Whether a Salt Dissolves Decides How You Have to Make It
Define a salt through neutralisation, tell normal, acid, basic and complex salts apart, learn the six laboratory methods for preparing normal salts, and use solubility to pick the right method and write its balanced equation for any salt you are asked to make.
Why can't every salt be made by the same method?
Suppose you need a sample of copper sulphate and a sample of barium sulphate. Both are sulphates, and both could in principle come from an acid meeting a base. Yet the routes you would actually take are completely different.
- Copper sulphate dissolves in water. You can make it in solution, filter off anything unwanted, and crystallise the salt out by evaporation
- Barium sulphate does not dissolve. If you tried the same route, the salt would stay mixed with the solid you meant to filter away — so instead you make it appear as a precipitate and filter the salt itself
So the first question in preparing any salt is simply: does it dissolve? The answer narrows six possible methods down to one or two, and a second question — is the metal reactive, and is its base soluble? — usually settles the rest.
Before choosing methods, this part pins down what a salt is. A salt is what remains when the hydrogen of an acid has been replaced by a metal or by the ammonium ion, and how much of that hydrogen is replaced — all of it or only part — gives different kinds of salt.
Salts are everywhere in daily life.
- Common salt, sodium chloride, on every table
- Baking soda, sodium hydrogen carbonate — an acid salt
- Blue vitriol, the blue copper sulphate crystals sold as a fungicide for crops
- Alum, used in many Indian homes to clear muddy water
- Plaster of Paris and washing soda
So this part moves from definition to classification to method to decision, and the final section is the one most examination questions are built on: given a salt, choose the method and write the equation.
This page covers the third part of the ICSE Class 10 Chemistry chapter on acids, bases and salts: neutralisation and the definition of a salt, the four classes of salt, the methods of preparing normal salts, and selecting a method by solubility.
- Copper sulphate dissolves in water. You can make it in solution, filter off anything unwanted, and crystallise the salt out by evaporation
- Barium sulphate does not dissolve. If you tried the same route, the salt would stay mixed with the solid you meant to filter away — so instead you make it appear as a precipitate and filter the salt itself
So the first question in preparing any salt is simply: does it dissolve? The answer narrows six possible methods down to one or two, and a second question — is the metal reactive, and is its base soluble? — usually settles the rest.
Before choosing methods, this part pins down what a salt is. A salt is what remains when the hydrogen of an acid has been replaced by a metal or by the ammonium ion, and how much of that hydrogen is replaced — all of it or only part — gives different kinds of salt.
Salts are everywhere in daily life.
- Common salt, sodium chloride, on every table
- Baking soda, sodium hydrogen carbonate — an acid salt
- Blue vitriol, the blue copper sulphate crystals sold as a fungicide for crops
- Alum, used in many Indian homes to clear muddy water
- Plaster of Paris and washing soda
So this part moves from definition to classification to method to decision, and the final section is the one most examination questions are built on: given a salt, choose the method and write the equation.
This page covers the third part of the ICSE Class 10 Chemistry chapter on acids, bases and salts: neutralisation and the definition of a salt, the four classes of salt, the methods of preparing normal salts, and selecting a method by solubility.
What is a salt, and how does neutralisation form NaHSO4 and Na2SO4?
A salt is formed when the replaceable hydrogen of an acid is partly or completely replaced by a metal ion or ammonium ion, and neutralisation is the reaction of an acid with a base that produces a salt and water only.
Definition of a salt. A salt is a compound formed by the partial or complete replacement of the replaceable hydrogen ions of an acid by a metallic ion or an ammonium ion.
Definition of neutralisation. The reaction between an acid and a base to form a salt and water only.
What actually happens in every neutralisation, whatever the acid and base:
The other ions are spectators that end up together as the salt. Neutralisation releases heat, which you can feel as a warm test tube.
Examples with different acids and bases:
Partial and complete replacement — sulphuric acid, basicity two.
Partial replacement: one of the two hydrogens is replaced, giving sodium hydrogen sulphate:
Complete replacement: both are replaced, giving sodium sulphate:
And the first salt can be taken on to the second by adding more alkali:
Worked check — the two routes to sodium sulphate agree. Adding the first and third equations:
**after cancelling from both sides — which is exactly the second equation. The partial route and the direct route use the same total amount of alkali.
Everyday neutralisations:
- An ant sting injects formic acid; rubbing on baking soda paste neutralises it and eases the pain
- An antacid neutralises excess hydrochloric acid in the stomach
- Slaked lime spread on acidic soil neutralises the acid so crops grow better
The boundary case to state precisely. Neutralisation does not always give a neutral solution. Ammonium chloride, from a weak base and a strong acid, gives a slightly acidic solution; sodium carbonate, from a strong base and a weak acid, gives an alkaline one.** The word describes the reaction between acid and base, not the pH of the product.
Definition of a salt. A salt is a compound formed by the partial or complete replacement of the replaceable hydrogen ions of an acid by a metallic ion or an ammonium ion.
Definition of neutralisation. The reaction between an acid and a base to form a salt and water only.
What actually happens in every neutralisation, whatever the acid and base:
The other ions are spectators that end up together as the salt. Neutralisation releases heat, which you can feel as a warm test tube.
Examples with different acids and bases:
Partial and complete replacement — sulphuric acid, basicity two.
Partial replacement: one of the two hydrogens is replaced, giving sodium hydrogen sulphate:
Complete replacement: both are replaced, giving sodium sulphate:
And the first salt can be taken on to the second by adding more alkali:
Worked check — the two routes to sodium sulphate agree. Adding the first and third equations:
**after cancelling from both sides — which is exactly the second equation. The partial route and the direct route use the same total amount of alkali.
Everyday neutralisations:
- An ant sting injects formic acid; rubbing on baking soda paste neutralises it and eases the pain
- An antacid neutralises excess hydrochloric acid in the stomach
- Slaked lime spread on acidic soil neutralises the acid so crops grow better
The boundary case to state precisely. Neutralisation does not always give a neutral solution. Ammonium chloride, from a weak base and a strong acid, gives a slightly acidic solution; sodium carbonate, from a strong base and a weak acid, gives an alkaline one.** The word describes the reaction between acid and base, not the pH of the product.
How are salts classified as normal, acid, basic and complex salts?
A normal salt has all its replaceable hydrogen replaced, an acid salt still has some, a basic salt still has hydroxyl groups from the base, and a complex salt contains a complex ion that stays intact in solution.
1. Normal salts. Formed by the complete replacement of the replaceable hydrogen of an acid. They contain neither replaceable hydrogen nor hydroxyl groups.
- Examples: , , , ,
2. Acid salts. Formed by the partial replacement of the replaceable hydrogen of a polybasic acid. They still contain replaceable hydrogen, and react with a base to form a normal salt.
- Examples: , , ,
3. Basic salts. Formed by the partial replacement of the hydroxyl groups of a polyacidic base by an acid radical. They still contain hydroxyl groups, and react with an acid to form a normal salt.
- Examples: basic copper nitrate , basic lead chloride , and basic copper carbonate
4. Complex salts. Contain a complex ion — a central metal ion bonded to molecules or ions around it — that does not break up into its parts in solution.
- Tetraamminecopper(II) sulphate gives the deep blue complex ion, not free copper ions and ammonia
- Potassium ferrocyanide gives the ferrocyanide ion, not free iron ions and cyanide ions
Worked check — charges in the complex salts. Tetraamminecopper(II): from copper, from each ammonia, so the ion is and one sulphate ion of balances it. Potassium ferrocyanide: .
A boundary case — complex salt versus double salt. Potash alum, , is a double salt: two simple salts crystallised together. In water it gives all its simple ions — potassium, aluminium and sulphate — whereas a complex salt keeps its complex ion intact. The test is what appears in solution.
Worked example — classify each salt.
- ** — contains replaceable hydrogen: acid salt
- — all hydrogen replaced: normal salt
- — contains a hydroxyl group: basic salt
- — contains a complex ion: complex salt
An everyday example. The green coating that forms on old copper vessels and copper roofs is largely basic copper carbonate, formed slowly by moisture and carbon dioxide from the air. The blue deep-coloured solution you will make when ammonia is added to copper sulphate in the analytical chemistry chapter is the complex salt above.
One trap in classification. Only a polybasic acid can form an acid salt. Hydrochloric acid is monobasic, so there is no acid salt of hydrochloric acid** — once its single hydrogen is replaced, nothing replaceable remains.
1. Normal salts. Formed by the complete replacement of the replaceable hydrogen of an acid. They contain neither replaceable hydrogen nor hydroxyl groups.
- Examples: , , , ,
2. Acid salts. Formed by the partial replacement of the replaceable hydrogen of a polybasic acid. They still contain replaceable hydrogen, and react with a base to form a normal salt.
- Examples: , , ,
3. Basic salts. Formed by the partial replacement of the hydroxyl groups of a polyacidic base by an acid radical. They still contain hydroxyl groups, and react with an acid to form a normal salt.
- Examples: basic copper nitrate , basic lead chloride , and basic copper carbonate
4. Complex salts. Contain a complex ion — a central metal ion bonded to molecules or ions around it — that does not break up into its parts in solution.
- Tetraamminecopper(II) sulphate gives the deep blue complex ion, not free copper ions and ammonia
- Potassium ferrocyanide gives the ferrocyanide ion, not free iron ions and cyanide ions
Worked check — charges in the complex salts. Tetraamminecopper(II): from copper, from each ammonia, so the ion is and one sulphate ion of balances it. Potassium ferrocyanide: .
A boundary case — complex salt versus double salt. Potash alum, , is a double salt: two simple salts crystallised together. In water it gives all its simple ions — potassium, aluminium and sulphate — whereas a complex salt keeps its complex ion intact. The test is what appears in solution.
Worked example — classify each salt.
- ** — contains replaceable hydrogen: acid salt
- — all hydrogen replaced: normal salt
- — contains a hydroxyl group: basic salt
- — contains a complex ion: complex salt
An everyday example. The green coating that forms on old copper vessels and copper roofs is largely basic copper carbonate, formed slowly by moisture and carbon dioxide from the air. The blue deep-coloured solution you will make when ammonia is added to copper sulphate in the analytical chemistry chapter is the complex salt above.
One trap in classification. Only a polybasic acid can form an acid salt. Hydrochloric acid is monobasic, so there is no acid salt of hydrochloric acid** — once its single hydrogen is replaced, nothing replaceable remains.
What are the methods of preparing a normal salt?
Six methods cover almost every normal salt: direct combination, displacement, precipitation, neutralisation of an insoluble base, titration of an alkali, and the action of dilute acids on carbonates and bicarbonates.
1. Direct combination, or synthesis. The metal combines directly with the non-metal.
Used for salts such as sulphides and for anhydrous chlorides like iron(III) chloride and aluminium chloride, which cannot be crystallised from water because they react with it.
2. Displacement — an active metal with a dilute acid.
Used for soluble salts of moderately reactive metals — magnesium, aluminium, zinc and iron. Not used for sodium, potassium or calcium, which react dangerously; nor for copper and silver, which lie below hydrogen and give no reaction; nor for lead, which gets coated with an insoluble salt.
3. Precipitation, or double decomposition. Two soluble salt solutions are mixed, and the ions exchange partners to form an insoluble salt.
Procedure: filter off the precipitate, wash it with distilled water to remove the soluble salt left with it, and dry it. Used for all insoluble salts.
4. Neutralisation of an insoluble base. A metal oxide, hydroxide or carbonate that does not dissolve is added to a warm dilute acid.
Procedure: add the base in excess so that all the acid is used up; filter off the unreacted base; evaporate the filtrate to the crystallisation point and cool to obtain crystals — blue in the first case.
5. Titration — a soluble base with an acid.
Procedure: run acid from a burette into a measured volume of alkali containing an indicator until the colour just changes; note the volume; repeat with the same volumes but no indicator, so the salt is not contaminated; evaporate and crystallise.
Why titration is needed here and not the excess method. An alkali dissolves, so any excess cannot be filtered off — the exact amount must be used. Used for sodium, potassium and ammonium salts.
6. Action of dilute acids on carbonates and bicarbonates.
An insoluble carbonate is used in excess and filtered, like an insoluble base; a soluble carbonate is titrated, like an alkali. The fizzing stops when the acid is used up, a convenient built-in signal.
Worked check — balancing the ammonium sulphate equation. Nitrogen ; hydrogen on the left and on the right; oxygen on the left and on the right; sulphur . Balanced.
One boundary case in precipitation. Lead chloride is only sparingly soluble in cold water but dissolves noticeably in hot water, so it is precipitated from cold solutions and washed with cold water, or much of the product is lost.
1. Direct combination, or synthesis. The metal combines directly with the non-metal.
Used for salts such as sulphides and for anhydrous chlorides like iron(III) chloride and aluminium chloride, which cannot be crystallised from water because they react with it.
2. Displacement — an active metal with a dilute acid.
Used for soluble salts of moderately reactive metals — magnesium, aluminium, zinc and iron. Not used for sodium, potassium or calcium, which react dangerously; nor for copper and silver, which lie below hydrogen and give no reaction; nor for lead, which gets coated with an insoluble salt.
3. Precipitation, or double decomposition. Two soluble salt solutions are mixed, and the ions exchange partners to form an insoluble salt.
Procedure: filter off the precipitate, wash it with distilled water to remove the soluble salt left with it, and dry it. Used for all insoluble salts.
4. Neutralisation of an insoluble base. A metal oxide, hydroxide or carbonate that does not dissolve is added to a warm dilute acid.
Procedure: add the base in excess so that all the acid is used up; filter off the unreacted base; evaporate the filtrate to the crystallisation point and cool to obtain crystals — blue in the first case.
5. Titration — a soluble base with an acid.
Procedure: run acid from a burette into a measured volume of alkali containing an indicator until the colour just changes; note the volume; repeat with the same volumes but no indicator, so the salt is not contaminated; evaporate and crystallise.
Why titration is needed here and not the excess method. An alkali dissolves, so any excess cannot be filtered off — the exact amount must be used. Used for sodium, potassium and ammonium salts.
6. Action of dilute acids on carbonates and bicarbonates.
An insoluble carbonate is used in excess and filtered, like an insoluble base; a soluble carbonate is titrated, like an alkali. The fizzing stops when the acid is used up, a convenient built-in signal.
Worked check — balancing the ammonium sulphate equation. Nitrogen ; hydrogen on the left and on the right; oxygen on the left and on the right; sulphur . Balanced.
One boundary case in precipitation. Lead chloride is only sparingly soluble in cold water but dissolves noticeably in hot water, so it is precipitated from cold solutions and washed with cold water, or much of the product is lost.
How do you choose the right method for preparing a given salt?
Decide first whether the salt is soluble; if it is insoluble, precipitate it; if it is soluble, choose between titration, an insoluble base or carbonate, and a metal with acid according to the metal.
The solubility rules you need:
- All sodium, potassium and ammonium salts are soluble
- All nitrates are soluble
- Chlorides are soluble except silver chloride and lead chloride, which is soluble in hot water
- Sulphates are soluble except barium sulphate and lead sulphate, with calcium sulphate only sparingly soluble
- Carbonates are insoluble except those of sodium, potassium and ammonium
- Oxides and hydroxides are insoluble except those of sodium, potassium and ammonium, with calcium hydroxide slightly soluble
The decision:
- Insoluble salt precipitation
- Soluble salt of sodium, potassium or ammonium titration of the alkali or soluble carbonate with the acid
- Soluble salt of a metal below hydrogen, such as copper insoluble base or carbonate with acid, in excess, then filter
- Soluble salt of a moderately reactive metal, such as zinc, magnesium or iron metal with dilute acid, or its insoluble base or carbonate
- Anhydrous chloride that reacts with water, such as iron(III) chloride direct combination with dry chlorine
Worked example 1 — lead chloride. Chlorides are soluble except silver and lead: insoluble, so precipitation.
Worked example 2 — sodium sulphate. A sodium salt: soluble, so titration.
Worked example 3 — copper(II) sulphate. Soluble, but copper does not react with dilute sulphuric acid, so use its insoluble oxide in excess:
Worked example 4 — zinc sulphate. Soluble, and zinc is above hydrogen: metal with dilute acid.
Worked example 5 — barium sulphate. Insoluble, so precipitation:
Worked example 6 — anhydrous iron(III) chloride. Reacts with water, so direct combination with dry chlorine over heated iron:
Worked example 7 — calcium carbonate. Carbonates of calcium are insoluble: precipitation.
Worked example 8 — ammonium sulphate. An ammonium salt: titration.
Three traps that the decision rules prevent.
- Iron with dilute hydrochloric acid gives iron(II) chloride, , not iron(III) chloride — which is why iron(III) chloride needs chlorine gas
- Lead sulphate cannot be made from lead and dilute sulphuric acid — the insoluble product coats the metal and the reaction stops, so precipitation is used
- Copper sulphate cannot be made from copper and dilute sulphuric acid — copper is below hydrogen
An everyday parallel. A cook separates boiled rice from its water by draining it, but recovers salt from brine by boiling the water away. Draining keeps the undissolved solid; evaporating keeps the dissolved one — exactly the difference between filtering a precipitate and crystallising a soluble salt.
The solubility rules you need:
- All sodium, potassium and ammonium salts are soluble
- All nitrates are soluble
- Chlorides are soluble except silver chloride and lead chloride, which is soluble in hot water
- Sulphates are soluble except barium sulphate and lead sulphate, with calcium sulphate only sparingly soluble
- Carbonates are insoluble except those of sodium, potassium and ammonium
- Oxides and hydroxides are insoluble except those of sodium, potassium and ammonium, with calcium hydroxide slightly soluble
The decision:
- Insoluble salt precipitation
- Soluble salt of sodium, potassium or ammonium titration of the alkali or soluble carbonate with the acid
- Soluble salt of a metal below hydrogen, such as copper insoluble base or carbonate with acid, in excess, then filter
- Soluble salt of a moderately reactive metal, such as zinc, magnesium or iron metal with dilute acid, or its insoluble base or carbonate
- Anhydrous chloride that reacts with water, such as iron(III) chloride direct combination with dry chlorine
Worked example 1 — lead chloride. Chlorides are soluble except silver and lead: insoluble, so precipitation.
Worked example 2 — sodium sulphate. A sodium salt: soluble, so titration.
Worked example 3 — copper(II) sulphate. Soluble, but copper does not react with dilute sulphuric acid, so use its insoluble oxide in excess:
Worked example 4 — zinc sulphate. Soluble, and zinc is above hydrogen: metal with dilute acid.
Worked example 5 — barium sulphate. Insoluble, so precipitation:
Worked example 6 — anhydrous iron(III) chloride. Reacts with water, so direct combination with dry chlorine over heated iron:
Worked example 7 — calcium carbonate. Carbonates of calcium are insoluble: precipitation.
Worked example 8 — ammonium sulphate. An ammonium salt: titration.
Three traps that the decision rules prevent.
- Iron with dilute hydrochloric acid gives iron(II) chloride, , not iron(III) chloride — which is why iron(III) chloride needs chlorine gas
- Lead sulphate cannot be made from lead and dilute sulphuric acid — the insoluble product coats the metal and the reaction stops, so precipitation is used
- Copper sulphate cannot be made from copper and dilute sulphuric acid — copper is below hydrogen
An everyday parallel. A cook separates boiled rice from its water by draining it, but recovers salt from brine by boiling the water away. Draining keeps the undissolved solid; evaporating keeps the dissolved one — exactly the difference between filtering a precipitate and crystallising a soluble salt.
Exam tip
What must a salt preparation answer include?
Name the method, justify it from solubility, write the balanced equation, and give the steps that separate the pure salt.
- State whether the salt is soluble before naming the method — the justification carries marks
- Name the method precisely: precipitation, titration, insoluble base in excess, displacement or direct combination
- Write a fully balanced equation with the correct formula of the salt
- For insoluble salts, say filter, wash with distilled water and dry
- For the excess-base method, say excess, filter off the unreacted solid, evaporate to the crystallisation point and cool
- For titration, mention the indicator run followed by a repeat without indicator
- Give the correct oxidation state: iron with dilute acid gives iron(II) salts
- Explain why a metal cannot be used when that is the point: copper below hydrogen, lead coated, sodium too violent
- Classify salts with a reason — replaceable hydrogen, hydroxyl group or complex ion
- Write complex ions in square brackets with their charge
The misconception to name. An acid salt does not always give an acidic solution. Sodium hydrogen carbonate is an acid salt, yet its solution is mildly alkaline — which is why baking soda can neutralise an acid. The name refers to the replaceable hydrogen in its formula, not to the pH of its solution.
A second trap. Choosing precipitation for a soluble salt. If both products of a double decomposition dissolve, nothing precipitates and nothing can be separated by filtering, so sodium nitrate, for example, can never be made that way.
- State whether the salt is soluble before naming the method — the justification carries marks
- Name the method precisely: precipitation, titration, insoluble base in excess, displacement or direct combination
- Write a fully balanced equation with the correct formula of the salt
- For insoluble salts, say filter, wash with distilled water and dry
- For the excess-base method, say excess, filter off the unreacted solid, evaporate to the crystallisation point and cool
- For titration, mention the indicator run followed by a repeat without indicator
- Give the correct oxidation state: iron with dilute acid gives iron(II) salts
- Explain why a metal cannot be used when that is the point: copper below hydrogen, lead coated, sodium too violent
- Classify salts with a reason — replaceable hydrogen, hydroxyl group or complex ion
- Write complex ions in square brackets with their charge
The misconception to name. An acid salt does not always give an acidic solution. Sodium hydrogen carbonate is an acid salt, yet its solution is mildly alkaline — which is why baking soda can neutralise an acid. The name refers to the replaceable hydrogen in its formula, not to the pH of its solution.
A second trap. Choosing precipitation for a soluble salt. If both products of a double decomposition dissolve, nothing precipitates and nothing can be separated by filtering, so sodium nitrate, for example, can never be made that way.
Did you know
Why do blue copper sulphate crystals turn white when heated?
Heat a few bright blue crystals of copper sulphate in a dry test tube and something surprising happens. Droplets of water appear on the cooler upper part of the tube, and the blue crystals crumble into a white powder. Let the tube cool and add a few drops of water to the powder, and it turns blue again — and the tube becomes noticeably warm.
The blue crystals are not pure copper sulphate. Each unit of the crystal holds five molecules of water of crystallisation, locked into the crystal structure:
That water is part of what makes the crystals blue. Drive it off and the anhydrous salt left behind is white; give it back and the colour returns.
How much of the crystal is water? The formula mass of the hydrated salt is , of which is water:
**So heating of blue crystals leaves only about**
of white powder — more than a third of the weight of what looked like a dry solid was water all along.
The reverse change has a practical use. Because white anhydrous copper sulphate turns blue with even a trace of water, it is used in the laboratory as a test for the presence of water in a liquid. The warmth you feel is the energy released as water molecules settle back into the crystal.
This is also why the last step of the excess-base method matters. Copper sulphate made from copper oxide and sulphuric acid is recovered by crystallisation, not by boiling the solution to dryness. Boiling to dryness would drive off the water of crystallisation and leave the white powder instead of the blue crystals the method is meant to produce.
The same idea explains the white crust that forms on washing soda crystals left open to air. Washing soda slowly loses some of its water of crystallisation to dry air and crumbles to a powder — a salt changing its appearance without any other substance being added at all.
The blue crystals are not pure copper sulphate. Each unit of the crystal holds five molecules of water of crystallisation, locked into the crystal structure:
That water is part of what makes the crystals blue. Drive it off and the anhydrous salt left behind is white; give it back and the colour returns.
How much of the crystal is water? The formula mass of the hydrated salt is , of which is water:
**So heating of blue crystals leaves only about**
of white powder — more than a third of the weight of what looked like a dry solid was water all along.
The reverse change has a practical use. Because white anhydrous copper sulphate turns blue with even a trace of water, it is used in the laboratory as a test for the presence of water in a liquid. The warmth you feel is the energy released as water molecules settle back into the crystal.
This is also why the last step of the excess-base method matters. Copper sulphate made from copper oxide and sulphuric acid is recovered by crystallisation, not by boiling the solution to dryness. Boiling to dryness would drive off the water of crystallisation and leave the white powder instead of the blue crystals the method is meant to produce.
The same idea explains the white crust that forms on washing soda crystals left open to air. Washing soda slowly loses some of its water of crystallisation to dry air and crumbles to a powder — a salt changing its appearance without any other substance being added at all.
Exam relevance
How does salt chemistry lead into JEE and NEET?
This is foundation work for Class 11 Equilibrium and Class 12 Coordination Compounds, both examined in JEE Main and NEET Chemistry, and for the inorganic preparations in JEE Main practical chemistry.
Where neutralisation leads. Class 11 explains why the solutions of salts are not all neutral through salt hydrolysis. The boundary cases on this page — ammonium chloride acidic, sodium carbonate alkaline, sodium hydrogen carbonate mildly alkaline — are exactly the categories of strong-acid weak-base and weak-acid strong-base salts, and predicting whether a salt solution is acidic, basic or neutral is a recurring question in both exams.
Where solubility and precipitation lead. Class 11 introduces the solubility product, which gives a precise condition for a precipitate to form. The solubility rules you learn here become the qualitative background for numericals on whether barium sulphate or silver chloride will precipitate when two solutions are mixed, and for the common ion effect.
Where complex salts lead. Class 12 Coordination Compounds is built on the distinction on this page between complex salts and double salts. Naming tetraamminecopper(II) sulphate and potassium hexacyanoferrate(II), counting the charge on the complex ion and working out the oxidation state of the metal are standard questions in both JEE Main and NEET.
Where the preparation methods lead. JEE Main practical chemistry includes the chemistry involved in preparing ferrous ammonium sulphate and potash alum by crystallisation — the same crystallisation step as the excess-base method here, and the double-salt idea from the classification section.
Where water of crystallisation leads. Formulae of hydrated salts and calculations of the percentage of water, as in the copper sulphate example, reappear in Class 11 Some Basic Concepts of Chemistry as mole-concept numericals.
Question types to expect. At this level: classification with reasons, method selection with balanced equations, and procedures. In competitive papers: hydrolysis and pH of salt solutions, solubility product numericals, naming and charges of complex ions, and double salt versus complex salt identification, often as match-the-column items.
The single trap that costs marks. Assuming an acid salt gives an acidic solution. Sodium hydrogen carbonate gives a mildly alkaline one, and both exams use it in assertion-reason questions precisely because the name suggests otherwise.
A second trap. Treating potash alum as a complex salt. It releases all its simple ions in water, which is the defining test of a double salt, and a question that asks how many ions one formula unit gives in solution separates the two immediately.
Board versus competitive emphasis. The ICSE paper marks the choice of method, its justification and the balanced equation; a competitive paper marks a pH prediction, a precipitation condition or a name. The transferable habit is asking what a salt does in water — dissolve, stay intact or react — because that one question drives preparation here and equilibrium and coordination chemistry later.
Where neutralisation leads. Class 11 explains why the solutions of salts are not all neutral through salt hydrolysis. The boundary cases on this page — ammonium chloride acidic, sodium carbonate alkaline, sodium hydrogen carbonate mildly alkaline — are exactly the categories of strong-acid weak-base and weak-acid strong-base salts, and predicting whether a salt solution is acidic, basic or neutral is a recurring question in both exams.
Where solubility and precipitation lead. Class 11 introduces the solubility product, which gives a precise condition for a precipitate to form. The solubility rules you learn here become the qualitative background for numericals on whether barium sulphate or silver chloride will precipitate when two solutions are mixed, and for the common ion effect.
Where complex salts lead. Class 12 Coordination Compounds is built on the distinction on this page between complex salts and double salts. Naming tetraamminecopper(II) sulphate and potassium hexacyanoferrate(II), counting the charge on the complex ion and working out the oxidation state of the metal are standard questions in both JEE Main and NEET.
Where the preparation methods lead. JEE Main practical chemistry includes the chemistry involved in preparing ferrous ammonium sulphate and potash alum by crystallisation — the same crystallisation step as the excess-base method here, and the double-salt idea from the classification section.
Where water of crystallisation leads. Formulae of hydrated salts and calculations of the percentage of water, as in the copper sulphate example, reappear in Class 11 Some Basic Concepts of Chemistry as mole-concept numericals.
Question types to expect. At this level: classification with reasons, method selection with balanced equations, and procedures. In competitive papers: hydrolysis and pH of salt solutions, solubility product numericals, naming and charges of complex ions, and double salt versus complex salt identification, often as match-the-column items.
The single trap that costs marks. Assuming an acid salt gives an acidic solution. Sodium hydrogen carbonate gives a mildly alkaline one, and both exams use it in assertion-reason questions precisely because the name suggests otherwise.
A second trap. Treating potash alum as a complex salt. It releases all its simple ions in water, which is the defining test of a double salt, and a question that asks how many ions one formula unit gives in solution separates the two immediately.
Board versus competitive emphasis. The ICSE paper marks the choice of method, its justification and the balanced equation; a competitive paper marks a pH prediction, a precipitation condition or a name. The transferable habit is asking what a salt does in water — dissolve, stay intact or react — because that one question drives preparation here and equilibrium and coordination chemistry later.
Key takeaways
What must you be able to do from this part?
One definition, four classes, six methods and a decision rule.
- A salt forms by partial or complete replacement of an acid's replaceable hydrogen by a metal ion or ammonium ion
- Neutralisation is acid + base giving salt and water only; the essential reaction is , and it releases heat
- Partial replacement:
- Complete replacement:
- Neutralisation need not give a neutral solution — ammonium chloride is acidic, sodium carbonate alkaline
- Normal salts have no replaceable hydrogen: ,
- Acid salts keep replaceable hydrogen and come only from polybasic acids: ,
- Basic salts keep hydroxyl groups: ,
- Complex salts keep a complex ion intact in solution: , ; double salts such as potash alum give all their simple ions
- Direct combination: , for anhydrous chlorides
- Displacement: , for Mg, Al, Zn, Fe
- Precipitation: , for insoluble salts — filter, wash, dry
- Insoluble base in excess: — filter, evaporate, crystallise
- Titration: for sodium, potassium and ammonium salts, because excess alkali cannot be filtered
- Carbonates and bicarbonates with acid: salt + water + , excess or titration depending on solubility
- Solubility rules: Na, K, NH4 salts and nitrates soluble; AgCl and PbCl2 insoluble; BaSO4 and PbSO4 insoluble; most carbonates, oxides and hydroxides insoluble
- Iron with dilute acid gives iron(II) salts; copper does not react; lead gets coated
- Crystallise, do not boil to dryness, or water of crystallisation is lost: is about water
The sharpest self-test is a list of eight salts. Write down silver chloride, potassium nitrate, magnesium sulphate, copper nitrate, lead sulphate, sodium chloride, zinc chloride and calcium carbonate, choose a method for each with a one-line reason, and write every equation — then check which of your choices the solubility rules would overturn.
- A salt forms by partial or complete replacement of an acid's replaceable hydrogen by a metal ion or ammonium ion
- Neutralisation is acid + base giving salt and water only; the essential reaction is , and it releases heat
- Partial replacement:
- Complete replacement:
- Neutralisation need not give a neutral solution — ammonium chloride is acidic, sodium carbonate alkaline
- Normal salts have no replaceable hydrogen: ,
- Acid salts keep replaceable hydrogen and come only from polybasic acids: ,
- Basic salts keep hydroxyl groups: ,
- Complex salts keep a complex ion intact in solution: , ; double salts such as potash alum give all their simple ions
- Direct combination: , for anhydrous chlorides
- Displacement: , for Mg, Al, Zn, Fe
- Precipitation: , for insoluble salts — filter, wash, dry
- Insoluble base in excess: — filter, evaporate, crystallise
- Titration: for sodium, potassium and ammonium salts, because excess alkali cannot be filtered
- Carbonates and bicarbonates with acid: salt + water + , excess or titration depending on solubility
- Solubility rules: Na, K, NH4 salts and nitrates soluble; AgCl and PbCl2 insoluble; BaSO4 and PbSO4 insoluble; most carbonates, oxides and hydroxides insoluble
- Iron with dilute acid gives iron(II) salts; copper does not react; lead gets coated
- Crystallise, do not boil to dryness, or water of crystallisation is lost: is about water
The sharpest self-test is a list of eight salts. Write down silver chloride, potassium nitrate, magnesium sulphate, copper nitrate, lead sulphate, sodium chloride, zinc chloride and calcium carbonate, choose a method for each with a one-line reason, and write every equation — then check which of your choices the solubility rules would overturn.