Difference between natural and synthetic indicators

The major difference between natural and synthetic indicators is that natural indicators are those that come from or are made of substances that occur naturally, whereas synthetic indicators are compounds that are manufactured in a lab from artificial substances.

Compounds that change colour when added to acidic or alkaline solutions are called indicators. Indicators are often used to determine pH. Colored indicators can be added to the reaction mixture to determine the endpoint or equivalence point of different titration methods, such as acid-base.

Indicators are either weak acids or weak bases that change colour when the amount of hydrogen ions in a solution or its pH changes. In water, the indicators dissociate slightly to form ions. In chemistry, there are primarily two types of indicators: natural indicators and artificial indicators.

Natural indicators:

Natural indicators are naturally occurring substances that can be used to determine whether a substance is acidic or basic. These naturally occurring indicators are used to identify hydrogen ions (H+) and hydroxyl ions (OH-) in a sample solution.

Examples of natural indicators: 

Turmeric, tomato, litmus, china-rose, clove oil, curry powder, red cabbage, vanilla extract, onion, grape juice, cherries, etc.

 

Natural indicator	Color in acid	Color in base
Turmeric	Yellow	Brick red
Beet root	Pink	Pale yellow
Red cabbage	Red	Green

Artificial indicators:

Indicators that do not occur naturally but are produced artificially in the laboratory or obtained through a chemical reaction are considered synthetic indicators. These synthetic indicators are used for titration, identifying acids and bases, and determining the pH of various solutions.

Examples of artificial indicators: 

Methyl orange, phenolphthalein, thymol blue, litmus paper, malachite green, etc.

Synthetic indicator	Color in acid	Color in base
Litmus	Red	Blue
Phenolphthalein	Colorless	Pink
Methyl Orange	Pink	Yellow

Keywords:
Define, 10 examples, methyl orange, litmus paper, class 9, class 10 students 

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Preparation of dimethyl yellow indicator solution

Learn the procedure for making a dimethyl yellow indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Dimethyl yellow (C14H15N3) is an organic compound chemically described as 4-(Dimethylamino)azobenzene, N,N-Dimethyl-4-(phenylazo)aniline, Butter yellow, Methyl yellow. It is used as an acid-base indicator, as well as an indicator for non-aqueous titration which changes from yellow to red through a pH range pH 2.9 to pH 4.00.

Dimethyl yellow comes in the form of a yellow to orange-brown powder that is soluble in alcohol, benzene, chloroform, ether, and petroleum but, insoluble in water.

How to prepare dimethyl yellow indicator mixture for titration:

• Accurately weigh 100 mg of calcein and pour it into a 100 mL volumetric flask with 50 ml of rectified spirit.
• Once it is dissolved, dilute it to 100.00 mL with rectified spirit.
• The concentration of the prepared solution is about 01 g/L.

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Preparation of calcein indicator solution

Learn the procedure for making a calcein indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Calcein (C30H26N2O13), which is also called fluorexon or fluorescein complex, is a fluorescent dye with wavelengths of 495 and 515 nm for excitation and emission, respectively. It is used as a complexometric indicator for the titration of calcium ions with EDTA and fluorometric determination of calcium. 

Calcein is a fluorescent metal indicator, chemically described as 2, 7 Bis[bis(carboxymethyl)aminomethyl]fluorescein. Calcein comes in the form of a yellowish-orange powder that is slightly soluble in water.

In acidic conditions, a calcein solution emits a yellowish-green fluorescence, whereas there is no fluorescence in basic conditions. However, under basic conditions, Calcein will emit fluorescence in the presence of metal ions such as Al, Ba, Ca, Cu, Mg, and Zn.

How to prepare calcein indicator mixture for titration:

To make a fine powder that can be used as an indicator, mix 1 gram of calcein with 100 gm of potassium chloride.

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Preparation of calcon indicator solution

Learn the procedure for making a calcon indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Calcon (C20H13N2NaO5S) is an azo dye chemically described as, 2-Hydroxy-1-(2-hydroxy-1-naphthylazo) naphthalene-4-sulfonic acid sodium salt that is used as an indicator (metal indicator) for complexometric titrations of calcium with ethylenediaminetetraacetic acid (EDTA) in the presence of magnesium.

Calcon comes in the form of a dark, blackish-brown powder that is soluble in water, ethanol, and methanol. It produces a purple color when calcium ions in an alkaline solution are present and a blue color when these ions are absent.

How to prepare calcon indicator for titration:

• Accurately weigh 100 mg of calcon and pour it into a 100 mL volumetric flask with 50 ml of water.
• Once it is dissolved, dilute it to 100.00 mL with water.
• The concentration of the prepared solution is about 01 g/L.

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Preparation of dithizone indicator solution

Learn the procedure for making a dithizone indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Dithizone (C13H12N4S) is a sulfur-containing organic compound chemically described as 1,5-diphenylthiocarbazone, which is used as a complexometry/metal indicator. It comes in the form of crystalline black powder which is soluble in ethanol but insoluble in water.

It is a good ligand and can form complexes with several toxic metals, including lead, thallium, and mercury. Additionally, it acts as a chelating agent for toxic metals including lead and mercury.

How to prepare dithizone indicator for titration:

• Accurately weigh 25 mg of dithizone and pour it into a 100 mL volumetric flask with 50 ml of ethanol.
• Once it is dissolved, dilute it to 100.00 mL with ethanol.
• The concentration of the prepared solution is about 01 g/L.

Dithizone has a pKa of 1:4.50;pK2:15 (25°C), When used as an indicator, it produces a pink color in a pH range of 4 to 5 when zinc ions are present, and a green color when these ions are absent.

References:

Wikipedia contributors. (2022, August 9). Dithizone. In Wikipedia, The Free Encyclopedia. Available Here:

Law.resource.org. 2022. [online] Available Here:

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Preparation of pyridylazo naphthol indicator solution

Learn the procedure for making pyridylazo naphthol indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Pyridylazonaphthol (C15H11N3O) is an orange-colored dye chemically described as 1-(2-Pyridylazo)-2-naphthol (PAN) which is used as an acid-base/complexometry/metal indicator. It comes in the form of orange powder which is soluble in ethanol but insoluble in water.

When copper ions are present in an acidic solution, it produces a yellow color. However, when these ions are absent, the solution becomes red. It is a useful indicator in complexometric titrations because it may form chelates with metal ions.

How to prepare pyridylazo naphthol indicator for titration:

• Accurately weigh 100 mg of pyridylazo naphthol and pour it into a 100 mL volumetric flask with 50 ml of methyl alcohol.
• Once it is dissolved, dilute it to 100.00 mL with methyl alcohol.
• The concentration of the prepared solution is about 01 g/L.

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Preparation of sodium alizarin sulfonate indicator solution

Learn the procedure for making sodium alizarin sulfonate indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. To find the endpoint or equivalence point of the titration, they are added to the reaction mixture.

Sodium alizarin sulfonate is a chemical compound, chemically described as sodium 1,2-dihydroxy-9, 10 anthraquinone-3-sulphonate, used as a metal indicator. It comes in the form of yellow-orange powder which is soluble in water and ethanol.

At a pH of about 4.0, it produces a bluish-red lake with thorium and aluminum ions, However, the solution turns yellow when these ions are absent.

How to prepare sodium alizarin sulfonate indicator for titration:

• Accurately weigh 100 mg of sodium alizarin sulfonate and pour it into a 100 ml volumetric flask with 50 ml of distilled water.
• Once it is dissolved, dilute it to 100.00 mL with distilled water.
• The concentration of the prepared solution is about 01 g/L.

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Preparation of alkali blue 6B indicator solution

Learn the procedure for making an alkali blue 6B indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. They are added to the reaction mixture to find the endpoint or equivalence point of the titration.

Alkali blue 6B (C37H29N3O3S) is a chemical compound, chemically described as sodium para-rosaline mono sulphonate.

Alkali blue 6B has been used as a pH indicator for the non-aqueous titration method, which changes color from blue to greenish-red and its pH range is 9.4 (Blue) to pH 14.0 (Red). Alkali Blue 6B is a solid that is a dark blue powder. It is soluble in ethanol, but not in water.

How to prepare dichlorofluorescein indicator for titration:

• Accurately weigh 2 g of Alkali blue 6B and pour it into a 100 ml volumetric flask with 50 ml of rectified spirit.
• Once it is dissolved, dilute it to 100.00 mL with rectified spirit.
• The concentration of the prepared solution is about 20.0g/L.

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Learn the procedure for making 1-Naphtholbenzein indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Most of the time, color indicators are used to measure pH. They are added to the reaction mixture to find the endpoint or equivalence point of the titration.

α-Naphtholphthalein (C28H18O4) is a dye, chemically describes as 4, 4′-(α-Hydroxybenzylidene)di-1-naphthol, p-naphtholbenzein

α-Naphtholbenzein has been used as a pH indicator during the acid-base titration method which changes color from colorless/reddish to greenish blue at pH 7.3–8.7. It changes from green (basic) to orange (neutral) to yellow (acidic) when used as an indicator for non-aqueous titrations.

Naphtholbenzein is a solid that is colorless to slightly reddish or blue-green. It dissolves in ethanol, acetic acid, and acetone, but not in water.

How to prepare dichlorofluorescein indicator for titration?

• Accurately weigh 0.2 g of 1-Naphtholbenzene and pour it into a 100 ml volumetric flask with 50 ml of acetic acid.
• Once it is dissolved, dilute it to 100.00 mL with acetic acid.
• The concentration of the prepared solution is about 2.0 g/L.

References:

1. Wikipedia contributors. (2022, April 3). Naphtholphthalein. In Wikipedia, The Free Encyclopedia. Available Here:
2. Law.resource.org. 2022. [online] Available Here:

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Preparation of dichlorofluorescein indicator solution

Learn the procedure for making dichlorofluorescein indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Color indicators are commonly used to measure pH and are added to the reaction mixture to identify the titration endpoint or equivalence point.

Dichlorofluorescein (C20H10Cl2O5) is an organic dye of the fluorescein family, which is chemically described as 2′, 7′-Dichloro-3, 6-fluorandiol.

In Fajan’s method, it is used as an indicator for argentometric titration and as adsoprtion indicator. The color of the titration reaction changes from colorless to pale pink on reaching the equivalence point.

Dichlorofluorescein usually comes in the form of crystalline powder, in the color of orange to red-brown powder, which is soluble in ethanol and methanol but insoluble in water.

How to prepare dichlorofluorescein indicator for titration:

• Weigh 0.1 gm of dichlorofluorescein accurately and pour it into a 100.00 ml volumetric flask containing 50.00 ml of rectified spirit.
• Once it is dissolved, dilute it to 100.00 ml with rectified spirit.
• The concentration of the prepared solution is about 1.0 g/L.

References:

1. Wikipedia contributors. (2022, June 30). Dichlorofluorescein. In Wikipedia, The Free Encyclopedia. Available Here:
2. Law.resource.org. 2022. [online] Available Here:

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Preparation of phenosafranine indicator solution

Learn the procedure for making phenosafranine indicator solution.

The compound that changes color when exposed to acidic or basic solutions is called an indicator. Color indicators are commonly used to measure pH and are added to the reaction mixture to identify the titration endpoint/equivalence point.

Phenosafranine (C18H15ClN4) is a compound or dye, chemically described as 3,7-Diamino-5-phenylphenazinium chloride.

Phenosafranine is a phenylphenazinium bacterial stain for microscopy that is used as an indicator in adsorption and redox titration.

It turns the precipitate red when used in the titration of chloride or bromide with silver nitrate in an acid solution. Once the precipitate is complete, the color of the precipitate changes to blue color.

Phenosafranine usually comes in the form of powder, crystals, or flakes, and its appearance of dark green powder which is soluble in water.

How to prepare phenosafranine indicator for titration:

• Weigh accurately 0.25 gm of phenosafranine and pour it into a 100.00 ml volumetric flask containing 50.00 ml of distilled water and mix it.
• Once it is dissolved, dilute it to 100.00 ml with distilled water.
• The concentration of the prepared solution is about 2.5 g/L.

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Why KCl is used for the calibration of conductivity meter

Potassium chloride, often known as KCl, is the most commonly used solution for calibrating conductivity meters due to its high stability and solubility.

If you want reliable and repeatable results from the conductivity meter, calibration using pre-prepared standard solutions is essential. The most common calibration solution is potassium chloride (KCl). Calibration of a conductivity meter is an important maintenance task. Regular calibration produces consistent and reliable readings.

Electrical conductivity (EC) is the measure of the ion concentration present in the sample. The ability of a substance to transmit electric current in a certain area is known as conductivity. Siemens (S) is the unit of measurement for electrical conductivity, and scientific instruments commonly show measured units as MilliSiemens per centimeter mS/cm or MicroSiemens per cm μS/cm.

Because of its solubility and stability, KCl is the solution most often used in the process of calibrating the conductivity meter. Conductivity standard solutions are made up of a KCl: Water ratio. The mixing ratio is determined by the desired ion concentration level of the standard solution.

Molarity/Concentration of KCl	Standard conductivity
0.001 M	147 µS/cm ±10% at 25°C
0.01 M	1.413 mS/cm ±10% at 25°C
0.1 M	12.88 mS/cm ±10% at 25°C

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How to prepare KCl solution for conductivity

Learn how to make potassium chloride solution for calibration of conductivity meter which is needed in many applications such as research, practical, pharmaceutical, chemical laboratory, industries, etc.

The conductivity meter is a very significant tool in several applications for measuring conductance, it measures the level of conductivity in solutions. To provide reliable readings, the conductivity cell must be in proper working condition.

A few factors that could shorten the lifespan of the conductivity cell include sample composition, temperature, and how the conductivity cell is stored and maintained. Care and preservation of conductivity cells ensure not just a longer lifespan but also more precise readings.

Requirements of glassware, chemicals, and apparatus:

Digital balance, beaker, pipette, pipette bulb, volumetric flask, measuring cylinder, glass rod, funnel, distilled water, AR/LR grade potassium chloride (KCl), etc.

How to prepare calibration solution for conductivity meter:

Calculation method:

Molar mass KCl = 39.1+35.45 = 74.55 g/mol,

Mol KCl needed to make 100 mL of 0.01 M solution:

mol = 100 mL / 1000 mL/L x 0.1 mol /L = 0.01 mol KCl

Mass KCl required = 0.01 mol * 74.55 g/mol = 0.7456 g KCl

Therefore, to make a standard KCl solution for conductivity (0.01 M KCl, 1411 μS at 25°C), weigh accurately 0.7456 gm of potassium chloride and dissolve it in 500 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 01 liter with distilled water, and properly mix it.

The standard value of conductance at coefficient-1.92, K=cell constant:

Molarity/Concentration of KCl	Standard conductivity
0.001 M	147 µS/cm ±10% at 25°C
0.01 M	1.413 mS/cm ±10% at 25°C
0.1 M	12.88 mS/cm ±10% at 25°C

Why do we use KCl for calibration of the conductivity meter?

The conductivity meter should be frequently calibrated using pre-prepared standard solutions to get accurate and reproducible results. Potassium chloride (KCl) is the most common solution to perform the calibration of the conductivity meter because it is soluble and very stable.

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Why KCl is used in pH meter

Why KCl is used in pH meter

The basic parts of a pH meter are the glass electrode, reference electrode, electrolyte solution, temperature sensor, etc. It is the most commonly used instrument in our lab for research, industrial or practical purposes.

To obtain accurate results, it is necessary to calibrate the device regularly and keep the pH electrode moist. Therefore a manufacturer advises storing the pH electrode in 4M KCl or pH 04 solutions and it does not require storing in distilled or deionized water.

Why KCl solution used in pH meter?

To understand the reason for the use of KCl as an electrolyte, it is required to know the role of the electrolytic solution in a pH meter. The pH meter operates by immersing a glass electrode and a reference electrode in an electrolytic solution, which is then connected to the test solution via a porous ceramic membrane. The voltage is then displayed on a voltmeter which provides a pH reading.

Therefore, we need a solution in an electrolytic solution that contains enough ions to complete the circuit and does not change the pH of the test solution, as it is in a way, connected to it.

For this reason, we use a solution of KCl because it contains a good source of ions in the form of chlorine (Cl) ions. Furthermore, because KCl is neutral, it does not affect or change the pH of the test solution. Because of these two factors, potassium chloride (KCl) is what is used in pH meters.

Why electrodes are only put in a KCl solution?

Potassium chloride (KCl) and sodium chloride (NaCl) are typical electrolytes in a salt bridge. Potassium chloride is preferred over NaCl because of its greater mobility of K+ ions. Transport number shows which ion would move faster than the other. As a result, KCl gives the electrolyte used in salt bridges the least chance of ionic imbalance, or the least chance of retardation.

Why is a calomel electrode kept in a KCl solution?

Because the calomel electrode is made of 1M KCl or a saturated KCl solution, maintaining the electrode in the KCl solution helps prevent the electrolyte to come out or change the concentration of the cell.

Why is the top layer of a calomel electrode filled with saturated KCl?

In the calomel electrode, KCl acts as a salt bridge. Electrolytes that have anion and cation mobility that is equal, or substantially equal, are used as salt bridges.

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How to prepare saturated KCL solution

Potassium chloride, also known as KCl or potassium salt, is a 1:1 mixture of potassium and chlorine ions in a metal halide salt. It looks like a white or colorless and odorless vitreous crystal.

The molecular weight of KCl is 74.555 g/mol−1, its density is 1.984 g/cm3, and it is soluble in water, highly soluble in alcohols, and insoluble in ether.

Requirements of glassware, chemicals, and apparatus:

Beaker, volumetric flask, glass rod, funnel, distilled water, AR/LR grade potassium chloride (KCl), spatula, etc.

How to prepare saturated solution of KCl

Using KCl and water, you can make a supersaturated solution by the following procedure:

• Take a 250 ml beaker; add 100 ml of distilled water to it, then add KCl and mix well to dissolve.
• Once it has completely dissolved, add more salt to the mixture and keep stirring.
• At a temperature of 20°C, it can be noticed that 35.00 g of potassium chloride can be completely dissolved, and the solution becomes saturated.
• If we increase the temperature, more amount of the solute will be able to dissolve in the solution.
• If we increase the temperature of the solution to 80 °C the excess solute separates out in the form of crystals.

What is the concentration of saturated KCl?

The temperature-dependent solubility of saturated KCl determines its concentration.

• KCl solubility at 20°C: 0,340 g.cm-3 in water = 340 g/l = 4.56 mol/l (Saturated KCl at 20°C)
• KCl solubility at 30°C: 0,374 g.cm-3 in water = 374 g/l = 5.02 mol/l (Saturated KCl at 30°C)

How to store saturated KCl solution?

The saturated and supersaturated KCl solutions must be kept at room temperature in a closed container. There is no need to be concerned about the growth of microorganisms in the solution as it has a very high osmotic pressure.

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How to prepare KCl solution for pH meter

Learn how to make potassium chloride solution for filling and storing electrodes of pH meter which are needed many applications such as research, practical, pharmaceutical, chemical laboratory, industries, etc.

The pH electrode is a very significant tool in several applications for measuring pH, it measures a solution's ionic potential, or how much charge is moving within a solution. This data is then converted into a pH value by the pH meter.

In order for the meter to provide reliable readings, the pH electrodes must be in proper working condition. The average lifetime of an electrode is between six months and two years, depending on its use and frequency of application.

A few factors that could shorten the lifespan of the pH electrode include sample composition, temperature, and how the electrode is stored and maintained. Care and preservation of pH probes ensure not just a longer lifespan but also more precise readings.

Requirements of glassware, chemicals, and apparatus:

Digital balance, beaker, pipette, pipette bulb, volumetric flask, measuring cylinder, glass rod, funnel, distilled water, AR/LR grade potassium chloride (KCl), etc.

How to prepare 3M KCl filling solution for the electrode of the pH meter:

Calculation method: We can calculate using the formula below.

Molar mass KCl = 39.1+35.45 = 74.55 g/mol,

Mol KCl needed to make 100 mL of 3 M solution:

mol = 100 mL / 1000 mL/L x 3 mol /L = 0.30 mol KCl

Mass KCl required = 0.30 mol * 74.55 g/mol = 22.635 g KCl

Therefore, to make a 3M solution of KCl, weigh accurately 22.635 gm of potassium chloride and dissolve it in 70 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 100 ml with distilled water, and properly mix it.

How to prepare pH electrode storage solution?

To prepare 4M potassium chloride electrode storage solution, weigh accurately 29.82 gm of KCl and dissolve it in 80 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 100 ml with distilled water, and properly mix it.

Tips:

• Always keep the pH electrode bulb moist to prevent clogging and for a fast response time.
• Keep the pH electrode bulb moist when not in use by pouring electrode storage solution into the cap and then placing the cap over the bulb.
• Frequently calibrate pH meter with standard buffer solutions such as pH 07, 04, and 09.20.
• Electrode manufacturers suggested that store the electrode in a solution of 4 M KCl solution if you don't have 4M KCl, you can use a pH 4 buffer solution.
• Do not store the electrode in distilled or deionized water, as this can cause ions to leach out from the glass bulb, rendering the electrodes useless.

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