Monday, November 15, 2021

Why do you need to calibrate a pH Meter

Regular calibration of pH meter is best to maintain in pH measurement since calibration maintains the readings reliable and accurate. You can calibrate the pH meter using the standard pH calibration solutions (pH 07.00, pH 04.00, and pH 09.20) and they are easily available.

What is a pH meter?

A pH meter is a scientific laboratory apparatus that determines the acidity or alkalinity of water-based sample solutions by measuring hydrogen-ion activity.

A membrane in a pH meter allows H+ ions to pass through, allowing current to flow and providing voltage. The voltage is measured by the meter and you tell it what standard buffer it should be in. To determine the pH of the sample solution, the pH meter compares the voltage of the unknown solution to that of the buffers.

Litmus paper (blue or red), pH paper, and universal indicator are some of the methods of pH determination that are available. However, a pH meter provides the most accurate results across the entire pH range.

What is calibration of pH meter?

The process of calibrating the pH meter by measuring liquids with a known pH value is known as pH calibration. This is important because the properties of your electrode can vary over time, and this needs to be compensated for. A calibration completes this by matching the current characteristics of the pH sensor to the pH meter.

What types of pH meter calibration are there?

There are single-point, two-point, and three-point or multi-point calibration options for pH meters. When need to test a consistent pH value with low variation, single-point calibration can be used. In, single-point calibration, only one buffer solution is used as a calibration reference.

In 3 point calibrations, three buffer solutions such as pH 04.00, 07.00, and 09.20 are used for the calibration procedure that covers the entire pH range (acidic, neutral, and basic). To maintain the accuracy of your pH meter at least three standards are required to create a calibration curve.

Why you need to perform pH meter calibration?

The calibration of the pH meter is important and it is performed to maintain the accuracy of results, to avoid drift, to account for differences, and to see the changes in their characteristics.

To maintain accuracy:

It is necessary to calibrate the pH meter to maintain the accuracy of the results. Calibrating the meter without standardized pH buffer solutions results will be useless and incorrect.

To avoid drift:

Most pH meters, as well as electrodes in general, are known to drift from their calibrated settings. Drift from calibrated settings cannot be avoided; however, it is significant to calibrate the pH meter frequently to make sure that you keep getting accurate results.

To account for differences:

When calibrating, using standardized buffers also helps eliminate differences between samples. Ionic strength differences and other membrane-related issues can be avoided with proper standards.

Changing characteristics:

The characteristics of pH electrodes can change over time due to aging and coating, and even the most stable electrodes cannot be made with the same characteristics. Calibration helps match the current characteristics of the pH meter to the pH sensor in use, compensates for any difference in the behavior of the pH electrode in theory and reality.

The electrodes are based upon the offset and slope, but, at the age, all the electrodes will change. Actually, at all times the electrodes will not behave as per the Nernst equation.

This is the reason, where the calibration arrives. When using a known buffer, proper calibration will be ready for the old electrode by determining the actual offset and slope to match and update the algorithm of the pH meter.

If the calibration results are not proper or not reproducible as per the standard solutions, it is a signal that the electrode of the pH meter is dirty, too old, or damaged and wants to be changed.



FAQ (Frequently Asked Questions):


What is the purpose of calibration of pH meter?
The purpose of calibration is to ensure that instrument readings are consistent and accurate.

Which solution is used to store glass electrodes?
The pH electrodes are recommended to be stored in a 04 M KCL solution, if 4 M KCL isn't available, use a pH 04.00 buffer solution to store them.

When should a pH meter be calibrated?
In the case of high-accuracy measures, the pH meter must be calibrated before each test, and for normal-accuracy measurements, it can be calibrated in a week or more.


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Saturday, November 13, 2021

Why are Buffer Solutions used to Calibrate the pH Meter

The buffer solution is used to calibrate the pH meter as they resist pH change. So we use a standard buffer solution with a known pH to set the pH meters to show those results. When using a pH meter to test pH, we want to confirm that the meter indicates pH = 7.00 and that the pH is 7.00.

We most likely used commercially available standard calibration buffers, which are sold in the form of tablets or capsules that dissolve in deionized water or ready-to-use solutions. The  calibration standard solutions used for the electrode calibration have a pH of 25°C.

Please keep in mind that high pH buffers are less stable since they absorb CO2 from the atmosphere, lowering their pH. We should only open the bottle to pour the buffer into the beaker during calibration; the bottle is never left open.


It is also worth noting that the pH of the buffer solution varies depending on the temperature. For example, at 80°C, the pH of potassium hydrogen phthalate (C8H5KO4) solution rises to 04.16. Several pH meters do not automatically adjust for these variations, even if they provide automatic temperature correction during measurement.

There is a membrane in the pH meter, which allows passing the H + ions, which permit the current to flow, creates the voltage. The voltage is estimated by the pH meter and you reveal to it which standard buffer, it is in, after that the pH meter compares the voltage of its unrecognized solution to the buffer to decide the pH of the samples. Standard buffer solutions are defined to be at a precise pH. They can generally be gotten from the manufacturer or self-prepared and these buffers are important for the proper functioning of the pH meter.

Why buffer solution is used for calibration of pH meter?

The buffer solution is used to calibrate the pH meter as they resist pH change. So we use a standard buffer solution with a known pH to set the pH meters to show those results. A buffer solution is an aqueous combination of a weak acid and its conjugate base, or vice versa. Buffer solutions are used in a variety of chemical applications as a means of keeping the pH at a nearly constant value.

Why buffers are the best for this purpose are as follows.

  • Buffer solutions are easy for the preparation for a particular pH.
  • Buffer solutions are stable for a long time.
  • If you add a little amount of acid or base or even water, then they resist modifying in pH.
  • Buffer solutions maintain the accuracy of the pH meter.
  • The prepared buffer solution can easily avail in the market and is cheap in rate.

Which buffer solution is used for the calibration of the pH meter?

Three types of buffer solutions pH 04.00, pH 07.00, and pH 9.20 are used to calibrate the pH meter. Prepared standard buffer solutions are available in the market or you can prepare them in the laboratory by the procedure given in the monograph.

These buffer solutions facilitate to display of the accurate pH value since when we use a pH meter we desire to be sure that the pH meter displays the correct measurement. Buffer solutions provide accurate results within ±0.01 pH at 25 °C. Prepared standard buffer solutions are available in the market or you can prepare them in the laboratory by the procedure given in the monograph.

Why 3 point calibration is used to calibrate the pH meter?

To ensure the accuracy of readings, it is necessary to calibrate with three-point calibration utilizing different buffer solutions such as pH 04.00, 07.00, and 09.20, which covers the whole pH range i.e. acidic, basic, and neutral.

A pH meter uses the Nernst equation to calculate the pH of the sample solution. A 2 or 3 point calibration, using 2 to 3 different buffer solutions is usually sufficient for calibration as the meter's electronic logic will calculate pH values in between.

What buffer solution is used for storing pH electrode and why it is chosen?

Manufacturers of pH electrodes often recommend storing them in a 04 M KCl solution. Use a pH 4 buffer solution if 4 M KCl isn't available. If you keep your electrode in distilled or deionized water, ions will leak out of the glass bulb, rendering it useless.

The pH-responsive electrode is normally made of glass, whereas the reference electrode is usually made of silver-silver chloride. The difference in potentials between the two sides of the glass electrode is measured by the pH electrode.


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Friday, November 12, 2021

What is pH paper and how does it work

A pH indicator is a chemical detector for hydrogen ions (H+) or hydronium ions (H3O+) in a sample solution. The pH (potential of hydrogen) value is used to express the amount of hydrogen in a solution. The negative logarithm of hydrogen ion concentration is related to the pH scale, which ranges from 0 to 14. Generally, the indicator causes the color of the solution to change depending on the pH.


In chemistry, there are three types of indicators: natural indicators, artificial indicators, and olfactory indicators, they are used in different types of titrations like acid-base, redox, precipitation, isometric, etc. The two theories that explain the acid-base indicator are Ostwald's theory and the Quinonoid theory. Depending on the application they use litmus paper (red or blue), pH paper, pH indicator, phenolphthalein, methyl orange, and potentiometer and pH meter, etc.

What is pH paper?

What is pH paper and how does it workA pH paper is a strip of paper that has been saturated with pH indicators and is used to determine the pH of a solution.
Whether there is an acid or a base present, pH paper can change color depending on the pH of the sample solution. It is a paper treated with a natural water-soluble dye from lichens. pH paper is also called a universal indicator, at different concentrations of hydrogen ions in a solution show different colors.

In an acidic media, the pH indicator will display colors such as yellow, orange, red, and orangish-yellow. In the basic medium, it will turn purple or blue-violet, and in the neutral medium, it will turn green. However, there are various universal pH indicators in the market, most of which are modifications of the formula patented by Yamada in 1933.

What is pH paper made of?

The pH paper is made of paper; the basic raw materials used to compose pH paper are wood cellulose, lichens, and additional compounds. 

The universal indicator is usually composed of water, phenolphthalein, methyl red, 1-propanol, thymol blue, bromothymol blue, sodium bisulfite and sodium hydroxide, etc. A pH paper is saturated with chemical compounds whose chromatophore is a different color at different pH concentrations.

How does pH paper work?

The pH paper is coated with a chemical indicator that changes color when hydrogen ions or hydroxide ions are present. A solution with a high concentration of hydrogen ions is acidic. A solution with a high concentration of hydroxide ions is basic or alkaline, and if the amount of hydrogen ions and hydroxide ions are equals that result in a neutral state.

pH indicators are weak acids or bases that change color at particular pH. Acids and bases are chemical compounds that exchange a proton. Acids donate a proton, whereas alkaline is substances is accept a proton. Acids are hydrogen donors and bases are hydrogen acceptors. Thus, the color changes when an acid or base accepts or donates a proton.

Types of pH paper:

Universal paper and litmus paper are the two most prevalent types of pH paper. Both types are widely used for acidity and basicity determination, which is depending on the type of measurement and degree of precision required. 

Universal paper can indicate a range of pH values, while red and blue litmus paper kind of ready-made test filter paper that roughly indicate whether a sample is acidic or basic. When the pH is alkaline, the red paper turns blue, and when the pH is acidic, the blue paper turns red.

Nowadays, there are many different types of Whatman pH indicator paper available, each of which provides a different range of pH depending on the application and provides reliable and accurate results.

How to use a pH paper strip:

To determine the pH of any liquid, put a drop of the sample solution on the pH strip with a dropper. Or hold one end of the pH strip and dip the other end into the sample solution, then remove it after the required amount of time has elapsed.

When a color change occurs, it is measured against the color comparison chart (01-14). It is not necessary to immerse the entire strip in the test sample, as the dye could potentially contaminate the sample.

Applications of pH paper:

  • The most common application of pH paper is to determine whether a solution is acidic, basic, or neutral.
  • The pH paper is used for research, health care, educational institutions, pharmaceuticals, chemical laboratories, agricultural fields, environmental science, and water quality testing, etc. where it is necessary to determine the pH level of substances.
  • It is also used in science classes (practicals/experiments) because it is rapid and easier to use than a pH meter.


FAQ (Frequently Asked Questions):


Is a pH paper strip reusable?
It is not recommended to reuse pH paper as the results may be inaccurate. It is not completely insoluble, and the paper tends to fall apart after a while, so it cannot be used indefinitely, although it will work a few more times.

What is pH paper coated with?
The pH paper is coated with a chemical indicator that changes color when hydrogen or hydroxide ions are present.

What is the pH value of a neutral solution give an example?
The pH value of any neutral solution is 7. Water and human blood are examples of neutral solutions.

What is the difference between litmus paper and universal pH paper?
The pH value of the sample solution can be determined using universal pH paper, whereas the acidity or alkalinity of the solution can be determined using litmus paper.
 
 
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Wednesday, November 10, 2021

Selection of indicators for a titration

An indicator is used to indicate the end of the titration as a noticeable pH change occurs near the equivalence point of the acid-base titration. While choosing an indicator for acid-base titration, select an indicator whose pH range falls within the pH change of the reaction.

Compounds that change color when exposed to acidic or alkaline solutions are known as indicators. Colored indicators are widely used to detect pH and can be added to the reaction mixture to determine the titration's endpoint or equivalence point.

In the laboratory of science classes for practical purposes, indicators like phenolphthalein and methyl orange are commonly used. A universal indicator, pH paper, litmus paper (Blue or Red), and pH meter are usually used to determine the pH of the substance.

In the titration process, a substance to be analyzed (Titrand) is placed in a conical flask, two to four drops of a suitable indicator are added, followed by a drop by drop addition of titrant of known strength from a burette until the completion of a chemical reaction.

History of indicators:

Sir Robert Boyle first described the use of a natural dye as an acid-base indicator in 1664 in his collection of articles "The Experimental History of Dyes." By using markers for the experimental classification of acids and bases, Boyle made significant contributions to the early theory of acids and bases.

The theory of acid-base indicators can be explained in two ways: Ostwald's theory and the Quinonoid theory. In chemistry, there are three types of indicators: natural indicators, artificial indicators, and olfactory indicators, however artificial and natural indicators are the two most commonly used.

What are the two criteria for an indicator to be used in titration?

The two general criteria for an indicator to be used in a titration are: the pH at the end of the titration should be close to the indicator's neutral point, and the indicator should indicate a sharp color change near the equivalence point of the titration point.

For example, when titrating a strong acid with a strong base, the pH quickly shifts from 03 to 11. Therefore, the phenolphthalein indicator, which has a pH range of 08 to 10, is a good choice for this type of titration.

Choice of indicators for various acid-base titrations:

Each pH indicator changes color over a defined range of pH, known as the indicator range. The indicator ranges are listed below for some of the indicators.

Titration of a strong acid against a strong base:

For example, hydrochloric acid (HCl) vs. sodium hydroxide (NaOH)
In this form of titration, the pH value at the endpoint changes about from 04 to 10. As a result, any indicator that changes color within this range can be used to titrate strong acid against a strong base. For this type of titration, phenolphthalein may be suitable as an indicator.

Titration of a weak acid against a strong base:

For example, Oxalic acid vs. sodium hydroxide (NaOH)
In this type of titration, the pH value changes slightly at the endpoint which is about 6.5 to 10. Since the working range of phenolphthalein is 8.3 to 10, it is a suitable indicator for this titration. In these cases, methyl orange is not an appropriate indicator, because it has a working range below pH 05.

Titration of strong acid against weak base:

For example, hydrochloric acid (HCl) vs. sodium carbonate (Na2CO3)
When HCl (strong acid) is titrated against sodium carbonate (a weak base), at the endpoint the pH varies from 3.5 to 7.5. Methyl orange, which changes color within this pH range, is the more suitable indicator for this type of titration.

Titration of weak acid against weak base:

For example, Acetic acid (CH3COOH) vs. ammonium hydroxide or ammonia solution (NH4OH)
pH value at the end of this form of titration there is no sharp change. Therefore, this no indicator is suitable, because it requires a vertical portion of the curve over two pH units. As a result, neither phenolphthalein nor methyl orange is useful.

Indicators for different types of titration:

There are four types of titration such as acid-base titrations, redox titrations, precipitation titrations, and complexometric titrations, etc. According to the pH range and reaction, each titration has a particular reaction mechanism and uses different indicators.

Indicators for acid-base titration:

Acid-base indicators are generally classified into listed three groups: phthaleins and sulphophthaleins, Azo indicators, Triphenylmethane indicators.

Examples: Phenolphthalein indicator, methyl orange indicator, and malachite green indicator, etc. can be used be as an indicator

Indicators for complexometric titration:

Complexometric indicators are water-soluble organic compounds that are used to detect the endpoint of a titration.

Examples: Eriochrome Black T, calgamite, arsenazo, xylenol orange, Eriochrome red B, fast sulphon black, and calcein, etc. can be used be as an indicator

Indicators for redox titration:

The redox indicators are indicators that show a reversible color change between oxidized and reduced forms that change color at a specific potential.

Examples: Potassium permanganate (self-indicator), phenanthroline blue, methylene blue, 1, 10 phenanthroline monohydrate, and safrannin-T, etc. can be used be as an indicator

Indicators for precipitation titration:

There are three types of indicators employed in precipitation titration: two of them are chloride ions, and one is cation silver.

Examples: In Mohr's titration, silver nitrate or potassium chromate can be used, In Volhard’s titration, the ferric ion can use, and In Fajan’s titration, the dye dichlorofluorescein can used be as an indicator.

Indicators for non-aqueous titration:

The ionized and unionized forms of indicators, as well as numerous resonant forms of indicators, are all generally useful for non-aqueous titration.

Examples: Crystal violet indicator, methyl red indicator, naphtholbenzein indicator, Quinaldine red, and thymol blue, etc. can be used be as an indicator

List of indicators:

The following table gives an approximate color of some indicators at different pH values, as well as the types of titrations for which they are useful.
Name of Indicator Range of pH Colour in acidic Colour in basic
Methyl Violet 0.0 to 1.6 Yellow Blue
Crystal Violet 0.0 to 1.8 Yellow Blue
Cresol red 0.2 to 1.8 Yellow Red
Malachite green 0.2 to 1.8 Yellow Green
Thymol Blue 1.2 to 2.8 Red Yellow
Methyl Yellow 2.9 to 4.0 Red Yellow
Bromophenol blue 3.0 to 4.6 Yellow Blue
Methyl Orange 3.1 to 4.4 Red Yellow
Bromocresol Green 3.8 to 5.4 Yellow Blue
Dichlorofluorescein 4.0 to 6.6 Colorless Green
Methyl Red 4.2 to 6.2 Red Yellow
Chlorophenol Red 4.8 to 6.4 Yellow Red
Bromocresol Purple 5.2 to 6.6 Yellow Purple
Bromothymol Blue 6.0 to 7.6 Yellow Blue
Phenol Red 6.4 to 8.0 Yellow Red
Cresol Purple 7.4 to 9.0 Yellow Purple
Thymol Blue 8.0 to 9.6 Yellow Blue
Phenolphthalein 8.0 to 9.8 Colorless Red
Thymolphthalein 9.3 to 10.5 Colorless Blue
Alizarin Yellow 10.1 to 12.0 Yellow Red
Indigo Carmine 11.4 to 13.0 Blue Yellow
- - - -
Litmus 5.0 to 8.0 Blue Red
Universal indicator 0.1 to 14.0 - -

FAQ (Frequently Asked Questions):


Why is universal indicator not used in titration?
A universal indicator is used to determine the approximate pH of a sample compound. As it changes color over a broad range of pH it is not used in the titration.

What is a natural and artificial indicator? Give some examples.
Natural indicators are types of indicators that occur naturally (e.g turmeric, red cabbage) and artificial indicators are those that are prepared artificially in the lab or produced through a chemical reaction (e.g. phenolphthalein, methyl orange), both can be used to identify whether a substance is acidic or alkaline.

What happens if you use the wrong indicator in titration?
If we use the wrong indicator in a titration, the endpoint will not be as expected, as the titrant solution may take more or less volume, making your calculations incorrect. As a result, the result may be wrong or incorrect.

References:

  1. Reactions of Acids and Bases in Analytical Chemistry. Hulanicki, A. and Masson, M.R. New York: Halsted Press, 1987.
  2. https://www.sciencecompany.com/ph_indicator_ranges.aspx
  3. https://en.wikipedia.org/wiki/Titration
  4. Aqueous Acid-Base Equilibria and Titrations. Levie, Robert De. New York : Oxford University Press, 1999.

 

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Monday, November 8, 2021

What is indicator in chemistry and its types

Learn the three types of indicators, natural, artificial, and olfactory which are used to determine whether a substance is acidic or basic. It is a weak acid or weak base that dissociates in a solution to form ions.

The concentration of hydrogen in a solution is expressed using the pH (potential of hydrogen) value. The negative logarithm of hydrogen ion concentration is related to the pH scale, which ranges from 0 to 14. An acid is a substance that gives hydrogen ions, so when a solution has more hydrogen ions than hydroxide ions, it is acidic. A base is a substance that accepts hydrogen ions, so when a solution has more hydroxide ions than hydrogen ions, it is alkaline.

What is an indicator in chemistry?

Indicators are compounds that change color when introduced to acidic or alkaline solutions. Colored indicators are commonly used to determine pH and can be added to the reaction mixture to determine the endpoint or equivalence point of the different types of titration.

In the laboratory (for practicals) of science classes, indicators such as litmus, phenolphthalein, and methyl orange are most commonly used. As well as to determine the pH of the compound a pH paper, universal indicator, litmus paper (Blue or Red), and pH meter are generally used.
Definition of the indicator: 
“Indicator is a substance that shows a color change when exposed to acids and bases.”
An indicator has a defined pH range in which it changes from an acid to a base form. For example, litmus paper is blue in a basic solution and red in an acidic solution. The range is from 0 to 14, with 07 being the neutral value. Acidic is indicated by a pH less than 07, while the base is indicated by a pH greater than 7. Indicators cannot work outside their pH range, as an indicator does not change color over a wide range of pH values.

Indicators are weak acids or bases (usually derived from plant pigments) that have unique colors in their ionized and non-ionized states and are related to the pH of the solution being analyzed.

For example, using phenolphthalein as an indicator for a titration with a strong base as the titrant and acid as the titrated ingredient would result in a color change in the conical flask from colorless to pink towards the equivalence point.

History of indicators:

In his collection of essays "The Experimental History of Dyes", Sir Robert Boyle first described the use of a natural dye as an acid-base indicator in 1664. Boyle made major contributions to the early theory of acids and bases by using markers for the experimental classification of acids and bases.

Theory of acid-base indicators:

The theory of the acid-base indicator is explained by two theories: Ostwald's theory and the quinonoid theory.

Ostwald's theory:

Ostwald's theory, which is based on the Arrhenius theory, was proposed by Ostwald in 1891. According to this theory, the acid-base indicator ionization causes the color change. The color of the unionized form varies from that of the ionized form.

Quinonoid theory:

According to the quinonoid theory, the color change of the acid-base indicator results from a structural change. An indicator is thought to exist in equilibrium as two tautomeric forms, the benzenoid and quinonoid forms. One form occurs in an acidic solution and the other form in an alkaline solution. The two forms have two different colors, and as the pH of the solution changes, the color of the solution containing the indicator changes.

How do pH indicators work?

Indicators are both weak acids and bases, by reacting with H+ and OH-, pH indicators detect the presence of H+ and OH-. The deprotonation produces a color change when an indicator is a weak acid that is colored and its conjugate base is a different color.

The color we observe is determined by the ratio of the concentration of the indicator, HInd, and its conjugate base, Ind-. According to the Henderson–Hasselbalch equation, this ratio depends on the pKa and pH of the indicator.

Characteristics of indicators in chemistry:

  • It should change color when the pH reaches a point close to the endpoint of the reaction.
  • It should be sensitive to the equivalence point of the titration.
  • The color change of the sample solution should be very sharp and clear.
  • The produced color should be brilliant and stable.
  • Two colors of an indicator must be contrasted so that they can be identified.
  • The indicator should be cheap and non-toxic in nature.

Different types of indicators:

Different types of indicators are used in different types of titrations like acid-base, redox, precipitation, complexometric, etc. Each of them has its pKa, pH range, and produces a particular color at different pH values.

There are three types of indicators in chemistry: natural indicators, artificial indicators, and olfactory indicators, although artificial and natural indicators are the two types of chemical indicators that are most commonly used.

Natural indicators:

Natural indicators are types of indicator that occurs naturally and can be used to identify whether a substance is acidic or alkaline. Some examples of natural indicators are litmus, turmeric, china-rose, curry powder, red cabbage, grape juice, cherries, tomato, and onion, etc.

To detect hydrogen ions (H+) and hydroxyl ions (OH-) in a sample solution, these natural indicators are used. Some examples of household pH indicators are red cabbage, rose-petals and turmeric.

Artificial indicators:

Artificial indicators are types of indicators that are prepared synthetically in the laboratory or obtained through a chemical reaction that does not occur naturally.

Some examples of artificial indicators are phenolphthalein, methyl orange, malachite green, and methyl red, etc. These synthetic indicators are used for titration as well as the identification of acids and bases.

Olfactory indicators:

Olfactory indicators are types of indicators that do not indicate a color change however a change in odor when added to acidic or a basic sample. Olfactory indicators, unlike other indicators, change their smell rather than their color to indicate whether a sample solution is acidic or basic.The odor varies depending on the nature of the medium.

Some examples of olfactory indicators are onion, garlic, vanilla extract, clove oil, etc.

Indicators pH range and color change:

Here are some common indicators are listed in table form along with their pH range, pKa value, and color change under acid and basic environments.

What is indicator in chemistry and its types

Universal indicator:

A universal indicator is a pH indicator composed of different compounds that changes color over a wide range of pH values to indicate the acidity or alkalinity of the sample solution. The new color that is formed is then compared with the pH chart.

Usually, it is formulated by, phenolphthalein, water, methyl red, 1-Propanol, sodium bisulfite, sodium hydroxide (NaOH), bromothymol blue, and thymol blue, etc. However, there are various universal pH indicators on the market, most of which are modifications of the formula patented by Yamada in 1933.

Litmus paper:

Litmus paper is a kind of indicator that is used to determine if a compound is acidic or basic. There are two types of paper strips available, blue and red, which are made by mixing water-soluble dyes that are absorbed into the filter paper strips. It is impregnated with lichens, giving it the capacity to change color when exposed to acids and bases.

Red litmus paper of a weak diprotic acid and changes color red to blue when exposed to a basic solution or gaseous sample. Blue litmus paper already consists of a blue conjugate base and it changes color blue to red when expose to the acidic solution or gaseous sample.

Phenolphthalein:

Phenolphthalein (C20H14O4) is a phthalein-family organic compound that is commonly used as an acid-base indicator. The ionic form of phenolphthalein is In-, while the molecular form is HIn. The molecule loses a hydrogen ion and becomes a negative ion in the basic solution. Phenolphthalein is the preferred indicator when titrating weak acids with strong bases.

It is a fine crystalline powder that is yellowish-white to pale orange and appears colorless in the liquid form up to pH 8.5, after which it changes from pink to dark red. It has a pKa value of 9.3, is slightly soluble in water, and is prepared by dissolving in alcohol as an indicator for acid-base titration experiments.

Methyl orange:

Methyl orange (C14H14N3NaO3S) is an organic dye commonly used as a pH indicator in titrations, since of its clear and different color variations at different pH values. In an acidic medium, methyl orange turns red, while in a basic condition, it turns yellow.

Methyl orange is preferred for titrations involving weak bases and strong acids, such as ammonia solution (NH4OH) and hydrochloric acid (HCl) because its pKa (3.4) lies in the acidic region of the pH scale.

Bromothymol blue:

Bromothymol blue is an ionic dye also known as bromothymol sulfone phthalein and BTB. Bromothymol blue (C27H28Br2O5S) is a pH indicator that can be used to detect weak acids and bases. It has a bluish-green color in neutral solution and turns blue at pH 7.6 and yellow at pH 6.0. BTB is a good choice for titrations that have an equivalence point near the neutral.

Phenol red:

Phenol red (C19H14O5S) is a water-soluble dye also known as phenolsulfonphthalein that changes color from yellow to red over pH 6.6 to 8.0, then turns a bright pink color above pH 8.1. Phenol red is a pH indicator dye that can be used in a variety of medical and cell biology studies.

Methyl red:

Methyl red (C15H15N3O2) is a dye used as an acid and base indicator in the laboratory and also used in textile industries. It is an azo dye, comes in a dark red colored crystalline powder; indicator is also available in liquid form.

At pH 4.4 and below, methyl red turns red, and when pH 6.2 is achieved, it turns yellow. The color orange is found between these color endpoints, in the pH range of 4.4 to 6.2.

Turmeric:

Turmeric is a yellow-colored natural indicator, also as well as household indicator. Turmeric solution or paper becomes reddish-brown by adding base and adding acid does not change its color.

Onion:

Onion indication is a kind of olfactory indicator in which the smell of onion paste or juice is lost when it is mixed with the base. It does not change its odor when exposed to acid. Red onion turns pale red in acidic solution and green in alkaline solution.

Uses of indicators in chemistry:

  • The most typical use of indicators is to determine the endpoints of a titration.
  • The role of indicators in chemistry is to determine whether a reaction has completed or undergone a chemical change.
  • It is used in laboratory experiments such as titration; it is one of the most primary experiments for secondary school of chemistry, analytical chemistry students, and it necessitates the use of a burette, pipette, and conical flask, etc.
  • Titratable acidity or alkalinity of the sample is determined using indicators.
  • Chemical indicators are used to detect the presence of certain compounds in water, such as pollutants.
  • It is also used in industrial chemical processes, such as pharmaceutical, food, agriculture, and environmental science, etc.


References:

1. Chemical Sensing Using Indicator Dyes: Wolfbeis, Otto F. Optical Fiber Sensing 1997, 4, 53-107.
2. Schwarzenbach, Gerold (1957). Complexometric Titrations. Harry (1st English ed.). London: Methuen & Co. pp. 29–46.
3. https://www.sciencedirect.com/topics/chemistry/acid-base-indicator
4. Modern reaction-based indicator systems. Dong-Gyu Cho and Jonathan L. Sessler Chem. Soc. Rev. 2009, 38, 1647-1662.
5. http://www.3rd1000.com/chem301/p00413.html
6. https://en.wikipedia.org/wiki/PH_indicator
 
 
 
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Monday, November 1, 2021

What is universal indicator and how does it work

Learn about the types of universal indicators, their uses, advantages, and how they work.

The pH is a significant factor that determines whether a solution is acidic or alkaline. It determines the concentration of hydrogen ions (H+) in the aqueous solution on a negative logarithmic scale from 0 to 14.

There are many activities in research, pharmaceuticals, science classes, and agriculture sectors, etc that involve pH testing, environmental science, titration in chemistry, water quality testing, and biological processes are among them. Depending on the application they use the pH paper, litmus paper (red and blue), pH indicators, and potentiometer or pH meter.

What is a universal indicator in chemistry?

A universal indicator is a pH indicator composed of a solution of a different compound that shows multiple color changes over a broad range of pH values ​​to indicate the solution’s acidity or alkalinity. To determine the pH value of the acid or base solution a sample is added to the universal indicator, which generates a new color that is used to match the color to the colors on a pH chart.
Definition: “A universal indicator is a combination of multiple indicators that produce distinct colors at various pH levels on the scale”

What is universal indicator and how does it workThe indicator will show colors such as orange, red, yellow, orangish-yellow in an acidic medium. It will change color to purple or blue-violet in basic medium and its color will change to green in neutral medium. However, several universal pH indicators are commercially available, most of which are variations of a formula patented in 1933 by Yamada.
Example: Solution or pH paper

History of indicators:

Sir Robert Boyle first mentioned the use of a natural dye as an acid-base indicator in his essays collection "Experimental History of Colors" in 1664. By using markers for the experimental classification of acids and bases, Boyle made significant contributions to the early theory of acids and bases.

Theory of indicators:

The two primary theories that lead to the working principle of acid-base indicators are the Ostwald theory and the quinonoid theory.

Ostwald's theory:

According to Ostwald's theory, due to the ionization of the acid-base indicator, the color changes. The color of the unionized form differs from that of the ionized form.

Quinonoid theory:

According to Quinonoid theory, an acid-base indicator exists in two tautomeric forms with different structures that are in equilibrium. One form is called the benzenoid form and the other is called the quinonoid form.

What is a universal indicator made of?

Generally, the universal indicator is usually composed of water (H2O), phenolphthalein (C20H14O4), 1-Propanol (C3H8O), methyl red (C15H15N3O2), sodium hydroxide (NaOH), thymol blue (C27H30O5S), sodium bisulfite (NaHSO3), and bromothymol blue (C27H28Br2O5S), etc.

Phenolphthalein, methyl red, thymol blue, and bromothymol blue are the major components of a universal indicator in the form of a solution. This mixture is significant because depending on the acidity or basicity of the solution being tested, each component loses or gains electrons.

Types of universal indicators:

A universal indicator can be a solution or a piece of paper that changes color in a solution and tells you how acidic or basic it is.

Solution:

Thymol blue, bromothymol blue phenolphthalein, and methyl red as a solution are the major components of universal indicators. This combination is crucial since each component loses or acquires protons which depend on the acidity or alkalinity of the solution being analyzed. In a colorless solution, this type of universal indicator is useful. This will improve the precision of the endpoint or indication.

There are several indicators available in the market where you can buy them, although a mixture of thymol blue, methyl red, bromothymol blue, and phenolphthalein is the most commonly used.

Paper form:

It's a piece of colored paper that turns blue when the solution is basic and turns red when the solution is acidic, e.g. litmus paper. To see the color change compared to the reference scale of 01–14, a drop of the sample solution can be placed on the pH paper strip or it may be immersed in the solution. When it comes into touch with the sample solution to be examined, it changes color depending on the pH of the sample solution.

Generally, both types of universal indicators are most commonly used in different classes or laboratory practical’s because they are a simple, fast, and reliable way of measuring acidity or alkalinity.

How does a universal indicator work?

The pH scale may be used to determine whether a solution is acidic or basic. The change of color of the pH indicator is due to the H+ ion being dissociated from the indicator itself. Remember that pH indicators include both natural dyes and weak acids. The weak acid indicator dissociates, causing the solution to change color.

Universal indicator color range or chart:
pH Range Category Colour
< 3 Strongly acidic Red
3 to 6 Weakly acidic Orange to Yellow
7 Natural Green
8 to 11 Weak alkali Blue
> 11 Strong alkali Indigo to Violet


How to perform universal pH test:

  • Liquid pH universal indicators are used by dropping a few drops of the solution into the substance being tested. Generally, it uses a 00.20 ml indicator for every 10.00 ml of sample to be tested.
  • For pH universal indicator papers (using a dropper) add a drop of the sample solution that can be placed on the pH paper strip or it may be immersed in the solution.
  • The acidity level of a substance is determined by the color that is produced as a result of the reaction.

Uses of universal indicators:

  • The basic use of universal indicators is to determine the strength of an acid or basic solution.
  • Universal indicators are more used compared to other indicators because of their wide range of results and applications.
  • It is commonly used in research, scientific laboratories, to teach in classes (practical) of science colleges, and pharmaceutical industry, etc.
  • It is used to determine the acidity or alkalinity of a wide range of substances, including food, sludge, soil, sewage, and water or beverage, etc.
  • It may be applied in different types of titration such as acid-base, redox, precipitation, and complexometric to determine whether a reaction (endpoint) is complete.

Advantages of universal indicator:

  • The advantage of a universal indicator is that it can indicate a wide range of colors on the pH scale of 01 to 14.
  • It works similarly to a pH meter in that there is no waiting for equilibrium and measurements are made quickly.
  • It is a simple, fast, and economical method that gives high accuracy results.


FAQ (Frequently Asked Questions):

What color would hydrochloric acid (pH 1) turn universal indicator?
Since it has a very strong pH, HCL causes the universal indicator to turn red.

What is the pH value of dilute HCL on the universal indicator?
Dilute hydrochloric acid has a pH value of 0-02 (acidic) which produces a red color when measured.

Difference between the indicator and universal indicator
Simple acid-base indicators can only be used to determine if a sample is a base or an acid, while the universal indicator is used to evaluate a sample's strength by indicating its pH value.

What are natural indicators?
Natural indicators are indications that can be acquired from plants. Examples: Turmeric, beetroot, red cabbage, cherries, and grape juice.

What is the color of universal indicator in neutral solution?
Green is the color of a universal indicator in a neutral solution.


References:


https://en.wikipedia.org/wiki/Universal_indicator
http://www.idc-online.com/technical_references/pdfs/Theory_of_indicators.pdf
Jorgensen,J.H., Pfaller , M.A., Carroll, K.C., Funke, G., Landry, M.L., Richter, S.S and Warnock., D.W. (2015) Manual., of Clinical Microbiology, 11th Edition. Vol. 1.
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