titration experiment pink

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Phenolphthalein Indicator

Phenolphthalein indicator is a pH indicator solution that is colorless under acidic and extremely alkaline conditions and fuchsia pink under ordinary alkaline conditions. Here is a look at the colors at different pH ranges, the chemistry of the color change, how to make phenolphthalein indicator, and uses of the chemical.

Phenolphthalein pH Range and Colors

The two main colors of phenolphthalein are colorless and fuchsia, which is a pink-purple-red color, depending on the indicator concentration. Phenolphthalein is colorless under acidic, neutral, and slightly basic conditions. It’s pink between pH values of 8.3 and 10.0 and then becomes colorless again at pH values greater than 10. At extremely low pH, like concentrated sulfuric acid, the color turns orange-red. However, the orange-red color is not seen with all strong acids.

There are at least four states of phenolphthalein in aqueous solution. Under extremely acidic conditions, the indicator is in an orange protonated form (HIn + ). Under strongly acidic, neutral, and slightly alkaline conditions, it presents the colorless lactone form (HIn). The purple-pink color forms under basic conditions, where the doubly-deprotonated or anion form of phenol (In 2- ) predominates. As alkalinity increases, the color slowly fades as the indicator converts into the In(OH) 3- form.

How to Make Phenolphthalein Indicator Solution

You can purchase pre-mixed phenolphthalein indicator solution, but it is easy to prepare yourself. Phenolphthalein is somewhat soluble in water (400 mg/L), but it is highly soluble in either ethanol or ether. So, usually you first dissolve the powdered dye in alcohol and then dilute it with water or else dissolve it in 50% alcohol.

Here is the recipe for 1% phenolphthalein indicator solution:

• Weigh out 1.0 grams of phenolphthalein.
• Dissolve it in 100 milliliters of 50% ethanol in water. Alternatively, dissolve it in a solution of 50 milliliters absolute ethanol and 50 milliliters water.

Another common formulation is a 0.5% solution:

• Weigh out 0.5 grams of phenolphthalein.
• Dissolve it in 50 milliliters of ethanol.
• Add distilled water to make a final volume of 100 milliliters.

Phenolphthalein Uses

One of the most common uses for phenolphthalein indicator is for finding the end point of an acid-base titration. However, the chemical has other purposes.

• pH indicator for acid-base titration
• pH indicator for concrete carbonation
• Disappearing ink
• Color change toys, such as “Hollywood Hair” Barbie
• Colored bubbles
• Laxative: The use of phenolphthalein as a laxative has been largely discontinued, due potential carcinogenicity.
• Phenolphthalin, the reduced form of phenolphthalein, is used in the Kastle-Meyer test for blood.
• Phenolphthalein is a component of universal indicator .
• Dunnick, J. K.; Hailey, J. R. (1996). “ Phenolphthalein Exposure Causes Multiple Carcinogenic Effects in Experimental Model Systems .” Cancer Research . 56 (21): 4922–4926.
• Schwarzenbach, Gerold (1957). Complexometric Titrations . Translated by Irving, Harry (1st English ed.). London: Methuen & Co.
• Wittke, Georg (1983). “Reactions of phenolphthalein at various pH values”. Journal of Chemical Education . 60 (3): 239. doi: 10.1021/ed060p239
• Zumdahl, Steven S. (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company.

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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?

More resources

Inspire learners and discover more ways chemists are making a difference to our world with our video job profiles .

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

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Introduction to Titration

Source: Laboratory of Dr. Yee Nee Tan — Agency for Science, Technology, and Research

Titration is a common technique used to quantitatively determine the unknown concentration of an identified analyte. 1-4 It is also called volumetric analysis, as the measurement of volumes is critical in titration. There are many types of titrations based on the types of reactions they exploit. The most common types are acid-base titrations and redox titrations. 5-11

In a typical titration process, a standard solution of titrant in a burette is gradually applied to react with an analyte with an unknown concentration in an Erlenmeyer flask. For acid-base titration, a pH indicator is usually added in the analyte solution to indicate the endpoint of titration. 12 Instead of adding pH indicators, pH can also be monitored using a pH meter during a titration process and the endpoint is determined graphically from a pH titration curve. The volume of titrant recorded at the endpoint can be used to calculate the concentration of the analyte based on the reaction stoichiometry.

For the acid-base titration presented in this video, the titrant is a standardized sodium hydroxide solution and the analyte is domestic vinegar. Vinegar is an acidic liquid that is frequently used as a culinary condiment or flavorings. Vinegar mainly consists of acetic acid (CH 3 COOH) and water. The acetic acid content of commercial vinegar can vary widely and the goal of this experiment is to determine the acetic acid content of commercial vinegar by titration.

The determination of acetic acid in vinegar is based on the principle of an acid-base titration method. The reaction between NaOH and CH 3 COOH is shown in Equation 1 :

CH 3 COOH (aq) + NaOH (aq)  → H 2 O (l) + NaCH 3 CO 2(aq)  ( 1 )

The standardized NaOH solution is progressively added to the vinegar with unknown acetic acid concentration until the end point is reached. During the acid-base titration, the pH can be plotted as a function of the volume of the titrant added. The inflection point on the curve, the point at which there is a stoichiometric equal amount of acid and base in a solution, is called the equivalence point. Most acids and bases are colorless, with no visible reaction occurring at the equivalence point. To observe when the equivalence point has been reached, a pH indicator is added. The endpoint is not the equivalence point but a point at which the pH indicator changes color. It is important to select an appropriate pH indicator so that the end point is as close to the equivalence point of titration as possible.

At the end point of this reaction, the conjugate base NaCH 3 CO 2 is slightly basic. Phenolphthalein indicator has a working pH range of 8.3–10.0, which is colorless in acidic solution and magenta above pH 8.2. Therefore, phenolphthalein is a preferred indicator as it will change from colorless to pink at this condition. When performing the experiment, it is best to keep the concentration of pH indicator low because pH indicators themselves are usually weak acids that react with base.

The volume of standardized NaOH solution added at end point can then be used to calculate the molar concentrations of acetic acid based on the stoichiometry of the above equation. In this experiment, the titrant NaOH is a strong alkaline and the analyte acetic acid is a weak acid.

Before performing the experiment, it is important to consider the hygroscopic nature of NaOH. This property requires its solution to be standardized with a stable primary standard such as potassium hydrogen phthalate (KHC 8 H 4 O 4 ). The exact molar concentration of NaOH solution can then be accurately determined after standardization. The reaction between the primary acid standard and NaOH is shown in Equation 2 :

KHC 8 H 4 O 4(aq)  + NaOH (aq) → H 2 O (l) + NaKC 8 H 4 O 4(aq)   ( 2 )

A detailed step-by-step titration protocol is presented in the following section.

1. Standardization of NaOH with Potassium Hydrogenphthalate (KHC 8 H 4 O 4 )

• To begin, the titrant, sodium hydroxide, must be standardized. Prepare a stock NaOH solution by dissolving about 4 g of NaOH pellet in 100 mL of deionized water. Note that NaOH is a hazardous chemical which is corrosive to skin and irritant to eye, be cautious and wear proper personal protection equipment (PPE) to avoid skin or eye contact.
• Make a 1:10 dilution of the sodium hydroxide solution by adding 25 mL of the stock sodium hydroxide solution to a 500-mL bottle. Sodium hydroxide absorbs carbon dioxide. It is important to prevent this by making sure to use boiled, deionized water, an oven-dried bottle, and to cap the bottle quickly. Make the solution up to 250 mL with the deionized water and shake to mix.
• Dry 4–5 g of the primary standard acid, KHC 8 H 4 O 4 at 110 °C for 4 h in a drying oven and then cool the solid in a desiccator for 1 h.
• Dissolve about 4 g of dried KHC 8 H 4 O 4 in 250 mL of deionized water. Record the mass accurately. Calculate the molar concentration of the KHC 8 H 4 O 4 solution.
• Pipette 25 mL of KHC 8 H 4 O 4 into a clean and dry Erlenmeyer flask. Add 2 drops of phenolphthalein, and swirl gently to mix well. Note that phenolphthalein is toxic and irritant, be cautious to avoid skin or eye contact.
• Clean a 50-mL burette and a funnel thoroughly with detergent and water. Flush the burette with water and rinse 3x with deionized water. Rinse the burette with the diluted NaOH solution 3x, making sure that the NaOH wets the entire inner surface and drain the waste through the tip. Mount the washed burette on a ringstand with a clamp and ensure that it stands vertically.
• Fill the clean burette with the diluted NaOH solution. It should be noted that the amount of the diluted NaOH needs not be exactly at the zero mark but should be within the scale and sufficient for at least one titration. Air bubbles may affect the accuracy of volume reading. Carefully check the burette for air bubbles, and gently tap the burette to free them and open the stopcock to let a few mL of titrant to flow through and at the same time releasing any trapped air. Read the volume by viewing the bottom of the meniscus after 10 s. Record this initial volume. Pay attention to the significant figures of the reading. Record the value to two decimal places in mL.
• Place the Erlenmeyer flask containing potassium hydrogen phthalate (KHC 8 H 4 O 4 ) under the burette and adjust the height of the burette properly. Titrate the KHC 8 H 4 O 4 solution by slowly adding NaOH solution in 1–2 mL increments using one hand to control the flow rate by adjusting the stopcock, and the other swirling the flask.
• When close to the endpoint, begin adding the titrant drop by drop. The endpoint is reached when the solution turns a faint, persistent pink color. Record the final volume of the diluted NaOH in the burette.
• Repeat the titration at least twice more to obtain consistent data. Calculate the molar concentration of the diluted NaOH solution.

2. Titration of Vinegar with Standardized Sodium Hydroxide Solution

• The sodium hydroxide solution is now standardized and can be used as a titrant to analyze vinegar. To reduce the pungent aroma of vinegar, dilute 10 mL of the vinegar solution to be tested in a 1:10 ration to a total volume of 100 mL.
• Pipette 25 mL of analyte, to a clean and dry Erlenmeyer flask (noted as V A ). Add 2 drops of phenolphthalein.
• Fill the burette with the standardized NaOH solution from the first part of the Procedure. Record the initial volume of titrant (V 1 ).
• Progressively add the standardized NaOH solution to the vinegar. When the volume of titrant approaches the expected value, adjust the stopcock to add the titrant drop by drop. Continue to swirl the flask with one hand and keep the other hand ready to close the stopcock. Once the analyte solution changes to light pink color, swirl for a few seconds to see whether the color will fade. If the color persists, the titration reaches the end point. Record the final volume of titrant (V 1 ' ). If the solution color fades, add one more drop of titrant. Wash the bottom tip of the burette using the wash bottle. Collect the washed mixture and watch the color change of the analyte solution. Continue the titration till the end point. Record the amount of titrant needed (V t1 = V 1 ' V 1 ).
• Repeat the titration at least twice until three concordant values that are within 0.1 mL of one another is obtained (V t2 and V t3 ).

Titration is a commonly applied method of quantitative chemical analysis used to determine the unknown concentration of a solution. A typical titration is based on a reaction between a titrant and an analyte. The titrant of known concentration is gradually added to a precise volume of an unknown analyte until the reaction reaches an endpoint.

At the endpoint, the moles of titrant and analyte are equal. By manipulating the equation relating volume and concentration, the concentration of analyte can be deduced.

This video will illustrate the principles behind titration, present a protocol to determine the amount of acetic acid in commercial vinegar, and finally explore some common applications of the method.

Titrations are classified based on the type of reaction carried out. For example, redox titrations make use of an oxidation-reduction exchange between reactants which involves the transfer of electrons from one reactant to another. Complexometric titrations rely on the formation of a largely undissociated complex. However, acid-base titrations, which exploit the neutralization of an acid with a base, are one of the most widely studied. To determine the concentration of acid in an analyte, a base, such as sodium hydroxide, is used. Sodium hydroxide is hygroscopic, that is, it has the property of absorbing moisture from the atmosphere. Before it can be used as a titrant, its exact concentration in solution must be standardized.

To do this, it is first titrated with the primary standard, potassium hydrogen phthalate. A primary standard should be pure, stable, non-hygroscopic, and have a high molecular weight. Because the amount of hydronium ions contributed by the primary standard is known to a high degree of accuracy, it is used to determine the exact concentration of the hydroxide ions in the titrant. During an acid-base titration, the pH can be plotted as a function of the volume of the titrant added. The inflection point on the curve, the point at which there is a stoichiometric equal amount of acid and base in a solution, is called the equivalence point.

Most acids and bases are colorless, with no visible reaction occurring at the equivalence point. To observe when the equivalence point has been reached, a pH indicator is added. This is a pH sensitive dye that changes color in different pH environments. Its important to note that endpoint is not equal to the equivalence point, but indicates when a particular pH value has been reached. For example, phenolphthalein changes color around a pH of 8 and is commonly used as an indicator for acid-base titrations with an equivalence point around pH 7. While an accurate indicator for the titration is one that changes color as close to the equivalence point as possible, the titration curve has a steep slope around the equivalence point, leading to an acceptable level of error. At the equivalence point, the moles of base added are equal to the moles of acid initially present. An equation that utilizes the molarity and volume of each component can be used. With the other three values known, the acid concentration can be calculated. Now that you understand the principles behind the procedure, lets take a look at an actual protocol to determine the percent acetic acid in a commercial vinegar sample by reacting it with a standardized sodium hydroxide solution.

Typically, a rough estimate titration is performed to approximate where the endpoint will be. To begin, the titrant, sodium hydroxide, must be standardized. First, dissolve roughly 4 g of sodium hydroxide into 100 mL of deionized water. Make a 1:10 dilution by adding 25 mL of this stock sodium hydroxide solution to a glass container. Bring the total volume to 250 mL with deionized water and shake to mix. As sodium hydroxide absorbs carbon dioxide, it is important to use boiled, deionized water and an oven-dried bottle, and to cap the bottle quickly.

Calculate the approximate molar concentration of sodium hydroxide. Then, weigh out 5 g of the standard acid, potassium hydrogen phthalate, and place it in a drying oven. Once dried, allow the solid to cool to room temperature in a desiccator.

Weigh out 4 g of the dried potassium hydrogen phthalate to a high degree of precision, and dissolve in 250 mL of deionized water. Calculate the molar concentration of the potassium hydrogen phthalate solution.

Using a volumetric pipette, transfer 25 mL of the potassium hydrogen phthalate solution into a clean, dry Erlenmeyer flask. Add 2 drops of phenolphthalein pH indicator. Gently swirl the flask to mix. Flush a clean 50-mL burette with water and rinse at least three times with deionized water. Following this, rinse again with the diluted sodium hydroxide solution three times, making sure that the sodium hydroxide wets the entire inner surface. Mount the washed burette on a ringstand with a clamp and ensure that it stands vertically.

Fill the burette with the diluted sodium hydroxide solution. Air bubbles can affect the accuracy of volumetric readings. Gently tap the burette to free any air bubbles present, and open the stopcock to allow a few mL of titrant to flow through to release any trapped air. Read the volume of sodium hydroxide, at the bottom of the meniscus.

Place the flask containing potassium hydrogen phthalate under the burette. Add the titrant from the burette in 1–2 mL increments using one hand to control the flow rate by adjusting the stopcock, and the other swirling the flask.

When close to the endpoint, begin adding the titrant drop by drop. The endpoint is reached when the solution turns a faint, persistent pink color. Record the volume in the burette.

Repeat the titration at least two more times for consistent data and calculate the molar concentration of the diluted sodium hydroxide solution used as shown in the text protocol.

The sodium hydroxide solution is now standardized and can be used as a titrant to analyze vinegar. To reduce the pungent aroma, dilute 10 mL to a total volume of 100 mL.

Pipette 25 mL of the diluted vinegar into an Erlenmeyer flask, and add 2 drops of phenolphthalein. Fill the burette with the standardized sodium hydroxide solution and record the initial volume. Similar to the previous titration, slowly add the titrant to the analyte in the flask while swirling until the solution turns a light pink color, and record the final volume of sodium hydroxide used.

In this experiment, the titration was performed in triplicate and the mean volume of sodium hydroxide dispensed to neutralize the acetic acid in vinegar was calculated. The concentration and volume of base was used to elucidate the moles of acetic acid in the vinegar. The volume and molar mass were then used to calculate the concentration. It was determined that the vinegar had a molarity of 0.7388. Converting to percent, it was 4.23% acetic acid by volume.

Titrations are robust and easily customizable methods commonly applied in research, industry, and healthcare.

Scientists often use the measure of dissolved oxygen in freshwater bodies as an indicator of overall health that ecosystem. This is done by a redox titration. Unlike acid-base neutralizations, these titrations are based on a reduction-oxidation reaction between the analyte and the titrant. Dissolved oxygen in the water sample is reduced with chemicals in a reaction that results in the production of iodine. The amount of iodine produced and thus the level of dissolved oxygen can be determined by titration using a starch indicator. Glucose in urine can be indicative of a pathological condition like diabetes. A test to quantify urine glucose level, called Benedict's Method, is another example of the importance of titration; in this case, in healthcare. In this titrimetric procedure, sugars from urine are first reacted with an alkali resulting in the formation of enediols with powerful reducing properties. These reduce copper two ions in Benedict's reagent to copper one, in a colorimetric reaction that correlates with the initial concentration of glucose present in the urine sample.

You've just watched JoVE's introduction to titration. You should now be familiar with the principles behind this method, know how to perform an acid-base titration, and appreciate some of the ways it is being applied in research and industry.

As always, thanks for watching!

Table 1. Titration results.

Sample calculations:

Mass of KC 8 H 5 O 4 = 4.0754 g

Molar mass of KC 8 H 5 O 4 = 204.22 g/mol

According to Equation 2,

Moles of NaOH dispensed = concentration of NaOH × mean volume of NaOH dispensed = 0.09928 mol/L × 18.60 mL = 1.847 × 10 -3 mol

According to Equation 1,

Number of moles of CH 3 COOH in 25.00 mL of diluted vinegar = 1.847 × 10 -3 mol

Hence concentration of undiluted vinegar = 10 × 7.388 10 2 mol/L = 0.7388 mol/L

The above steps are presented to illustrate the calculation procedure; we can simply apply Equation 3 to obtain the concentration of undiluted vinegar in one step.

Therefore 1.000 L of undiluted vinegar contains 0.7388 mol of CH 3 COOH.

Applications and Summary

Titration is an important chemical method that is frequently applied in current chemistry research. For example, acid base titration is applied to determine amine or hydroxyl value of a sample. The amine value is defined as the number of milligrams of KOH equivalent to the amine content in one gram of sample. To determine the hydroxyl value, the analyte is first acetylated using acetic anhydride then titrated with KOH. The mass in milligrams of KOH then corresponds to hydroxyl groups in one gram of sample. 13 Another example is the Winkler test, a specific type of redox titration used to determine the concentration of dissolved oxygen in water for water quality studies. Dissolved oxygen is reduced using manganese(II) sulfate, which then reacts with potassium iodide to produce iodine. Since the iodine released is directly proportional to the oxygen content, the oxygen concentration is determined by titrating iodine with thiosulfate using a starch indicator. 14

Besides applications in basic chemical research, titration has also been widely adopted in industrial and everyday use. In biodiesel industry, waste vegetable oil (WVO) must first be neutralized to remove free fatty acids that would normally react to make undesired soap. A portion of WVO is titrated with a base to determine the sample acidity, so the rest of the batch could be properly neutralized. 15 Benedict's method, a test for quantification of urine glucose level, is another example showing the importance of titration in healthcare. In this titration, cupric ions are reduced to cuprous ions by glucose, which then react with potassium thiocyanate to form a white precipitate, indicating the endpoint. 16

• Llc, B. Titration: Ph Indicator, Thermometric Titration, Nonaqueous Titration, Equivalence Point, Acid-Base Titration, Amperometric Titration . General Books LLC, (2010).
• Jacobsen, J. J., Jetzer, K. H., Patani, N., Zweerink, G. & Zimmerman, J. Titration Techniques. J. Chem. Educ.  72 , 612, doi:10.1021/ed072p612 (1995).
• Harris, D. C. Quantitative Chemical Analysis . 8th edn, W. H. Freeman, (2010).
• Mattock, G., Taylor, G. R. & Paul, M. A. pH Measurement and Titration. J. Electrochem. Soc.  110 , 31C, doi:10.1149/1.2425702 (1963).
• De Levie, R. Aqueous Acid-base Equilibria and Titrations . Oxford University Press, (1999).
• Cannan, R. K. The Acid-Base Titration of Proteins. Chem. Rev.  30 , 395-412, doi:10.1021/cr60097a005 (1942).
• Michalowski, T. & Lesiak, A. Acid-Base Titration Curves in Disproportionating Redox Systems. J. Chem. Educ.  71 , 632, doi:10.1021/ed071p632 (1994).
• Waser, J. Acid-base Titration and Distribution Curves. J. Chem. Educ.  44 , 274, doi:10.1021/ed044p274 (1967).
• Gorbikova, E. A., Vuorilehto, K., Wikström, M. & Verkhovsky, M. I. Redox Titration of All Electron Carriers of Cytochrome c Oxidase by Fourier Transform Infrared Spectroscopy. Biochemistry  45 , 5641-5649, doi:10.1021/bi060257v (2006).
• Silverstein, T., Cheng, L. & Allen, J. F. Redox Titration of Multiple Protein Phosphorylations in Pea Chloroplast Thylakoids. Biochim. Biophys. Acta (BBA)-Bioenerg.  1183 , 215-220, doi:10.1016/0005-2728(93)90022-8 (1993).
• Lenghor, N., Jakmunee, J., Vilen, M., Sara, R., Christian, G. D. & Grudpan, K. Sequential Injection Redox or Acid-Base Titration for Determination of Ascorbic Acid or Acetic Acid. Talanta  58 , 1139-1144, doi:10.1016/S0039-9140(02)00444-7 (2002).
• Mitchell, P., Moyle, J. & Smith, L. Bromthymol Blue as a pH Indicator in Mitochondrial Suspensions. Eur. J. Biochem.  4 , 9-19, doi:10.1111/j.1432-1033.1968.tb00166.x (1968).
• Perkins, E. G. Analyses of Fats, Oils and Derivatives . AOCS press, (1993).
• Spellman, F. R. Handbook of Water and Wastewater Treatment Plant Operations . 2 edn, CRC Press, (2009).
• Purcella, G. Do It Yourself Guide to Biodiesel: Your Alternative Fuel Solution for Saving Money, Reducing Oil Dependency, Helping the Planet . Ulysses Press, (2007).
• Nigam. Lab Manual Of Biochemistry . Tata McGraw-Hill Education (2007).

Place the flask containing potassium hydrogen phthalate under the burette. Add the titrant from the burette in 1–2 mL increments using one hand to control the flow rate by adjusting the stopcock, and the other swirling the flask.

Scientists often use the measure of dissolved oxygen in freshwater bodies as an indicator of overall health that ecosystem. This is done by a redox titration. Unlike acid-base neutralizations, these titrations are based on a reduction-oxidation reaction between the analyte and the titrant. Dissolved oxygen in the water sample is reduced with chemicals in a reaction that results in the production of iodine. The amount of iodine produced and thus the level of dissolved oxygen can be determined by titration using a starch indicator. Glucose in urine can be indicative of a pathological condition like diabetes. A test to quantify urine glucose level, called Benedict’s Method, is another example of the importance of titration; in this case, in healthcare. In this titrimetric procedure, sugars from urine are first reacted with an alkali resulting in the formation of enediols with powerful reducing properties. These reduce copper two ions in Benedict’s reagent to copper one, in a colorimetric reaction that correlates with the initial concentration of glucose present in the urine sample.

You’ve just watched JoVE’s introduction to titration. You should now be familiar with the principles behind this method, know how to perform an acid-base titration, and appreciate some of the ways it is being applied in research and industry.

Titration Curves & Equivalence Point Calculations

Core Concepts

In this tutorial, you will learn about titration curves, titration analysis and the steps required to perform one. You will also learn about equivalence points and endpoints, and titration calculations.

Topics Covered In Other Articles

• Acid and Base Chemistry
• Properties of Acids and Bases
• Strong Acids and Bases
• Lab Safety and Tips

Introduction to Titration

What is the definition of titration.

Titration is a laboratory technique used to determine the concentration of a solution by reacting it with a known volume and concentration of another solution. This is done by carefully measuring out a specific volume of the first solution, called the titrant, and adding it to the second solution, called the analyte, until the reaction is complete. The point at which the reaction is complete is called the endpoint, and it can be detected using a variety of methods, such as a color change or a change in electrical conductivity. The concentration of the analyte solution can then be calculated based on the volume and concentration of the titrant that was used.

Titration is a useful tool for determining the concentration of a wide range of substances, including acids, bases, and other chemicals. Scientists employ titration to assess the quantity of a constituent, or analyte, of a given sample. They accomplish this by adding a known quantity of another substance, called a standard solution, to the sample.

To learn how to create a known concentration of a solution, check out this article on molarity . This standard solution contains the titrating reagent, or titrant. When titrant is added, it reacts with the analyte in a known proportion (according to the reaction equation) with the titrant, revealing its concentration and amount.

Types of Titrations

1. acid-base titrations.

This kind of titration determines the concentration of an acid or the concentration of a base. The former investigation uses a basic standard solution to neutralize the sample, while the latter uses an acidic standard solution. You can also determine the strength of the acid or base by viewing the shape of the titration curve.

2. Redox (Oxidation-Reduction) Titrations

This type of titration involves a redox chemical reaction between the analyte and the titrant, during which these substances transfer electrons. 

3. Precipitation Titrations

This titration variety uses the formation of a precipitate to signal the completion of the reaction between the analyte and the titrant. 

4. Complexometric Titrations

During this type of titration, a larger, undissociated complex forms. This generates less room for interpretation errors than normal precipitation titrations. 

5. Thermometric Titrations

This type measures the temperature of the solution to determine when the reaction is complete. It relies on the heat generated or absorbed during the reaction of the analyte and titrant.

Equivalence Point versus Endpoint 

But how do scientists determine the point at which chemically equivalent quantities of analyte and titrant exist in the solution, so that they can quantify the analyte? They look for an “equivalence point,” the point at which enough titrant has combined with the analyte to neutralize it. At this specific point, the amount of titrant in the system reveals the amount of analyte in the system; the moles of both species equal one another. 

However, this equivalence point differs from the “endpoint” of a titration analysis. The endpoint indicates the end of the reaction; it denotes the amount of reactant titrant needed to facilitate a complete chemical reaction with the reactant analyte. A color change in the system of interest signals that it has reached this endpoint. Materials called “indicators,” which undergo these color changes, can be added to the system to designate the endpoint. 

Different indicators exhibit different endpoints due to their varying chemical compositions. Scientists tend to choose indicators whose endpoints roughly equal their equivalence points. When this occurs, the color transition denotes both the endpoint and the equivalence point, revealing the amount of titrant needed to equal the amount of analyte and thus the quantity of analyte in the system. Some substances, such as polyprotic acids, possess multiple equivalence points, but for a given indicator there is generally only one endpoint.

Below is a chart summarizing the characteristics differentiating equivalence points and endpoints:

Titration Curves 

Titration curves provide graphical depictions of titration analyses. Scientists first measure a property of the analyte solution (usually pH for acid-base titrations). Then, they plot this property against the volume of the titrant added. 

Using a titration curve, scientists can often identify an endpoint closer to the equivalence point, compared with using an indicator to find the endpoint.

In addition to pH, several other properties can reveal the endpoint of a reaction. These include the absorptivity, temperature, conductivity, or electrical current generated. Some of these are only applicable to certain types of reactions.

Acid-Base Titration Curves

Acid-base titration, the most frequently employed method of titration, adheres to certain curve trends established by the acid and the base involved. The pH of the solution at the equivalence point depends on the characteristics of the analyte and the titrant. The following combinations produce: 

Below is an Acid-Base Titration curve used to quantify the oxalic acid analyte using NaOH as a titrant. The curve consists of a plot displaying pH. The equivalence point of a system occurs during a vertical increase in its pH values, while the endpoint of a visual indicator tends to occur during the last slight increase in pH value. The endpoint shown below is based on the indicator, and cannot be determined from the curve itself.

How to Perform a Simple Titration

The following steps describe the methods involved in a titration analysis using a visual indicator:

• Fill your burette with titrant solution, using the graduations on the instrument to precisely gauge the amount added. 
• Clamp the burette to your burette stand. Make sure the stand clamps adhere tightly to the burette to minimize movement during the analysis.   
• Fill a beaker with your analyte solution, again ensuring precise measurements using the graduations on the glassware . 
• Using a pipette or dropper, add a few drops of indicator to the analyte solution. The type and amount of indicator required will depend on your titrant and analyte. 
• Next, agitate your analyte solution by placing it on a magnetic stir plate. This thoroughly mixes the components. 
• Place the beaker, containing the analyte, directly under the burette, containing the titrant. Gently turn the stopcock to allow the titrant to drip from the burette to the beaker. 
• Continue adding titrant solution to analyte solution until you observe a color change, indicating the endpoint of the solution. If the indicator’s endpoint equals its equivalence point, you can infer that the amount of analyte in the system equals the amount of titrant at this point. 
• Determine how much titrant you added to the analyte by subtracting the final volume in the burette from the starting volume.
• Quantify the amount of analyte in the system using the amount of titrant added; these amounts should be equivalent. 

Titration Calculations

Automatic Titration Apparatus

Although the example discussed above is a manual titration, there are also many automatic titrators available today. These can often inject far more precise amounts of titrant than a standard burette would, and they can reduce waste by using extremely small quantities of reagents. Below is an image of a standard automatic titrator, which shows the physical apparatus as well as the software interface displaying a titration curve in the lower left corner.

Applications of Titration in Today’s World 

This analytical technique has many applications, ranging from medical, to cosmetic, and even to environmental science.

In the food industry, workers use titration and equivalence points to determine the quantity of different salts, sugars, vitamins, and fatty acids in foods. Similarly, other major industries like cosmetics and cleaning use titration to add appropriate and safe concentrations of chemicals to their products. Pharmaceutical companies rely on titration during the medication developing process. Even medical labs use this type of analysis to study blood and urine samples from patients.

In addition to industry, there are many uses for titration in academic research. For one, environmental chemists conduct titrations using rainwater, melted snow, and other liquids to assess their composition. For analysis of water content in an analyte, there is a special technique known as Karl-Fischer titration. In these and many other fields, titration also serves to determine the exact concentration of reagents, a process known as standardization.

As you can see, titration has an astounding range of uses and applications. As a result, it is one of the most frequently conducted methods of chemical analysis worldwide, both in the classroom and in the workplace.  

Titration Calculation Practice Problems

Acetic acid and NaOH react according to the following equation:

H 2 O 2 and KMnO 4 react according to the following equation:

Titration Calculation Practice Problem Solutions

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Lesson Explainer: Titration Experiments Chemistry • Third Year of Secondary School

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In this explainer, we will learn how to describe acid–base titration methods and their use in determining acid and base concentrations.

When we want to determine the concentration of an acid or a base, we can perform a titration experiment. In a titration experiment, a solution with a known concentration is added to an exact volume of a solution with an unknown concentration in the presence of an indicator.

The apparatus necessary for the experiment is shown below.

A buret is used in this experiment because we are unsure of exactly what volume of a solution with a known concentration will be necessary. The stopcock or faucet at the bottom of the buret allows us to easily add a little amount of solution at a time. An Erlenmeyer flask, also known as a conical flask, is used because the flask can easily be swirled without spilling. Notice that the buret and Erlenmeyer flask are placed over the base of the retort stand to prevent the apparatus from toppling over. The white tile is optional but can help us to more easily identify subtle color changes.

A standard solution, also called a titrant, will be placed in the buret. A standard solution is a solution that has a known concentration.

Definition: Standard Solution (Titrant)

It is a solution with a known concentration.

If the solution that we are trying to determine the concentration of is an acid, the standard solution should be a base and vice versa. Standard solutions can be purchased but are often easily made in the laboratory.

How To: Making a Standard Solution from a Solid Solute

• Use a small amount of water to rinse any remaining solute on the weigh boat or weight paper into the flask.
• Stopper the flask and invert several times to mix thoroughly.

How To: Making a Standard Solution from a Liquid Solute

Titration experiments are often performed four or more times, so enough standard solution must be prepared to account for all of the trials.

Example 1: Recalling the Meaning of the Term Standard Solution

When undertaking a titration, the solution in the buret is usually a standard solution of acid or base. What does the term standard solution refer to?

A titration experiment is performed when we wish to determine the concentration of an acid or a base. During a titration, a solution with a known concentration is added to an exact volume of a solution with an unknown concentration in the presence of an indicator. The solution with an accurately known concentration is called a standard solution. The term standard solution refers to an acid or base of a known concentration.

When adding a standard solution to the buret, it is important to ensure that the stopcock at the bottom of the buret is in the closed position, perpendicular to the column.

A buret should always be filled at eye level to prevent spilling corrosive acids or bases into your face or eyes. A funnel can be placed into the top of the buret to further help prevent spillage when filling. To fill the buret, we slowly add titrant, filling to near the 0 mL mark. Then, we remove the funnel. It is not necessary to begin a titration at exactly 0.00 mL .

We should then check the buret for air bubbles, gently tapping to remove them. Next, we place a waste container under the buret and open the stopcock to allow a few milliliters of liquid to drain. We should use a wash bottle to rinse the tip of the buret with deionized water. We can then record the starting volume of titrant. We must remember to read the volume from the bottom of the meniscus.

Example 2: Understanding the Risks Associated with Filling a Buret

Why should a buret always be filled at eye level and never above?

• To reduce the risk of splashing acid or base onto the eyes or face
• To know when the beaker or measuring cylinder is empty
• To make sure the solution is poured into the buret and not on the floor
• To watch the solution move down the buret
• To allow the buret to be filled while sitting down

A buret is frequently used when performing an acid–base titration and is therefore often filled with an acid or base. Acids and bases are corrosive. They can irritate or burn the eyes and skin and can cause respiratory distress. To minimize the potential for spilling such substances on the eyes or face, a buret should never be filled above eye level. The correct answer is choice A.

The solution of unknown concentration is called the titrand or analyte.

Definition: Titrand (Analyte)

It is a solution with an unknown concentration.

An exact volume of titrand should be added to an Erlenmeyer flask using a volumetric pipet. The volume of titrand used is sometimes called an aliquot.

Definition: Aliquot

It is a known volume of solution.

Example 3: Choosing Which Solution to Place into an Erlenmeyer Flask and Buret in an Acid–Base Titration

A student wants to use titration to determine how much acid is needed to neutralize a known volume of base. They set up the experiment as shown. At the start of the experiment, which solution should go into the Erlenmeyer flask and which should be used to fill the buret?

The Erlenmeyer flask is filled using a pipet. A pipet is a piece of glassware used to deliver a specific volume of liquid. A buret is a piece of glassware that is filled with a liquid that can then be dispensed. Burets are used when the volume of liquid necessary is not known as the stopcock at the bottom of the buret can easily be opened or closed to control the volume of the liquid delivered.

In this titration experiment, a student is using a known volume of base but does not know the total volume of acid needed. Therefore, the base should go into the Erlenmeyer flask and the acid into the buret.

During a titration experiment, the titrant will be added to the titrand until the acid and base have completely neutralized one another. This is called the equivalence point.

Definition: Equivalence Point

It is the point at which an acid and base have completely neutralized one another.

The approximate volume of titrant needed to reach the equivalence point can be visually determined by adding an acid–base indicator to the Erlenmeyer flask. An indicator is a weak acid or base that undergoes a color change over a specific pH range.

Definition: Indicator

It is a weak acid or base that undergoes a color change over a specific pH range.

There are a variety of indicators that can be chosen for a titration experiment depending on the acid and base used as the titrant and titrand. Near the equivalence point, addition of a single drop of titrant will cause the pH of the solution in the Erlenmeyer flask to change drastically. A good acid–base indicator will exhibit an abrupt color change during this point in the experiment.

The equivalence point of the reaction between a strong acid and strong base will occur at a pH of 7. However, a drastic pH change with the addition of single drops of titrant will occur between a pH of approximately 3.5 and 10.5. Therefore, we should choose an indicator that changes color within this range.

There are two indicators that are commonly used for strong acid–strong base titrations. They are phenolphthalein and methyl orange. We commonly say that phenolphthalein is colorless in acid and pink in base; although technically, this color change occurs between a pH of 8.2 and 10. We commonly say that methyl orange is red in acid and yellow in base; although technically, this color change occurs between a pH of 3.2 and 4.4.

A universal indicator is commonly found in the chemistry laboratory and is often used to test a solution to determine an approximate pH.

However, the universal indicator is not used in titrations because it changes color over a wide range of pH values and the color changes are less abrupt than other indicators.

Example 4: Determining the Color of an Indicator at One pH Value and Which Indicator to Use When the pH Changes from 8 to 11

The table below shows the color range for several different indicators.

• Which indicator is blue at a pH of 5?
• Which indicator would be best to show that the pH of a solution has changed from 8 to 11?

The graph shows the pH range in which each indicator changes color. For example, bromophenol blue changes color from yellow to blue between a pH of 3 and 4.6. Below a pH of 3, bromophenol blue will be yellow and above a pH of 4.6 bromophenol blue will be blue.

At a pH of 5, bromophenol blue will be blue, bromothymol blue will be yellow, cresolphthalein will be colorless, and alizarine yellow will be yellow. The indicator that is blue at a pH of 5 is bromophenol blue.

The indicator that would be best to show that the pH of a solution has changed from 8 to 11 will be an indicator that undergoes a color change in this pH range. Bromophenol blue and bromothymol blue will both remain blue as the pH changes from 8 to 11. Alizarine yellow will be yellow between a pH of 8 and 10 and will begin to turn orange as the pH gets closer to 11. There is color change; however, it occurs at too high of a pH. Cresolphthalein is colorless at a pH of 8 and purple at a pH of 10. As cresolphthalein undergoes a drastic color change in the desired pH range, it would be the best indicator to use for this experiment.

Let us consider a titration experiment where a base titrant, acid titrand, and phenolphthalein indicator are used. At the start of the experiment, phenolphthalein and the acid titrand will be in the Erlenmeyer flask. As phenolphthalein is colorless in acid, the flask will appear colorless.

When the base titrant is added, a bright-pink spot will appear in the Erlenmeyer flask as the phenolphthalein reacts with the base and turns pink. However, the color will quickly disappear as the acid titrand in the flask reacts with the base, neutralizing it. It is important to continue to swirl the flask, to allow the acid and base to react completely during this process.

As more base is added to the flask, the pink color will remain for longer periods of time. This is because as the experiment progresses, there is less acid in the solution and the acid–base neutralization reaction takes longer to occur. Once all of the acid has been neutralized, any additional base added will remain in the solution and can react with the phenolphthalein. When adding a drop of base causes the solution to turn faintly pink, the end point of the experiment has been reached.

Definition: End Point

It is the point in a titration experiment where the indicator has just changed color without reverting.

It is important to recognize that the end point of the experiment and the equivalence point are often not the same. The equivalence point occurs when all of the acid and base have been neutralized, but the end point occurs when the indicator has changed color.

The final volume of titrant in the buret should be recorded when the end point has been reached. If the Erlenmeyer flask appears a dark-pink color, too much base has been added and the experiment will need to be repeated.

The table below shows the desired end point for the phenolphthalein indicator and methyl orange indicator.

Example 5: Recalling the Colors of Methyl Orange

Fill in the blanks: Methyl orange is a useful indicator, particularly for titrations against acids. It is in strongly acidic solutions and in basic solutions.

Methyl orange is an indicator that is red in a solution with a pH less than 3.1 and yellow in a solution with a pH greater than 4.4. Solutions are acidic when the pH is less than 7, neutral when the pH is equal to 7, and basic when the pH is greater than 7.

This means that methyl orange will appear yellow in all basic solutions and red in strongly acidic solutions. We should fill in the first blank with the word red and the second blank with the word yellow.

It is very easy to add too much titrant during a titration. We can get an estimate of the amount of titrant needed to reach the end point by first performing a rough titration.

How To: Performing a Rough Titration

• Properly fill a buret with titrant.
• Record the initial volume of titrant in the buret.
• Use a pipet to transfer an exact volume of titrand to an Erlenmeyer flask.
• Add a few drops of the desired indicator to the Erlenmeyer flask and swirl.
• Place the Erlenmeyer flask under the tip of the buret.
• Open the stopcock of the buret to allow the titrant to quickly enter the flask, and swirl the flask continuously.
• When the Erlenmeyer flask has undergone a permanent color change, close the stopcock.
• Record the final volume of the titrant in the buret.
• Subtract the initial volume of titrant from the final volume of titrant to determine the volume of titrant used in the titration.

Subtracting five milliliters from the volume of titrant used in a rough titration tells us how much titrant we can safely add quickly during a good titration. Never use the rough titration titrant volume in any titration calculations.

How To: Performing a Good Titration

• Open the stopcock of the buret, add five milliliters less of titrant than was used during a rough titration. Close the stopcock.
• Open the stopcock partway, allow the slow addition of titrant to the flask, swirling continuously.
• When flashes of a different color (change in the indicator color) begin to appear in the flask, adjust the stopcock to allow only one drop to flow at a time.
• Continue adding one drop of titrant at a time and swirling until the indicator color change takes longer to fade. At this point, close the stopcock.
• Rinse the tip of the buret with deionized water into the flask and swirl.
• If the color change remains, the end point has been reached and the final volume should be recorded.
• If the color change does not remain, repeat adding a single drop of titrant, rinsing the tip of the buret, and swirling until the color change persists.
• Record the final volume of titrant in the buret.
• Repeat the entire titration two or three more times and average the volumes of the titrant used.

There are several common errors that can occur when performing a titration experiment. Measurement errors include inaccurate measurement of the titrand or misreading the initial or final buret volumes. Air bubbles in the buret will affect the volume readings as well. We can remove air bubbles by gently tapping the buret. In addition, if extra unreacted titrant is on the buret tip or in the neck of the flask, we can record a titrant volume that is too large. We can eliminate this error by rinsing the tip of the buret and the neck of the flask with deionized water when we are close to the end point of the titration.

If the indicator never changes color during the titration, the wrong indicator may have been selected or the same solution may have been placed in both the buret and Erlenmeyer flask. If the Erlenmeyer flask remains colorless during a titration involving phenolphthalein, double-check that the indicator was added to the flask.

• An acid–base titration experiment is used to determine the concentration of a solution.
• A standard solution or titrant is placed in the buret.
• An aliquot of titrand or analyte is placed in the Erlenmeyer flask.
• Indicators are added to the flask to signify the end point.
• A phenolphthalein indicator is colorless in an acid and pink in a base, while a methyl orange indicator is red in a strong acid and yellow in a base.
• When nearing the end point of a titration, a titrant should be added dropwise just until the indicator has changed color.

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Classic acid and base titration in pink

Susan Schmidt, Somerville High School, Somerville, Massachusetts

Department of Chemistry 200 University Ave. W Waterloo, Ontario, Canada N2L 3G1

[email protected]

What Is Titration?

What is the principle of titration, how is titration carried out, titration formula: how is concentration of the analyte calculated, types of titration, titration utility: practical applications and examples.

Titration is a technique of determining the concentration of an unknown solution by using a solution of a known concentration. The known solution (titrate) is added in drops to the analyte (unknown solution) until the endpoint is reached.

Is lemonade your favorite drink? Or do you occasionally add it to your diet for weight loss? Whatever the reason may be, do you know what concentration of lemon juice you’re consuming in the form of lemonade? Aren’t you worried about causing damage to your own stomach if the concentration exceeds the limit?

First of all…Relax! You don’t have to worry about the concentration of lemon juice in your lemonade.

However, when it comes to certain drugs, food preservatives, the pH level of your fish tank, and in various industrial procedures, the concentration of different constituents in solutions does play a major role.

This is where titration comes into play! Titration determines the concentration of a solution by following certain strict procedures. This article explains the principle and methodology behind titration.

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Titration is a quantitative analysis to determine the concentration of an unknown solution by adding a solution of known concentration in a drop at a time. By this process, the acid or base of a known concentration completely neutralizes the acid or base of the unknown concentration. The actual reaction that takes place during neutralization is between a hydrogen ion from acid and hydroxide ion from base that combine to form water. It is important to maintain a uniform pH throughout the process. Hence analyte and titrate should be of equal concentration. Involving strong titrate and a diluted analyte or vice versa will affect the physical of the analyte. The process is generally monitored by pH electrodes or indicators.

The solution of known concentration is the titrate and the solution whose concentration is to be determined is the analyte. The equivalence point of this process is obtained when the titrate completely neutralizes the base or acid in the analyte.

Also Read: Buffer Capacity: Definition And Method Of Calculation

The analyte is measured and filled in a beaker. A few drops of an acid base indicator, such as methyl orange or phenolphthalein, is added to the analyte.

Phenolphthalein remains colorless in acid, but turns pink in the presence of a base. Similarly, methyl orange is a reddish-orange powder that turns to deep red in an acid and pales out to yellow in a base.

A standard solution is taken in a burette to start the process. The standard solution is allowed to drip slowly into the beaker containing the analyte. The process is carried out until the color of the analyte changes, indicating the arrival of an end point. This means that all of the base or acid in the analyte has been completely neutralized by the titrate.

The volume change of the standard solution at which the end point occurs is noted. This volume indicates the amount of  titrate that has been used to neutralize the analyte. These values are then used in further calculations.

Also Read: Why Are pH Values Only In A Range Of 0-14?

Like every other quantitative analysis, titration has a standard formula to determine the unknown concentration. The generally used formula is:

Volume of titrate * molarity of titrate = Volume of analyte * molarity of analyte

Let’s consider that 45.6ml of 1.25 M sodium hydroxide has been used to neutralize 20ml of hydrochloric acid. The concentration of HCl is then given by,

45.6  X  1.25 = 20.0 * molarity of analyte.

Solving the above equation will yield the concentration of the analyte.

Based on the type of reaction involved, titraion is classified into 4 different types.

• Acid-base titration: This is based on the neutralization between an acid or base and the analyte.

• Complexometric titration: This is widely used to determine the number of metal ions involved in the reaction. EDTA is used as an indicator. At the end of this titration, an undissociated complex is formed that marks the equivalence point. For example:

Ag +  + 2CN –  → [Ag(CN) 2 ] –

• Redox titration: In this type, the transfer of electrons takes place in the reacting ions of the solution. It is further classified into 3 types based on the type of reagent used.

Permanganate:    MnO 4 – + 8H + 5e → Mn 2+ + 4H 2

Dichromate­:       K 2 Cr 2 O 7 + 4H 2 SO 4  → K 2 Cr 2 (SO 4 ) + 4H 2 O + 3[O]

Iodimetric and iodometric:    l 2  + 2e → 2l –

• Precipitation titration: This reaction forms an insoluble precipitate when the two reacting substance are brought into contact.

AgNO 3  + NaCl → AgCl + NaNO 3

Titration is presumably applied in one way or another in our day-to-day activities. Various clinical tests, such as blood tests and urine tests, use titration to determine the concentration of chemicals of interest. It is also used in the food industry to determine the amount of certain chemicals in food.

Often, it is used to determine carbohydrate, fat and vitamin content. Titration is also widely applied in the medical field. Pregnancy tests, blood glucose level tests and other urinalysis applications use titration. It is also included as a part of academics for high school students to analyze their practical aptitude.

Sea water properties, such as the concentration of ammonia, nitrates and pH level are calculated and the environment is modified accordingly to maintain optimum conditions.

Concentration—commonly expressed in terms of molarity (number of moles of solute per liter in solution)—is of great importance in chemistry because it determines the rates of reaction and the conditions at equilibrium. Titration is the most prominent and widely used method to determine unknown concentration.

• What is a Titration?. Purdue University
• ChemTeacher - chemteacher.chemeddl.org

Janani Anand is pursuing a Bachelor of Engineering in Biomedical Engineering at Rajalakshmi Engineering College, Chennai. She is a voracious reader and a creative writer. She is also involved in social activities, enjoys watching cricket and a great fan of MS Dhoni.

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AP®︎/College Chemistry

Course: ap®︎/college chemistry   >   unit 4.

• Acid–base titrations

Worked example: Determining solute concentration by acid–base titration

• Redox titrations
• Introduction to titration

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Video transcript