Practical Chemistry is a vital part of the chemistry syllabus for competitive exams like JEE and NEET, as well as for board examinations. It emphasizes the application of theoretical concepts through hands-on experiment and real-life chemical processes. The topic “Principles Related to Practical Chemistry” covers detection of extra elements in organic compounds, detection of functional groups in organic compounds, inorganic salt analysis, titration, etc. Whether preparing for exams or pursuing further studies in science, a solid grasp of practical chemistry principles is essential for success and confidence in the laboratory.
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Practical Chemistry Syllabus
- Detection of extra elements (Nitrogen, Sulphur, Halogens) in organic compounds; Detection of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketones) carboxyl, and amino groups in organic compounds.
- The chemistry involved in the preparation of the following: Inorganic compounds (Mohr’s salt, potash alum), Organic compounds (Acetanilide, p-nitro acetanilide, aniline yellow, iodoform).
- The chemistry involved in the titrimetric exercises– Acids, bases and the use of indicators, oxalic acid vs KMnO4, Mohr’s salt vs KMnO4
- Chemical principles involved in the qualitative salt analysis: Cations (Pb2+, Cu2+, Al3+, Fe3+, Zn2+, Ni2+, Ca2+, Ba2+, Mg2+, NH4+), Anions (CO32−, S2−, SO42−, NO3−, NO2−, Cl−, Br−, I− (Insoluble salts excluded).
- Chemical principles involved in the following experiments: Enthalpy of solution of CuSO4, Enthalpy of neutralization of strong acid and strong base, Preparation of lyophilic and lyophobic sols, Kinetic study of the reaction of iodide ions with hydrogen peroxide at room temperature.
Qualitative Analysis of Organic Compounds: Analysis involving detection of all elements present in an organic compound. It involves the detection of nitrogen, halogen and sulphur by Lassigne’s test.
Detection of Elements (Nitrogen, Sulphur, Halogens) in Organic Compounds
Organic compound generally consists of carbon and hydrogen. Some other elements which may be present inorganic compounds are nitrogen, sulphur, oxygen and halogens.
Lassigne’s test is mostly preferred to confirm the presence of elements like N, S, Cl, Br and I. In this test, salt is used with sodium metal to prepare its water soluble salt called soda extract or Lassigne’s extract.
Preparation of Lassigne’s Extract
To prepare soda extract, an small amount of organic compound is heated with dry sodium in a fusion tube. The red hot tube is crushed in distilled water and filtered, the filtrate is known as Lassigne’s extract.
The chemical reactions of various elements with Na are as:
Na + C + N → NaCN
2 Na + S → Na2S
Na + X → NaX (X may be Cl, Br or I)
If nitrogen and sulphur both are present in the compound then sodium thiocyanate is formed.
Na + C + N + S → NaSCN
Formation of sodium thiocyanate is possible only when sodium metal is present in little amounts.
Detection of Nitrogen
(a) Sodium extract in a test tube is added with freshly prepared saturated solution of ferrous sulphate. A green precipitate of ferrous hydroxide is formed. Acidify the solution with dilute sulphuric acid and warm ferric chloride solution is added. A Prussian blue precipitate indicates the presence of nitrogen.
Fe4[Fe(CN)6]3 is the compound named Iron (III) hexacyanidoferrate (II) which shows Prusian Blue colour.
Detection of Sulphur
(a) The sodium fusion extract is added with sodium nitroprusside, then appearance of violet colour confirms the presence of S.
(b) The soda extract is added with acetic acid and lead acetate is added to it. The appearance of black precipitate of lead sulphide confirms the presence of S.
Detection for Nitrogen and Sulphur when present together
When both nitrogen and sulphur are present in an organic compound, sodium thiocyanate may be formed on fusion with sodium metal. Soda extract is acidified with dilute HCl, 2-3 drops of FeCl3 is added to it. Appearance of blood red colouration confirms the presence of both nitrogen and sulphur in the organic compound.
Detection of Halogens
To test halogens, first nitrogen and sulphur are removed in the form of HCN and H2S. For this soda extract is added with few drops of conc. HNO3.
(a) Now, AgNO3 is added to the solution of a white precipitate is formed which dissolved in excess of NH3 (l) then the presence of chlorine is confirmed.
(b) If a yellow precipitate partically soluble in ammonium hydroxide solution is formed, then the presence of bromine is confirmed in the organic compound.
(c) Iodine ion formd dark yellow precipitate that is insoluble in ammonia solution.
Layer Test for bromine and iodine
Soda extract is added with CCl4 or CS2 and chlorine water, appearance of orange coloured layer confirms the presence of bromine and violet coloured layer confirms the presence of iodine.
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Detection of Functional Groups
In this detection of functional group, we will be going through the chemical reactions involved for the detection of various functional groups present in the organic compounds, like -COOH, -OH, -CO-, -NH2, etc.
Test for Alcoholic Group [R-OH]
Ceric Ammonium Nitrate Test
Alcoholic compounds on reaction with ceric ammonium nitrate give a red colouration due to the formation of a complex.
Lucas Test
Lucas reagent (anhydrous ZnCl2 + conc. HCl) converts alcohols into the corresponding alkyl chlorides. Zinc chloride (a lewis acid) increases the reactivity of alcohols with HCl.
(i) Tertiary alcohol: Turbidity appears immediately.
(ii) Secondary alcohol: Turbidity appears after 5 minutes.
(iii) Primary alcohol: Turbidity appears after heating.
ZnCl2 is not required on reaction with tertiary alcohol. Turbidity is due to the formation of alkyl chloride.
Iodoform Test
This test is given by alcohols and carbonyl compounds having alpha-methyl groups such as CH3CH(OH)- or CH3CO- groups respectively.
Compound is heated with I2 + NaOH or I2 + Na2CO3 or NaIO3 to give iodoform.
A yellow precipitate of iodoform confirms the test.
Test for Phenolic Group [Ar-OH]
Ferric Chloride Test
Phenols form violet coloured complex with freshly prepared FeCl3 solution.
Resorcinol, o-cresol, m-cresol and p-cresol give violet or blue colouration, catechol gives green colour which rapidly darkens. 1-Naphthol and 2-Naphthol do not give characterstic colour.
Phthalein Dye Test
Phenol condenses with phthalic anhydride in the presence of concentrated H2SO4. Phenol condenses to give phneolphthalein which gives a dark pink colour with NaOH solution. This is called as phthalein dye test.
Test for Aldehydic and Ketonic Functional Group [-CHO and -CO-]
2,4-Dinitrophenylhydrazine Test (2,4-DNP Test)
This test is performed by both aldehydes and ketones. 2,4-Dinitrophenylhydrazine reacts with aldehydes and ketones to give 2,4-dinitrophenylhydrazone. The colour of the product is yellow and red for aldehydes and ketones respectively.
Tests to distinguish Aldehydes and Ketones
Schiff’s test, Fehling’s test and Tollen’s test are given by aldehydes only. So, the carbonyl compounds which take part in 2,4-DNP test as well as positive Fehling test, Schiff’s test and Tollen’s test contain aldehydic group while ketones give negative results for Fehling test, Schiff’s test and Tollen’s test.
Fehling Test
Fehling reagent is a mixture of Fehling A and Fehling B solution. Fehling A is an aqueous solution of copper sulphate while Fehling B is an alkaline solution of sodium potassium tartarate (Rochelle’s salt).
We mix the two solution just before the reaction as Fehling reagent (mixture) is not stable while separately Fehling A and Fehling B are quite stable. Fehling reagent (mixture) contains complex of copper with tartarate ions as follows:
Cupric ions oxidises aldehyde into carboxylic acid and brick red precipitate of Cu2O is formed while the ketones remain unaffected. The important point to note here is that aromatic aldehydes do not give positive Fehling Test.
Benedict solution is the modified form of Fehling solution where a single solution is used. It is an alkaline solution containing a mixture of copper sulphate and sodium citrate (2Na3C6H5O7.11H2O). It is more stable than Fehling solution and can be stored for a longer duration.
Schiff’s Test
Schiff’s reagent is prepared by decolourising an aqueous solution of p-rosaniline hydrochloride dye by adding sodium sulphite to it or by passing SO2 gas in it. A pink colouration on adding Schiff;s reagent to the compound indicates the presence of aldehyde group.
Tollen’s Test
In the freshly prepared silver nitrate solution (~2%) a few drops of sodium hydroxide solution is added to obtain a dark brown precipitate of silver oxide. Further ammonium hydroxide solution is added to it in drop-wise manner to dissolve the precipitate. The obtained solution is called as Tollen’s reagent.
An aqueous or an alcoholic solution of the organic compound is added to the Tollen’s reagent and heated. Formation of a layer of silver metal on the inner surface of the test tube (silver mirror) indicates the presence of an aldehyde.
Tollen’s test is also known as silver mirror test.
Test for Carboxyl Group [-COOH]
Litmus Test
Blue litmus turns red when a 2-3 drops of carboxylic acid is added to the litmus paper or litmus solution.
Sodium hydrogencarbonate Test
If an organic compound containing carboxylic acid group is added to the aqueous solution of sodium hydrogen carbonate, the carbon dioxide gas is evolved with brisk effervescence.
Ester Test
Organic compound containing carboxylic acid group when added to alcohols in acidic medium then a compound with fruity or sweet smell is formed.
Test for Amino Group [-NH2]
Carbylamine Test
When a primary aliphatic or aromatic amine is heated with chloroform in the presence of alkali, a foul smelling compound (isocyanide/carbylamine) is formed.
Note: Secondary and Tertiary amines do not give positive carbylamine test.
Solubility Test
Due to the basic nature of amino group, amines react with acids to form salts which are soluble in water, hence, it can be used as a test for the amino group.
Azo Dye Test
It is a well known test to identify primary aromatic amines and commonly used confirmatory test for it. The test involves two steps, namely diazotisation and coupling.
Diazotisation: Reaction of primary aromatic amine with HNO2 at 0-5oC to produce diazonium chloride.
Coupling reaction: On addition of phenol, aniline and beta-naphthol to the diazonium salt, coupling reaction takes place to form complex having colour orange, yellow and scarlet red respectively.
The formation of respective coloured complex confirms the formation of aromatic primary amine.
Chemistry involved in the preparation of Inorganic Compounds
Here the preparation of two double salts, i.e., potash alum and mohr’s salt, are given below:
Potash alum
Crystallisation of an equimolar mixture of potassium sulphate and aluminium sulphate results in the formation of potash alum, also known as Potassium Aluminium Sulphate (PAS).
Procedure
- Aluminium sulphate with dilute sulphuric acid is dissolved in 10 ml of distilled water.
- A small amount of powdered potassium sulphate is added and mixture is heated and then cooled to room temperature.
- White crystals of potash alum gets separated on cooling.
- Crystals are washed with mixture of cold water and alcohol and then dried.
Mohr’s Salt
Crystallisation of an equimolar mixture of ferrous sulphate and ammonium sulphate results in the formation of Mohr’s salt, also know as Ferrous Ammonium Sulphate (FAS).
Procedure
- Ferrous sulphate and ammonium sulphate are taken in 2:1 ratio and dissolved in minimum amount of distilled water.
- Few drops of dilute sulphuric acid is added to the solution and it is concentrated by heating till the crystallisation point is reached.
- On cooling, light green crystals of ferrous ammonium sulphate are obtained which are washed with a mixture of cold water and alcohol followed by drying.
Chemistry involved in the preparation of Organic Compounds
Here the preparation of Acetanilide, p-nitroacetanilide, Aniline yellow and Iodoform are given below:
Acetanilide
Acetanilide is prepared from aniline by replacement of one H-atom from -NH2 group with acetyl group (-CO-CH3) using glacial acetic acid. Acetylation of aniline can be done using acetic anhydride or acetyl chloride.
This reaction is carried out in the presence of strong base than amine like pyridine, which removes HCl so formed and shifts the equilibrium towards forward direction.
Procedure
- Aniline is taken in a round bottom flask and mixed with acetic anhydride and glacial acetic acid (acetylating mixture).
- An air condenser on the mouth of the round bottom flask is fitted afetr adding pumice stones and refluxing the mixture gently on a sand bath.
- Next step is cooling the reaction mixture and pouring it slowly in ice cold water and stirring.
- After cooling next step is filtering the solid, washing it with cold water and recrystallizzing a small amount of sample from hot water containing a few drops of methanol and ethanol.
Para-Nitro Acetanilide
Formation of p-nitro acetanilide follows electrophilic aromatic substitution reaction in which nitronium ion (NO2+) acts as an electrophile. In this reaction, the acetanilide reacts with nitrating reagent (mixture of concentrated nitric acid and concentrated sulphuric acid) to get p-nitro acetanilide as major product.
-NHCOCH3 group of acetanilide acts as an electron donating group to benzene ring through +M effect which directs the incoming electrophile to the ortho and para position.
Formation of an electrophile
NO2+ is produced by transfer of proton (from sulphuric acid) to nitric acid acid as follows:
Acetanilide is itself prepared using acetylation of aniline.
Procedure
- Acetanilide is dissolved in glacial acetic acid and taken in a beaker.
- The above mixture becomes hot and clear solution is obtained when conc. H2SO4 is added.
- Cold mixture of conc. HNO3 and conc. H2SO4 is added to the above mixture drop by drop with constant stirring.
- After removing the beaker from ice bath mixture is allowed to attain room temperature.
- The mixture is stirred and the compound obtained is filtered.
- After drying, recrystallise a small amount of the pale yellow solid from alcohol.
- Colourless crystals of p-nitroacetanilide are obtained.
Aniline Yellow
The reaction of diazonium salt with aniline forms p-aminoazobenzene (yellow dye). This reactions follows electrophilic substitution reaction and known as coupling reaction.
Procedure
- Finely powdered diazoaminobenzene and aniline are dissolved in a conical flask.
- Powdered aniline hydrochloride is added to the above mixture.
- The mixture is heated on a water bath for approximately 1 hour.
- Temperature of mixture is lowered in the next step by allowing it to stand at room temperature.
- Excess aniline is converted to its acetate by adding glacial acetic acid and water to the mixture.
- Allow the mixture to stand for 15 minutes with occasional stirring.
- In the next step, p-aminoazobenzene is filtered, washed with little cold water and dried.
- In the last step, a small portion of crude p-aminoazobenzene is recrystallised from carbon tetrachloride.
Iodoform
Iodoform (CHI3) is a pale yellow crystalline solid having a characterstic odour. It is used as a mild antiseptic and disinfectant. It can be prepared by treating any organic compound containing CH3CH(OH)- group, e.g., (ethanol, 2-propanol, 2-butanol) or CH3CO- group (e.g., propanone, 2-butanone, 2-pentanone) with iodine in the presence of sodium hydroxide. In the laboratory, it is usually prepared from either ethanol or propanone. The chemical reactions involved are:
(a) With Ethanol
(b) With Propanone
Qualitative Analysis of Inorganic Salts: Analysis done to understand the nature of the substance and to identify its consitutnets is termed as Qualitative Anslysis. Constituents of inorganic salts are cations and anions. In the qualitative analysis of organic salts, we identify cations and anions present in salt or in the mixture of salts.
Qualititative analysis is carried out through the reactions in which we can easily identify the changes by some of the physical observations like:
(a) Evolution of gas
(b) Change in colour
(c) Formation of a precipitate
For the systematic analysis of an inorganic salt, we follow the following steps:
(a) Preliminary examination of solid salt and its solution.
(b) Determination of anions by reactions carried out in solution (wet tests) and confirmatory tests.
(c) Determination of cations by reactions carried out in solution (wet tests) and confirmatory tests.
Preliminary examination of a salt is not conclusive but generally it gives quite important hints for the presence of certain anions or cations and simplifies the process of analysis. Preliminary tests are termed as ‘dry test’ and include to observe the general appearance and physical properties like colour, odour, solubility, etc.
The first step in the qualitative analysis of a salt includes identification of anions present in it and we proceed with preliminary tests followed by confirmatory analysis. Gases evolved in the preliminary tests with dilute sulphuric acid and concentrated sulphuric acid help to identify the presence of anions (acid radicals).
Systematic Analysis of Anions – Acid Radicals
In the systematic analysis of anions, we will analyze the negative part of the inorganic salt present. The anionic part of the salt comes from acid (e.g., HCl, H2SO4, HNO3, HI, etc.), so the anions are known as acidic radicals.
The analysis of anions include two steps, namely Preliminary test and Confirmatory test.
Preliminary tests are performed under two group reagents, namely dilute Sulphuric Acid and concentrated Sulphuric Acid.
Preliminary test with dilute Sulphuric acid
Dilute Sulphuric acid is the group reagent for the following anions:
- Carbonate ion (CO32-)
- Sulphide ion (S2-)
- Sulphide ion (SO32-)
- Nitrite ion (NO2–)
In this test, the given inorganic salt is treated with dilute sulphuric acid, then the observations obtained are as follows:
| Observation | Gas Evolved | Inference |
|---|---|---|
| Evolution of colourless gas with brisk effervescence | CO2 | CO32- may be present |
| Evolution of colourless gas with smell of rotten eggs | H2S | S2- may be present |
| Evolution of colourless, pungent gas with smell of buring sulphur | SO2 | SO32- may be present |
| Evolution of brown fumes | NO2 | NO2– may be present |
Preliminary test with concentrated Sulphuric acid
Concentrated Sulphuric acid is the group reagent for the following anions:
- Chloride ion (Cl–)
- Bromide ion (Br–)
- Iodide ion (I–)
- Nitrate ion (NO3–)
In this test, the given inorganic salt is treated with concentrated sulphuric acid, then the observations obtained are as follows:
| Observation | Gas Evolved | Inference |
|---|---|---|
| Evolution of pungent smelling colourless gas | HCl | Cl– may be present |
| Evolution of pungent smelling reddish brown gas | Br2 | Br– may be present |
| Evolution of violet coloured gas or violet sublimate deposit | I2 | I– may be present |
| Evolution of brown fumes which gets dense with Cu turnings | NO2 | NO3– may be present |
Chemistry of Confirmatory Tests
Test for Carbonate [CO32-] ion
On adding dilute sulphuric acid to the solid salt, there is evolution of a colourless and odourless gas with brisk effervescence. This indicates the presence of carbonate ion. Evolved gas turns lime water milky due to the formation of CaCO3.
Milkiness disappears on passing excess of CO2 through lime water due to the formation of soluble Ca(HCO3)2.
Test for Sulphide [S2-] ion
(a) With warm dilute sulphuric acid, sulphide salt gives hydrogen sulphide gas which can easily be identified by its rotten egg like smell.
When this gas is placed on a filter paper dipped in lead acetate then the paper turns black due to the formation of lead sulphide.
(b) Check solubility of salt in water:
If salt is water soluble then tale its aqueous solution, make it alkaline by adding ammonium hydroxide solutions and add sodium nitroprusside solution to it.
If salt is water insoluble then take its sodium carbonate extract and add sodium nitroprusside solution to it.
Purple or violet coloured complex compound Na4[Fe(CN)5(NOS)] is formed which confirms the presence of sulphide ion in the salt.
Sodium Carbonate Extract: To make this extract, mix salt with 3-part sodium carbonate and add distilled water to it. Boil the mixture for 10 minutes and filter. Filtrate is known as sodium carbonate extract.
Test for Nitrite [NO2–] ion
(a) When solid salt containing nitrite anion is heated with dilute sulphuric acid, reddish brown fumes of NO2 gas are evolved.
In the salt solution, if potassium iodide followed by freshly prepared starch solution is added and acidification with acetic acid is done then blue colouration confirms nitrite ion. The same test can also be done by taking a filter paper moistened with potassium iodide and starch solution and a few drops of acetic acid. When evolved gas comes in contact with that filter paper, it turns blue due to the interaction of liberated iodine with starch.
(b) Sulphanilic acid-1-naphthylamine reagent test (Griss-Ilosvay test):
Take nitrite salt and acidified with acetic acid and add it to the sulphanilic acid and 1-naphthylamine. Acidification of nitrite salt will form nitrous acid which leads to diazotisation of sulphanilic acid. Diazotised acid couples with 1-naphthylamine to form red coloured azo dye.
The test solution should be very dilute. In concentrated solution, reaction does not proceed post diazotisation.
Test for Chloride [Cl–] ion
(a) Evolution of a colourless gas with pungent smell, on treatment with warm concentrated sulphuric acid. When evolved gas is passed in ammonia solution then it gives dense white fumes.
(b) Evolution of light greenish yellow gas with effervescence on treatment with MnO2 and conc. H2SO4.
(c) Acidify the salt solution with dilute nitric acid and add silver nitrate to it. A curdy white precipitate which is soluble in ammonium hydroxide solution is obtained.
(d) Chromyl Chloride Test:
A little amount of salt is mixed with an equal amount of solid potassium dichromate and a few drops of conc. sulphuric acid. Evolved gas is passed through sodium hydroxide solution. Obtained yellow coloured solution is divided into two parts:
(i) First part is acidified and lead acetate is added to it. Formation of yellow precipitate of lead chromate confirms the presence of chloride in salt.
(ii) Second part is acidified with dilute sulphuric acid. Further addition of amyl alcohol followed by 10% hydrogen peroxide solution turns the organic layer blue.
Reason for Blue colouration: CrO42- ions formed in the reaction mixture react with NaOH and H2O2 to form chromium pentoxide (CrO5). Formed CrO5 gets dissolved in amyl alcohol to give blue colour.
Test for Bromide [Br–] ion
(a) Presence of bromide ions can be identified by the evolution of brown fumes of bromine on heating the salt with conc. sulphuric acid. Evolution of fumes gets intense on addition of MnO2 to reaction mixture.
Starch paper turns yellow when comes in contact with these bromine vapours.
(b) Take the salt solution in water or take its sodium carbonate extract and neutralize it with dilute HCl. Prepare a mixture of carbon tetrachloride or chloroform and freshly prepared chlorine water and add to the water extract or sodium carbonate extract of salt in dropwise manner. Brown colouration in organic layer shows the presence of bromide ions.
(c) Acidified sodium carbonate extract of the salt is taken and silver nitrate solution is added to it. Formation of yellow precipitate (soluble in ammonium hydroxide) shows the presence of bromide ions.
Test for Iodide [I–] ion
(a) Evolution of violet coloured vapours of iodine with a pungent smell on heating the salt with conc. sulphuric acid shows the presence of iodide ions. Addition of MnO2 to the reaction mixture intensify the evolution of violet vapours. These vapours turn starch paper blue and leave violet sublimate on the sides of the test tube.
Along with iodine vapours some HI, SO2, H2S and Sulphur is also formed.
(b) Take the salt solution in water or take its sodium carbonate extract and neutralize it with dilute HCl. Prepare a mixture of carbon tetrachloride or chloroform and freshly prepared chlorine water and add to the water extract or sodium carbonate extract in dropwise manner. Violet colouration in organic layer shows the presence of iodide ions (Iodide is soluble in organic solvents).
(c) Acidified sodium carbonate extract of the salt is taken and silver nitrate solution is added to it. Formation of yellow precipitate (insoluble in ammonium hydroxide) shows the presence of iodide ions.
Test for Nitrate [NO3–] ion
(a) When nitrate salt is heated with conc. sulphuric acid, light brown fumes are evolved. Excess of brown fumes are evolved when salt is heated with few copper turnings and conc. sulphuric acid. The colour of the solution also turns blue due to the formation of copper sulphate.
(b) Brown Ring Test:
Small amount of aqueous solution of the salt is mixed with freshly prepared ferrous sulphate solution and conc. sulphuric acid is added slowly to it along the sides of the test tube. It results in the formation of a dark brown ring at the junction of the two solutions due to the formation of nitroso ferrous sulphate.
Alternatively, first conc. sulphuric acid can be mixed with aqueous salt solution in a dropwise manner and add freshly prepared ferrous sulphate solution to it along the side of the test tube post cooling the mixture. It will also form a brown coloured ring.
Test for Sulphate [SO42-] ion
If salt neither reacts with dilute sulphuric acid nor with conc. sulphuric acid, then it may contain sulphate anions. The following confirmatory tests are done in this case.
(a) A white precipitate of barium sulphate is formed when an aqueous solution or sodium carbonate extract of the salt is acidified with acetic acid and mixed with barium chloride. Precipitate of barium sulphate is insoluble in conc. HCl or conc. HNO3.
(b) A white precipitate of lead sulphate is formed when an aqueous solution or sodium carbonate extract of the salt is neutralised with acetic acid and treated with lead acetate solution.
Systematic Analysis of Cations – Basic Radicals
In the systematic analysis of cations, we will analyse the positive part of the inorganic salt present. The cationic part of the salt comes from base (e.g., Al(OH)3, Pb(OH)2, Mg(OH)2, NH4OH, etc.), so the cations are known as basic radicals.
Preliminary Tests of the Salt for the Identification of Cations
Colour Test
Colour of the salts give useful information about metal ions. Characteristic colours of some metal ions are given as:
| Colour | Cations |
|---|---|
| Light green, Yellow, Brown | Fe2+, Fe3+ |
| Blue | Cu2+ |
| Bright green | Ni2+ |
| Blue, Red. Violet, Pink | Co2+ |
| Light Pink | Mn2+ |
Dry Heating Test
Observe colour of the salt in hot and cold conditions.
| Colour when cold | Colour when hot | Inference |
|---|---|---|
| Blue | White | Cu2+ |
| Green | Dirty white or yellow | Fe2+ |
| White | Yellow | Zn2+ |
| Pink | Blue | Co2+ |
Flame Test
Metal salts impart characterstic colour to the flame test. Platinum wire is used to perform the test.
| Colour of the flame observed by naked eye | Colour of the flame observed through blue glass | Inference |
|---|---|---|
| Green flame with blue center | Same colour as observed without glass | Cu2+ |
| Crimson red | Purple | Sr2+ |
| Apple green | Bluish green | Ba2+ |
| Brick red | Green | Ca2+ |
Borax Bead Test
Borax reacts with metal salts to form metal borates which have characterstic colour. Coloured salts are considered for this test.
| Colour of salt bead when heating in oxidising (non-luminous) flame [In hot] | Colour of salt bead when heating in oxidising (non-luminous) flame [In cold] | Inference |
|---|---|---|
| Green | Blue | Cu2+ |
| Violet | Reddish brown | Ni2+ |
| Light violet | Light violet | Mn2+ |
| Yellowish brown | Yellow | Fe3+ |
| Colour of salt bead when heating in reducing (luminous) flame [In hot] | Colour of salt bead when heating in reducing (luminous) flame [In cold] | Inference |
|---|---|---|
| Colourledd | Red opaque | Cu2+ |
| Grey | Grey | Ni2+ |
| Colourless | Colourless | Mn2+ |
| Green | Green | Fe3+ |
Charcoal Cavity Test
On heating metal carbonates in a charcoal cavity, the carbonates are converted to metal oxide, which are coloured residues. Sometimes metal oxides are converted to metal on carbon reduction within the charcoal cavity.
| Observations | Inference |
|---|---|
| Yellow residue when hot and grey metal when cold | Pb2+ |
| White residue with the odour of garlic | As3+ |
| Brown residue | Cd2+ |
| Yellow residue when hot and white when cold | Zn2+ |
Cobalt Nitrate Test
If the residue in the charcoal cavity is white, cobalt nitrate test is performed.
(i) The residue is treated with few drops of cobalt nitrate solution.
(ii) Heat the mixture with the help of a blow pipe and observe the colour of the residue.
On heating, cobalt nitrate decomposes into CoO, which gives a characterstic colour on reaction with metal oxide present in the charcoal cavity.
Wet Test for Identification of Cations
In the first step, prepare the clear aqueous solution of the salt. This is called original solution. Cations of different groups with group reagent have been depicted in the table given below:
| Group | Cations | Group Reagent |
|---|---|---|
| Group Zero | NH4+ | None |
| Group I | Pb2+ | Dilute HCl |
| Group II | Pb2+, Cu2+, As3+ | H2S gas in the presence of dil. HCl |
| Group III | Al3+, Fe3+ | NH4OH in the presence of NH4Cl |
| Group IV | Co2+, Ni2+, Mn2+, Zn2+ | H2S in the presence of NH4OH |
| Group V | Ba2+, Sr2+, Ca2+ | (NH4)2CO3 in presence of NH4OH |
| Group VI | Mg2+ | None |
To identify metal cations, these are precipitated from the original solution by using the group reagents according to the flow chart given below:
Analysis of Group-0 Cation
Small amount of salt is treated with few drops of sodium hydroxide solution and then heated. If there is smell of NH3, this indicates the presence of ammonium ions.
Confirmatory Tests for NH4+ ion
(a) Ammonia gas evolved by the action of sodium hydroxide on ammonium salts reacts with hydrochloric acid to give ammonium chloride, which is visible as dense white fumes.
(b) On passing the gas through Nessler’s reagent, a brown precipitate of basic mercury (II) amido-iodine is formed.
Analysis of Group-I Cation
In original solution, add dilute HCl, white precipitate of PbCl2 forms indicate the presence of Pb2+ ions. The precipitate is soluble in hot water.
Confirmatory Tests for Pb2+ ion
(a) To the hot solution, potassium iodide solution is added which results in a yellow precipitate of lead iodide. This confirms the presence of Pb2+ ions.
This yellow precipitate is soluble in boiling water and reappears on cooling as shining crystals. This yellow precipitate is also soluble in excess of KI and forms K2[PbI4].
(b) To the hot solution, potassium chromate solution is added which results in a yellow precipitate of lead chromate. This confirms the presence of Pb2+ ions.
Lead chromate is soluble in hot sodium hydroxide solution.
(c) Addition of alcohol followed by treatment with dilute H2SO4 solution gives white precipitate of PbSO4.
Lead sulphate is soluble in ammonium acetate solution due to the formation of tetraacetoplumbate (II) ions.
Analysis of Group-II Cation
If Group-I metal cation is absent, and few ml of water to the same test tube. Warm the aqueous solution and pass H2S gas for few minutes and shake the test tube. If a precipitate is observed, this indicates the presence of Group-II cations. Pass excess of H2S gas to the solution for complete precipitation.
Confirmatory test for Pb2+ ion
Lead sulphide precipitate is dissolved in dilute HNO3. To it dilute H2SO4 and a few drops of alcohol is added to solution, a white precipitate of lead sulphate appears. This indicates the presence of lead ions.
Confirmatory test for Cu2+ ion
(a) Copper sulphide precipitate dissolves in nitric acid due to the formation of copper nitrate.
On heating the reaction mixture, sulphur obtained is oxidised to sulphate resulting in the formation of copper sulphate and the solution turns blue. Addition of small amount of ammonium hydroxide, precipitates basic copper sulphate which is soluble in excess of ammonium hydroxide solution due to the formation of deep blue tetraamminecopper(II) sulphate.
(b) Addition of acetic acid to the blue solution followed by treatment with K4[Fe(CN)6] gives a chocolate colouration due to the formation of copper ferrocyanide.
Analysis of Group-III Cation
If Group-II is absent, take original solution and add few drops of conc. nitric acid and heat to oxidise Fe2+ ions to Fe3+ ions. After cooling add a small amount of solid ammonium chloride and excess of ammonium hydroxide solution and shake the test tube till small amount of NH3 arises. If a brown or white precipitate is formed, this indicates the presence of Group-III metal cations.
Group-III metal cations are precipitated as their hydroxides, which dissolve in dilute hydrochloric acid due to the formation of corresponding metal chlorides.
Confirmatory test for Al3+ ion
Aqueous solution of AlCl3 is treated with sodium hydroxide, a white gelatinous precipitate of aluminium hydroxide is formed which is soluble in excess of sodium hydroxide solution due to the formation of sodium aluminate.
Confirmatory test for Fe3+ ion
Ferric hydroxide, the reddish brown precipitate, dissolved in HCl and ferric chloride is formed. Divide the FeCl3 solution into two parts.
(a) To the first part of ferric chloride solution, add potassium ferrocyanide solution, which results in the formation of blue precipitate. It is ferric ferro cyanide, which is prussian blue in colour. It confirms the presence of Fe3+ ion.
(b) To the second part of ferric chloride solution, add potassium thiocyanate solution, which results in the formation of blood red colouration. It confirms the presence of Fe3+ ion.
Analysis of Group-IV Cation
If Group-III is absent, H2S gas is passed in the solution of Group-III for a few minutes. If a precipitate appears (white or black) then the Group-IV cations are present.
Group-IV cations are precipitated as their sulphides. A white colour of the precipitate indicates the presence of Zn2+ ions while black colour indicates Ni2+ ions.
Confirmatory test for Zn2+ ion
The precipitate, ZnS, dissolves in HCl to form zinc chloride.
(a) Aqueous solution of ZnCl2 on reaction with NaOH solution gives a white precipitate of zinc hydroxide, which is soluble in exvess of NaOH solution on heating. This confirms the presence of Zn2+ ions.
(b) When potassium ferrocyanide solution is added to the solution after neutralisation by NH4OH solution, a white precipitate of zinc ferrocyanide is formed. This confirms the presence of Zn2+ ions.
Confirmatory test for Ni2+ ion
Nickel sulphide dissolves in aqua regia.
When dimethyl glyoxime (dmg) is added to the aqueous solution of nickel chloride and is made alkaline by adding NH4OH solution, a brilliant red precipitate [Ni(dmg)2] is obtained.
Analysis of Group-V Cation
If Group-IV cations are absent then take original solution and add a small amount of solid NH4Cl and an excess of NH4OH solution followed by solid ammonium carbonate. Appearance of white precipitate indicates the presence of Group-V cations.
Group-V cations are precipitated as their carbonates which dissolve in acetic acid due to the formation of corresponding acetates.
Confirmatory test for Ba2+ ion
(a) Precipitate BaCO3, on treatment with acetic acid followed by addition of K2CrO4 results in yellow precipitate of barium chromate.
(b) Flame test: It gives apple green colour to the flame test.
Confirmatory test for Ca2+ ion
(a) Group-V precipitate is dissolved in acetic acid which is then tretaed with ammonium oxalate. A white precipitate is obtained which confirms Ca2+ ion.
(b) Flame test: Calcium imparts brick red colour to the flame.
Analysis of Group-VI Cation
If Group-V is absent, the solution may contain magnesium carbonate, which is soluble in water in the presence of ammonium salts because the equilibrium is shifted towards the right side.
The concentration of carbonate ions required to produce a precipitate is not sufficient to the given solution. Disodium hydrogenphosphate solution is added and the inner wall of the test tube is scratched with a glass rod, appearance of white crytalline precipitate of magnesium ammonium phosphate confirms the presence of Mg2+ ions.
Acids and Bases
Acids and Bases are popular chemicals which react with each other to form salt and water.
Acidic substances are usually identified by their sour taste. An acid is basically a molecule which can donate a H+ ion and can remain energetically favourable after a loss of H+ ion. Acids are known to turn blue litmus red.
Bases, on the other hand, are characterised by abitter taste and a slippery texture. A base that can be dissolved in water is referred to as an alkali. Bases are known to turn red litmus blue.
Theories of Acid and Bases:
- The Arrhenious theory of acids and bases states that “an acid generates H+ ions in its aqueous solution whereas a bases produces an OH– ion in its aqueous solution”.
- The Bronsted-Lowry theory of acids and bases defined “an acid as a proton donor and a base as a proton acceptor”.
- The Lewis Acid theory of acids and bases describes “acids as electron pair acceptors and bases as electron pair donors.
pH of Acids and Bases
In order to find the numeric value of the level of acidity or basicity of a substance, the pH scale can be used. The pH scale is the most common and trusted way to measure how acidic or basic a substance is. A pH scale measure can vary from 0 to 14 at 25 0C, where 0 is the most acidic and 14 is the most basic a substance can be. Another way to check if a substance is acidic or basic is to use litmus paper.
There are two types of litmus paper available that can be used to identify acids and bases – red litmus paper and blue litmus paper. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic or alkaline conditions.
Indicators in Acid Base Titration
Acid base indicators are sensitive to pH change. For most acid base titrations, it is possible to select indicators which exhibit colour change at pH close to the equivalence point. We will discuss here about only two indicators – Phnolphthalein and Methyl Orange.
Phenolphthalein
Phenolphthalein is a weak acid, therefore it does not dissociate in the acidic medium and remains in the unionised form, which is colourless.
Ionised and unionised forms of phenolphthalein are given below:
In acidic medium, equilibrium lies to the left. In the alkaline medium, the ionisation of phenolphthalein increases considerably due to the constant removal of H+ ions released from HPh by the OH– ions from the alkali. So the concentration of Ph- ion increases in the solution, which imparts pink colour to the solution.
For a weak acid vs strong base titration, phenolphthalein is the most suitable indicator. This is so because the last drop of added alkali brings the pH of the solution in the range in which phenolphthalein shows sharp colour change.
Methyl Orange
Methyl orange is a weak base and is yellow in colour in the unionised form. Sodium salt of methyl orange is represented as:
The anion formed from the indicator is an active species, which on accepting a proton (i.e., acting as Bronsted base) changes from the benzenoid form to the quinoid form. The quinoid form is deeper in colour and thus is responsible for the colour change at the end point. This is illustrated in the following manner:
pH range of indicators
| Name of indicator | Colour in acidic medium | Colour in basic medium | pH range |
|---|---|---|---|
| Methyl orange | Orange red | Yellow | 3.1 – 4.5 |
| Methyl red | Red | Yellow | 4.2 – 6.2 |
| Phenol red | Yellow | Red | 6.2 – 8.2 |
| Phenolphthalein | Colourless | Pink | 8.2 – 10.2 |
Acid – Base Titration
An acid-base titration is an experimental technique used to acquire information about a solution containing an acid or base.
Important point of titration
- Titration is always possible in two opposite nature solutions, i.e, one is acidic and other is basic.
- For any titration only that indicator is suitable if their working pH range is in pH range of titration.
- At the equivalence point of titration, number of equivalents of acids and bases are equal.
- At the equivalence point, nature of solution depends on the type of titration.
Strong acid/Strong base – Neutral (pH = 7) at 25 oC
Strong acid/Weak base – Acidic
Weak acid/Strong base – Basic
Weak acid/Weak base – Acidic/Basic
(i) Strong acid – Strong base titration: In the titration of HCl with NaOH, the equivalence point lies in the pH range of 3-11. Thus methyl orange, methyl red and phenolphthalein will be suitable indicators.
(ii) Weak acid – Strong base titration: In the titration of CH3COOH with NaOH, the equivalence point lies in the pH range 7-11. Thus, phenolphthalein will be the suitable indicator.
(iii) Weak base – Strong acid titration: In the titration of NH4OH against HCl, the equivalence point lies in the pH range 3-7. Thus, methyl orange and methyl red will be the suitable indicators.
(iv) Weak acid – Weak base titration: In the titration of the weak acid with weak base, the pH at the equivalence point is about 7, i.e., lies between 6.5 and 7.5, but no sharp change in pH is observed in these titrations. Thus, no simple indicator can be employed for the detection of the equivalence point.
Titrimetric Analysis (Redox Titration)
A reaction which involves simultaneous oxidation and reduction is called redox reaction. The titrations involving redox reactions are called redox titrations. In acid -base titrations, indicators which are sensitive to pH change are used to identify the end point. Similarly, in redox titrations, there is a change in oxidation potential of the system. The indicators used in redox reactions are sensitive to change in oxidation potential intermediate between the values for the solution being titrated and the titrant and these show sharp readily detectable colour change.
- Titration: It is a process of quantitative analysis which is carried out by determining the volume of standard solution which is required to react quantitatively with known volume of a solution to be determined.
- Titrant: A reagent of known concentration (standard solution) in the solution.
- Titrand: The substance being titrated is termed as titrand or analyte.
- Titration curve: A plot of pH vs millilitres of titrant showing the manner in which pH changes vs millilitres of titrant during acid-base titration.
- Equivalence point: The point at which chemical reaction is complete and number of equivalents of titrant is added to number of equivalents of analyte.
- End point: Point at which indicator changes colour which usually occurs after few milliseconds after equivalence point.
Titration of Oxalic Acid (H2C2O4) against Potassium Permanganate (KMnO4)
The titration of KMnO4 vs Oxalic acid is a good exampleof the oxidation – reduction reaction. KMnO4 solution is considered a strong standard oxidising agent and H2C2O4 as strong reducing agent.
For the experiment , let 0.1 M standard solution of oxalic acid is to be prepared. The formula of hydrated crystalline oxalic acid is (COOH)2 . 2H2O and ts molar mass is 126 g. If 126 g of oxalic acid is present in one liter of the solution, it is known as one molar (1.0 M) solution. For the preparation of one liter of 0.1 M oxalic acid solution, we require 12.6 g of hydrated oxalic acid.
Following reactions occur during the titration:
Here MnO4– is reduced to Mn2+ and C2O42- is oxidised to CO2. The oxidation number of carbon in C2O42- changes from +3 to +4.
The strength of the unknown solution in terms of molarity may be determined by the following equation:
n1M1V1 = n2M2V2 …(i)
For oxalic acid vs potassium permanganate titration:
n1 = 2 (the number of electrons lost per formula unit of oxalic acid in the balanced equation of half cell reaction)
n2 = 5 (the number of electrons gained per formula unit of potassium permanganate in the balanced equation of half cell reaction)
M1 and M2 are molarities of oxalic acid and potassium permanganate solutions used in the titration.
V1 and V2 are the volumes of oxalic acid and potassium permanganate solutions.
On putting the values of n1 and n2 in the equation (i), we get
Number of equivalents of oxalic acid = Number of equivalents of KMnO4
2M1V1 = 5 M2V2
M2 = (2M1V1) / (5V2)
We can calculate the molarity of potassium permanganate solution by the above equation, and the strength of the solution is given by:
Strength = Molarity x Molar Mass
If one of the reactants has some characteristic intense colour, then no external indicator is added for the indication of completion of reaction. For example, KMnO4 which has an intense violet colour. As soon as the reaction complete, one extra drop of KMnO4 from burette is capable of changing the colour of the solution. Titration of oxalic acid solution against potassium permanganate solution.
Titration of Ferrous Ammonium Sulphate (FeSO4.(NH4)2SO4.6H2O) against Potassium Permanganate (KMnO4)
Ferrous Ammonium Sulphate (FAS) also acts as a reducing agent in the titration against potassium permanganate.
To prepare 0.05 M standard solution of FAS, following steps are followed:
(i) Weigh 4.90 g of FAS and transfer it into a 250 ml measuring flask through a funnel.
(ii) Transfer the solid sticking to the funnel with the help of distilled water into the flask and add dilute H2SO4 into the flask drop wise to get the clear solution and shake the flask till the substance dissolves.
Following reaction occur during the titration:
The oxidation number of iron in FAS (Mohr’s salt) is +2. Iron isoxidised during the reaction and its oxidation number changes from +2 to +3.
From the overall balanced chemical equation, it is clear that 1 moles of potassium permanganate react with 5 moles of Mohr’s salt.
Therefore,
Chemical Principles Involved in Some Experiments
Enthalpy of Solution of CuSO4
Generally reactions in lab are carried out at atomspheric pressure so heat changes observed during these reactions are enthalpy changes. The relationship between enthalpy change and temperature is given by formula.
ΔH = qp
ΔH = mCpΔT
ΔH = VdCpΔT
where V = Volume of the solution, d = Density of the solution, Cp = Heat capacity, ΔT = Change in temperature
Calorimeters are the vessels that are used to measure the heat changes. Thermochemical reactions are generally carried out in beakers placed in thermos flask or thermally insulated boxes. Use of metallic calorimeters is not preferred for measuring thermochemical changes because of tendency of metals to react with substances. So stainless steel or gold plated copper calorimeters can be used. The amount of heat absorbed by calorimeter, thermometer and stirrer during the measurement of heat change is called calorimeter constant.
Glass is a material with low thermal conductivity, for such materials calorimeter constant for that part is found which is in direct contact with reaction mixture. It has been observed that when calorimeter is made up of low thermal conductivity material, maximum heat is absorbed by area in direct contact with liquid.
In order to determine the calorimeter constant, specified volume of hot water at particular temperature is mixed with specified volume of water contained in the calorimeter at room temperature. Heat absorbed by calorimeter and cold water is equal to heat released by hot water, this is according to the law of conservation of energy.
So the equation can be written as-
ΔH1 + ΔH2 = – ΔH3
where, ΔH1 = Enthalpy change of calorimeter, stirrer and thermometer (calorimeter constant), ΔH2 = Enthalpy change of cold water, ΔH3 = Enthalpy change of hot water.
In thermochemical measurement, mostly aqueous solutions are mixed. So water act as medium for which temperature change is due to chemical reactions taking place in solution.
The sum of enthalpy changes taking place as either gain or loss in energy in the calorimeter must be zero. Therefore,
ΔH1 + ΔH2 + ΔH3 + ΔH4 = 0
where. ΔH4 = Enthalpy change of reaction
In these reactions, the product of density and heat capacity of solutions, dCp = 4.1814 J ml-1 K-1. Enthalpy of solution is defined as the amount of heat released or absorbed when one mole of a solute is dissolved in such a large quantity of solvent that there is no heat change on further dilution.
Enthalpy of Neutralisation of Strong Acid and Strong Base
During the neutralisation reaction, H+ ions furmished by acid combines with OH– furnished by base combines to form H2O. This reaction is always exothermic as it involves bond formation between H+ and OH–. Enthalpy of neutralisation is defined as amount of heat released when 1 mole H+ ions combines with 1 mole of OH– and forms water.
H+ (aq) + OH– (aq) → H2O (l), Enthalpy of nuutralisation (ΔneutH) = -ve
If both the acid and base undergoing neutralisation reaction are strong, then 57 kJ of heat is released for the formation of 1 mole of H2O. So, ΔneutH for strong acid and base is 57 kJ/mol.
If any one of the acid or base is weak or if both of them are weak, then some amount of heat released is used for ionisation of the acid or base or both of them and heat released is less than 57 kJ/mol.
Preparation of Lyophilic and Lyophobic Sols
A colloidal solution is a heterogenous system in which one phase (dispersed phase) is dispersed as very fine particles in another system (dispersion medium). Colloidal particles are so small that they remain suspended in medium. Generally, size of colloidal particles ranges from 1-1000 nm. A colloidal system having solid and liquid as dispersed phase and dispersion medium respectively are called sol.
Based on the nature of interaction between dispersed phase and dispersion medium, there are two types of colloidal sol. First is lyophilic sol means liquid loving sol and second is lyophobic sol means liquid hating sol. Lyophilic sol can be prepared by mixing the two phases together. Example of lyophilic sol are egg albumin, starch and gum. Lyophobic sol cannot be prepared directly by mixing. They can only be prepared by special methods. Examples of lyophobic sold are freshly prepared ferric hydroxide, aluminium hydroxide, arsenic sulphide, etc.
Lyophilic sols are more stable than lyophobic sols. Lyophilic sols are also called reversible sols because if dispersed phase and dispersion medium has separated, the sol can be reconstituted simply by mixing. This is not the case with lyophobic sol, so lyophobic sol are called irreversible sol. Therefore, lyophobic sol need stabilising agents for their preservation. Two factors responsible for the stability of sol are charge and solvation of the colloidal particles by the solvent.
Lyophilic sols are stabilised due to solvation of colloidal particles by dispersion medium while stability of lyophobic sol is due to charge on colloidal particles. The charge on the colloidal particles prevents coagulation of colloidal particles and helps them to remain in medium. The charges can be negative or positive. Starch and arsenious sulphide are negatively charged sol. Addition of FeCl3 to excess of hot water leads to formation of positively charge sol while negatively charged sol of hydrated ferric oxide is formed when ferric chloride is added to NaOH solution.
Kinetic Study of the reaction of Iodide ions with Hydrogen Peroxide at room temperature
The reaction between hydrogen peroxide and iodide ion in acidic medium can be represented as:
I– has been oxidised to I2 so the above reaction shows oxidising nature of H2O2. On addition of calculated amount of sodium thiosulphate in the presence of starch solution, the released I2 reacts with thiosulphate and is reduced back to I2, till all the thiosulphate ions are oxidised to tetrathionate ions.
When all the thiosulphate ions are consumed, the concentration of iodine released in the reaction of H2O2 with I- increases rapidly to a point where iodine forms intense blue complex with starch. The time required to consume a fixed amount of thiosulphate can be reproduced. As the time for the appearance of colour is noted, the reaction is also referred as clock reaction.
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FAQs
What is Practical Chemistry?
Practical Chemistry refers to the hands-on experimental side of chemistry – where theoretical concepts are applied in the laboratory through tests, preparation of compounds, titrations, qualitative analyses, etc. It’s a key part of the syllabus for classes like 12 and important in competitive exams.
Why is the topic “Principles Related to Practical Chemistry” important for students?
Because it bridges theory and practice. Understanding the underlying principles helps students not only perform experiments correctly but also interpret results, avoid errors, and succeed in board exams (and competitive exams like JEE/NEET) as noted in this article.
What are the key components or sections covered in this topic?
The article outlines several major sections:
-Detection of elements (Nitrogen, Sulphur, Halogens) in organic compounds.
-Detection of functional groups (alcoholic, phenolic, carbonyl, carboxyl, amino groups).
-Preparation of inorganic compounds (e.g., potash alum, Mohr’s salt) and organic compounds (e.g., acetanilide, p-nitro acetanilide, aniline yellow, iodoform) with their chemistry.
-Qualitative analysis of salts: systematic analysis of anions and cations.
-Titration (acid-base, redox) and related chemical principles like enthalpy, kinetics, sols.
What is the procedure for detection of nitrogen, sulphur and halogens in organic compounds?
In brief:
-You prepare the “Lassigne’s extract” (sodium fusion extract) when analysing an organic compound for N/S/halogen.
-For nitrogen: the extract is treated with ferrous sulphate then acidified and reacted with ferric chloride to give Prussian blue if N is present.
-For sulphur: tests like sodium nitroprusside (violet colour) or lead acetate (black PbS) from the extract show S.
-For halogens: after removal of N, S, the extract is treated (e.g., with AgNO₃, ammonium hydroxide) and or a layer test with CCl₄ or CS₂ + chlorine water to spot Cl, Br, I.
What functional group tests are covered?
Yes — the article covers many:
-Alcoholic group (R–OH): tests like ceric ammonium nitrate, Lucas test, iodoform test for α-methyl alcohols.
-Phenolic group (Ar–OH): ferric chloride test, phthalein dye test.
-Aldehydic & Ketonic (–CHO / –CO–): 2,4-DNP test; distinguishing tests (Fehling’s, Schiff’s, Tollen’s) for aldehydes.
-Carboxyl (–COOH): litmus test, sodium hydrogen carbonate test, ester test.
-Amino group (–NH₂): carbylamine test, solubility test, azo-dye test.
What preparations of compounds are included and why are they significant?
This article includes:
-Inorganic compounds: e.g., potash alum (potassium aluminium sulphate) and Mohr’s salt (ferrous ammonium sulphate). These show crystallisation, double salt formation & link to theory.
-Organic compounds: e.g., acetanilide (acetylation of aniline), p-nitroacetanilide (nitration), aniline yellow (azo coupling dye), iodoform (from certain alcohols/ketones). These illustrate reaction mechanisms, etc.
These are significant because they help students understand preparation, purification, reaction chemistry and link to exam-questions.
What does qualitative analysis of salt involve?
Qualitative analysis refers to identifying the constituent ions (anions and cations) in inorganic salts, through a sequence of preliminary (dry/wet) tests followed by confirmatory tests.
-Anions: e.g., CO₃²⁻, S²⁻, NO₂⁻, Cl⁻, Br⁻, I⁻, NO₃⁻, SO₄²⁻. The article shows tests using acid treatments, precipitates, colours.
-Cations: using group reagents (-NH₄⁺, Pb²⁺, Cu²⁺, Al³⁺, Fe³⁺, Zn²⁺, Ni²⁺, Ba²⁺, Ca²⁺, Mg²⁺) and tests like flame test, borax bead test, charcoal cavity, precipitation groups.
How does titrimetric analysis feature in the article?
Titrimetric analysis (such as acid-base titrations and redox titrations) is explained:
For acid-base: concepts like indicators, equivalence point, working pH of indicators.
Examples: Oxalic acid vs KMnO₄, Mohr’s salt vs KMnO₄.
Understanding these helps students perform titrations and interpret results properly.
What additional “chemical principles” are covered beyond the basic tests and preparations?
Yes — the article also includes: reaction kinetics (e.g., iodide ions with hydrogen peroxide), sol systems (lyophilic/lyophobic sols), thermochemistry (enthalpy of solution, enthalpy of neutralisation) in the context of experiments. This adds depth by linking experimental evidence to underlying theory.
How should a student best use this article for exam preparation?
Here are some suggestions:
Read and understand each test and preparation: know what the observation means and why it happens (not just rote).
-Make short notes or flashcards for each detection/preparation/test: what reagent is used, what observation (colour/fume/precipitate), what inference.
-Practice writing clear step-wise procedures (preparation of salts/compounds) as often exam questions ask for “write the steps” or “give the chemistry involved”.
-For qualitative analysis of salts: understand the flow-chart/grouping logic (preliminary → confirmatory) and remember key distinguishing features (e.g., S²⁻ gives rotten egg smell; CO₃²⁻ gives effervescence with dilute acid).
-For titrations: be sure you understand equivalence point, indicator choice, pH ranges.
-Use the article’s diagrams/images (or your own sketches) to visualise tests: this helps recall.
-Regularly revisit and quiz yourself (or have a friend ask observations).
-Lastly, practice past-year questions (if available) referencing this syllabus to apply knowledge.
Are there any common mistakes students make with practical chemistry and how does the article help avoid them?
Yes — some frequent mistakes:
-Mixing up reagent and outcome (e.g., what colour precipitate means what ion).
-Skipping the “why” behind the observation and just memorising.
-Ignoring the sequence of tests (especially for salts) which can lead to wrong inference.
-Using wrong indicator or misunderstanding equivalence point in titrations.
-Not writing neat procedure or missing key steps (like washing crystals, igniting, cooling).
The article helps by providing structured steps, clear observations, and linking to theoretical rationales, thereby helping students avoid pure memorisation and favour understanding.
Does the article cover safety precautions or laboratory best-practices?
From the content reviewed, the focus is on the principles, tests, preparations and analyses. While many procedures are described (e.g., heating, acid treatment), specific safety precautions may not be emphasised in depth. Students should always follow standard laboratory safety norms (like wearing goggles, gloves, working under fume hood where required, disposing chemicals properly) alongside this material.
How does this article integrate with the rest of the chemistry syllabus (for Class 12)?
This topic forms the ‘practical’ component of the curriculum alongside theoretical chapters. It complements topics like organic/inorganic/physical chemistry by offering experimental verification of concepts (e.g., functional groups, salt solubility, reaction kinetics). For board exams and competitive exams (which often have practical‐based questions), mastering this section is critical. The article outlines the full syllabus in its “Practical Chemistry Syllabus” section.
Is this article suitable for self-study by students?
Yes — especially for motivated students preparing for exams like JEE, NEET, or board exams. The content is structured, covers theory plus procedures plus observations, and includes many visuals (in the original article) that help. However, students should also attempt hands-on practice or virtual labs, and solve problems, to supplement.
What are the most important topics in “Principles Related to Practical Chemistry” for exams like JEE Main / NEET?
Key focus areas include:
-Detection of elements (e.g., nitrogen, sulphur, halogens) in organic compounds.
-Detection of functional groups in organic compounds.
-Preparation of inorganic and organic compounds.
-Volumetric (titrimetric) analysis.
-Qualitative analysis of salts (anions & cations).
How many marks or questions does this chapter generally carry in JEE/NEET?
According to recent reviews, this chapter usually contributes 1 question in the JEE Main exam in each of the last several years and 1-2 question in NEET exam. So while the weightage may seem small, the level of difficulty can be moderate — making it strategic to prepare well.
What are some common MCQ themes students should prepare for in this chapter?
Examples include:
-Choosing proper indicator for a titration (strong acid vs strong base etc.).
-Which glassware or apparatus gives more accurate measurement (e.g., burette vs beaker etc.).
-What reagent is used for detecting particular ion (e.g., AgNO₃ for Cl⁻).
-What principle underlies a separation technique (e.g., distillation is based on difference in boiling points).
Why is understanding the “principle” behind an experiment so important?
Because in exams and in practicals you’re not just asked what you did, but why you did it. For example:
Knowing that in distillation you separate liquids due to differences in boiling points – helps you apply the method correctly.
In qualitative salt analysis, knowing why certain precipitates form or why gas evolves helps avoid rote mistakes.
What mistakes do students frequently make in this chapter and how to avoid them?
Common mistakes:
-Memorising observations without linking them to underlying principle → leads to confusion when variations appear.
-Not practising varied MCQs so unfamiliar phrasing confuses them.
-Ignoring apparatus and procedural details (e.g., using wrong glassware, skipping standardisation) → leads to procedural errors in practicals or viva.
-Treating this chapter as purely “easy” because weightage is low → leads to inadequate revision.
Tips to avoid: Focus on linking tests/preparations to principle, practise MCQs regularly.
How should one prepare this chapter for the practical exam (viva + experiment) for boards?
Here’s a suggested plan:
List all experiments/tests in the syllabus (element detection, functional groups, salt analysis, titrations, etc.).
For each, write: objective → principle → reagents/apparatus → observations → conclusions.
Practice procedural writing: steps should be logically ordered, with safety notes.
Mock viva: Ask yourself “Why did I add this reagent?”, “What is the function of this apparatus?”, “Why did I wash the precipitate?”, etc. (Such questions are trending for viva preparation).
Revise MCQs from this chapter (to cover board theory + practical-based questions).
Time-management: practice answering quickly since in practical exams time is limited.
Which experiments/tests from this chapter are most likely to appear in board/viva/practical?
Based on trend:
-Detection of N, S, halogens in organic compounds.
-Functional group tests (e.g., –OH, –COOH, >C=O, –NH₂).
-Preparation of one inorganic and one organic compound.
-Volumetric titration of an acid–base or redox type.
-Qualitative salt analysis (one cation + one anion).
These have high visibility in both board practicals and competitive exam lists.
How can students make quick revision notes for this chapter?
Some suggestions:
-Create tables of “Reagent → Test → Observation → Conclusion” for element/functional group detection.
-Create flow-charts for salt analysis (preliminary tests → confirmatory tests).
-List key principles (e.g., titration: equivalence point vs end point; volumetric analysis principle).
-Make a cheat sheet of important preparations: compound, reagents, diagram, yield, key reactions.
-Use flashcards for MCQ-type facts (e.g., “Which apparatus gives most accurate volume?”).
What is qualitative analysis in chemistry?
It is the process of identifying ions (cations and anions) in a salt sample using chemical reactions and observations.
What is quantitative analysis in chemistry?
It is the determination of the exact amount of a substance present through titration and chemical measurements.
Why do we do titration in chemistry?
Titration helps determine the concentration of an unknown solution using a known standard solution.
What are indicators in titration?
Indicators are chemicals that change colour at a specific pH to show the end point of a titration (e.g., phenolphthalein, methyl orange).
What are common errors in acid-base titration?
Parallax error, air bubbles in burette, improper washing of apparatus, and incorrect reading of meniscus.
What is meant by end point in titration?
It is the point at which the indicator changes colour, signalling the completion of the reaction.
What is meant by equivalence point?
The exact point where stoichiometric amounts of acid and base react — close to but not always equal to end-point.
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