LABORATORY REAGENTS ACCORDING TO SCIENCE LABORATORY TESTS AND STANDARD PROCEDURES

Laboratory reagents

Laboratory reagents are substance used to carry out a laboratory test. They can be classified based on:

1. Purity.

• Technical Grade Chemicals: used for low-grade applications like commercial or industrial purposes. Due to the present impurities, it isn’t utilized for drug, food, or medicinal purposes.

• Synthesis Grade Chemicals: The primary purpose of this involves organic synthesis and preparative tasks.

• Lab Grade Chemicals: found in educational or teaching labs. Though their purity levels are high, the precise impurity levels remain anonymous.

• AR grade Chemicals: These are used for high precision work since trace impurities are restricted to the lowest possible limits for high precision. Such reagents get used mainly for analytical applications, research, and quality control.

• General Reagent (GR): reagent that meets or exceeds AR grade specifications.

• Extra Pure Grade Chemicals: suitable for laboratory accreditations and also work requiring compliance with pharmacopoeia standard requirements.

2. Applications.

• Electronic Grade Chemicals: have very stringent limits for metallic impurities as required for use in electronic component industry as such as below ppt or ppb levels

• HPLC Grade Chemicals: adequately pure ion-pair reagents, solvents, and buffers used as the mobile phase of High-Performance Liquid Chromatography (HPLC).

• Spectroscopy Grade Chemicals: includes solvents of high purity, low residue on boiling, and having absorption blank in the wavelength region of interest. Spectroscopy grade salts consist of alkali metal salts having transparency in the IR region such as potassium bromide, sodium chloride etc.

Methods of preparing solutions

• Solution: A uniform homogeneous mixture of two or more substances. The individual substances may be present in varying amounts.

• Solute: The substance which is dissolved, or has gone into solution (typically a solid).

• Solvent: The substance which does the dissolving (typically a liquid, such as water or alcohol). Must be greater than 50% of the solution.

• Standard Solution: A very precise solution, usually to 3–4 significant figures, used in quantitative analysis or an analytical procedure.

• Saturated Solution: A solution that contains the maximum amount of a particular solute that will dissolve at that temperature.

• Supersaturated Solution: A solution that contains more solute than equilibrium conditions allow; it is unstable and the solute may precipitate upon slight agitation or addition of a single crystal.

Solutions can be prepared from the following:

1. Simple inorganic salts

The guidelines on solubility of inorganic salts are:

• Nitrates (NO3–): All nitrates are soluble.

• Acetates (C2H3O2–): All acetates are soluble; silver acetate is moderately soluble.

• Bromides (Br–) Chlorides (Cl–) and Iodides (I–): Most are soluble except for salts containing silver, lead, and mercury.

• Sulfates (SO42–): All sulfates are soluble except barium and lead. Silver, mercury(I), and calcium are slightly soluble.

• Hydrogen sulfates (HSO4–): The hydrogen sulfates (aka bisulfates) are more soluble than the sulfates.

• Carbonates (CO32–), phosphates (PO43–), chromates (CrO42–), silicates (SiO42–): All carbonates, phosphates, chromates, and silicates are insoluble, except those of sodium, potassium, and ammonium. An exception is MgCrO4, which is soluble.

• Hydroxides (OH–): All hydroxides (except lithium, sodium, potassium, cesium, rubidium, and ammonium) are insoluble; Ba (OH)2, Ca(OH)2 and Sr(OH)2 are slightly soluble.

• Sulfides (S2–): All sulfides (except sodium, potassium, ammonium, magnesium, calcium and barium) are insoluble. Aluminum and chromium sulfides are hydrolyzed and precipitate as hydroxides. Sodium (Na+), potassium (K+), ammonium (NH4+): All sodium, potassium, and ammonium salts are soluble (Except some transition metal compounds.)

• Silver (Ag+): All silver salts are insoluble. Exceptions: AgNO3 and AgClO4; AgC2H3O2 and Ag2SO4 are moderately soluble.

A solvent will only dissolve a limited quantity of solute at a definite temperature. However, the rate at which the solute dissolves can be accelerated by the following methods:

• Pulverize or grind up the solid to increase the surface area of the solid in contact with the liquid.

• Heat the solvent. This will increase the rate of solution because the molecules of both the solvent and the solute move faster.

• Stir vigorously.

2. Acid solutions

Methodology

When making acid solutions, always ADD ACID (AA) to water since a great amount of heat is liberated when acid is added to water. The temperature of the solution will rise rapidly.

The temperature control can be controlled by immersing your mixing vessel in a bucket of ice. Always add the acid to water very slowly while stirring continuously.

Safe storage of acids

Corrosive chemicals, such as strong acids and bases, must be isolated from other chemicals to prevent accidental contact and hazardous reaction conditions. The best way to isolate your corrosive chemicals is to store them in an approved corrosive storage cabinet e.g. an acid cabinet. Locked storage cabinets also provide security against theft and vandalism.

3. Base solutions

In diluting alkalis, sodium hydroxide, a great amount of heat is liberated. The temperature of the solution may rise very rapidly and possibly spatter a hot, caustic solution.

Immerse the flask or beaker in an ice–water bath to control the solution temperature and add ingredients slowly with continuous stirring.

Safety equipment needed during solution preparation:

• Eyewash/Body Drench

• Spill Control and Clean-up Materials

• Fire Extinguisher

• Chemical-resistant Gloves and Aprons

• Chemical Splash Goggles

• Chemical First Aid Kit

Equipment needed during solution preparation:

• Electronic Balance

• Magnetic Stirrers

• Volumetric Flasks

• Graduated Cylinders

• Water Purification System

• Bottles

• Labels

Preparing solutions

Definitions

• Concentration: The relative amount of solute and solvent in a solution.

• Molality: A concentration unit (m); defined as the number of moles of solute divided by the number of kilograms of solvent.

• Molar Mass: The mass of a mole of any element or compound.

• Molarity: A concentration unit (M); defined as the number of moles of solute divided by liters of solution.

• Normality: A concentration unit (N); defined as the number of equivalents of solute per liter of solution. (e.g., 1 M H2SO4 = 2 N H2SO4)

1. Molarity vs Normality

Molarity: number of moles of solute per liter of solution. For example, a 1 M solution of H2SO4 contains 1 mole of H2SO4 per liter of solution.

Normality is a measure of concentration that is equal to the gram equivalent weight per liter of solution. Gram equivalent weight is a measure of the reactive capacity of a molecule.

Example:

H2SO4 dissociates into H+ and SO4- ions in water. For every mole of H2SO4 that dissociates in solution, 2 moles of H+ and 1 mole of SO4- ions are formed.

For acid reactions, a 1 M H2SO4 solution will have normality (N) of 2 N because 2 moles of H+ ions are present per liter of solution. For sulfide precipitation reactions, where the SO4- ion is the most significant factor, the same 1 M H2SO4 solution will have a normality of 1 N.

Converting From Molarity to Normality

N = M*n where n is the number of equivalents

2. Percent solutions

• Mass percent solutions are defined based on the grams of solute per 100 grams of solution.

Example: 20 g of sodium chloride in 100 g of solution is a 20% by mass solution.

• Volume percent solutions are defined as milliliters of solute per 100 mL of solution.

Example: 10 mL of ethyl alcohol plus 90 mL of H2O (making approx. 100 mL of solution) is a 10% by volume solution.

• Mass-volume percent solutions are also very common. These solutions are indicated by w/v% and are defined as the grams of solute per 100 milliliters of solution.

Example: 1 g of phenolphthalein in 100 mL of 95% ethyl alcohol is a 1 w/v% solution.