solution to another clean conical flask. Add 2-3 drops of the phenolphthalein indicator solution.
• Carefully titrate this solution to the end-point, adding the alkali drop by drop when you think the colour is about to change.
• Repeat steps 5, 6 and 7 at least twice more.
• Empty the burette and wash it carefully immediately after the titration, especially if it has a ground glass tap.

You should record burette readings to the nearest 0.05 cm 3  (approximately one drop). Consecutive titrations should agree to within 0.10 cm 3  and, strictly, you should repeat the titration until this is achieved. However, you may have neither the time nor the materials to do this. With practice, your technique will improve so that it is not necessary to do more than four titrations.

Calculate the mean of the two (or preferably three) closest consecutive readings and quote this also to the nearest 0.05 cm 3

Note that this does not introduce a fourth significant figure; it merely makes the third figure more reliable.

Table of Results (PDF)

• What effect would each of the errors described below have on the calculated value of the concentration of sodium hydroxide?

(a) The burette is not rinsed with the sodium hydroxide solution.

(c) The tip of the burette is not filled before titration begins.

(d) The conical flask contains some distilled water before the addition of potassium hydrogen phthalate.

• In using phenolphthalein as an indicator, we prefer to titrate from a colourless to pink solution rather than from pink to colourless. Suggest a reason for this.
• Why is it advisable to remove sodium hydroxide from the burette as soon as possible after the titration.

Share this:

• Click to share on Twitter (Opens in new window)
• Click to share on Facebook (Opens in new window)
• Click to print (Opens in new window)
• Click to email a link to a friend (Opens in new window)

Type your email…

• Already have a WordPress.com account? Log in now.
• Subscribe Subscribed
• Copy shortlink
• Report this content
• View post in Reader
• Manage subscriptions
• Collapse this bar

Master Chemistry

Acid base titration-Working principle, Process, types and indicators

Learning objectives.

Table of Contents

In this article, the  author  has explained Acid base titration, its working principle types and process of acid base titration

Acid-base reactions are used to find the amount of acidic or basic substances. The solution with unknown molarity is the analyte. The analyte will react with the solution with known molarity.

First lets start with a simple definition of titration.

What is titration?

Titration is a chemical analysis method used to determine the concentration of a particular analyte.

Titration, also known as volumetric analysis, is a method in which the titrant is added from a burette until the reaction is complete, and an indicator is usually employed to mark the endpoint of the reaction.

Theory of titration

Titration is a method used to determine the unknown concentration of a substance in a solution. It involves the gradual addition of a known concentration of a solution, called the titrant , to a known volume of the substance being analyzed, called the analyte or titrand. The point at which the two solutions are chemically balanced is called the endpoint of the titration.

Types of titration

There are several types of titrations, including acid-base titrations, redox titrations, and complexometric titrations. Acid-base titrations involve the reaction between an acid and a base to form a salt and water. Redox titrations involve the transfer of electrons from one species to another. Complexometric titrations involve the formation of a complex between the analyte and the titrant.

How endpoint of titration is measured?

The endpoint of the titration can be determined using a variety of methods, including visual indicators, conductivity measurements, or pH measurements. The endpoint is typically indicated by a change in the color of the solution, a change in the electrical conductivity of the solution, or a change in the pH of the solution.

Titrations are widely used in various fields, including chemistry, biology, and environmental science, to determine the concentration of a variety of substances in solutions.

important terms in titration

I have also written few titration related other articles, I am sure you would learn from these as well.

• What are precautions during titration process in lab
• Uses of titration in industries
• Why is titration important in chemistry?

What is acid-base titration?

An acid-base titration is a quantitative analysis method used to determine the concentration of an acid or base by neutralizing the acid or base with a known concentration standard solution.

The concentration of a solution can be determined by knowing the acid and base dissociation constant. If the solution concentration is known, a titration curve can be used.

Principle of acid-base titration

In the theory of acid-base titration, the principle involves using a burette and pipette to determine the concentration of an acid or basic.

An indicator is a dye added to a solution to change its color. It is dissolved in the sample solution and can be used to detect the end of the titration.

Importance of acid base titration

Acid-base titrations are important because they allow for the precise determination of the concentration of an acid or a base in a solution. This information is useful in a variety of fields, including chemistry, biology, and environmental science.

In chemistry, acid-base titrations are used to determine the concentration of acids and bases in solutions, which is important for understanding chemical reactions and for the preparation of standard solutions.

In biology, acid-base titrations are used to determine the pH of solutions, which is important for understanding the behavior of enzymes and other biological molecules.

In environmental science, acid-base titrations are used to determine the acidity or basicity of water, which is important for understanding the impact of acid rain on aquatic ecosystems.

Overall, acid-base titrations are a powerful tool for understanding and quantifying the concentration of acids and bases in solutions, and they have many important applications in various fields.

Acid base titration and volumetric analysis

Acid-base titrations are a type of volumetric analysis, which is a method of chemical analysis that involves the measurement of volume in order to determine the concentration of a substance in a solution. In an acid-base titration, a solution of known concentration (called the titrant) is gradually added to a known volume of the substance being analyzed (called the analyte). The point at which the two solutions are chemically balanced is called the endpoint of the titration.

Indicators used in acid-base titration:

The acid strength of the indicator is important in determining the pH range. The indicator changes color from acid to base when it’s in the range of pH values.

The acid form can only be seen at the highest pH, the base form can only be seen at the lowest. Since the indicator doesn’t change color at certain pH levels, it’s not sensitive to changes outside of its range.

Selection of indicator in acid base titration

There are several factors to consider when selecting an indicator for an acid-base titration:

• The pH range of the titration: The indicator should have a color change within the pH range of the titration. For example, if the titration involves a strong acid and a strong base, the pH range will be wide, and an indicator with a wide range, such as bromothymol blue, can be used. If the titration involves a weak acid and a strong base, the pH range will be narrow, and an indicator with a narrow range, such as phenolphthalein, can be used.
• The desired endpoint: The indicator should undergo a color change at the desired endpoint of the titration. For example, if the endpoint of the titration is the point at which the acid and base are neutralized, an indicator with a pK value close to 7, such as bromocresol green, can be used. If the endpoint of the titration is the point at which the acid and base are in a specific ratio, an indicator with a pK value close to the desired ratio, such as methyl red, can be used.
• The sensitivity of the indicator: The indicator should undergo a noticeable color change at the endpoint of the titration. Some indicators, such as thymol blue, have a sharp color change at the endpoint, while others, such as methyl orange, have a more gradual color change.

Overall, the selection of an indicator for an acid-base titration is based on the pH range of the titration, the desired endpoint, and the sensitivity of the indicator.

Classification of acid-base titration indicators

Acid-base indicators are generally classified into below listed three groups.

1. The phthaleins and sulphophthaleins:  example-phenolphthalein indicator

2. Azo indicators:  example- methyl orange indicator

3. Triphenylmethane indicators: for example- malachite green indicator

Specific indicator for different acid-base titration:

Different indicators are used in acid-base titrations. The selection of indicators depends on the type of titration and the range of the reaction.

• Strong acid-strong base: Phenolphthalein is generally preferred due to color change seen more easily.
• Weak acid-strong base:  Phenolphthalein is more proffered for this titration because it changes sharply at the equivalence point.
• Strong acid-weak base:  Methyl orange is more proffered for this titration because it changes sharply at the equivalence point.
• Weak acid-weak base:  Because a vertical portion of the curve above two pH units is required, there is no indication is suitable for this titration.

Types of acid-base titration with their examples:

There are four different types of acid-base titration, which include strong acid-strong base, weak acid-strong base, strong acid-weak base, and weak acid-weak base.

Strong acid-strong base:

It’s one of the easiest titrations to perform among the four forms of acid-base titrations. It involves the dissociation of a strong acid and a strong base in water, which results in a strong acid-strong base neutralization reaction. The equivalency point is reached when the moles of acid and base are the same and the pH is zero.

Weak acid-strong base

The direct transfer of the protons from the weak acid to the hydroxide ion is possible in this type of titration. The acid and base react in a one-to-one ratio when reacting a weak acid with a strong base. At the equivalent point of a weak acid–strong base titration, the pH is larger than 07.

Strong acid-weak base

The acid and base will react to form an acidic solution. A conjugate acid is formed which reacts with water to form a hydronium ion. At the point of a strong acid-weak base titration, the pH is less than 7.

Weak acid-weak base

The shape of a weak acid’s or base’s titration curve depends on the acid’s or base’s identity and the associated acid ionization constant. In the titration of a weak acid or a weak base, the pH gradually changes around the equivalence point, which is greater or less than 7.

Below are examples of these titrations.

• Hydrochloric acid (HCl) and sulphuric acid (H 2 SO 4 ) are two examples of strong acids.
• Acetic acid (CH 3 COOH) and formic acid (CH 2 O 2 ) are two examples of weak acids.
• Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are two examples of strong bases.
• Ammonia and methylamine are two examples of weak bases.

Apparatus and Process of titration

First of all we need to arrange the following apparatus to perform acid-base titration.

• Conical flask
• burette stand
• wash bottle
• unknown solution
• Standard solution

Then we need to clean all the apparatus using distilled water.

After then we need to fill the burette with a standardized solution, accurately measure the volume of the analyte, and add in the conical flask, also add a few drops of indicator using the pipette.

Titrate it with the standardized solution until the indicator changes color. When the indicator permanently changes the color, the endpoint reaches.

Repeat the titration at least three more times and record the initial and final readings in the observation table and calculate the value.

How does titration work?

Titration is a method used to determine the concentration of a substance in a solution. It involves adding a known concentration of a solution (titrant) to a known volume of the substance (analyte) and measuring the volume of titrant at the point of chemical balance (endpoint).

What is the major purpose of acid base titration?

The major purpose of acid-base titration is to determine the concentration of an acid or a base in a solution.

Which indicator is used in acid base titration?

An indicator is chosen for an acid-base titration based on the pH range of the titration, the desired endpoint, and the sensitivity of the indicator. Some commonly used indicators include bromothymol blue, phenolphthalein, bromocresol green, methyl red, thymol blue, and methyl orange.

is titration only for acids and bases?

No, titration can be used to determine the concentration of various substances, including acids, bases, and other types of molecules. There are several types of titrations, including acid-base, redox, and complexometric.

How do acid base indicators work?

Acid-base indicators change color at a specific pH and are used to indicate the endpoint of a titration. They are chosen based on the desired endpoint and the pH range of the titration, and their color change at the endpoint is used to determine the concentration of the acid or base in the solution.

Why is phenolphthalein an appropriate indicator for a weak acid-strong base titration?

Phenolphthalein is a suitable indicator for a weak acid-strong base titration because it has a narrow pH range and undergoes a sharp color change at the endpoint.

Find Study Materials for

• Explanations
• Business Studies
• Combined Science
• Computer Science
• Engineering
• English literature
• Environmental Science
• Human Geography
• Macroeconomics
• Microeconomics
• Social Studies
• Browse all subjects
• Textbook Solutions
• Read our Magazine

Create Study Materials

• Flashcards Create and find the best flashcards.
• Notes Create notes faster than ever before.
• Study Sets Everything you need for your studies in one place.
• Study Plans Stop procrastinating with our smart planner features.

Chapter 18: Problem 10

Describe acid-base titration. What is the equivalence point?

Short answer, step by step solution, introduction to acid-base titration, understanding the equivalence point, detection of the equivalence point, key concepts.

These are the key concepts you need to understand to accurately answer the question.

Equivalence Point

Ph indicators, one app. one place for learning..

All the tools & learning materials you need for study success - in one app.

Most popular questions from this chapter

Describe the solubility of CaF \(_{2}\) in each solution compared to its solubility in water. a. in a 0.10 \(\mathrm{M} \mathrm{NaCl}\) solution b. in a 0.10 \(\mathrm{M}\) NaF solution c. in a 0.10 \(\mathrm{M}\) HCl solution

The pH at the equivalence point of the titration of a strong acid with a strong base is \(7.0 .\) However, the \(\mathrm{pH}\) at the equivalence point of the titration of a weak acid with a strong base is above \(7.0 .\) Explain.

What is molar solubility? How do you obtain the molar solubility of a compound from \(K_{\mathrm{sp}}\) ?

What is the pH range of human blood? How is human blood maintained in this pH range?

Is each compound more soluble in acidic solution or in pure water? Explain. a. \(\mathrm{Hg}_{2} \mathrm{Br}_{2}\) b. \(\mathrm{Mg}(\mathrm{OH})_{2}\) c. \(\mathrm{CaCO}_{3}\) d. \(\mathrm{AgI}\)

Recommended explanations on Chemistry Textbooks

Nuclear chemistry, making measurements, chemistry branches, ionic and molecular compounds, inorganic chemistry, the earths atmosphere.

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.

Privacy overview.

• school Campus Bookshelves
• menu_book Bookshelves
• perm_media Learning Objects
• login Login
• how_to_reg Request Instructor Account
• hub Instructor Commons

Margin Size

• Download Page (PDF)
• Download Full Book (PDF)
• Periodic Table
• Physics Constants
• Scientific Calculator
• Reference & Cite
• Tools expand_more
• Readability

selected template will load here

This action is not available.

7.2: Lab - Titrations

• Last updated
• Save as PDF
• Page ID 352566

\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

\( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

\( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

\( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

\( \newcommand{\Span}{\mathrm{span}}\)

\( \newcommand{\id}{\mathrm{id}}\)

\( \newcommand{\kernel}{\mathrm{null}\,}\)

\( \newcommand{\range}{\mathrm{range}\,}\)

\( \newcommand{\RealPart}{\mathrm{Re}}\)

\( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

\( \newcommand{\Argument}{\mathrm{Arg}}\)

\( \newcommand{\norm}[1]{\| #1 \|}\)

\( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

\( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

\( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

\( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

\( \newcommand{\vectorC}[1]{\textbf{#1}} \)

\( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

\( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

\( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

Learning Objectives

Goals: 

• Understand that parts of an acid-base titration
• be able to determine the K a  or K b  from pH data associated with the titration of a weak acid or base
• be able to determine the molar mass of a solid monoprotic acid from titration data
• be able to calculate K a1  and K a2  for a polyprotic acid

By the end of this lab, students should be able to:

• design , develop and perform  acid base titrations .

Prior knowledge:

• Stoichiometry of Acid-Base Titrations in quantitative analysis (s ection 4.7.2 )

Concurrent Reading & Additional Resources

• pH and pOH  ( section 16.2 )
• pH of Weak Acids and Bases ( section 16.3 , 16.5.4 and 16.5.5 )
• pH of salts ( sections 16.4 , 16.5.6 and 16.5.7 )
• buffers ( section 17.2 )
• E ye protection is mandatory in this lab, and you should not wear shorts or open toed shoes. 

The following image shows the damage 4 M sodium hydroxide can cause.  By following proper procedures and using proper PPE (Personal Protective Equipment) the risk to damage can be reduced to near zero. There are two primary factors that determine the extent of injury due to exposure to corrosive chemicals like NaOH, the concentration of the chemical and the time of contact (exposure to the chemical).  For this reason we do not perform titrations with concentrated acids and bases, but dilute ones.  If you ever have an acid or base spill you need to immediately inform your instructor, who will clean it up or instruct you on what to do.  If you spill it on your body you should immediately wash it off with copious amounts of water.

• Sodium hydroxide
• Acetic Acid
• Check the cord on the hotplate, inform the instructor if it is frayed.
• Make sure the electrical cord never touches the surface of the hotplate
• label all containers
• All solutions should be considered harmful and care should be taken to avoid contact with your skin or other body tissues. 
• In event of contact with reagents you should flush contacted area with water and notify instructor immediately.
• All waste is placed in the labeled container in the hood and will be recycled when the lab is over.  Contact your instructor if the waste container is full, or about full.

Equipment and materials needed

Indicator Titrations

The next several labs will involve laboratory techniques that take into account the equilibrium chemistry associated with the stoichiometry of chemical reactions. The first will involve titrations of acid base reactions that may include the equilibria of weak acids or bases and the second will deal with the formation of complex ions.  The stoichiometry of acid base titrations were introduced in the first semester ( section 4.7 ) as an analytical technique to determine the concentration of an unknown ( analyte ) by adding a standard of known concentration ( titrant ) until they were in stoichiometric proportions ( the equivalence point) .  At the equivalence point the analyte has been neutralized and converted to its salt (conjugate form). The volume of titrant required to neutralize the analyte could be quickly determined through the use of an appropriate indicator, where titrant was added until the solution changed color, which was at a volume known as the endpoint of the titration.  For this to work the pH at which an indicator changes color must be the same as that of the salt of the analyte being neutralized. The analyte may be weak or strong, but the titrant must be strong and typically is monoprotic. Figure  \(\PageIndex{2}\) shows the experimental setup for an indicator based titration

Figure \(\PageIndex{6}\) of the experimental section is a chart of the colors and pHs various indicators change at and it is important that you pick an appropriate indicator based on the acidity or basicity of the neutralized analyte.  If the analyte is a strong acid or base the indicator should change color around a pH of 7.  If the analyte is a weak acid the indicator should change color in a basic solution, and if it is a weak base it should change color in an acidic solution.  The calculations for determining this are in the expermintal section of this lab.

pH Titrations

In a pH titration you measure the pH as a function of the volume of titrant added and determine the equivalence point as the point in where there is an inflection in the slope of the curve.  Figure \(\PageIndex{2}\) shows the four common types of titrations.  Initially the pH is that of the pure analyte. Before the equivalence point the titrant is neutralizing the analyte and converting it to its salt, but since there is an excess of the analyte it is not completely consumed and so a buffer is formed, which is a mixture of the analyte and its salt. As the equivalence point is approached the pH changes rapidly as the analyte is being consumed and so there is nothing to react with the titrant (which must be a strong acid or base) and after the equivalence point the pH stabilizes as it is effectively determined by the pure titrant, which is both strong and in excess, and a change is only due to dilution.  Lets look at two examples

• Titration of Acetic Acid with Sodium hydroxide (analagous to figure \(\PageIndex{3}\)c.  Initially the pH is due to pure acetic acid . As sodium hydroxide is added it reacts with the acetic acid forming its conjugate base,  the salt sodium acetate.  This is an acetic acid/acetate buffer and the pH is determined by the ratio of the un-neutralized to neutralized acetic acid. \[C_2H_3O_2(aq) + NaOH(aq)  \rightarrow NaC_2H_3O_2(aq) + H_2O(l) \] At the equivalence point all the acetic acid  has been neutralized  and converted to its conjugate base acetate, so the pH is determined by the concentration of its salt, sodium acetate.  After the equivalence point there is no more acetic acid to react with the sodium hydroxide and so it accumulates, with the pH being dictated by the amount of excess sodium hydroxide.
• Titration of Ammonia with Hydrochloric Acid  (analagous to figure \(\PageIndex{3}\)d.  Initially the pH is due to pure ammonia  As HCl is added it reacts with the ammonia forming its salt, ammonium chloride.  This is an ammonia/ammonium buffer and the pH is determined by the ratio of the un-neutralized to neutralized ammonia. \[NH_3(aq) + HCl(aq) \rightarrow NH_4Cl(aq)  \] At the equivalence point  all the ammonia has been neutralized and converted to its conjugate acid ammonium, and so the pH is dictated by the concentration of the ammonium chloride salt.  After the equivalence point there is no more ammonia to react with the HCl and so it accumulates and the pH is dictated by the excess HCl.

Equivalence Point Determination

A closer look at figure \(\PageIndex{3}\) indicates that the steepest part of the titration curve is the equivalence point and that there is an inflection in the slope of the line as the solution goes from excess analyte to excess titrant.  Lets look at parts (a) and (c) of figure \(\PageIndex{3}\).  As you approach the equivalence point the slope increases and then after the equivalence it decreases.  This can be shown with a first derivative plot of the curve as in figure \(\PageIndex{4}\) \(\left ( lim \;\Delta V \to 0 \;  \frac{\Delta pH}{\Delta V} \right )\) .  What is clear in the first derivative plot is that the line is going higher and higher and then reverses direction and goes lower and lower.  The point where it changes from increasing to decreasing is the inflection point, and this can identified where the second derivative plot goes through zero \(\left ( lim \;\Delta V \to 0 \;  \frac{\Delta^2 pH}{\Delta V^2} \right )\) . Note, if there is noise in your data over the flat portion of the curve you will have a lot of false inflection points and so you do not need to take the second derivitive plot over all the data, just in the region around the equivalence point

Titration of a weak monoprotic acid.

Figure \(\PageIndex{4}\) shows the four "regions" of the titration curve for the titration of a weak acid with a strong base. It should be noted that region two is a buffer because there is excess acid (analyte) and so only part of it been neutralized by the base and converted to it's salt (the acid's conjugate base). This is one of the two ways to make a buffer (see section 17.2.3 ).

Weak Acid + Strong Base -> Salt +Water

where the salt is the conjugate base of the acid.

The four parts of the titration curve are described below and you should look to the approriate text section to see how they are treated.

• Pure Acid (0 ml of base is added, section 17.3.2.1 )
• K a  can be calculated from the pH at half equivalence ( section 17.3.2.2.1 )
• If the acid was a solution you can determine its molarity from he volume titrated.
• If the acid was a solid you can determine its molar mass from the mass titrated.
• Excess Base (you have added more base than there was acid, section 17.3.2.4 )

Figure  \(\PageIndex{5}\): Four parts of the titration curve for a weak acid being titrated with a strong base. Notice that two parts are points (1 & 3) and two parts are regions (2 & 4).

Be sure to go over the four parts of the titration curve in section 17.3.2 as that material is not being repeated here.  If you are titrating a weak base with a strong acid you should go over section 17.3.4

Titration of Weak Diprotic Acid

A diprotic acid has two titratable protons and if K a1 >1000K a2  there are two clearly defined equivalence points.

Determination of K a

K a  can be determined by reading the pH at half equivalence (when half of the acid has been neutralized and converted to its salt).  This is in the buffer region and uses the Henderson Hasselbach equation

\[pH=pK_a+\log \dfrac{[A^-]}{HA}\]

Since at half equivalence [HA]=[A - ] pH = pK a  , at half equivalence

\[K_a =10^{-pH\text{, at half equivalence}} \]

So you find the equivalent point on the titration curve and read the value of the curve at half of that volume.  For this reason you need to collect data half way along the curve (red circle).

In this experiment we will use a Ph probe, which is an electronic device that measures the pH. These are very common and they should always be checked against standard solutions of known pH and calibrated if they read incorrectly. The pH probe is an electrochemical cell and we will cover these in chapter 19 , sections 19.3-19.5 and 19.7. The following YouTube from Oxford Press does an excellent job of describing how a pH probe works.  It is imperative that you test your probe in a buffer to be sure it is reading accurately and if it is not, you will need to calibrate it.

Video \(\PageIndex{1}\) 2:30 YouTuve describing the operation of a pH probe developed by Oxford University Press ( https://youtu.be/aIn4D2QXUy4 ).

The pH reading is not accurate until the probe stabilizes, so when you change the pH you need to wait until the reading becomes steady before recording the value.

Exploratory Run

Before running a pH titration you should make a trial run with an indicator ( section 17.3.4.2 ), which is a chemical that undergoes a color change at a specific pH.  You want an indicator that indicates when the titrant and analyte have been added in stoichiometric proportions, (the equivalence point), which is when the analyte has been converted to its salt. From section 17.3.3.2 we see that for the titration of a weak acid

\[[OH^-]=\sqrt{\left (\frac{K_w}{K_a} \right )[A^-]_e} \]

where K w =water ionization constant (10 -14 ), K a =acid ionization constant and [A - e ]=the salt concentration at the equivalence point (when all the acid is neutralized). 

So  \[pOH =-log\sqrt{\left (\frac{K_w}{K_a} \right )[A^-]_e} \]

and \[pH=14-pOH=14+ \sqrt{\left (\frac{K_w}{K_a} \right )[A^-]_e} \]

In the first experiment we are neutralizing 25.00 mL of 0.100M acetic acid with 0.100M NaOH, and so when 25.00 mL of NaOH has been added all the acetic acid will be converted to acetate ions, but the volume has doubled and so the concentration is now 0.05M A - .  This gives a pH of 

\[pH=14-pOH=14+ \sqrt{\left (\frac{10^{-14}}{1.8x10^{-5}} \right )[0.05M]_e}= 8.72\]

From figure \(\PageIndex{4}\) we see that phenolphthalein would be a good indicator for a weak acid like acetic acid as it is clear up until just below a pH of 9, when it turns pink.

That is, you want an indicator that changes color at the pH of the salt of the acid or base that you are titrating, and that way you can tell when you have completely neutralized it.

Experimental Procedures

Before running a pH titration we are going to make a quick exploratory run with an indicator.  The endpoint of an indicator titration is when the indicator changes color and if we choose an indicator that changes color at the pH of the salt of the analyte, it gives us a bearing on the equivalence point, which is when that analyte and titrant have been added in stoichiometric proportions (moles acid = moles base for monoprotic acid being titrated with a monoprotic base).  You will then use the results of the exploratory run to design the pH titration.  

Designing the pH titration

The challenge is that the pH probes are old and it takes a while for their readings to stabilize.  If you do not let the reading stabilize there will be a lot of noise in your data.

The goal of the exploratory run is to figure out where the equivalence point is.  You then need to record data dropwise for about 3/4 a mL before and after the endpoint and collect data around half equivalence.  You need to be sure to record the inital pH (pure acid) and extend your data at least 5 mL beyond the equivalence point.

Tips on getting good data

If you look at YouTube you will see many titrations where people are creating a vortex with magnetic stirrers.  This can increase the rate at which gasses dissolve and there are a class of non-metal oxides called the acid anhydrides that form acids when they combine with water.  The carbon dioxide you exhale is an acid anhydride and the following youtube shows what happens if you breath over a beaker that is rapidly being stirred.

Video \(\PageIndex{2}\): 1:10 minute video showing effect of cavitation and breathing on a slightly basic solution ( https://youtu.be/4RiftqpXI8c , Belford)

In the above video a slightly basic solution with phenolphthalein indicator is pink.  If you listen carefully you can hear someone breathing above it and due to the vortex the carbon dioxide they exhale reacts with the water to form carbonic acid and the solution turns clear.

\[CO_2(g) +H_2O(l) \leftrightharpoons H_2CO_3(aq)\]

The goal of the exploratory run is to give you a feeling for the volume of actual titrant you will need to neutralize 25 mL of your analyte. Using a volumetric pipette 25 mL of acetic acid and a few drops of phenolphthalein were added to the Erlenmeyer flask. Then it was titrated with 0.1M NaOH and the volume of NaOH needed to neutralize the acetic acid was quickly determined. This video shows how to quickly do this, and we are not using this to measure the concentration, but to get a quick bearing on how to design the pH titration.

Video \(\PageIndex{3}\): 2:01 minute video showing a quick exploratory run with an indicator.  Note how as the titration proceeds the color takes longer to disappear as the solution approaches the end point. 

In this experiment we will hook up the Vernier pH probe to a $35 Raspberry Pi microcomputer that transmits the data to your Google Sheet in real time. We will run two python programs on the Raspberry Pi.  The first program we will run from the command line and it gives you the pH readings every 10 seconds, and you use this to decide when to upload data to your Google Sheet.  The second program you will run from the Thonny IDE (Interactive Development Environment), and this program will allow you to input your volumes and pH to your Google Sheet.  These programs were developed during the COVID pandemic to allow instructors to stream data to students in real time, and we have decided to let students use them directly, as data science skills are important for today's students to learn.  Further information may be obtained in the Internet of Science Things course at UALR.  Please do not get water on the Raspberry Pis as you will kill them.

COVID 19 Protocols

The second reason we have decided to use the Raspberry Pis is that we feel the lab can be run safer in a pandemic than using the normal equipment.  We ask students to take up roles for each experiment, and change the roles when they perform different titrations.  Note, at the discretion of your instructor these roles may be modified by the number of people in your group.  Each group will have two stations on opposite sides of the bench.  On one side is the Raspberry Pi, keyboard and monitor, and on the other side is the titration setup.  All students need to work together, make sure the lab is run safely and that you get the best data possible.

• Titrator.  Only one person handles the buret (opens and closes the stopcock). 
• burette reader. This person assists the titrator and reads the volume. This may be the person running the titration.
• Titration supervisor. This person coordinates with the titrator and burette reader to determine the approriate volumes for when they should make a measurement and for communicating with the data supervisor. 
• Pi operator.  This is the only person who touches the keyboard and mouse.  This person runs the python programs.
• Data supervisor. This person assists the pi operator in determining when the pH is stable enough to upload to the Google Sheet, and is responsible for communicating with the titration supervisor.

IOT Enhanced Titration

The following image shows the setup for the titration lab. Note this Raspberry Pi is being run in "headless mode" through VNC viewer and connected to a laptop.  Your Pi may be run in desktop mode, where it is connected to a monitor and keyboard.  Note the tip of the pH probe is submerged and magnetic stirrer is set where it is under the burette and does not touch the probe as it spins.  A gentle spin is all you need.

  Be sure to add enough water to submerge the pH probe and take the dilution effect of this water into account when determining the initial concentration of the acid. 

The following image shows all the programs on the desktop of a Raspberry Pi (remotely accessed via VNC viewer).  Step (1) of the blue commands show you how to run the "current pH" program in command line, where (2) shows the display with the values being streamed every 10 seconds.  The green tabs show how to activate and run the progam that sends data to the Google Sheet (ph_Venier_sheets.py).  (1) Open the folder pH_lab on the desktop and then (2) clicking the python program pH_Veneir_sheets.py opens that program in the Thonny. The shell of Thonny (3) allows you to input the volume in mL (do not include units) and when you hit <enter> that volume and the current pH are uploaded to the Google Sheet.  You need to use both progams concurrently.   

Stepwise procedures

• Add titrant
• Record the new volume of titrant added to analyte in Thonny Shell (running pH_Venier_sheets.py)
• Observe pH in command line (running current_pH.py)
• When pH is stable, hit <enter> on Thonny
• repeat above steps with a new volume of titrant

Note, you do not need to run the Google Sheet, but it would be nice to be sure your data is being uploaded (everytime you hit enter on Thonny your sheet should be updated).  You can also run the Google Sheet on a separate laptop.

Printable Instructions

Experiments

Experiment 1: weak acid.

• Using a  volumetric pipette transfer 25 mL 0.1M Acetic Acid to a 250 mL Erlenmeyer flask
• Add a few drops of phenolphthalein
• Attach burette to a ring stand and fill with 0.1 M NaOH
• Titrate to endpoint when solution turns pink and use this value when decided what volumes to take data at in the pH titration
• Place probe in buffer solution
• In command line (The black box to the left of screen) you will see
• Sudo python /home/Desktop/pH_lab/current_ph.py
• First time running it will sometime get error, just hit up arrow on keyboard then enter
• After calibrating before pressing ok, go to storage tab
• Select save calibration to sensor, then press ok twice when prompted 
• Using a  volumetric pipette transfer 25 mL 0.1M Acetic Acid to a 500 mL beaker.
• Add 50 mL water to make sure the pH probe will be fully submerged. If not, add more water and record total volume of water added
• Set up titration station like the demo station in the lab. 
• Make sure heat is OFF. 
• Carefully transfer 0.100 M NaOH to the burette
• Gently turn on the magnetic stirrer and make sure it does not hit the sides of the beaker or the probe.
• Press Green play button to start
• Enter y to make your group's tab
• Enter a name for your group 
• Look at the  black box (current ph reading) and wait for the numbers to stabilize
• When ready Enter volume of titrant (first will be zero) then press enter
• Add appropriate amount of base (see image on back side of handout)
• Record total volume base added in the Thonny shell (do not hit enter)
• Observe the pH in command line and when it is stable, click enter in Thonny shell
• Repeat above steps adding the next increment of base, while recording the total amount of base added in the Thonny shell.
• Be sure to collect data at least 5 mL beyond the equivalence point
• Check data on the data tab of lab report copy

Experiment 2: Solid Acid

• Weigh around 0.5 grams of unknown acid.
• Transfer acid to 500 mL beaker.
• Add around 50 mL of water to beaker and dissolve solid acid
• Titrate to the endpoint when the solution turns pink and use this value to design the pH titration and choose volumes to take data at.
• Place probe in buffer solution 
• If you need to calibrate then plug the probe into the lab quest and calibrate as normal
• Select save calibration to sensor, then press ok twice when prompted
• Reattach pH probe to Raspberry Pi via Go!Link
• Set up a titration station like the demo station in the lab.
• Weigh around 0.5 grams of unknown acid and record value to the precision of the balance on data sheet
• Add 50 mL water and make sure the pH probe tip is fully submerged. If not, add more water.
• Turn on the magnetic stirrer and make sure it does not hit the sides of the beaker or the probe. Avoid cavitation

Experiment 3: Diprotic Acid

The challenge with the drop counter is to set the flow rate slow enough so that the pH readings stabilize between drops. The burette has two valves, the top one to adjust the flow rate and the bottom one to turn on/off the device.  Before calibrating the drop counter you need to adjust the flow rate by filling the burette with about 40 mL of water, 

• Open bottom valve and close top valve of burette and place on stand.
• place 100 mL beaker under burette.
• Add around 40 mL water to burrette
• Slowly open top valve until a drop falls every 2 to 4 second.
• When desired rate is achieved close bottom valve.

Video \(\PageIndex{4}\):  

• Place 10 mL graduated cylinder under burette 
• Align drop detector so drops break cross IR beam path as they fall from burette to graduated cylinder
• Choose calibrate
• Choose calibrate now
• Open valve and the counter should increment one value every time a drop falls.
• After you have around 9 mL you can stop the calibration and the device will know record the volume of each drop

• Set up calibrated drop counter and burette over a 500 mL beaker containing 50 mL of water
• Connect pH sensor to LabQuest Analog port, it should read the pH

• Add magnetic stirrer, set up to a slow stir rate and make sure it does not hit the probe
• Add 15 mL of the diprotic acid to a 500 mL beaker
• Submerge pH probe into solution, add water if probe not submerged
• Switch to graphing mode so you can observe the titration curve grow in real time.
• complete for about 5 mL in the excess base region after the second equivalence point (there should be two buffer regions, do not mistake the second buffer for being excess base.

Data Analysis

Experiment 1 (Titration of Amino Acids) Updated Protocol (24M)