Mix Two Clear Solutions and a Bright Yellow Solid Appears
Learn to classify a reaction as combination, decomposition, displacement or double decomposition, write the three standard precipitation equations and explain the ion exchange, tell exothermic from endothermic, and name the conditions a reaction needs.
Why does a yellow solid appear when two clear liquids meet?
Take a colourless solution of lead nitrate and a colourless solution of potassium iodide. Neither shows a trace of yellow. Pour one into the other and a bright yellow solid forms at once and settles.
Nothing was added. The yellow came out of the two clear liquids.
What happened is an exchange of partners. Lead nitrate in water exists as and ions; potassium iodide exists as and . Mix them and four kinds of ion are swimming about together — and one pairing, with , makes a compound that will not dissolve. It drops out of the solution as the yellow precipitate lead iodide.
The potassium and nitrate simply stay dissolved, as they were.
That is one of four patterns every reaction in this chapter falls into. Recognising which pattern you are looking at is what lets you predict the products, write the equation and explain the colour — and that is what this page is about.
It covers the ICSE Class 9 Chemistry chapter on chemical changes and reactions — the four types, the ion exchange in double decomposition, heat changes, and the conditions a reaction needs.
Nothing was added. The yellow came out of the two clear liquids.
What happened is an exchange of partners. Lead nitrate in water exists as and ions; potassium iodide exists as and . Mix them and four kinds of ion are swimming about together — and one pairing, with , makes a compound that will not dissolve. It drops out of the solution as the yellow precipitate lead iodide.
The potassium and nitrate simply stay dissolved, as they were.
That is one of four patterns every reaction in this chapter falls into. Recognising which pattern you are looking at is what lets you predict the products, write the equation and explain the colour — and that is what this page is about.
It covers the ICSE Class 9 Chemistry chapter on chemical changes and reactions — the four types, the ion exchange in double decomposition, heat changes, and the conditions a reaction needs.
What are the four types of chemical reaction?
Combination joins, decomposition splits, displacement replaces one element, and double decomposition swaps partners.
Direct combination, or synthesis — two or more substances join to form one product.
The test is one product. If two or more things go in and exactly one comes out, it is a combination.
Decomposition — one compound splits into two or more products. It is named after whatever supplies the energy.
- Thermal decomposition, by heat:
- Electrolytic decomposition, by electricity:
- Photochemical decomposition, by light:
Displacement, or single displacement — a more reactive element pushes out a less reactive one from its compound.
Worked observation. Drop a strip of zinc into blue copper sulphate solution. The blue fades as leaves the solution, and a red-brown deposit of copper appears on the zinc. Zinc is more reactive than copper, so it takes the sulphate and the copper is turned out.
Double decomposition, or double displacement — two compounds exchange their radicals, giving two new compounds.
Displacement and double decomposition are easy to confuse and easy to separate. In a displacement one of the reactants is an element — zinc, iron, chlorine — and it ends up as part of a compound. In a double decomposition both reactants are compounds and both products are compounds, with no free element anywhere.
So the quickest classification is to count the free elements. Two reactants and one product means combination; one reactant means decomposition; a free element on the left means displacement; and no free element at all with two compounds each side means double decomposition.
Direct combination, or synthesis — two or more substances join to form one product.
The test is one product. If two or more things go in and exactly one comes out, it is a combination.
Decomposition — one compound splits into two or more products. It is named after whatever supplies the energy.
- Thermal decomposition, by heat:
- Electrolytic decomposition, by electricity:
- Photochemical decomposition, by light:
Displacement, or single displacement — a more reactive element pushes out a less reactive one from its compound.
Worked observation. Drop a strip of zinc into blue copper sulphate solution. The blue fades as leaves the solution, and a red-brown deposit of copper appears on the zinc. Zinc is more reactive than copper, so it takes the sulphate and the copper is turned out.
Double decomposition, or double displacement — two compounds exchange their radicals, giving two new compounds.
Displacement and double decomposition are easy to confuse and easy to separate. In a displacement one of the reactants is an element — zinc, iron, chlorine — and it ends up as part of a compound. In a double decomposition both reactants are compounds and both products are compounds, with no free element anywhere.
So the quickest classification is to count the free elements. Two reactants and one product means combination; one reactant means decomposition; a free element on the left means displacement; and no free element at all with two compounds each side means double decomposition.
Formula
How do ions exchange in a double decomposition?
The two positive radicals swap their negative partners, and the reaction proceeds because one of the new pairings is insoluble.
Worked example 1 — sodium sulphate and barium chloride.
Check the balance: Na ; S ; O ; Ba ; Cl .
The ion exchange. In solution the four ions present are , , and . Of the two new pairings:
- with gives barium sulphate, which is insoluble — it appears as a white precipitate
- with gives sodium chloride, which is soluble and stays in solution
Worked example 2 — silver nitrate and sodium chloride.
Already balanced: Ag ; N ; O ; Na ; Cl .
pairs with to give silver chloride, an insoluble white curdy precipitate, while sodium nitrate remains dissolved. This is the standard test for a chloride — add silver nitrate solution and a white precipitate appears.
Worked example 3 — lead nitrate and potassium iodide.
Check the balance: Pb ; N ; O ; K ; I .
pairs with to give lead iodide, an insoluble bright yellow precipitate — the reaction this page opened with. Potassium nitrate stays in solution.
What actually drives the reaction. All four ions were free to move before mixing, and they are still free afterwards — except for the pair that formed an insoluble solid. Those ions have been removed from the solution, so the exchange cannot run backwards.
A double decomposition happens only if a product leaves the solution. Mix sodium chloride and potassium nitrate and all four possible compounds are soluble, so nothing precipitates, no ions are removed, and no reaction occurs at all — the beaker simply holds four kinds of ion. So the insoluble product is not a by-product of the reaction; it is the reason there is a reaction, and that is why the arrow in these equations always carries a precipitate.
The other way a product can leave is as a gas. Marble and hydrochloric acid exchange partners and the carbonic acid formed breaks up at once, so carbon dioxide escapes:
A precipitate and a gas do the same job — both remove a product from the mixture, which is why neutralisation, precipitation and gas evolution all belong to this one family.
Worked example 1 — sodium sulphate and barium chloride.
Check the balance: Na ; S ; O ; Ba ; Cl .
The ion exchange. In solution the four ions present are , , and . Of the two new pairings:
- with gives barium sulphate, which is insoluble — it appears as a white precipitate
- with gives sodium chloride, which is soluble and stays in solution
Worked example 2 — silver nitrate and sodium chloride.
Already balanced: Ag ; N ; O ; Na ; Cl .
pairs with to give silver chloride, an insoluble white curdy precipitate, while sodium nitrate remains dissolved. This is the standard test for a chloride — add silver nitrate solution and a white precipitate appears.
Worked example 3 — lead nitrate and potassium iodide.
Check the balance: Pb ; N ; O ; K ; I .
pairs with to give lead iodide, an insoluble bright yellow precipitate — the reaction this page opened with. Potassium nitrate stays in solution.
What actually drives the reaction. All four ions were free to move before mixing, and they are still free afterwards — except for the pair that formed an insoluble solid. Those ions have been removed from the solution, so the exchange cannot run backwards.
A double decomposition happens only if a product leaves the solution. Mix sodium chloride and potassium nitrate and all four possible compounds are soluble, so nothing precipitates, no ions are removed, and no reaction occurs at all — the beaker simply holds four kinds of ion. So the insoluble product is not a by-product of the reaction; it is the reason there is a reaction, and that is why the arrow in these equations always carries a precipitate.
The other way a product can leave is as a gas. Marble and hydrochloric acid exchange partners and the carbonic acid formed breaks up at once, so carbon dioxide escapes:
A precipitate and a gas do the same job — both remove a product from the mixture, which is why neutralisation, precipitation and gas evolution all belong to this one family.
How do you tell an exothermic reaction from an endothermic one?
An exothermic reaction gives out heat and the vessel warms up; an endothermic one takes heat in and the vessel cools or must be heated continuously.
Exothermic reactions — energy is released.
- Slaking quicklime — add water to quicklime and the vessel becomes hot enough to be uncomfortable to hold
- Burning anything: ; burning magnesium also gives a dazzling white light
- Neutralisation:
- Diluting concentrated sulphuric acid — a great deal of heat is released, which is why the acid is added to water and never water to acid
- Respiration in living cells, which releases the energy of glucose
Endothermic reactions — energy is absorbed.
- Thermal decomposition of any kind — the heating must continue or the reaction stops
- , which is why this happens in a lightning flash and not at room temperature
- Dissolving ammonium chloride or ammonium nitrate in water — the beaker becomes noticeably cold
- Photosynthesis, which absorbs light energy
- Electrolysis of water, which absorbs electrical energy
The energy need not be heat. The three forms named in this chapter all appear:
- Heat — slaking lime releases it, decomposing limestone absorbs it
- Light — burning magnesium gives out light; silver bromide decomposing takes in light
- Electricity — a cell's reaction releases electrical energy; electrolysis absorbs it
Worked observation to tell them apart. Put a thermometer in the beaker. Slaking lime drives the reading up; dissolving ammonium nitrate drives it down. That single measurement classifies the reaction, and it is the practical test behind the definitions.
A reaction that needs heating is not automatically endothermic. Burning coal needs a match to start it and then releases far more heat than the match supplied — it is strongly exothermic. Decomposing limestone needs heating throughout, and stops the moment the heating stops.
So the question is not whether heat is needed to begin, but whether heat is given out or taken in overall. The energy used to start a reaction is its activation, and an exothermic reaction pays that back with interest. Confusing the starting energy with the overall energy change is the commonest error here, and a fire is the clearest counter-example.
Exothermic reactions — energy is released.
- Slaking quicklime — add water to quicklime and the vessel becomes hot enough to be uncomfortable to hold
- Burning anything: ; burning magnesium also gives a dazzling white light
- Neutralisation:
- Diluting concentrated sulphuric acid — a great deal of heat is released, which is why the acid is added to water and never water to acid
- Respiration in living cells, which releases the energy of glucose
Endothermic reactions — energy is absorbed.
- Thermal decomposition of any kind — the heating must continue or the reaction stops
- , which is why this happens in a lightning flash and not at room temperature
- Dissolving ammonium chloride or ammonium nitrate in water — the beaker becomes noticeably cold
- Photosynthesis, which absorbs light energy
- Electrolysis of water, which absorbs electrical energy
The energy need not be heat. The three forms named in this chapter all appear:
- Heat — slaking lime releases it, decomposing limestone absorbs it
- Light — burning magnesium gives out light; silver bromide decomposing takes in light
- Electricity — a cell's reaction releases electrical energy; electrolysis absorbs it
Worked observation to tell them apart. Put a thermometer in the beaker. Slaking lime drives the reading up; dissolving ammonium nitrate drives it down. That single measurement classifies the reaction, and it is the practical test behind the definitions.
A reaction that needs heating is not automatically endothermic. Burning coal needs a match to start it and then releases far more heat than the match supplied — it is strongly exothermic. Decomposing limestone needs heating throughout, and stops the moment the heating stops.
So the question is not whether heat is needed to begin, but whether heat is given out or taken in overall. The energy used to start a reaction is its activation, and an exothermic reaction pays that back with interest. Confusing the starting energy with the overall energy change is the commonest error here, and a fire is the clearest counter-example.
What conditions does a reaction need before it will happen?
Bringing the reactants together is often not enough. Most reactions need one or more of heat, light, pressure, a catalyst, a solvent or electricity.
Heat. Many reactions will not start, or will not continue, without it.
Iron filings and sulphur powder can be stirred together indefinitely with nothing happening; heat them and they combine to form iron sulphide.
Light. Some reactions are driven by light alone.
- Photosynthesis in a green leaf needs sunlight
- , which is why silver salts are kept in dark bottles
- Hydrogen and chlorine mixed in the dark are stable, and in bright sunlight they combine explosively
Pressure. A reaction between gases can need a high pressure to push the molecules together.
Catalyst. A substance that changes the rate of a reaction without being used up in it.
- Manganese dioxide speeds the decomposition of potassium chlorate and of hydrogen peroxide
- Iron is the catalyst in the manufacture of ammonia
- The catalyst can be recovered unchanged at the end, which is what distinguishes it from a reactant
Solvent — usually water. This is the condition most often forgotten.
Dry hydrogen chloride gas and dry sodium hydroxide, mixed together dry, do not react. Dissolve each in water and they react instantly. Likewise dry litmus paper is not turned red by dry hydrogen chloride gas, and is turned red at once if either is moist.
The water is not a spectator. It allows the compounds to split into their ions, and it is the ions that react. No water means no ions and no reaction, which is why the moisture condition is a real requirement and not a laboratory nicety.
Electricity. Some decompositions happen only when a current is passed.
Two more requirements that are easy to overlook. The reactants must be in contact — which is why a gas and a solid react at the solid's surface only. And a finely divided reactant reacts faster than a lump of the same mass, because more of it is exposed.
A condition is written above the arrow, not among the reactants. Heat, light, a catalyst and a pressure all go above or below the arrow, because they are not consumed and do not appear in the products. Writing a catalyst as a reactant is a marked error — it would have to appear on the right as well, since it is recovered, and the equation would be cluttered with something that never changed.
And the conditions are part of the answer to "what are the limitations of an equation". As the previous chapter noted, a bare equation tells you nothing about the conditions — so the conditions written over the arrow are the chemist's way of repairing exactly that gap.
Heat. Many reactions will not start, or will not continue, without it.
Iron filings and sulphur powder can be stirred together indefinitely with nothing happening; heat them and they combine to form iron sulphide.
Light. Some reactions are driven by light alone.
- Photosynthesis in a green leaf needs sunlight
- , which is why silver salts are kept in dark bottles
- Hydrogen and chlorine mixed in the dark are stable, and in bright sunlight they combine explosively
Pressure. A reaction between gases can need a high pressure to push the molecules together.
Catalyst. A substance that changes the rate of a reaction without being used up in it.
- Manganese dioxide speeds the decomposition of potassium chlorate and of hydrogen peroxide
- Iron is the catalyst in the manufacture of ammonia
- The catalyst can be recovered unchanged at the end, which is what distinguishes it from a reactant
Solvent — usually water. This is the condition most often forgotten.
Dry hydrogen chloride gas and dry sodium hydroxide, mixed together dry, do not react. Dissolve each in water and they react instantly. Likewise dry litmus paper is not turned red by dry hydrogen chloride gas, and is turned red at once if either is moist.
The water is not a spectator. It allows the compounds to split into their ions, and it is the ions that react. No water means no ions and no reaction, which is why the moisture condition is a real requirement and not a laboratory nicety.
Electricity. Some decompositions happen only when a current is passed.
Two more requirements that are easy to overlook. The reactants must be in contact — which is why a gas and a solid react at the solid's surface only. And a finely divided reactant reacts faster than a lump of the same mass, because more of it is exposed.
A condition is written above the arrow, not among the reactants. Heat, light, a catalyst and a pressure all go above or below the arrow, because they are not consumed and do not appear in the products. Writing a catalyst as a reactant is a marked error — it would have to appear on the right as well, since it is recovered, and the equation would be cluttered with something that never changed.
And the conditions are part of the answer to "what are the limitations of an equation". As the previous chapter noted, a bare equation tells you nothing about the conditions — so the conditions written over the arrow are the chemist's way of repairing exactly that gap.
Exam tip
Exam tip: count the free elements to classify a reaction
Classify by counting. One product means combination; one reactant means decomposition; a free element on the left means displacement; two compounds each side with no free element means double decomposition.
Displacement has an element; double decomposition never does. That single test separates the pair that gets confused.
Name the kind of decomposition — thermal, electrolytic or photochemical — after what supplies the energy.
Mark the precipitate with a downward arrow and the gas with an upward one. and carry marks.
Give the colour of the standard precipitates: barium sulphate white, silver chloride white and curdy, lead iodide bright yellow.
Explain the ion exchange in two steps — which ions are present, and which new pairing is insoluble.
Say why the reaction happens: a product leaves the solution as a precipitate or a gas, so the exchange cannot reverse.
Check every equation element by element and write that check out.
Needing heat to start is not the same as endothermic. Burning coal needs a match and is strongly exothermic.
Write conditions ABOVE the arrow, never as a reactant — a catalyst is recovered unchanged.
And remember the water condition: dry hydrogen chloride and dry sodium hydroxide do not react, because there are no ions without a solvent.
Displacement has an element; double decomposition never does. That single test separates the pair that gets confused.
Name the kind of decomposition — thermal, electrolytic or photochemical — after what supplies the energy.
Mark the precipitate with a downward arrow and the gas with an upward one. and carry marks.
Give the colour of the standard precipitates: barium sulphate white, silver chloride white and curdy, lead iodide bright yellow.
Explain the ion exchange in two steps — which ions are present, and which new pairing is insoluble.
Say why the reaction happens: a product leaves the solution as a precipitate or a gas, so the exchange cannot reverse.
Check every equation element by element and write that check out.
Needing heat to start is not the same as endothermic. Burning coal needs a match and is strongly exothermic.
Write conditions ABOVE the arrow, never as a reactant — a catalyst is recovered unchanged.
And remember the water condition: dry hydrogen chloride and dry sodium hydroxide do not react, because there are no ions without a solvent.
Did you know
Why nothing at all happens when you mix two soluble salts
Mix sodium chloride solution with potassium nitrate solution. Stir. Wait.
Nothing happens. No precipitate, no gas, no colour, no warmth.
And yet the ions did exactly what they do in every double decomposition — they moved about freely and met every possible partner. met ; met . All four new pairings were formed and unformed constantly.
The trouble is that all four possible compounds are soluble. Sodium chloride, potassium nitrate, sodium nitrate and potassium chloride would every one of them stay dissolved. So nothing ever settles out, nothing is removed, and the ions simply keep mingling.
Evaporate the water and you get a mixture of crystals — but that is the water leaving, not a reaction.
So whether a double decomposition "happens" is not a question about the ions at all. It is a question about solubility, and the whole reaction depends on one of the four possible products refusing to dissolve.
That gives a genuinely useful prediction rule. Look at the four compounds that could form, and ask whether any of them is insoluble. If one is, you have a reaction and a precipitate; if none is, you have a mixture and nothing more.
It also explains why the same three or four reactions appear in every textbook. Barium sulphate, silver chloride and lead iodide are reliably insoluble, so they precipitate every time and make clean demonstrations. The reactions are famous because the products are stubborn, not because the chemistry is special — and that stubbornness is what turns them into the standard laboratory tests for a sulphate, a chloride and an iodide.
Nothing happens. No precipitate, no gas, no colour, no warmth.
And yet the ions did exactly what they do in every double decomposition — they moved about freely and met every possible partner. met ; met . All four new pairings were formed and unformed constantly.
The trouble is that all four possible compounds are soluble. Sodium chloride, potassium nitrate, sodium nitrate and potassium chloride would every one of them stay dissolved. So nothing ever settles out, nothing is removed, and the ions simply keep mingling.
Evaporate the water and you get a mixture of crystals — but that is the water leaving, not a reaction.
So whether a double decomposition "happens" is not a question about the ions at all. It is a question about solubility, and the whole reaction depends on one of the four possible products refusing to dissolve.
That gives a genuinely useful prediction rule. Look at the four compounds that could form, and ask whether any of them is insoluble. If one is, you have a reaction and a precipitate; if none is, you have a mixture and nothing more.
It also explains why the same three or four reactions appear in every textbook. Barium sulphate, silver chloride and lead iodide are reliably insoluble, so they precipitate every time and make clean demonstrations. The reactions are famous because the products are stubborn, not because the chemistry is special — and that stubbornness is what turns them into the standard laboratory tests for a sulphate, a chloride and an iodide.
Exam relevance
How are reaction types tested in JEE Main and NEET?
Because the four types become the framework for redox and ionic chemistry, and the precipitation reactions become the whole of qualitative analysis.
This is the foundation for Class 11 Chemistry Redox Reactions and Equilibrium, and Class 12 Electrochemistry, all examined in JEE Main and NEET. Displacement reactions become redox reactions: when zinc displaces copper, the zinc is oxidised and the copper reduced, and Class 11 tracks that with oxidation numbers and half-equations. The reactivity ordering that decides which metal displaces which becomes the electrochemical series, and Class 12 attaches a numerical electrode potential to every entry so that the prediction becomes a calculation.
Double decomposition becomes ionic equilibrium. Class 11 Equilibrium and Class 12 introduce the solubility product, a number for each sparingly soluble salt, and a precipitate forms exactly when the ion concentrations exceed it. The rule discovered on this page — that a reaction happens only if a product is insoluble — becomes a quantitative test, and solubility-product numericals are recurring JEE Main material.
The three named precipitations are the basis of qualitative analysis. The practical syllabus for both boards and the NEET and JEE practical component identify basic and acidic radicals by exactly these tests: silver nitrate for a chloride, barium chloride for a sulphate, lead nitrate for an iodide. Match-the-column questions pairing a reagent with a precipitate colour are standard, which is why the colours are worth learning alongside the equations.
Heat changes become thermochemistry. Class 11 Thermodynamics gives the enthalpy change a sign: negative for exothermic, positive for endothermic, written as . The thermometer test on this page is that sign measured, and Hess's law and bond-enthalpy calculations follow. The distinction drawn here between the energy needed to start a reaction and the overall change becomes activation energy against enthalpy change in Class 12 Chemical Kinetics — and that pair is a standard assertion-reason item.
Catalysts get their own treatment. Class 12 Chemical Kinetics explains that a catalyst lowers the activation energy without altering the enthalpy change or the position of equilibrium, and Class 12 Surface Chemistry classifies catalysts and explains adsorption. The point made here — that a catalyst is recovered unchanged — is what distinguishes it from a reactant there too.
The solvent condition becomes ionisation. Class 11 explains that water's high polarity separates an ionic lattice into free ions, which is why the dry-gas experiment fails. For NEET Biology, enzymes are biological catalysts and respiration and photosynthesis are the standard exothermic and endothermic pair, examined directly.
What the questions look like. For board work, expect classify a given reaction with its balanced equation, write the three precipitation equations and explain the ion exchange, distinguish exothermic from endothermic with examples, and identify the conditions from a reaction description. Colours and arrows carry marks. For JEE Main and NEET, expect redox balancing, electrochemical-series predictions, solubility-product calculations, enthalpy signs and kinetics.
How board and competitive emphasis differ. A board paper rewards the named type with the equation and the colour. A competitive paper assumes the classification and asks whether a given displacement will occur at all, which needs the electrode potentials.
The single trap that costs the most marks. Calling a reaction endothermic because it needs heating. Burning coal needs a match and is strongly exothermic; decomposing limestone needs heating continuously and is endothermic. The defence is to ask whether the heating can be stopped once the reaction is going — if the reaction keeps itself going, it is exothermic, and if it stops the moment the flame is removed, it is endothermic.
This is the foundation for Class 11 Chemistry Redox Reactions and Equilibrium, and Class 12 Electrochemistry, all examined in JEE Main and NEET. Displacement reactions become redox reactions: when zinc displaces copper, the zinc is oxidised and the copper reduced, and Class 11 tracks that with oxidation numbers and half-equations. The reactivity ordering that decides which metal displaces which becomes the electrochemical series, and Class 12 attaches a numerical electrode potential to every entry so that the prediction becomes a calculation.
Double decomposition becomes ionic equilibrium. Class 11 Equilibrium and Class 12 introduce the solubility product, a number for each sparingly soluble salt, and a precipitate forms exactly when the ion concentrations exceed it. The rule discovered on this page — that a reaction happens only if a product is insoluble — becomes a quantitative test, and solubility-product numericals are recurring JEE Main material.
The three named precipitations are the basis of qualitative analysis. The practical syllabus for both boards and the NEET and JEE practical component identify basic and acidic radicals by exactly these tests: silver nitrate for a chloride, barium chloride for a sulphate, lead nitrate for an iodide. Match-the-column questions pairing a reagent with a precipitate colour are standard, which is why the colours are worth learning alongside the equations.
Heat changes become thermochemistry. Class 11 Thermodynamics gives the enthalpy change a sign: negative for exothermic, positive for endothermic, written as . The thermometer test on this page is that sign measured, and Hess's law and bond-enthalpy calculations follow. The distinction drawn here between the energy needed to start a reaction and the overall change becomes activation energy against enthalpy change in Class 12 Chemical Kinetics — and that pair is a standard assertion-reason item.
Catalysts get their own treatment. Class 12 Chemical Kinetics explains that a catalyst lowers the activation energy without altering the enthalpy change or the position of equilibrium, and Class 12 Surface Chemistry classifies catalysts and explains adsorption. The point made here — that a catalyst is recovered unchanged — is what distinguishes it from a reactant there too.
The solvent condition becomes ionisation. Class 11 explains that water's high polarity separates an ionic lattice into free ions, which is why the dry-gas experiment fails. For NEET Biology, enzymes are biological catalysts and respiration and photosynthesis are the standard exothermic and endothermic pair, examined directly.
What the questions look like. For board work, expect classify a given reaction with its balanced equation, write the three precipitation equations and explain the ion exchange, distinguish exothermic from endothermic with examples, and identify the conditions from a reaction description. Colours and arrows carry marks. For JEE Main and NEET, expect redox balancing, electrochemical-series predictions, solubility-product calculations, enthalpy signs and kinetics.
How board and competitive emphasis differ. A board paper rewards the named type with the equation and the colour. A competitive paper assumes the classification and asks whether a given displacement will occur at all, which needs the electrode potentials.
The single trap that costs the most marks. Calling a reaction endothermic because it needs heating. Burning coal needs a match and is strongly exothermic; decomposing limestone needs heating continuously and is endothermic. The defence is to ask whether the heating can be stopped once the reaction is going — if the reaction keeps itself going, it is exothermic, and if it stops the moment the flame is removed, it is endothermic.
Key takeaways
Reaction types, ion exchange and heat changes: quick revision
- Combination: two or more substances give one product. ; ; .
- Decomposition: one compound gives two or more. Thermal — , ; electrolytic — ; photochemical — .
- Displacement: a more reactive element pushes out a less reactive one. — the blue fades and red-brown copper deposits.
- Also ; ; .
- Double decomposition: two compounds exchange radicals.
- — white precipitate.
- — white curdy precipitate, the test for a chloride.
- — bright yellow precipitate.
- To classify, count the free elements: one product means combination, one reactant means decomposition, a free element on the left means displacement, and two compounds each side means double decomposition.
- The reaction happens because a product LEAVES the solution — as a precipitate or as a gas, so the exchange cannot reverse.
- Mixing two soluble salts gives no reaction at all if all four possible products are soluble.
- — a gas doing the same job as a precipitate.
- Exothermic gives out energy: slaking quicklime, all burning, neutralisation, diluting concentrated sulphuric acid (add acid to water), respiration.
- Endothermic takes energy in: all thermal decomposition, , dissolving ammonium nitrate (the beaker goes cold), photosynthesis, electrolysis.
- The energy can be heat, light or electricity — magnesium burning emits light, silver bromide absorbs it; a cell releases electrical energy, electrolysis absorbs it.
- A thermometer classifies the reaction — up for exothermic, down for endothermic.
- Needing heat to start is not endothermic. Coal needs a match and is exothermic; limestone needs heating throughout.
- Conditions: heat (limestone, iron and sulphur), light (photosynthesis, silver salts, hydrogen with chlorine), pressure (ammonia synthesis), catalyst (manganese dioxide, iron), solvent (water), electricity (electrolysis).
- Dry hydrogen chloride and dry sodium hydroxide do NOT react — water is needed to give ions, and dry litmus is not reddened by dry hydrogen chloride gas.
- A catalyst changes the rate and is recovered unchanged, so it is written above the arrow, never as a reactant.
- Reactants must be in contact, and a finely divided solid reacts faster than a lump.
Mix a little washing soda solution with a pinch of Epsom salt dissolved in water and watch for a cloudiness — then work out which of the four possible products refused to dissolve.
- Decomposition: one compound gives two or more. Thermal — , ; electrolytic — ; photochemical — .
- Displacement: a more reactive element pushes out a less reactive one. — the blue fades and red-brown copper deposits.
- Also ; ; .
- Double decomposition: two compounds exchange radicals.
- — white precipitate.
- — white curdy precipitate, the test for a chloride.
- — bright yellow precipitate.
- To classify, count the free elements: one product means combination, one reactant means decomposition, a free element on the left means displacement, and two compounds each side means double decomposition.
- The reaction happens because a product LEAVES the solution — as a precipitate or as a gas, so the exchange cannot reverse.
- Mixing two soluble salts gives no reaction at all if all four possible products are soluble.
- — a gas doing the same job as a precipitate.
- Exothermic gives out energy: slaking quicklime, all burning, neutralisation, diluting concentrated sulphuric acid (add acid to water), respiration.
- Endothermic takes energy in: all thermal decomposition, , dissolving ammonium nitrate (the beaker goes cold), photosynthesis, electrolysis.
- The energy can be heat, light or electricity — magnesium burning emits light, silver bromide absorbs it; a cell releases electrical energy, electrolysis absorbs it.
- A thermometer classifies the reaction — up for exothermic, down for endothermic.
- Needing heat to start is not endothermic. Coal needs a match and is exothermic; limestone needs heating throughout.
- Conditions: heat (limestone, iron and sulphur), light (photosynthesis, silver salts, hydrogen with chlorine), pressure (ammonia synthesis), catalyst (manganese dioxide, iron), solvent (water), electricity (electrolysis).
- Dry hydrogen chloride and dry sodium hydroxide do NOT react — water is needed to give ions, and dry litmus is not reddened by dry hydrogen chloride gas.
- A catalyst changes the rate and is recovered unchanged, so it is written above the arrow, never as a reactant.
- Reactants must be in contact, and a finely divided solid reacts faster than a lump.
Mix a little washing soda solution with a pinch of Epsom salt dissolved in water and watch for a cloudiness — then work out which of the four possible products refused to dissolve.