Class 10 Science – Chapter 2: Acids, Bases and Salts

1. Introduction – Acids, Bases and Salts

Every drink, food item and cleaning product around you is classified in chemistry by its nature: acidic, basic or neutral. Substances like lemon juice and vinegar taste sour, while substances like soap solution feel slippery. These everyday observations are linked to chemical properties.

How acids and bases differ

In water, an acid releases hydrogen ions (H⁺), while a base releases hydroxide ions (OH⁻). These two ions react together to form water, which is why acids and bases neutralise each other. Along with their different tastes and litmus tests, this ionic behaviour is the fundamental difference between them.

Everyday examples of acids and bases

2. Acids and Bases in Terms of Ions

When acids and bases dissolve in water, they break apart (dissociate) into positively and negatively charged particles called ions. This ionic behaviour in aqueous solution is what gives acids and bases their characteristic properties.

What are ions?

An ion is a charged particle that forms when a neutral atom or a group of atoms gains or loses electrons. A positively charged ion is a cation (like H⁺), and a negatively charged ion is an anion (like Cl⁻).

The hydrogen ion (H⁺) and hydronium ion (H₃O⁺)

When an acid dissolves in water, it produces hydrogen ions (H⁺). A hydrogen ion is simply a hydrogen atom that has lost its electron. Because it is a bare proton, H⁺ does not remain free in water — it quickly combines with a water molecule to form the hydronium ion (H₃O⁺).

\[ H^+ + H_2O \rightarrow H_3O^+ \]

The hydroxide ion (OH⁻)

When a base dissolves in water, it releases hydroxide ions (OH⁻). Sodium hydroxide, for example, dissociates completely in water.

Why aqueous solutions matter

Pure HCl gas and pure H₂SO₁₀ show acidic behaviour only in the presence of water. It is the dissociation into H⁺ (as H₃O⁺) that gives an acid its acidic character, and dissociation into OH⁻ that gives a base its basic character. In the absence of water (for example, in dry benzene or in the dry gaseous state), these ions are not formed and the substances do not show acidic or basic behaviour.

3. Chemical Properties of Acids

A. Reaction of acids with metals

Acids react with active metals (metals above hydrogen in the reactivity series) to form a salt and hydrogen gas.

General pattern: Acid + Metal → Salt + Hydrogen gas

Testing for hydrogen gas

Bring a burning splinter (or a lighted match) near the mouth of the test tube. If hydrogen is present, it burns with a characteristic pop sound, because hydrogen is a highly combustible gas.

B. Reaction of acids with metal carbonates

Acids react with metal carbonates to form a salt, carbon dioxide gas and water.

General pattern: Acid + Metal carbonate → Salt + Water + Carbon dioxide

Testing for carbon dioxide

Pass the gas through lime water (calcium hydroxide solution). Carbon dioxide turns lime water milky because insoluble calcium carbonate is formed.

\[ Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l) \]

C. Reaction of acids with metal hydrogencarbonates

Acids react with metal hydrogencarbonates (bicarbonates) to form a salt, carbon dioxide gas and water.

General pattern: Acid + Metal hydrogencarbonate → Salt + Water + Carbon dioxide

D. Reaction of acids with bases (neutralisation)

Acids react with bases to form a salt and water. This reaction is called neutralisation, because the acidic (H⁺) and basic (OH⁻) ions combine to form neutral water.

\[ H^+(aq) + OH^-(aq) \rightarrow H_2O(l) \]

E. Reaction of acids with metal oxides

Metal oxides are basic oxides: in water they form bases (for example, CaO in water gives Ca(OH)₂). That is why they react with acids like a base would, to form a salt and water.

General pattern: Acid + Metal oxide → Salt + Water

4. Chemical Properties of Bases

A. Reaction of bases with acids (neutralisation)

Bases neutralise acids to form a salt and water, exactly as in the acid + base section above (the same reaction viewed from the base side).

B. Reaction of bases with non-metal oxides

Non-metal oxides such as CO₂ and SO₂ are acidic oxides — they form acids with water. That is why they react with bases to form a salt and water.

General pattern: Base + Non-metal oxide → Salt + Water

C. Reaction of bases with certain metals

Some metals react with strong alkalis like sodium hydroxide to form a salt and hydrogen gas. The metals that do this are amphoteric metals such as zinc and aluminium, which can react with both acids and bases.

5. What do All Acids and All Bases have in Common?

All acids, though they are different compounds, behave alike as a class. Same for bases. The reason lies in the ions they produce in water.

What all acids have in common

Every acid produces H⁺ ions when dissolved in water. In water these exist as hydronium ions (H₃O⁺). It is these hydrogen ions that:

• give the sour taste of acids,
• turn blue litmus red,
• make the solution conduct electricity, and
• react with bases to form water.

What all bases have in common

Every base produces OH⁻ (hydroxide) ions when dissolved in water. These hydroxide ions:

• give the bitter taste of bases,
• turn red litmus blue,
• make the solution soapy/slippery, and
• combine with H⁺ to form water.

6. Indicators

An indicator is a substance whose colour (or smell) changes depending on whether it is in an acidic or a basic medium. Indicators help us detect acids and bases without tasting them.

Types of indicators

• Natural indicators – obtained from nature, e.g. litmus, turmeric, red cabbage juice, china rose (hibiscus) petals.
• Synthetic indicators – made in the laboratory, e.g. methyl orange, phenolphthalein.
• Olfactory indicators – substances whose smell changes in acidic or basic medium, e.g. onion, vanilla essence, clove oil.

Litmus

Litmus is extracted from lichens. It is available as blue and red litmus papers. Blue litmus turns red in an acid; red litmus turns blue in a base. Litmus gives no colour change in a neutral solution.

Methyl orange

Methyl orange is red in an acidic medium and yellow in a basic medium.

Phenolphthalein

Phenolphthalein is colourless in an acidic and neutral medium and turns pink in a basic medium.

Turmeric

Turmeric is yellow in an acidic and neutral medium and turns red in a basic medium. This is why turmeric stains on cloth turn red when washed with soap (a base).

Indicator colour table

Olfactory indicators

Onion juice has a pungent smell. Onion juice loses its smell in a basic medium but keeps its smell in an acidic medium. Vanilla essence loses its characteristic smell in a basic medium but retains it in an acidic medium. Because the change is detected by smell, these are called olfactory indicators.

7. The pH Scale

Litmus and indicators tell us only whether a solution is acidic or basic. To measure how strong an acid or base is, we use the pH scale.

The pH scale

• pH = 0 to 6 – acidic solution. The smaller the pH, the stronger the acid.
• pH = 7 – neutral solution (pure water).
• pH = 8 to 14 – basic solution. The larger the pH, the stronger the base.

Everyday applications of pH

• Farmers test soil pH to decide whether to add a base (lime) or an acid to soil.
• pH of skin-care products is kept close to the natural pH of skin.
• pH indicators (pH paper) are used in water-testing and swimming pools.
• Mild base solutions are used to neutralise acid spills; mild acid solutions neutralise base spills.

8. Importance of pH in Everyday Life

A. pH in our digestive system

Our stomach produces hydrochloric acid (gastric juice) which helps digest food. This acid has a very low pH (about 1.0–2.0). Excess acid causes acidity or heartburn. To relieve it, people take antacids — mild bases such as magnesium hydroxide (milk of magnesia, Mg(OH)₂) or sodium hydrogencarbonate (baking soda, NaHCO₃) that neutralise the excess stomach acid.

\[ Mg(OH)_2(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + 2H_2O(l) \]

B. pH and tooth decay

The mouth has bacteria that break down the sugar in food into acids. These acids lower the pH of the mouth below 5.5. Tooth enamel (made of calcium phosphate) is attacked by this acid, and decay starts. Brushing after meals removes the food and the acid; using a toothpaste that is basic neutralises the excess acid, protecting the enamel.

C. Self-defence by plants and animals

Many organisms use acids and bases to defend themselves:

• A bee sting injects methanoic (formic) acid into the skin, causing pain. Applying a mild base such as baking soda (NaHCO₃) neutralises the acid.
• A wasp sting is basic. Applying a mild acid such as vinegar (acetic acid) neutralises it.
• Nettle plants release methanoic acid on their stinging hairs; rubbing a dock leaf (a mild base) or baking soda paste gives relief.

D. pH of soil

Plants grow best in soil of a suitable pH (generally near neutral, about 6.5–7). If soil is too acidic, farmers add a weak base like lime (calcium oxide/hydroxide) to neutralise it. If soil is too basic, organic manure or a suitable acid treatment is used. This is why fertilisers and soil testing are important in agriculture.

9. Salts

Formation of salts – neutralisation

Salts are most commonly formed by the neutralisation of an acid with a base: acid + base → salt + water.

Types of salts (Class 10 level)

• Normal salts – formed when all the hydrogen of the acid is replaced. Example: NaCl, Na₂SO₁₀, KNO₃.
• Acid salts – formed when only part of the hydrogen of a polyprotic acid is replaced. Example: sodium hydrogencarbonate (NaHCO₃), sodium bisulphate (NaHSO₁₀).
• Basic salts – contain hydroxide groups as well as a metal. Example: basic copper carbonate CuCO₃·Cu(OH)₂.
• Salt hydrates (hydrated salts) – salts containing water of crystallisation. Example: blue vitriol CuSO₁₀·5H₂O, washing soda Na₂CO₃·10H₂O.

10. Common Salts and their Preparation

A. Sodium chloride (common salt)

Formula: NaCl

Common salt is found dissolved in sea water and in underground salt deposits (rock salt). Sea water is evaporated in large shallow beds to obtain salt. Rock salt is mined.

Sodium chloride is the starting material for the manufacture of sodium hydroxide, sodium carbonate and bleaching powder.

Important uses:

• Flavouring and preserving food.
• Used for making sodium hydroxide, washing soda and bleaching powder.
• Used on icy roads in winter to lower the freezing point.

B. Sodium hydroxide (chlor-alkali process)

Formula: NaOH

When an electric current is passed through an aqueous solution of sodium chloride (brine), it decomposes to give chlorine, hydrogen gas and sodium hydroxide. This is called the chlor-alkali process.

\[ 2NaCl(aq) + 2H_2O(l) \xrightarrow{electricity} 2NaOH(aq) + Cl_2(g) + H_2(g) \]

The products are highly useful:

• Chlorine (Cl₂) – used for water purification, bleaching, and making PVC, insecticides and HCl.
• Hydrogen (H₂) – used for making ammonia (Haber process) and as a fuel.
• Sodium hydroxide (NaOH) – used for making soap, paper, artificial rayon and in degreasing.

C. Bleaching powder

Formula: CaOCl₂ (calcium oxychloride)

Bleaching powder is prepared by passing chlorine gas over dry slaked lime (calcium hydroxide).

\[ Ca(OH)_2 + Cl_2 \rightarrow CaOCl_2 + H_2O \]

Properties: bleaching action, oxidising power, reacts with acids to release chlorine, used to disinfect water.

Uses: bleaching of cotton and linen, disinfecting water, removing stains, and as an oxidising agent.

D. Baking soda

Chemical name: Sodium hydrogencarbonate (sodium bicarbonate)

Formula: NaHCO₃

Baking soda is prepared by passing carbon dioxide through brine saturated with ammonia (the Solvay-type process).

\[ NaCl(aq) + NH_3(aq) + H_2O(l) + CO_2(g) \rightarrow NaHCO_3(s) + NH_4Cl(aq) \]

Properties and everyday use:

• It is used in baking (batter rises as CO₂ is released).
• It is a mild base and is used in antacids.
• It develops a heating effect: on heating it decomposes to give sodium carbonate, water and carbon dioxide.

\[ 2NaHCO_3(s) \xrightarrow{\Delta} Na_2CO_3(s) + H_2O(l) + CO_2(g) \]

E. Washing soda

Chemical name: Sodium carbonate decahydrate

Formula: Na₂CO₃·10H₂O

Washing soda is obtained by heating baking soda (the above decomposition), which gives anhydrous sodium carbonate, then dissolving it in water and recrystallising it (adding water of crystallisation).

\[ Na_2CO_3(s) + 10H_2O(l) \rightarrow Na_2CO_3 \cdot 10H_2O(s) \]

Uses: washing of clothes, making of glass and soap, softening of hard water, and as a laboratory reagent.

11. Crystallisation

Crystallisation is used to:

• separate a pure solid from a solution, for example common salt from sea water or crystals of copper sulphate from its solution;
• purify solids that are contaminated (e.g. impure salt) — the impurities, being more soluble or less soluble, remain in the mother liquor;
• obtain crystals with a definite shape and water of crystallisation (e.g. blue vitriol, potash alum).

12. Water of Crystallisation

Examples

Effect of heating

On heating, hydrated salts lose their water of crystallisation and become anhydrous. For example:

• Hydrated copper sulphate (blue) loses water to become white anhydrous copper sulphate.

\[ CuSO_4 \cdot 5H_2O(s) \xrightarrow{\Delta} CuSO_4(s) + 5H_2O(g) \]

• Adding a few drops of water to the white anhydrous copper sulphate turns it blue again (test for water).
• Washing soda and other hydrates similarly lose water on heating.

13. Important Chemical Equations (Quick Reference)

14. Formulae Quick-Reference Table

15. Comparison Tables

A. Acids vs Bases

B. Strong acids vs Weak acids

C. Acidic vs Neutral vs Basic solutions

D. Natural vs Synthetic indicators

E. Baking soda vs Washing soda

F. Salts and their uses

16. Application-Based Learning

Here are everyday situations that use the ideas of this chapter. Work out the answer yourself before reading it.

17. Competency and Reasoning Examples

18. Case-Based Questions

19. Solved Examples

Basic level

Moderate level

Application level

Competency level

Case-based level

20. Common Mistakes to Avoid

1. 1
   Wrong formulae. Learn by heart: HCl, H₂SO₁₀, HNO₃, NaOH, KOH, Ca(OH)₂. A wrong formula costs marks even when the concept is right.
2. 2
   Unbalanced equations. Balance atoms on both sides. For example, H₂SO₁₀ + 2NaOH → Na₂SO₁₀ + 2H₂O (the "2" on NaOH and H₂O is easy to forget).
3. 3
   Confusing baking soda and washing soda. Baking soda = NaHCO₃ (sodium hydrogencarbonate); washing soda = Na₂CO₃·10H₂O (sodium carbonate decahydrate).
4. 4
   Writing "acid gives hydrogen gas" instead of hydrogen ions. Acids release H⁺ ions (as H₃O⁺) in water; H₂ gas appears only in particular reactions like acid + active metal.
5. 5
   Forgetting the pop test. H₂ burns with a pop; CO₂ turns lime water milky. Mixing up the two tests is a classic error.
6. 6
   Indicator colour mix-ups. Turmeric red in base; phenolphthalein pink in base; methyl orange red in acid, yellow in base. Frequently inverted by students.
7. 7
   pH of pure water written as "water is pH 0". Pure water is pH 7 (neutral).
8. 8
   Thinking the bee sting should be treated with an acid. Bee sting is acid — treat with a mild base. Wasp sting is base — treat with a mild acid.
9. 9
   Chlor-alkali products. Remember the three products of electrolysing brine: NaOH (cathode side), Cl₂ and H₂ (anode side). Do not drop any.
10. 10
    Water of crystallisation. CuSO₁₀·5H₂O loses 5H₂O on heating (blue → white). The number "5" is part of the formula — keep it in the equation.
11. 11
    Metal oxide vs non-metal oxide. Metal oxides are basic (react with acids); non-metal oxides are acidic (react with bases). Mixing this reverses the reaction products.
12. 12
    Not writing states or conditions. Add (s), (aq), (g), Δ (heat) or "electricity" where asked; the exam rewards complete, precise equations.
13. 13
    Calling all bases soluble. A base that dissolves in water is an alkali; not all bases dissolve (e.g. Cu(OH)₂ is sparingly soluble).
14. 14
    Antacid direction. Antacids are bases, not acids. They lower acidity in the stomach.

21. Board Exam Tips

• Read the whole question before writing; many errors come from answering only part of a 3-mark question.
• Write balanced equations with physical states (s), (aq), (g) and the arrow conditions (Δ, electricity).
• Structure long answers: define, give the equation, then the observation. For example, for "acid + metal" write the pattern, one equation, and the pop test.
• Learn the formulae table fully — it directly answers 1-mark formula questions.
• pH questions: state pH = 7 neutral, < 7 acidic, > 7 basic, then justify with ions (H⁺ vs OH⁻).
• Everyday-life applications (bee sting, tooth decay, soil pH, antacids, baking soda in cooking) come up frequently — summarise each in one or two lines.
• Common name ↔ chemical name ↔ formula must be a reflex: quicklime CaO, slaked lime Ca(OH)₂, blue vitriol CuSO₁₀·5H₂O, etc.
• Practise case-based questions: the pattern is a short situation + 3 questions; answer each independently using the chapter concepts.
• Time management: attempt 1-mark questions first, then the equation-based questions, leaving the descriptive ones for last with neat headings.
• Underline key terms like hydronium ion, water of crystallisation, neutralisation — markers reward precision.

22. Last-Minute Revision

1. Acids: sour taste, blue litmus → red, produce H⁺/H₃O⁺ in water.
2. Bases: bitter taste, soapy feel, red litmus → blue, produce OH⁻ in water.
3. Hydronium ion: H⁺ + H₂O → H₃O⁺ — this is the acid ion actually present in water.
4. Acid + metal: salt + H₂ (pop test: burning splinter pops).
5. Acid + carbonate / hydrogencarbonate: salt + H₂O + CO₂ (CO₂ turns lime water milky).
6. Neutralisation: acid + base → salt + water; H⁺ + OH⁻ → H₂O.
7. Metal oxides are basic; non-metal oxides are acidic.
8. Indicators (acid → base): methyl orange red → yellow; phenolphthalein colourless → pink; turmeric yellow → red; litmus red ↔ blue.
9. Olfactory indicators: onion, vanilla, clove oil (smell changes in base).
10. pH scale 0–14: < 7 acidic (lower = stronger), 7 neutral, > 7 basic (higher = stronger).
11. Stomach: HCl ~pH 1–2; antacid = mild base (Mg(OH)₂, baking soda).
12. Tooth decay: acid below pH 5.5 attacks enamel; basic toothpaste neutralises.
13. Bee sting: methanoic acid → treat with baking soda; wasp sting: basic → treat with vinegar.
14. Soil pH: too acidic → add lime; too basic → add manure/mild acid.
15. Chlor-alkali: brine electrolysis → NaOH + Cl₂ + H₂.
16. Bleaching powder CaOCl₂: chlorine + dry slaked lime; bleaches and disinfects.
17. Baking soda NaHCO₃: on heating → Na₂CO₃ + H₂O + CO₂.
18. Washing soda Na₂CO₃·10H₂O: from Na₂CO₃ + 10H₂O.
19. Water of crystallisation: CuSO₁₀·5H₂O → (heating) CuSO₁₀ + 5H₂O; blue ↔ white test.
20. Crystallisation: pure crystals from hot saturated solution; used to purify salts and obtain hydrated crystals.

23. Self-Practice Questions

Attempt these on your own, then check the answers below.

Conceptual (6)

1. What is an alkali?
2. Why do all bases have a bitter taste and a soapy feel?
3. State the ion present in all acids in aqueous solution.
4. Why does dry HCl gas not conduct electricity?
5. What is meant by neutralisation?
6. Name any two olfactory indicators.

Equation based (6)

7. Write the balanced equation for zinc + dilute hydrochloric acid.
8. Write the equation for the reaction of sodium carbonate with HCl.
9. Write the chlor-alkali process equation.
10. Write the equation for the preparation of bleaching powder.
11. Write the decomposition of baking soda on heating.
12. Write the reaction between H₂SO₁₀ and KOH.

pH and indicators (5)

13. A solution has pH 2 and another has pH 12. Which is more basic and which is more acidic?
14. What is the colour of phenolphthalein in a basic solution?
15. What is the pH of pure water?
16. Why does turmeric turn red when soap is rubbed on it?
17. Name the indicator used to test hydrogen gas (pop test).

Application (5)

18. Why does a person with acidity take antacid?
19. What treats a bee sting better: vinegar or baking soda paste?
20. How does brushing teeth after a meal prevent decay?
21. Why does a farmer add lime to acidic soil?
22. A cake made without an acid ingredient tastes soapy. Why?

Competency (5)

23. How would you prove that the blue colour of copper sulphate crystals comes from its water of crystallisation?
24. Two powders look alike. One is baking soda and the other washing soda. How would you distinguish them chemically?
25. Red litmus turns blue in solution X, and blue litmus turns red in solution Y. What are X and Y?
26. A gas turns lime water milky but burns when a burning splinter is brought. Explain whether this can happen.
27. Why should zinc not be used in the laboratory reaction with H₂SO₁₀ to prepare H₂ gas collection without care?

Case based (3)

28. A gardener notices his plants yellowing; soil test shows pH 3.5. Suggest the remedy and the balanced equation of the remedy with soil acid.
29. A student passes CO₂ through lime water. It turns milky, then clears on passing more gas. Explain with equations.
30. A street food vendor cleans his utensils with washing soda water. The next day a customer's fingers turn slippery and bitter. Relate this to the nature of washing soda and its pH.

Answers

1. A base that dissolves in water.
2. Because they release hydroxide ions (OH⁻) in solution, which produce bitter taste and soapy/slippery feel.
3. H⁺ ions (present as hydronium ions H₃O⁺).
4. Dry HCl has no free ions / does not dissociate; it has no H⁺ ions to conduct electricity.
5. The reaction of an acid with a base to give salt and water: acid + base → salt + water.
6. Onion juice and vanilla essence (also clove oil).
7. Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g).
8. Na₂CO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g).
9. 2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + Cl₂(g) + H₂(g).
10. Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O.
11. 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(l) + CO₂(g).
12. H₂SO₁₀(aq) + 2KOH(aq) → K₂SO₁₀(aq) + 2H₂O(l).
13. pH 12 is more basic; pH 2 is more acidic.
14. Pink.
15. 7 (neutral).
16. Soap solution is basic; turmeric is a natural indicator that turns red in a base.
17. None — a burning splinter is brought near H₂ to prove it (pop). Litmus is the indicator for acid/base tests.
18. Antacids are mild bases that neutralise the excess hydrochloric acid in the stomach.
19. Baking soda paste (mild base) is better for a bee sting — it neutralises the methanoic acid.
20. Brushing removes food and the acid formed by bacteria, keeping the mouth pH above 5.5 and protecting enamel.
21. Lime (CaO / Ca(OH)₂) is a base that neutralises the soil acid, raising pH toward neutral.
22. Without an acid, baking soda leaves Na₂CO₃ (a base) in the cake, giving a soapy/bitter taste.
23. Heat blue CuSO₁₀·5H₂O — it turns white (water lost), and adding a few drops of water turns it blue again.
24. Add dilute HCl to each: baking soda (NaHCO₃) gives brisk, short effervescence of CO₂, washing soda (Na₂CO₃·10H₂O) also fizzes but is a stronger base; a simple check is to compare pH or the water-of-crystallisation behaviour on heating (washing soda loses water, baking soda decomposes).
25. X is basic; Y is acidic.
26. This cannot happen with one gas: CO₂ (which makes lime water milky) does not burn as H₂ does (pop). The milkiness from CO₂ and the pop from H₂ are different gases.
27. Hydrogen is flammable; collect it over water or use a safety container — uncontrolled mixing of H₂ with air can cause a loud pop or explosion risk. Careful collection technique prevents accidents.
28. Add lime (CaO or Ca(OH)₂). Equation with soil acidity — lime reacts with acid: Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O (improving pH toward 6–7).
29. Initially CO₂ forms insoluble CaCO₃ (milky): Ca(OH)₂ + CO₂ → CaCO₃ + H₂O. Excess CO₂ converts the solid to soluble Ca(HCO₃)₂: CaCO₃ + H₂O + CO₂ → Ca(HCO₃)₂, so the milkiness clears.
30. Washing soda is sodium carbonate decahydrate, a base with pH > 7, which is why the customer's fingers feel slippery and taste bitter (typical base properties). It should be rinsed off thoroughly after cleaning utensils.