Qualitative analysis of salts lab manual MANIK

Laboratory Manual Course Code: 2114

Course Title: Inorganic Pharmacy II Lab

Prepared ByShadiduzzaman and Md. Imran Nur Manik

Lecturer

Department of Pharmacy

Northern University Bangladesh

Edited BySomaia Haque Chadni

LecturerDepartment of Pharmacy

Northern University Bangladesh

Inorganic Pharmacy II Lab

Prepared By: and

Edited By:

Lecturer; Department of Pharmacy; Northern University Bangladesh (NUB).

INDEX

Sl.

No. Date Name of the experiment

Page

No.

01. Conversion of different water insoluble or sparingly soluble

substances into water soluble form. 01 – 06

02.

Safety First 07 – 07

Qualitative analysis (identification) of anions from inorganic salt

solutions. 08-11

Qualitative analysis (identification) of group I & group II cations. 12 – 15

Qualitative analysis (identification) of group III cation. 16 – 20

Qualitative analysis (identification) of group IV & group V cations. 21 – 23

Md. Imran Nur Manik


Prepared By: Shadiduzzaman and Md. Imran Nur Manik Edited By: Somaia Haque Chadni Page 1


Experiment No. 01

Name of the Experiment: Conversion of different water insoluble or sparingly soluble

substances into water soluble form.

Principle:

Solubility, the phenomenon of dissolution of solute in solvent to give a homogenous

system, is one of the important parameters to achieve desired concentration of drug in

systemic circulation for desired (anticipated) pharmacological response. Low aqueous

solubility is the major problem encountered with formulation development of new

chemical entities as well as for the generic development. More than 40% new chemical

entities developed in pharmaceutical industry are practically insoluble in water. Any drug

to be absorbed must be present in the form of solution at the site of absorption. Various

techniques are used for the enhancement of the solubility of poorly soluble drugs which

include physical and chemical modifications of drug and other methods like particle size

reduction, crystal engineering, salt formation, solid dispersion, use of surfactant,

complexation, and so forth. Selection of solubility improving method depends on drug

property, site of absorption, and required dosage form characteristics.

Terminologies:

A solution may be defined as a homogeneous mixture of two or more components that

form a single phase.

The component that determines the phase of the solution is termed the solvent and

usually constitutes the largest proportion of the system. The other components dispersed

as molecules or ions throughout the solvent are termed solutes.

The transfer of molecules or ions from a solid state into solution is known as dissolution.

The extent to which the dissolution proceeds under a given set of experimental conditions

is referred to as the solubility of the solute in the solvent.

Solution Process In order to dissolve a solid, the particles of the solvent must be able to separate the particles of the solute and occupy the intervening space. This happens when the positive end of the solvent molecule approaches the negative end of the solute molecule. The solute molecule is pooled in the solution when the force between the solvent solute molecules overcomes the attractive forces between the solute molecules themselves. Therefore, the solute will be soluble in the solvent if the solute-solvent attraction is stronger than the solvent-solvent attraction and solute-solute attraction.

Figure: A molecular view of the solution process. The

solute molecules and solvent molecules first being

spread apart, and then being mixed together. The

relative strength of forces holding solvent molecules

togetherH1, solute particles together H2, and the

forces between solvent and solute molecules H3 in the

solution are important in determining the solubility.

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An ionic solid dissolves in water because the number of water molecules around the surface is greater than the number of other ions of the solid. The attraction between polar water molecules and a charged ion enables the water molecules to pull ions away from the crystal, a process called dissolving.

A General Solubility Rule:

Polar molecules are soluble in polar solvents e.g. C2H5OH in H2O. Thus polar compounds, like table sugar (C12H22O11), are soluble in polar solvents and insoluble in non-polar solvents. On the other hand non-polar molecules are soluble in non-polar solvents e.g. CCl4 in C6H6.

Accordingly non-polar compounds, like naphthalene (C10H8), are soluble in non-polar solvents and insoluble in polar solvents. This it the like dissolves like rule.

Factors Affecting Solubility of a Solute

1. Temperature Generally in many cases solubility increases with the rise in temperature and decreases with the fall of temperature but it is not necessary in all cases. However we must follow two behaviours: In endothermic process solubility increases with the increase in temperature and vice versa. For example: solubility of potassium nitrate increases with the increase in temperature. In exothermic process solubility decrease with the increase in temperature. For example: solubility of calcium oxide decreases with the increase in temperature.

2. Effect of pressure

The effect of pressure is observed only in the case of gases. An increase in pressure increases of solubility of a gas in a liquid. For example CO2 is filled in cold drink bottles (such as coca cola, Pepsi 7up etc.) under pressure.

3. Chemical natures of the solute and solvent A polar solute will dissolve in a polar solvent but not in a non-polar solvent. The adage "like dissolves like" is very useful. Example: Alcohol (polar substance) dissolves in water (polar substance). Water (polar substance) does not dissolve in oil (non-polar substance)

4. Stirring With liquid and solid solute, stirring brings fresh portions of the solvent in contact with the solute, thereby increasing the rate of dissolution.

5. Amount of solute already dissolved

When there is little solute already in solution, dissolving takes place relatively rapidly. As the solution approaches the point where no solute can be dissolved, dissolving takes place more slowly.

6. Molecular size The larger the molecules of the solute are, the larger is their molecular weight and their size. If the pressure and temperature are the same than out of two solutes of the same polarity, the one with smaller particles is usually more soluble.

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Methods for Expression of Concentration

Expression Symbol Definitions

Mole Fraction X Ratio of moles of one constituent of a solution to the

total moles of all constituents (solute & solvent)

Mole percent Moles of one constituent in 100 moles of the solution.

It is obtained by multiplying X by 100

Percent by weigh %W/W Gram of solute in 100 gm solution

Percent by volume %V/V ml of solute in 100 ml of solution

Percent by weight-in-volume

%W/V Gram of solute in 100 ml solution

Milligram percent mg of solute in 100 ml of solution

Parts By Mass Factor tionMultiplicaSolventofMass

SoluteofMass

Parts per million ppm Multiplication Factor =106

Parts per billion ppb Multiplication Factor =109

Solubility Definitions

Description Parts of Solvent required to dissolve 1 part of Solute

Very Soluble Less than 1 part

Freely Soluble from 1 to 10 parts

Soluble from 10 to 30 parts

Sparingly Soluble from 30 to 100 parts

Slightly Soluble from 100 to 1000 parts

Very Slightly Soluble from 1000 to 10,000 parts

Practically Insoluble more than 10,000 parts

Approximate quantity of solvent by volume for one part of soluble by weight. For example, 1g of a very soluble substance dissolves in less than 1ml of solvent. Importance of Solubility Solubility is one of the important parameters to achieve desired concentration of drug in systemic circulation for achieving required pharmacological response. Poorly water soluble drugs often require high doses in order to reach therapeutic plasma concentrations after oral administration. Most of the drugs are either weakly acidic or weakly basic having poor aqueous solubility. These poorly water soluble drugs having slow drug absorption leads to inadequate and variable bioavailability and gastrointestinal mucosal toxicity. For orally administered drugs solubility is the most important one rate limiting parameter to achieve their desired concentration in systemic circulation for pharmacological response. Solubility also plays a major role for other dosage forms like parenteral formulations as well.

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Apparatus

1. Test tubes

2. Pipette

3. Dropper

4. Bunsen Burner

5. Watch glass

Reagents

1. Solid Sample a. Salicylic acid b. Benzoic acid

2. Chemicals a. Sodium Carbonate (Na2CO3) b. Sodium Hydroxide (NaOH)

3. Distilled water Salicylic acid

Solubility: 1 gm in 460m ml water

Conversion Reaction:

2 +Na2CO3 2 +CO2 + H2O

Procedure:

1. Take a small amount of salicylic acid (0.5 g) in a dry test tube.

2. Add a small volume of water (1 ml) and shake well. Observe that whether it is

dissolving or not.

3. Now add small amount of Sodium carbonate (Na2CO3) and shake well.

4. Observe that whether it is dissolving or not.

5. Continue the shaking for complete dissolution.

Pharmaceutical uses of Salicylic Acid (Salicylic Acid Topical)

Uses

1. This medication is used on the skin to treat common skin and foot

(plantar) warts. Salicylic acid helps cause the wart to gradually peel off.

2. This medication is also used to help remove corns and calluses.

3. Topical salicylic acid is used to help clear and prevent pimples and skin blemishes in people who have acne.

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http://www.webmd.com/drugs/index-drugs.aspx

http://www.webmd.com/skin-problems-and-treatments/picture-of-the-skin

http://www.webmd.com/skin-problems-and-treatments/warts

http://www.webmd.com/drugs/2/drug-76393/salicylic+acid+in+collodion+topical/details

http://www.webmd.com/skin-problems-and-treatments/guide/understanding-corns-calluses-basics




4. Topical salicylic acid is also used to treat skin conditions that involve scaling or overgrowth of skin cells such as psoriasis (a skin disease in which red, scaly patches form on some areas of the body), ichthyoses (inborn conditions that cause skin dryness and scaling), dandruff, corns, calluses, and warts on the hands or feet.

5. Topical salicylic acid should not be used to treat genital warts, warts on the face, warts with hair growing from them, warts in the nose or mouth, moles, or birthmarks.

6. Salicylic acid is in a class of medications called keratolytic agents.

7. Topical salicylic acid treats acne by reducing swelling and redness and unplugging

blocked skin pores to allow pimples to shrink.

8. It treats other skin conditions by softening and loosening dry, scaly, or thickened skin so that it falls off or can be removed easily.

Brand names of Salicylic Acid

Sl. Brand Name Company

Benzoic acid

Solubility: 1 gm in 300 ml water

Conversion Reaction:

Procedure:

1. Take a small amount of benzoic acid (0.5 g) in a dry test tube.

2. Add a small volume of water (1 ml) and shake well. Observe that whether it is

dissolving or not.

3. Now add small amount of Sodium hydroxide (NaOH) and shake well.

4. Observe that whether it is dissolving or not.

5. Continue the shaking for complete dissolution.

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Pharmaceutical uses of Benzoic Acid

Uses

1. Benzoic acid helps prevent infection caused by bacteria.

2. Benzoic acid and salicylic acid topical (for the skin) is a combination medicine used to treat skin irritation and inflammation caused by burns, insect bites, fungal infections, or eczema.

3. As an inactive ingredient in the pharmaceutical industry, it is used as a. Antimicrobial preservative, b. Antifungal, and c. Tablet and capsule lubricant.

4. Benzoic acid has been used in combination with salicylic acid, as in Whitfield's

ointment, for use as an antifungal for athlete's foot and ringworm.

Brand names of Benzoic Acid

Sl. Brand Name Company

Precautions

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SAFETY FIRST! In this lab, observe all precautions, especially the

ones listed below. If you see a safety icon beside a

step in the Procedure, refer to the list below for its

meaning.

01: Wear your safety goggles. (All steps.)

02: Never cover the opening of a test tube with your finger when mixing

chemicals in the tube. To mix the contents, “flick” the tube. Do not shake

carelessly.

03: Silver, barium, and oxalate compounds are poisonous. Avoid contact with

these chemicals.

04: Silver nitrate will stain skin and clothing.

05: Never pick up a dropper bottle by its cap. Always hold a dropper with the tip lower than the

rubber bulb so that the liquid does not run into the bulb.

06: Exercise care when working with a hot water bath.

07: Nitric acid, HNO3, is a strong acid, and is a corrosive and toxic substance which

can cause severe burns and discoloration of the skin. Avoid contact with your skin,

eyes, and clothing. If you do spill some on yourself, wash immediately with plenty of

cold water. Notify your instructor.

08: Sulfuric acid, H2SO

4, is a strong acid, and is a corrosive and toxic substance

which can cause severe burns. Avoid contact with your skin, eyes, and clothing. If

you do spill some on yourself, wash immediately with plenty of cold water. Notify

your instructor.

09: Silver nitrate, AgNO3, is a toxic substance which can cause burns and

discoloration of the skin. Avoid contact with your skin, eyes, and clothing. If you

do spill some on yourself, wash immediately with plenty of cold water. Notify

your instructor.

10: Return or dispose of all materials properly.

11: Wash your hands thoroughly after completing this experiment.

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Experiment No. 02

1. Name of the Experiment Qualitative analysis (identification) of anions from inorganic salt solutions.

2. Principle

Salts are compounds which are formed from the reaction between acids and bases.

An inorganic salt is just a salt which has no organic part in its molecular structure.

Any salt is composed of two parts – cation and anion. The cation part usually comes from

the base and the anion part usually comes from the acid. As a result, most of the common

anions are also called acid radicals.

Some of the most common acid radicals are sulfate (SO42–), chloride (Cl–), bromide (Br–),

iodide (I–) and nitrate (NO3–) which come from H2SO4, HCl, HBr, HI, HNO3 respectively.

Identification of these radicals in an unknown salt sample is done by taking advantage of

the fact that, not all salts of the radicals are highly soluble in water. For example, BaSO4

is poorly water-soluble but BaCl2 has good water solubility. So, if sulfate ion is present in

a sample, addition of BaCl2 will cause formation of precipitate.

Test Principle Reactions

Sulphate test

principle

Sulphate salt of barium is insoluble in water and forms white precipitate, but other salts of barium are soluble in water. So, if addition of BaCl2 causes formation of precipitate, then sulphate is present.

SO42 + BaCl2 BaSO4 + 2 Cl

(soluble) (white ppt.) HCl is required to remove any carbonate ion that may be present (because carbonate ion also produce white PPT with BaCl2). Instead of HCl, HNO3 can be used but not H2SO4.

Halide test

principle

Silver salts of chloride, bromide and iodide are insoluble in water, with solubility decreasing from chloride salt to iodide salt. So, addition of silver nitrate will cause formation of precipitate (white/cream/yellow). When ammonium hydroxide

is added, some of the silver ions will form soluble diamminesilver(I) ions. Thus the insoluble salt will become soluble. But silver iodide salt is so insoluble that even concentrated ammonium hydroxide can’t dissolve it.

(soluble) (ppt.)X + AgNO3 AgX + NO3

Chloride, bromide, and iodide ions will give

white, pale cream, and pale yellow PPT.

will be formed. When ammonium hydroxide is added, the following reaction moves to the right:

Ag++NH4OH ⇌ [Ag(NH3)2]

+ As a result, the amount of free ions in the solution will decrease. So, the following reaction will move to the right.

AgX Ag+ + X

(insoluble) (soluble) As a result, the salt will be solubilized. But, in case of silver bromide, concentrated ammonium hydroxide is required to significantly dissolve the salt. And for silver iodide, it is not possible to dissolve it even with conc. ammonium hydroxide.

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Nitrate test

This is known as the Ring test of Nitrate. In this test, a brown colored ring is formed between two layers of acid and ferrous sulphate solution in a test tube if nitrate ion is present.

NO3- + 3 Fe2+ + 4 H+ NO + 3 Fe3+ + 2H2O

[Fe(H2O)6]2+SO42- + NO [Fe(H2O)5NO]2+SO4

2- + H2ONitrosoferrous sulfate

(brown) Addition of acid (such as sulphuric acid or hydrochloric acid) will reduce the nitrate to NO. This NO will then attack the ferrous ions to form nitrosoferrous complex which resides in the interface of the acid layer and ferrous sulphate layer.

Acetate

ion (Ferric

chloride test)

When ferric chloride solution is added to the salt solution reddish brown colour (brick red) forms.

On the addition of dil. HCl the color disappears (PPT forms). Boiling the brick red colored solution so formed, followed by the addition of water reddish brown PPT forms.

The reddish color is due to the formation of (CH3COO)3Fe.

The reddish brown precipitate is due to the formation of (CH3COO)(OH)2Fe.

3. Required apparatus

1. Test tubes

2. Pipette (5mL)

3. Water bath or ice bath

4. Required reagents

1. For sulphate identification

a. Dilute (6M) Hydrochloric acid (HCl)

b. 0.1M Barium chloride (BaCl2) or Barium nitrate {Ba(NO3)2} solution

2. For halide ions identification

a. Dilute nitric acid (2M)

b. 5% Silver nitrate (AgNO3) solution

c. Dilute ammonium hydroxide solution (2M)

d. Concentrated ammonium hydroxide solution (~9M)

3. For nitrate ion identification

a. Concentrated Hydrochloric acid (37%) {/Conc. Sulphuric acid }

b. 5% Ferrous sulphate (FeSO4) solution

4. For Acetate ion identification

a. Dilute Hydrochloric acid

b. 0.2M Ferric chloride (FeCl3) solution

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5. Procedure

5.1 Test for sulphate ion

1. Take 3-4mL of the supplied salt solution in a test tube.

2. Add 2mL of dilute (6M) hydrochloric acid solution to the test tube.

CAUTION: Take special care when handling hydrochloric acid, it is extremely

corrosive. Add slowly to the side wall of the test tube.

3. Add 0.5mL of 0.1M Barium chloride or Barium nitrate solution slowly using the

pipette. Check to see if white precipitate is formed.

4. Add more Barium salt solution if precipitate doesn’t form initially or only small

precipitate is obtained initially.

5.2 Test for halide ions

1. Take 2mL of the supplied salt solution in a test tube.

2. Add 2mL of dilute nitric acid to the test tube.

CAUTION: Be careful when handling nitric acid. Add slowly to the side wall of the

test tube.

3. Add 0.5mL of 5% silver nitrate solution. Check to see if white/Cream/Pale Yellow

precipitate is formed.

4. Add more silver nitrate solution if precipitate doesn’t form initially or only small

precipitate is obtained initially.

5. If precipitate is formed, add dilute ammonium hydroxide solution. Check to see if

precipitate dissolves.

6. If precipitate doesn’t dissolve, add concentrated ammonium hydroxide solution.

Check to see if precipitate dissolves.

5.3 Test for nitrate ion


2. Add concentrated hydrochloric acid as such that salt solution: Conc. HCl is 2:1.

3. Cool it.

NOTE: You can use a water bath or ice bath for this.

4. Add 5mL of 5% ferrous sulphate solution very carefully with a pipette to the side

wall of the test tube.

5. The ferrous sulphate solution should form a layer above the acidic solution.

6. Check to see if a brown ring is formed at the interface of the two layers.

5.3 Test for acetate ion


2. Add 2M ferric chloride solution drop-wise.

3. Reddish color (brick red) forms.

4. (If PPT appears) Filter and divide the filtrate into two portions.

5. To one part, add dilute hydrochloride acid solution

6. Reddish color disappears.

7. To second part, add water and boil.

8. Reddish brown precipitate appears. This confirms the presence of acetate ion.

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6. Inference

Name of

the test

Observation Inference

Sample 1 Sample 2 Sample 3

Sulphate

Halide

Nitrate

Acetate

7. Result

8. Precautions

Specimen: Writing inference

Confirmation test for Acid Radicals

Experiment Observations Inference

2 ml of the supplied sample solution was taken in

a test tube. Then 0.5 ml 5% AgNO3 was added to

the sample solution, followed by 2ml dil. HNO3.

White PPT was formed.

Reaction: Cl–(aq)+Ag

+(aq)AgCl(s)

White PPT Cl

– Present

Then excess aq. NH3 (2M) solution was added to

it.

PPT dissolved in excess 2M NH3.

AgCl(s)+2NH3+(aq) ⇌ [Ag(NH3)2]

+(aq)+ Cl

–

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Experiment No. 02

1. Name of the Experiment

Qualitative analysis (identification) of group I & group II cations.

2. Principle




the base and the anion part usually comes from the acid.

Based on solubility characteristics, the common inorganic cations (usually metal in

nature) are classified into five groups. Cation of each group is precipitated by a common

reagent.

So, by step-wise addition of those reagents, we can first identify which group the cation

belongs to. Then by using specific reagents we can confirm the cation.

Group Cations Specific precipitating agents

I Hg+, Pb2+, Ag+ Dilute HCl

II Bi3+, Cu2+, Hg2+, Cd2+, Pb2+, Sn2+, Sn4+, Sb3+, Sb5+, As3+, As5+

H2S

III Fe2+, Fe3+, Al3+, Cr3+, Zn2+, Ni2+, Co2+, Mn2+ NH4Cl + excess NH4OH (NaOH + NH4OH)

IV Ca2+, Ba2+, Sr2+ (NH4)2CO3 (& Na2CO3)

V Na+, K+, Mg2+, NH4+ -

The whole process of group identification can be shown using the following flow-chart.

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The cations of group I are separated when dilute HCl is added because of the formation of

insoluble chloride salts. Thus, to identify whether the cation belongs to group I or not, one

should add dilute HCl. If precipitate is formed, then the specific cation must be identified.

Ag+ + Cl– AgCl(s)

Pb2+ + 2Cl– PbCl2

For the confirmation of the specific cation of group I following process should be utilized.

Confirmation test

Description Reactions

Ag+ test When NH3 solution is added to silver chloride, the white precipitate dissolves.

Ag++ NH4OH ⇌[Ag(NH3)2]+

Pb2+ test

Potassium chromate is added to the stock solution. A bright yellow precipitate confirms Pb2+ ion.

Pb2+(aq) + K2CrO4(aq) → PbCrO4(s)

If no precipitate is formed when dilute HCl is added to the stock solution, then group I

cations are not present. Then we have to add H2S which will form insoluble sulfide salts of

group II cations. H2S is not directly produced, rather thioacetamide is added which will

produce H2S in solution.

Hg2+ + H2S HgS (black) + 2H+ + 2Cl–

Cu2+ + H2S CuS (black) + 2H+ + 2Cl–

If precipitates are formed upon addition of thioacetamide, the next step is to confirm the

cation.

Confirmation

test

Description Reaction

Hg2+ test

When HNO3 is added, HgS doesn’t dissolve, but sulfides of other cations dissolve.

Then HCl is added (aqua regia), which will cause formation of soluble HgCl2 (mercuric chloride) salt.

Addition of SnCl2 will cause formation of Hg2Cl2 and elemental Hg both which form precipitates.

HgS +HNO3 X

HgS + HNO3 + 6Cl– + 8H+ 3HgCl2 + 2NO + 3S + 4H2O

Hg2Cl2 + Sn2+ Hg + Hg2Cl2 + Sn4+

Cu2+ test

When Sodium hydroxide is added to a

solution of cupric (Cu2+) ions, blue gelatinous precipitate will form.

When aqueous ammonia is added to a solution of cupric (Cu2+) ions, blue precipitate will form. Excess addition of ammonia will dissolve the blue precipitate forming a deep blue color.

Cu2+ + NH4OH Cu(OH)2

(blue PPT)

Cu2+ + NH4OH Cu(OH)2

(blue PPT)

Cu(OH)2 + NH4OH [Cu(NH3)4]2+ (blue sol.)

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4. Test tubes 5. Pipette (5mL) 6. Water bath and/or ice bath


A. For group I cations a. Dilute hydrochloric acid (6M) b. Ammonia solution (9M) c. Potassium chromate (1M)

B. For group II cations

a. Ferrous sulfide (FeS) b. Concentrated sulfuric acid c. Nitric acid (6M) d. Stannous chloride (0.1M) e. Concentrated Ammonia solution (9M) f. Dilute Ammonia solution (2M) g. Dilute Sodium hydroxide solution (2M)

5. Procedure

5.1 Test for group I cations

1. Take 3-4mL of the supplied salt solution in a test tube. 2. Add 2mL of dilute (6M) hydrochloric acid solution to the test tube.

CAUTION: Take special care when handling hydrochloric acid, it is extremely corrosive. Add slowly to the side wall of the test tube.

3. Check to see if precipitate is formed. 4. If precipitate is formed, go for confirmation tests. 5. If precipitate is not formed, go for test of group II cations.

5.1.1 Confirmation test of Ag+

1. Take the solution containing the precipitate. 2. Add 2-3mL of ammonia solution. If the precipitate is dissolved, then Ag+ is

confirmed. 5.1.2 Confirmation test of Pb2+

1. Take 2mL of the stock solution in a test tube. 2. Add 0.5mL of 1M potassium chromate solution. If precipitate starts to form, add

more 1M potassium chromate (up to 2mL). 3. Bright yellow precipitate confirms Pb2+ ions.

5.2 Test for group II cations

7. From 5.1, take the solution and add H2S from the apparatus. 8. Check to see if any precipitate is formed. If black precipitate is formed, first go for

Hg2+ or Cu2+ confirmation test.

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5.2.1 Confirmation test of Hg2+

1. First separate the black precipitate from the solution by filtration. 2. Take the precipitate in another test tube and add 1mL distilled water. 3. Then add 1mL 6M Nitric acid & Shake vigorously. 4. If precipitate doesn’t dissolve, it is indicative of the presence of Hg2+. 5. Add 1mL 6M HCl and 1mL 6M HNO3 to the test tube. Heat in a boiling water bath. 6. Cool the solution in an ice bath. 7. Add 0.5mL of 0.1M SnCl2 drop wise. 8. Grey to black precipitates confirms the presence of Hg2+.

5.2.2 Confirmation test of Cu2+

1. Take 2mL of the stock solution in a test tube. 2. Add 0.5 mL dilute NaOH solution drop wise. 3. Check to see if light blue gelatinous precipitate is formed. 4. If precipitate is observed, then add more NaOH solution. Precipitate insoluble in

excess alkali.

5. Again take 2mL of the stock solution in a test tube

6. Add 0.5 mL dilute ammonia solution drop wise. 7. Check to see if light blue precipitate is formed. 8. If precipitate is observed, then add more ammonia solution. A deep blue colored

solution will be formed. This confirms the presence of Cu2+.

6. Inference

Group test

Confirmation test



Group I

-

Hg+ test

Ag+ test

Pb2+ test

Group II

-

Hg2+ test

Cu2+ test

7. Result 8. Precautions

Specimen: Writing inference Confirmatory test for Group II cations

Experiment Observations Inference 2ml stock solution was taken in a test tube and

2M NaOH solution was added to it.

Blue gelatinous PPT was formed.

Reaction: Cu2+

(aq)+2OH–(aq) Cu(OH)2(s)

Blue PPT Cu2+

Present Then excess 2M NaOH solution was added to it. The Precipitate was insoluble in excess NaOH.

Again, 2ml stock solution was again taken in a

test tube and ammonia solution was added to it.

Blue gelatinous PPT formed.

Reaction: Cu2+

(aq)+2NH3(aq)+2H2O Cu(OH)2(s)+ 2NH4

+(aq)

Blue gelatinous PPT

Cu2+

Confirmed Afterwards, excess ammonia solution was added

to it.

Precipitate was soluble in excess 2M NH4OH.

Excess aq. NH3

Cu(OH)2 (s)+ 4NH3(aq)

[Cu(NH3)4]2+

(aq) + 2OH–(aq)

Tetraamminecopper(II) ion

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Experiment No. 02

1. Name of the Experiment Qualitative analysis (identification) of group III.

2. Principle






nature) are classified into five groups. Cations of each group is precipitated a common

reagent. So, by step-wise addition of those reagents, we can first identify which group the

cation belongs to. Then by using specific reagents we can confirm the cation.

The common inorganic cations can be classified into five groups based on solubility

characteristics. Cations of each class are precipitated by a specific reagent.




H2S



V Na+, K+, Mg2+, NH4+ -

Cations of group III are not precipitated by dilute HCl or H2S, but they can be precipitated

from a salt solution by the addition of NH4Cl in the presence of excess NH4OH.

With this reagent, hydroxide or sulphide salts of cation will be precipitated. If precipitate

is formed, then confirmation tests should be performed to identify specific cation.

Confirmation test Description Reactions

Fe2+ test

If the precipitate from the above reaction is greenish, then it may be ferrous. If dilute sodium hydroxide is added to a solution containing ferrous ions (Fe2+), a green gelatinous precipitate of iron (II) hydroxide will form. Excess addition of sodium hydroxide will not dissolve the precipitate. When aqueous ammonia (dil. /2M) is added to a solution of ferrous (Fe2+) ions, a green gelatinous precipitate will form. Excess addition of ammonia will not dissolve the precipitate. To confirm, potassium ferricyanide is added to the stock solution, a deep blue precipitate is obtained.

Fe2+(aq) +2OH–(aq) Fe(OH)2(s) Green precipitate

Fe2+(aq) +2NH3(aq)

Fe(OH)2(s) +2NH4+(aq)

Green precipitate

Fe2+ + K3[Fe(CN)6] KF[Fe(CN)6] (blue PPT)

Md. Imran Nur Manik





Fe3+ test

If the precipitate from the above reaction is brownish, then it may be ferric. If dilute sodium hydroxide is added to a solution containing ferric ions (Fe3+), a brownish red gelatinous precipitate of iron (III) hydroxide will form. Excess addition of sodium hydroxide will not dissolve the precipitate. When aqueous ammonia (dil. /2M) is added to a solution of ferric (Fe3+) ions, a brownish red gelatinous precipitate will form. Excess addition of ammonia will not dissolve the precipitate. To confirm, potassium ferricyanide is added to the stock solution, a red (greenish) precipitate is obtained. In another test, addition of potassium thiocyanate to stock solution will produce blood red solution.

Fe3+(aq) +3OH–(aq) Fe(OH)3(s) Reddish brown precipitate

Fe3+(aq) +3NH3(aq)

Fe(OH)3(s) +3NH4+(aq)

Reddish brown precipitate

Fe3+ + K3[Fe(CN)6] F[Fe(CN)6] (red PPT.) Fe3+ + KCNS + H2O [Fe(H2O)5(CNS)]2+ + K+


Zn2+ test

When dilute sodium hydroxide is added to a solution containing zinc ions (Zn2+), a gelatinous white precipitate is formed. Addition of concentrated sodium hydroxide (excess dil. NaOH) produces a soluble complex resulting in a clear solution. When ammonia (dil. /2M) is added to a solution of zinc (Zn2+) ions, white gelatinous precipitate will form. Excess addition of ammonia will dissolve the precipitate forming a water soluble colourless tetraaminezinc ion.

Zn2+ + 2NaOH

Na+ + Zn(OH)2(white PPT.) Zn(OH)2 + 2NaOH Na2[Zn(OH)4] (SodiumZincate.)

Zn2+(aq)+2NH3(aq)+2H2O Zn(OH)2(s)+ 2NH4

+(aq) White gelatinous PPT Zn (OH)2 (s)+ 4NH3(aq) [Zn (NH3)4]2+ (aq)+ 2OH–(aq) Tetraamminezinc(II) ion

Al3+ test

When dilute sodium hydroxide is added to a solution containing aluminium ions (Al3+), a white gelatinous precipitate is formed. Addition of concentrated sodium hydroxide (excess dil. NaOH) produces a soluble complex resulting in a clear solution. When aqueous ammonia (dil. /2M) is added to a solution of aluminium (Al3+) ions, white gelatinous precipitate will form. Excess addition of ammonia will not dissolve the precipitate.

Al3+ + 3OH– Al(OH)3

(white PPT.) Al(OH)3 (s)+ OH– (aq) [Al(OH)4]–(aq) Tetrahydroxoaluminate ion

Al3+(aq)+3NH3(aq)+3H2O Al(OH)3(s)+ 3NH4

+(aq) White gelatinous PPT

Md. Imran Nur Manik





1. Test tubes

2. Pipette (5mL)

3. H2S generating apparatus


1. For group III cations

a. Ammonium chloride

b. Ammonia solution (9M)

c. Ammonia solution (2M)

d. Potassium ferricyanide (K3[Fe(CN)6]) [250 mg in 10 ml]

e. Potassium thiocyanate [250 mg in 10 ml]

f. Sodium hydroxide (2M)

g. Sodium hydroxide (10M)

5. Procedure

5.1 Test for group III cations





3. Check to see if precipitate is formed.

4. If no precipitate is formed, add H2S in the test tube. Check again to see if


5. If no precipitate is formed above, add 2mL concentrated NH4OH and 1gm of NH4Cl.

Check to see if precipitate is formed.

6. If precipitate is formed then go for confirmation tests. If no precipitate is formed, go

for ‘Test for group IV cations’.

5.1.1 Confirmation test of Fe2+ cation

1. Take 2mL of the stock solution in a test tube.

2. Add 0.5 mL dilute NaOH solution drop wise.

3. Check to see if light green gelatinous precipitate is formed.

4. If precipitate is observed, then add more NaOH solution. Precipitate insoluble in

excess alkali.


6. Add 0.5 mL dilute ammonia solution drop wise.

7. Check to see if green gelatinous precipitate is formed.

8. If precipitate is observed, then add more ammonia solution.

Precipitate insoluble in excess aqueous ammonia.

Md. Imran Nur Manik




To further confirm the presence of Fe2+

1. Take 2ml of the stock the solution in a test tube.

2. Add 2mL potassium ferricyanide (just one drop) solution. Check to see if precipitate

is formed.

1. Light blue precipitate confirms the presence of Fe2+.

5.1.2 Confirmation test of Fe3+ cation


2. Add 0.5 mL dilute NaOH solution drop wise.

3. Check to see if light brownish red gelatinous precipitate is formed.

4. If precipitate is observed, then add more NaOH solution. Precipitate insoluble in

excess alkali.


6. Add 0.5 mL dilute ammonia solution drop wise.

7. Check to see if brownish red gelatinous precipitate is formed.

8. If precipitate is observed, then add more ammonia solution.

Precipitate insoluble in excess aqueous ammonia.

To further confirm the presence of Fe3+

1. Take 2ml of the stock the solution in a test tube.

2. Add 2mL potassium ferricyanide (just one drop) solution. Check to see if precipitate

is formed.

3. Brownish (Green) colour indicates Fe3+.

4. To confirm Fe3+, you can also use potassium thiocyanate.

5. Take 2ml of the stock the solution in a test tube and add 2mL of potassium

thiocyanate (just one drop) solution to stock solution, if the colour change to blood

red then Fe3+ is confirmed.

5.1.3 Confirmation test of Zn2+


1. First add 2M sodium hydroxide drop wise very slowly.

2. Check to see if white precipitate has been formed. If precipitate is formed, add

concentrated NaOH (10M) {excess 2M NaOH} to check if the precipitate dissolves.

3. If the precipitate dissolves, then Zn2+ is confirmed.

4. Again take 2mL of the stock solution in a test tube.

5. Now add ammonia (dil. /2M) drop wise very slowly.

6. Check to see if white precipitate has been formed. If precipitate is formed, add

concentrated ammonia (9M) {excess 2M ammonia} to check if the precipitate

dissolves.

7. If the precipitate dissolves, then Zn2+ is further confirmed.

(10dilution may be needed for the appearance and disappearance of PPT)

Md. Imran Nur Manik




5.1.4 Confirmation test of Al3+


2. First add 2M sodium hydroxide drop wise very slowly.

3. Check to see if white gelatinous precipitate has been formed. If precipitate is

formed, add concentrated NaOH (10M) {excess 2M NaOH} to check if the precipitate

dissolves.

4. If the precipitate dissolves, then Al3+ is confirmed.

5. Again take 2mL of the stock solution in a test tube. 6. Now add ammonia (dil. /2M) drop wise very slowly. 7. Check to see if white precipitate has been formed. If precipitate is formed, add

concentrated ammonia (9M) {excess 2M ammonia} to check if the precipitate dissolves.

8. If the precipitate does not dissolves, then Al3+ is further confirmed.

6. Inference

Group test

Confirmation test



Group III

-

Fe2+ test

Fe3+ test

Zn2+ test

Al3+ test

7. Result

8. Precautions

Specimen: Writing inference

Analysis for Group III cations

Experiment Observations Inference 2ml stock solution was taken in a test

tube and 2M NaOH solution was

added to it.

Green gelatinous PPT was formed.

Reaction: Fe2+

(aq)+2OH–(aq) Fe(OH)2(s)

Green jelly like PPT Fe2+

Present

Then excess 2M NaOH solution was

added to it. Precipitate was insoluble in excess NaOH.

2ml stock solution was again taken in

a test tube and ammonia solution was

added to it.

Green gelatinous PPT formed.

Reaction: Fe2+

(aq)+2NH3(aq)+2H2O Fe(OH)2(s)+ 2NH4+(aq)

Green gelatinous PPT Fe2+

Present

Afterwards, excess ammonia solution

was added to it. Precipitate was insoluble in excess aqueous ammonia solution.

Again 2mL of the stock solution was

taken in a test tube and one drop

potassium ferricyanide solution was

added to it.

Light blue precipitate was formed. Fe

2+

Confirmed

Md. Imran Nur Manik




Experiment No. 02

1. Name of the Experiment

Qualitative analysis (identification) of group IV & group V cations.

2. Principle






nature) are classified into five groups. Cations of each group is precipitated a common

reagent. So, by step-wise addition of those reagents, we can first identify which group the

cation belongs to. Then by using specific reagents we can confirm the cation.

The common inorganic cations can be classified into five groups based on solubility

characteristics. Cations of each class are precipitated by a specific reagent.




H2S



V Na+, K+, Mg2+, NH4+ -

Cations of group IV are not precipitated by dilute HCl or H2S, as well as by the addition of

NH4Cl in the presence of excess NH4OH.

If no precipitate is formed with the addition of NH4OH and NH4Cl to the stock solution,

then (NH4)2CO3 is added to it. If precipitate is formed, then group IV cations are present.

In that case, confirmation tests are carried out as follows:

Confirmation test

Description Reaction

Ca2+ test When the sample salt (solid form or solution) is heated in Bunsen burner, a brick-red flame confirms presence of Ca2+.

NH4+ test

If Nessler’s reagent is added to a solution containing ammonium ions (NH4

+), a brown

precipitate is formed. When sodium hydroxide is added to a solution containing ammonium ions (NH4

+), and the mixture a colourless gas will be produced which turns moist red litmus paper blue.

NH4+ + 2[HgI4]2− +

4OH− →HgO·Hg(NH2)I(s)↓

+ 7I− + 3H2O (Yellow or brown)

NH4+ (aq)+ OH–

NH3(g) + H2O(l)

Ammonia

Md. Imran Nur Manik





a) Test tubes b) Pipette (5mL) c) H2S generating apparatus d) Platinum wire

4. Required reagents 1. For group IV cations

a) Ammonium carbonate b) Conc. hydrochloric acid

2. For group V cations

a) Litmus paper (red)

b) Nessler's reagent [a 0.09 mol/L solution of potassium tetraiodomercurate(II) (K2[HgI4]) in 2.5 mol/L potassium hydroxide]

5. Procedure

5.1 Test for group IV cations





3. Check to see if precipitate is formed.

4. If no precipitate is formed, add H2S in the test tube. Check again to see if


5. If no precipitate is formed above, add 2mL concentrated NH4OH and 1gm of NH4Cl.

Check to see if precipitate is formed. If precipitate is formed then go for

confirmation tests.

6. If no precipitate is formed, then add (NH4)2CO3. If precipitate is formed, then go for

‘Test for group IV cations’.

7. If no precipitate is formed, then go for ‘Test for group V cations’.

5.2 Test for group IV cations

1. From 5.1, take the solution and add 5% ammonium carbonate solution.

2. Check to see if any precipitate is formed. If precipitate is formed,

go for confirmation test.

5.2.1 Confirmation test of Ca2+

1. Take a platinum or nichrome wire and dip it in concentrated hydrochloric acid.

2. Then heat the wire in Bunsen burner until flame is no longer visible.

3. Then take the wire and dip it in the sample salt solution. Conversely, you can dip it

in a solid sample of the salt.

4. Place the wire in the Bunsen burner. Observe the flame colour.

5. Brick-red flame confirms the presence of Ca2+ in the sample.

Md. Imran Nur Manik




5.2.2 Confirmation test of NH4+


2. Add Nessler's reagent drop wise.

3. Check to see if brown precipitate has been formed. If precipitate is not forming or

disappears as quickly as it forms then add more Nessler’s reagent until the

precipitate forms.

4. Add more Nessler’s reagent to check if the precipitate dissolves.

5. If the precipitate does not dissolve, then NH4+ is confirmed.

For further confirmation


2. Add 3ml NaOH solution to it.

3. Then heat the solution carefully in Bunsen burner until a colourless gas evolves.

4. Place a moist red litmus paper in the mouth of the test tube.

5. Check to see if the red litmus paper turns into blue.

6. If the colourless gas turns the red litmus paper into blue then NH4+ is further

confirmed.

Hold the tube containing the ammonium nitrate solution with a test-tube

holder. Gently warm the tube along its sides using a back-and-fort

motion through a burner flame. Do not allow the solution to boil.

CAUTION: At all times, make sure that the opening of the tube is

pointed away from other people. Hold a moistened piece of red litmus

paper near the mouth of the test tube, as shown in Figure. The test will

be spoiled if the solution contacts the litmus paper. Record the changes

you observe. Fan the vapours coming out of the tube toward your nose

with your hand. Cautiously sniff the vapours. Record your observations.

6. Inference

Group test

Confirmation test



Group IV

-

Ca2+ test

Group V

-

NH4+ test

7. Result

8. Precautions

Md. Imran Nur Manik

Laboratory Techniques

The objective of qualitative analysis is the identification of the ions present in a sample. The usual procedure is

to run a carefully selected series of reactions on a solution containing all the ions of a particular group. This

solution is called a “known” solution or “control.” The behavior of an unknown, which contains one or more of

the ions of the group, is then compared with that of the known. The occurrence of certain reactions and the

absence of others allow us to determine which ions are present in the unknown. Such identification of ions

requires careful laboratory work.

These procedures and techniques are generally used throughout the qualitative analysis scheme and will not be

described in detail when they occur.

1. Cleanliness: If a known or unknown sample comes in contact with any surface that has not been thoroughly

cleaned, it will be contaminated and the analytical results will be unreliable. It is important that all glassware be

kept clean (not necessarily dry) during lab. Label two beakers, “clean” and “dirty”, and use these for stirring

rods, droppers, test tubes, etc. Always rinse glassware twice with deionized water and use only deionized water

whenever instructions call for water. When a sample is contaminated it should be discarded.

2. Labeling: Label all solutions and precipitates immediately with the designations used in the lab procedure;

i.e., 2A, 3B, etc. Use these same labels in the lab notebook.

3. Dispensing Solutions: Volumes of solutions can usually be estimated. You should have a plastic dropper

with a 1 mL mark etched into it for your convenience. It is also helpful to determine how many drops from

your glass droppers are equal to 1 mL (usually around 20 drops). You should also be aware of the level of a

milliliter of liquid in your 10 x 75 mm test tubes so that you can readily estimate that volume.

When adding reagents from dropper bottles (either your own or those in the hoods) be careful not to

contaminate the dropper: do not put the tip into your solution or touch the inside of the test tube with the

dropper tip. And always place the dropper back into the reagent bottle immediately. Never lay the dropper on

a contaminated bench top.

If a reagent bottle does not have a dropper, pour a small amount into a clean test tube and use one of your own

droppers to dispense it from the test tube. Never place your own dropper into a reagent bottle that is used by

others.

4. Testing Acidity/Basicity: Usually when instructions call for the addition of acid or base to a solution it is

important to add enough to change the solution from acidic to basic or vice versa. A litmus test is used to detect

this change. Do not dip litmus paper into your solution in the test tube since this may contaminate your

sample. Use the stirring rod to transfer a drop of the solution to a small piece of litmus paper and note when the

color change occurs.

5. Heating Solutions: Keep an enough precautions during heating of your solutions.

6. Precipitation of Solids: To accomplish a precipitation, add the indicated amount of precipitating agent to

the solution. Some precipitates form slowly and must be given time to form completely.

Md. Imran Nur Manik

Preparation of stock solution

1. For all in general, 1 gm salt in 10 ml DW 2. Filter if required 3. For ZnSO4, 500 mg in 10 ml (make further dilution if required) 4. 0.2M FeCl3 (324 mg in 10 ml)

5. For NH4Cl, 250 mg in 10 ml

Halide ion

(2ml)

2 ml HNO3 then 0.5 ml

5% AgNO3

Action of dilute NH3 on

product (excess 2M NH3)

Action of conc NH3 on

product (excess 9M NH3)

Cl– White ppt

(Sun lightturns grey quickly)

White ppt dissolves White ppt dissolves

Br– Pale Cream ppt

(Sun lightturns grey slowly)

No reaction Cream ppt dissolves

slightly

I– Pale Yellow ppt

(Sun lightno change)

Insoluble insoluble

SO42-

Add barium chloride and mix followed by dilute nitric or

HCl acid

A white precipitate insoluble in

dilute acid

NO3-

Add a small amount of freshly prepared ferrous sulphate

solution.

Gently pour conc. H2SO4 down the side of the ppt tube

such that the solutions do not mix

A brown ring may be formed at the

interface/junction of the two

solutions.

Cation Reaction with NaOH (aq) Reaction with aqueous ammonia (NH3)

Ba

2+

White precipitate formed

Precipitate insoluble in excess alkali


Precipitate insoluble in excess aqueous ammonia

Ca

2+





Mg

2+





Al3+ White precipitate formed

Precipitate soluble in excess alkali



NH4+ No white precipitate

When heated a gas which turns moist red

litmus paper blue is evolved

No white precipitate

When heated a gas which turns moist red litmus

paper blue is evolved.

Zn2+




Precipitate soluble in excess aqueous ammonia

Fe2+

Green precipitate formed


Green precipitate formed


Fe3+ Brown precipitate formed


Brown precipitate formed


Md. Imran Nur Manik

Cation Reaction with NaOH (aq) Reaction with aqueous ammonia (NH3)

Pb2+ White precipitate formed




Cu2+ Blue precipitate formed


Blue precipitate formed

Precipitate soluble in excess aqueous ammonia

forming a deep blue solution.

Test Observations Inferences

Add dilute HCl

to solution

- White ppt formed

- No white ppt is formed

Pb2+,

Hg2+

or Ag+ ions to be present

Pb2+,

Mg2+

and Ag+ absent

Add dilute

H2SO4

- White ppt formed

- No white ppt formed

Pb2+

, Ca2+

or Ba2+

present

Pb2+

, Ca2+

and Ba2+

ions absent

SO42-

Add barium chloride and mix followed by

dilute nitric or HCl acid

A white precipitate insoluble in dilute acid

NB: the precipitate dissolves due to the formation of complex ions e.g.

Amphoteric hydroxides of Aluminiu. Zinc and lead dissolves in excess NaOH.

• Pb(OH)2 (s) + 2OH- (aq) Pb(OH)4

2- (aq) + 2Na

+

• Zn(OH)2 (s) + 4NH3(aq) {Zn (NH3)4}2+

(aq) + 2OH-(aq)

• Al(OH)3 (s) + OH- (aq) Al(OH)4

-(aq) + Na

+(aq)

Zinc and copper hydroxide dissolve in excess ammonia solution to form colourless tetra amine zincate and deep

blue tetra amine copper respectively.

• Zn(OH)2 (s) + 4NH3(aq) { Zn (NH3)4}2+

(aq) + 2OH-(aq)

Tetramine zinc (II) ion.

• Cu(OH)2 (s) + 4NH3(aq) { Cu(NH3) }2+

(aq) + 2OH-(aq)

Tetramine Copper (II) ion

If on adding few drops of NaOH or ammonia solutions, white precipitate is formed, conclude by writing;

• Either; Ca2+

, Mg2+

,Al3+

, Zn2+

,Pb2+

or Ba2+

ions present.

• If green precipitate is formed, Fe2+

ions present.

• If brown precipitate is formed, Fe3+

ions present.

• If blue precipitate is formed, Cu2+

ions present.

• If on adding excess NaOH solution white precipitate persisted (insoluble in excess NaOH), conclude; Ca2+

, Mg2+ or Ba2+

ions present.

• If on adding excess NaOH solution white precipitate dissolves (soluble in excess NaOH) conclude.

Al3+

,

Zn2+ or Pb2+

ions present.

• If on adding excess ammonia solution white precipitate persisted (insoluble in excess), conclude; Either Ba2+

, Ca2+ , Mg2+

, Al3+ , or Pb2+

ions present.

• If on adding excess ammonia solution white precipitate dissolves (soluble in excess) conclude; Zn2+

ions confirmed present

Md. Imran Nur Manik

For cations which exhibit similar properties in both dilute NaOH and aqueous ammonia. A confirmatory test

has to be performed to distinguish them e.g. Al3+

ions and Pb2+

ions which form white precipitates which are

soluble in excess NaOH but insoluble in excess aqueous ammonia.

To distinguish them, two to three drops of either sodium chloride (NaCl), sodium sulphate (NaSO4) or

potassium iodine KI are added. Lead ions form insoluble chloride, sulphate or iodide respectively.

Flame test for some cations

This method is based on the colours the cations render a flame when heated.

Procedure.

• Dip a nichrome wire into concentrated HCl acid.

• Heat the end that has been dipped in the acid using a flame until it cease to impart colour on the flame.

This is done to clean the wire.

• Dip the cleaned wire into the solution containing the test ion and then heat it in the hottest part of the

flame. The test ion imparts its characteristics colour.

NB: platinum wire, metallic spatula or glass rod can be used instead of the nichrome wire.

Cation Colour of the flame produced

Lithium (Li+) Crimson (deep red)

Sodium (Na+) Yellow

Potassium (K+) Lilac (Purple)

Calcium (Ca2+

) Red

Copper (Cu2+

)) Blue-green

Barium (Ba2+

)) Apple green

Acetate ion

1. Take 2 ml Salt solution (1gm Lead Acetate in 10 ml DW) in attest tube

2. Add 0.2M FeCl3 drop wise

3. Check to see if Brick red colour if is produced

4. Now add dil. HCl

5. Check to see if Brick red colour disappears

6. Add water and boil the colourless solution so produced

7. Check to see if Reddish brown ppt appears

8. If ppt appears then acetate is confirmed.

Md. Imran Nur Manik

Qualitative analysis of salts lab manual MANIK

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Transcript of Qualitative analysis of salts lab manual MANIK