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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Preparation of α-Naphtholbenzein indicator solution

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

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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Preparation of KCl solution

Learn how to make different concentrations of molar and normal potassium chloride solutions, which are needed for many applications such as research, practical, pharmaceutical, chemical laboratory, industries, etc.

Potassium chloride, often referred to as KCl or potassium salt, is a metal halide salt composed of potassium and chlorine anions in the ratio of 1:1. It has the appearance of a vitreous crystal that is either white or colorless and odorless.

Generally, it is mostly used in the agriculture industry for the production of crop fertilizers, as well as is also applied in the field of medicine (Pharmaceutical), food processing, and in scientific applications.

Generally, a solid form of potassium chloride (KCl) in different-sized packs is supplied in the market by vendors in the form of white crystalline solid powder.

The molecular weight of KCl is 74.555 g·mol⁻¹

The melting point of KCL is 770°C

The density of KCl is 1.984 g/cm3

The solubility of KCl: Soluble in water, highly soluble in alcohols, and insoluble in ether

Table of Contents: Preparation of 0.02M potassium chloride Preparation of 0.05M potassium chloride Preparation of 0.1M potassium chloride Preparation of 0.2M potassium chloride Preparation of 0.25M potassium chloride Preparation of 0.5M potassium chloride Preparation of 1M potassium chloride Preparation of 2M potassium chloride Preparation of 3N potassium chloride Preparation of 3M potassium chloride Preparation of 2% potassium chloride Preparation of 5% potassium chloride Preparation of 7% potassium chloride Preparation of 10% potassium chloride Preparation of 20% potassium chloride

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.

Calculation method:

Example: To prepare 1 m KCl with a final volume of 250 ml, we can calculate using the formula below.

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

Mol KCl needed to make 250 mL of 1 M solution:

mol = 250 mL / 1000 mL/L x 1 mol /L = 0.25 mol KCl

Mass KCl required = 0.25 mol * 74.55 g/mol = 18.6375 g KCl

Therefore, to make a 01 M solution of KCl, dilute 18.63 gm of potassium chloride in 250 ml of distilled water.

Similarly, we can use the above method to determine the weight of KCl that is required to prepare any molarity/normality solution, as 0.1 N KCl is the same as 0.1 M KCl.

How to prepare a 0.02M KCl solution

Weigh accurately 01.49 gm of potassium chloride and dissolve in 250 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare a 0.05M KCl solution

Weigh accurately 3.73 gm of potassium chloride and dissolve in 250 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare 0.1M KCl solution

Weigh accurately 07.45 gm of potassium chloride and dissolve in 500 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare a 0.2M KCl solution

Weigh accurately 14.90 gm of potassium chloride and dissolve in 500 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare a 0.25M potassium chloride solution

Weigh accurately 18.63 gm of KCl and dissolve it in 750 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1 liter with distilled water, and properly mix it.

How to prepare 0.5M KCl solution

Weigh accurately 09.31 gm of KCl and dissolve it in 150 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 250 ml with distilled water, and properly mix it.

How to prepare 1M KCl solution

Weigh accurately 74.55 gm of potassium chloride and dissolve it in 100 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare 2M KCl solution

Weigh accurately 14.91 gm of KCl and dissolve it in 50 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 3M KCl solution

Weigh accurately 22.36 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 4m KCl solution

Weigh accurately 29.82 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 do you make a 2% KCl solution?

Weigh accurately 02.00 gm of potassium chloride and dissolve it in 50 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 a 5% KCl solution

Weigh accurately 05.00 gm of potassium chloride and dissolve it in 50 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 do you make a 7% potassium chloride solution?

Weigh accurately 07.00 gm of KCl solid powder and dissolve in 50 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 make 10% KCl solution

Weigh accurately 10.00 gm of KCl and dissolve 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.

How do you make a 20% solution of potassium chloride?

Weigh accurately 20.00 gm of potassium chloride and dissolve in 50 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.

PRECAUTIONS:

• KCl crystals are hygroscopic in nature and should be stored in a tightly closed, cool, and dry place.
• Stir a little amount of KCl into a large volume of water at a time, and then dilute the solution.
• When making KCl solutions, it is recommended that always use distilled water.
• Wear protective gloves, clothing, eye protection, and face protection.
• Wash your hands/skin thoroughly after handling.
• Handled the chemicals with care, follow laboratory safety measures (SOP), and please use extreme caution when preparing the solution concentrations.

References:

• Wikipedia contributors. (2022, August 27). Potassium chloride. In Wikipedia, The Free Encyclopedia. Available Here:
• USP Pharmacopoeia: USP29–NF24 Page 1762, Available Here:

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Preparation of sodium bicarbonate solution

Learn how to make different concentrations of molar and normal sodium bicarbonate solutions, which are needed for many applications such as research, practical, pharmaceutical, chemical laboratory, industries, etc.

Sodium bicarbonate is a chemical substance with the formula NaHCO3, also known as sodium hydrogen carbonate, monosodium salt, baking soda, or bicarbonate of soda. Sodium bicarbonate is a white, crystalline substance that often appears as a fine powder, which is a salt composed of bicarbonate anion (HCO3-) and sodium cation (Na+). Its taste is somewhat alkaline and salty, like washing soda (sodium carbonate:Na2CO3).

The molecules of sodium bicarbonate include one sodium cation and one bicarbonate anion. Here, an ionic bond is formed between the sodium ion, which has a positive charge, and the oxygen ion, which has a negative charge.

Generally, a solid form of sodium bicarbonate (NaHCO3) in different-sized packs is supplied in the market by vendors in the form of white crystals, and odorless powder.

The molecular weight of NaHCO3 is 84.0066 g mol−1

The melting point of NaHCO3 is 50°C.

The density of NaHCO3 is 2.20 g/cm3

The solubility of NaHCO3: soluble in water and insoluble in alcohol

Table of Contents: Preparation of 0.05M sodium bicarbonate Preparation of 0.1M sodium bicarbonate Preparation of 0.2M sodium bicarbonate Preparation of 0.25M sodium bicarbonate Preparation of 0.5M sodium bicarbonate Preparation of 1M sodium bicarbonate Preparation of 2M sodium bicarbonate Preparation of 3M sodium bicarbonate Preparation of 2% sodium bicarbonate Preparation of 5% sodium bicarbonate Preparation of 7% sodium bicarbonate Preparation of 10% sodium bicarbonate Preparation of 20% sodium bicarbonate

Requirements of glassware, chemicals, and apparatus:

Digital balance, beaker, pipette, pipette bulb, volumetric flask, measuring cylinder, glass rod, funnel, distilled water, AR/LR grade sodium bicarbonate (NaHCO3), etc.

Calculation method:

Depending on the volume you want to prepare, you can prepare 01 M of NaHCO3 and use the following formula:

M = g / V x molar mass

For example, if you want to make 1M in 0.5L of NaHCO3, you need to determine the amount of sodium bicarbonate in grams.

g = M x V x m.m

g = 01 x 0.5 x 84

g =42.0 g

Therefore, to make a 01 M solution of sodium bicarbonate, dilute 42.00 gm of NaHCO3 in 500 ml of distilled water.

How to prepare a 0.05M sodium bicarbonate solution

Weigh accurately 04.20 gm of NaHCO3 and dissolve in 250 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare 0.1M sodium bicarbonate solution

Weigh accurately 08.40 gm of NaHCO3 and dissolve in 500 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare a 0.2M sodium bicarbonate solution

Weigh accurately 16.80 gm of NaHCO3 and dissolve in 500 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare a 0.25M NaHCO3 solution

Weigh accurately 21.00 gm of sodium bicarbonate and dissolve it in 750 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1 liter with distilled water, and properly mix it.

How to prepare 0.5M sodium bicarbonate solution

Weigh accurately 10.50 gm of sodium bicarbonate and dissolve it in 150 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 250 ml with distilled water, and properly mix it.

How to prepare 1M sodium bicarbonate solution

Weigh accurately 84.00 gm of NaHCO3 and dissolve it in 100 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 1000 ml with distilled water, and properly mix it.

How to prepare 2M sodium bicarbonate solution

Weigh accurately 16.80 gm of sodium bicarbonate and dissolve it in 50 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 3M sodium bicarbonate solution

Weigh accurately 63.00 gm of NaHCO3 and dissolve it in 100 ml of distilled water in a volumetric flask. Once it has completely dissolved, make up the volume to 250 ml with distilled water, and properly mix it.

How do you make a 2% sodium bicarbonate solution?

Weigh accurately 02.00 gm of NaHCO3 and dissolve it in 50 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 5% sodium bicarbonate solution

Weigh accurately 05.00 gm of NaHCO3 and dissolve it in 50 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 do you make a 7% sodium bicarbonate solution?

Weigh accurately 07.00 gm of NaHCO3 solid powder and dissolve in 60 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 make 10% sodium bicarbonate solution

Weigh accurately 10.00 gm of sodium bicarbonate and dissolve 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.

How do you make a 20% solution of sodium bicarbonate?

Weigh accurately 20.00 gm of NaHCO3 and dissolve in 50 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.

PRECAUTIONS:

• Sodium bicarbonate is non-hygroscopic in nature it should be stored in a tightly closed, cool, and dry place.
• Stir a little amount of sodium bicarbonate into a large volume of water at a time, and then dilute the solution.
• When making sodium bicarbonate solutions, it is recommended that always use distilled water.
• Wear protective gloves, clothing, eye protection, and face protection.
• Wash your hands/skin thoroughly after handling.
• Handled the chemicals with care, follow laboratory safety measures (SOP), and please use extreme caution when preparing the solution concentrations.

References:

• Indian Pharmacopoeia 1996
• Wikipedia contributors. (2022, August 23). Sodium bicarbonate. In Wikipedia, The Free Encyclopedia. Available Here:
• ‘Preparation of Standard Solution of Sodium bicarbonate - Chemistry Practicals Class 11’. BYJUS, Available Here:
• ‘How to Make a .5 Solution of Sodium Bicarbonate’. Sciencing, Available Here:

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