experiment standardization of sodium hydroxide solution

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Standardizing a solution of sodium hydroxide.

Experiment #6 from Advanced Chemistry with Vernier

Introduction

It is often necessary to test a solution of unknown concentration with a solution of a known, precise concentration. The process of determining the unknown’s concentration is called standardization .

Solutions of sodium hydroxide are virtually impossible to prepare to a precise molar concentration because the substance is hygroscopic. In fact, solid NaOH absorbs so much moisture from the air that a measured sample of the compound is never 100% NaOH. On the other hand, the acid salt potassium hydrogen phthalate, KHC 8 H 4 O 4 , can be measured out in precise mass amounts. It reacts with NaOH in a simple 1:1 stoichiometric ratio, thus making it an ideal substance to use to standardize a solution of NaOH.

In this experiment, you will

• Prepare an aqueous solution of sodium hydroxide to a target molar concentration.
• Determine the concentration of your NaOH solution by titrating it with a solution of potassium hydrogen phthalate, abbreviated KHP, with an exact molar concentration.

Sensors and Equipment

This experiment features the following sensors and equipment. Additional equipment may be required.

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This experiment is #6 of Advanced Chemistry with Vernier . The experiment in the book includes student instructions as well as instructor information for set up, helpful hints, and sample graphs and data.

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6: Titration of an Unknown Acid

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• To perform an analytical titration.
• To standardize a basic solution.
• To determine the molecular mass of an unknown acid.

Chemical reactions between acids and bases are important processes. These reactions can be used to synthesize new substances or to analyze the quantity of a pure substance or of a compound in a mixture. In this experiment, you will first determine the concentration of a base, NaOH, and then use this standardized base to determine the molar mass of an unknown weak acid. Based on the molar mass you will determine the identity of the unknown weak acid.

In water strong acids produce hydronium ions and strong bases produce hydroxide ions. The reaction of a strong acid with a strong base is represented with the chemical reaction shown in Equation 1. Chemists regard this as a reaction that goes to completion; that is all of the reactants are converted to products.

\[\ce{H_{3}O^{+}} + \ce{OH^{-}} \rightarrow 2\ce{H_{2}^O} \label{1}\]

The reaction between weak acids and strong bases is represented by Equation 2.

\[\ce{HA} + \ce{OH^{-}} \rightarrow \ce{H_{2}O} + \ce{A^{-}} \label{2}\] 

This reaction also goes to completion. In this experiment you will use the reaction in Equation 3 to determine the molar concentration of the strong base NaOH using the weak, monoprotic acid potassium hydrogen phthalate (KHP), where P - is the phthalate group.

\[\ce{HP^{-}} + \ce{OH^{-}} \rightarrow \ce{P^{2-}} +\ce{H_{2}O} \label{3}\]

The method of analysis will be a titration of acid with base. In the first part of the experiment, you will standardize (determine the exact concentration of) your sodium hydroxide solution. We will determine the concentration of the solution by titrating a known mass of a known acid with your sodium hydroxide solution, using an acid-base indicator to find the endpoint of the titration. In the second part of the experiment, we will determine the identity of an acid by titrating a known mass with out standardized NaOH solution and determining its molecular weight.

Pre Lab Video

Safety and Waste Disposal

• The solutions can be disposed of down the drain. 

Part A: Standardization of a Sodium Hydroxide Solution

In this part of the experiment you will prepare and standardize a sodium hydroxide solution.

Step 1

Record the initial volume to two decimal places. (It need not be exactly 0.00.)

Record the final volume of NaOH in your buret to two decimal places. The difference between the final measurement and the initial is the volume of NaOH used.

Determine the concentration of NaOH as shown in the Calculations section.

Calculate the molarity of the NaOH solution to 4 significant figures.

Repeat the titration with fresh samples of KHP until you have two concentrations that agree within 1.5 %.

Part B: Determining the Molecular Mass of an Unknown Acid

In this part of the experiment you will use your standardized sodium hydroxide solution to titrate an unknown acid and determine its molecular mass.

Calculate the molecular weight of the unknown acid. (see Calculations).

Repeat the determination.

Weigh out two more samples of your unknown, increasing or decreasing the mass to have the volume of NaOH used be around 25mL.

Titrate these samples and record your results and any observations.

Calculate the molecular weight of the acid. Your results should agree to within 2%.

Determine the identity of your unknown based on your average molecular mass. 

Part A – Standardization of a Sodium Hydroxide Solution

Calculate the moles of \(\ce{KHP}\).

Calculate the volume of \(\ce{NaOH}\).

Calculate the molarity of the \(\ce{NaOH}\) solution to 4 significant figures.

Determine if your result of the standardization value for \(\ce{NaOH}\) is after two titrations, is within the tolerance. Calculate the percent difference of two  titrations. 

\[ \%\: difference = \frac{|M_{1} - M_{2}|}{M_{2}} \times 100 \label{4}\]

*If necessary

• Average molarity of \(\ce{NaOH}\) solution: ___________________ M

In the space below, clearly show all calculations for your Trial 1 data only:

Part B – Determining the Molecular Mass of an Unknown Acid

Unknown Number of Solid Acid: ________________

Calculate the moles of base used to reach the endpoint

These are monoprotic weak acids so the moles of base equal the moles of acid present at the equivalence point.

Using the mass of the unidentified acid you measured and the moles of acid you calculated, determine the MW of the unidentified acid. 

\[MW=\frac{grams\: in\: sample}{moles\: in\: sample} = \frac{g}{mol} \label{5}\]

*Only three trials are required, but space for a fourth is given if needed. Average equivalent mass of unknown acid: ___________________ \(g·eq^{–1}\) In the space below, clearly show all calculations for your Trial 1 data only:

Unknown Mass Identification:                                      

Unknown Possibilities:

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Titration Lab: NaOH with Standardized solution of KHP

• Titration Lab: NaOH with Standardized…

To standardize a sodium hydroxide (NaOH) solution against a primary standard acid [Potassium Hydrogen Phthalate (KHP)] using phenolphthalein as an indicator.

• Two 100 cm 3 Beakers (One for making the KHP solution, one for pouring NaOH into the burette)
• 1 Digital Balance (up to 2 decimal places accuracy)
• 1 Stirring rod
• 100 cm 3 volumetric flask
• 100 ml distilled water
• Conical flask
• Phenolphthalein
• 100 cm 3 burette
• Sodium Hydroxide

Data Processing

Qualitative data.

Potassium Hydrogen Phthalate ( referred to in the experiment as KHP) was a brittle, white, crystalline substance. The crystals required intense stirring before they could be dissolved in water. The resultant Acidic solution was transparent, with a small amount of undissolved granules of KHP. As the transparent NaOH solution came into contact with transparent phenolphthalein in the KHP solution, it turned pink which on shaking became transparent. However, as NaOH was added further, there came a point when no amount of stirring changed the pink colour.

Uncertainties

• % Uncertainty of KHP Mass = (0.01/m KHP ) x 100
• % Uncertainty of (aq) KHP in Volumetric Flask = (0.1/100) x 100
• % Uncertainty of (aq) KHP in Pipette = (0.1/10) x 100
• % Uncertainty of NaOH Volume = (0.1/ V NaOH ) x 100

Stoichiometric Calculations

Moles (n vf ) of KHP in volumetric flask = m KHP /M KHP where M KHP is the Molar Mass of KHP (204.22 g)

[c] KHP in 100 cm 3 volumetric flask = n/V

Moles of KHP in 10 cm 3 of solution in where V is a given volume of water

conical flask (n cf ) = [c] KHP x V

General Formulae

n = m/M mol

[c] 1 x V 1 = [c] 2 x V 2

The volume of NaOH added = Final Volume – Initial Volume

n is the number of moles of KHP

m is the mass of KHP in grams

M is the Molar mass in grams

V is the volume in cm 3

[c]KHP = (n/V) mol dm -3 = (0.00974/0.1) mol dm -3 = 0.0974 mol dm -3

Where [c]KHP is the concentration of KHP Acid .

Sample Calculations

Number of moles of KHP in 2.00 grams = (m/M) = (2/204.22) mol = 0.00979 mol

[c] KHP = n/V = (0.00979/0.1) mol dm -3

Number of moles of KHP in 0.01 dm 3 of solution in conical flask = [c] x V

= 0.0979 x 0.01 = 9.79 x 10 -4 mol.

KHC 8 H 4 O 4 + NaOH → H 2 O + NaKC 8 H 4 O 4

Mole ratio = 1 KHP:1NaOH

From the mole ratio, the number of moles of NaOH = 0.00979 mol.

[c] NaOH = n/V = (0.00979/0.0950) = 0.103 mol dm -3 (cm 3 is converted into dm 3 )

Lab 1: Preparation of KHP Acid

Weight of weighing boat before adding KHP = 2.67 g

Weight of weighing boat with KHP = 4.67 g

Weight of weighing boat after transfer = 2.68 g

Molar Mass (M) of KHP = 204.22 g

Volume (V) of water = 100 cm3 = 0.1 dm3

Mass of KHP Transfer = Weight of weighing boat with KHP – Weight of weighing boat after transfer

=    (4.67 − 2.68) g = 1.99 g.

Lab 2: Standardisation of NaOH solution

*Initial volume is the initial reading of the burette and final volume is the reading after adding NaOH solution

Calculations

Mole ratio = 1 KHP: 1NaOH

From mole ratio, number of moles of NaOH = 0.00974 mol

Uncertainty Calculations

Data Analysis and Conclusion

One experimental flaw which resulted in readings inconsistent with the literature value was due to human error. This flaw was due to allowing excess sodium hydroxide to flow, causing the KHP solution to become pinker than it should have. This might have caused some deviations because the volume of sodium hydroxide added was excess. The theoretical value of NaOH to be poured was 9.50 cm 3 , and more or less than 0.1 cm 3 of that value. However, the amount I added on an average was 10.4 cm 3 , which suggests why the solution became unusually dark pink as supposed to light pink.

The theoretical value of the Sodium Hydroxide that was expected to be used was 9.50 cm 3 . However, there has been a deviation of 0.9 cm 3 , which is significant, but not high. The resulting percentage error out of this deviation is:

There is almost a 1% deviation. Taking the value of 9.50 cm 3 and mass of 2 grams, the concentration of NaOH should have been 0.103 mol, but the value I obtained due to the excessive deviation gave me 0.0937 mol. The percent error that has resulted in:

9.03% is by far a significant error that has resulted from a small error in the volume. Another error was caused by the deviation in the mass of KHP. 2.00 grams was the amount expected to be taken, but the experimental amount was 1.99 grams. However, this, being only 0.01 grams of the expected value, could have only constituted a very small portion of the error. This means that due to systematic error, my accuracy has fallen by 9.03%, which, although not high, is quite a deviation inaccuracy.

The deviation in the volume, however, is not the only indicator of noticeable systematic errors. The volumes of NaOH used up show significant fluctuations. For example, in trial 1, I used 11.0 cm 3 of NaOH, which is 1.50 cm 3 off 9.50 cm 3 , and in my rough trial, the volume used was 9.9 cm 3 and in trial 2, the amount used up was 10.4 cm 3 . The difference between these sets of data indicates that the systematic error of allowing the KHP solution to become too pale resulted in strange fluctuations. These fluctuations caused the 0.95% error.

The percentage uncertainty calculated of the concentration of NaOH was 2.57%, which indicates that the level of precision, although not low, could have been better. The expected % uncertainty that was expected was 0.500%, and the uncertainty I obtained was 0.503%. Also, the % uncertainty of the volume of NaOH was ±1.05%, taking the value of 9.50 cm 3 . Therefore, due to flaws in raw data values taken from systematic errors, there has been a deviation in uncertainty too, indicating the impact of methodical flaws. The uncertainty of 2.57% indicates that my values were accurate up to within ±2.57%.

Overall, the data obtained, although not completely inaccurate, is not as accurate as it could have been. These errors were avoidable.

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This is an awesome source of information, Thank you !

Hello..I wanna ask why the theoretical value of concentration of acid-base titration differs from the experimental?

Due to Uncertainity

What I know is due to various errors which tend to happen when conducting an experiment such as contamination of the sample used (impurities) also external factor like temperature and humidity which results the sample to react with the atmosphere (air)

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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?

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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Preparation and standardization of sodium hydroxide

It is an acid-base titration. The sodium hydroxide is an alkali whose strength changes over time and it can be effectively standardized utilizing primary standard viz. oxalic acid. Sodium hydroxide reacts with oxalic acid in presence of phenolphthalein indicator. The color changes from colorless to pink at the end point. 1

Hence, based on the above theory our aim is to prepare and standardize sodium hydroxide using oxalic acid.

REQUIREMENTS

            40 g of NaOH dissolved in 1000 ml water = 1N NaOH

            31.5 g of oxallic acid dissolved in 1000ml water = 0.5 N oxalic acid

Calculation

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19 Buffers, Indicators, and Solution Standardization

To explore the properties of buffers and indicators, and to review experimental techniques associated with acid-base titration

Expected Learning Outcomes

After completing this experiment, students are expected to be able to

• Calibrate and use a pH meter.
• Prepare and recall the properties of a buffer solution.
• Select appropriate indicators for use in a titration.
• Prepare and standardize a solution of sodium hydroxide.

Textbook Reference

From Tro,  Chemistry: Structures and Properties , 2nd Ed:

• Buffers are discussed in Ch. 17.2-3.
• Indicators are discussed in Ch. 17.4.
• The preparation and standardization of acidic and basic solutions was conducted previously in CHEM-C 125 .

Introduction

Buffers are formed by mixing together similar amounts of weak acids and their conjugate base, such that

[latex]0.1 < \frac{[\mbox{base}]}{[\mbox{acid}]} < 10[/latex]

The pH of a buffer solution can be calculated using the  Henderson-Hasselbalch equation . [1]

\begin{equation} \mbox{pH} = \mbox{p}K_a + \log_{10} \left(\frac{[\mbox{base}]}{[\mbox{acid}]}\right) \label{414:HH} \end{equation}

Indicators are weak acids that absorb different colors in their acid and conjugate base forms. In environments with different pHs, they have different colors. [2]

Phenol red changes color from yellow at pH 6 to red at pH 8:

As part of this experiment, you will examine the colors of different indicators at different pHs. You may have used indicators before for standardization of acid and base solutions.

Standardization of Acids and Bases

Often, it is vital that the concentration of solutions are critical for human life. These are often determined using various titrimetric and other analytical techniques. In these cases, it is vital that reference solutions be prepared accurately.

The concentration of chloride ions in water is typically determined by titration with silver, using the precipitation of silver chloride as a standard.

\begin{equation} \ce{Ag+}(aq) + \ce{Cl-} (aq) \rightleftharpoons \ce{AgCl}(s) \end{equation}

When this is performed, it is vital that the concentration of silver ions in the analyte (we typically use the water sample as the titrant) is known accurately.

For acid-base experiments, the challenge is even greater. The solid form of many acids and bases are hygroscopic; they will absorb water from the ambient environment. This is true in particular for sodium hydroxide, which will form swollen pellets that may clump together as it absorbs moisture from air:

For this reason, specific primary standards  of acids and bases are used to determine the concentration of other acids and bases using titrimetric methods rather than simply by dissolving the appropriate mass of solution.

In this experiment, you will prepare a solution of approximately 0.1 M sodium hydroxide and standardize this solution using potassium hydrogen phthalate (KC 8 H 5 O 4 , often referred to as “KHP”; molar mass = 204.2 g/mol) as the primary standard, as you did previously in the CHEM-C 125 experiment  Analysis of Acids and Bases by Titration .  Please review the discussion about acid-base titration in the lab manual for that class.

• You will perform this experiment in pairs.  However, for the solution preparation and standardization part, if your lab partner is absent you should conduct this part of the experiment individually.
• Review the directions on using a pH meter in  Using Laboratory Equipment before coming to lab.
• All solutions are properly labeled.
• Volumes are measured accurately using volumetric glassware when preparing solutions. The exact volume for the pH measurement, however, does not matter. You will need to ensure that you have at least about 10 mL of solution to take good measurements such that the minimum liquid level on the pH probe is exceeded.
• You must label and turn in your solution of sodium hydroxide. In the next lab period, you will use this solution as a secondary standard } to perform acid-base titration experiments.

The Action of Buffers

• Turn on your device and open the Vernier Graphical Analysis app. Connect the pH probe following the directions in this manual . Calibrate the pH probe using pH 4, 7 and 10 buffers following the directions here . [3]
• Prepare a buffer solution by mixing 15.0 mL of 1.0 M acetic acid (CH 3 COOH) with 15.0 mL of 1.0 M sodium acetate (CH 3 COONa) in a graduated cylinder. Mix these solutions thoroughly and measure its pH in a test tube.
• Measure out 5 mL of this solution into each of three different test tubes.
• 5 mL deionized water
• 5 mL of 0.05 M sodium hydroxide
• 5 mL of 0.05 M hydrochloric acid.
• Repeat steps 3 and 4 using deionized water instead of the buffer solution you made in step 2.

This part of the experiment will be done individually; each student will be assigned a different pH buffer.

As you may know different indicator have different colors depending on their pH.  These are often summarized on charts like this:

You will be assigned one of the buffers provided (pH 4, 7, 10) to study.  The goal of this experiment is to use  two different indicators from this list to identify the pH as best you can.  You will be given the following indicator solutions:

• alizarin yellow
• bromothymol blue
• methyl orange
• phenolphthalein

Add a small amount of your chosen indicator solutions into separate samples of your assigned buffer solution.  Record the colors of each of these solutions.

Preparation of a Sodium Hydroxide Solution

In this part of the experiment, you will prepare using solid sodium hydroxide approximately 500 mL of 0.1 M standardized sodium hydroxide. To determine the pH, you will prepare a standard potassium hydrogen phthalate (KHP) solution as a primary reference.

• Calculate the mass of solid sodium hydroxide required to prepare 500 mL of 0.1 M sodium hydroxide solution.
• Your names/initials.
• The identity of the solution (including the approximate concentration).
• Today’s date.
• The hazard class ( look up from an SDS sheet )

Standardization of Sodium Hydroxide

Following the directions in the  Analysis of Acids and Bases by Titration experiment, prepare a sodium hydrogen phthalate (KHP) solution and use this to determine (accurately) the concentration of sodium hydroxide solution you prepared.

The bottle of sodium hydroxide you prepared should be placed in your drawer for use next week.

Waste Disposal

All waste from this experiment can be disposed of in the drain with plenty of water.

• This is equivalent in assumption to neglecting x in ICE tables. ↵
• Universal indicator is a mixture of different indicators designed to have different colors at different pHs. ↵
• These may be substituted by any two or three pH values.  Be sure to pay attention to the pH values on bottle. ↵

IU East Experimental Chemistry Laboratory Manual Copyright © 2022 by Yu Kay Law is licensed under a Creative Commons Attribution-NonCommercial 4.0 International License , except where otherwise noted.

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• Inorganic Chemistry

Experiment (1) Standardization of sodium hydroxide NaOH solution

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