Acids, Bases And Salts

Introduction

Three clear liquids can look identical and behave very differently. One changes blue litmus to red; another reverses that change; a third may do neither. Chemistry turns these observations into an explanation about particles in water. The aim of this chapter is to connect what you observe, which ions are involved, and which reaction accounts for the change.

This guide follows Class 10 NCERT Science, Acids, Bases and Salts, Chapter 2, in the 2026–27 reprint. It covers indicators, reactions, ions in water, pH, salts and the preparation and uses of the named everyday chemicals. The current CBSE curriculum lists these topics and does not require the logarithmic definition of pH. Check your school’s current assessment guidance and your own board’s prescribed text; this is not a claim of coverage across every board. NCERT chapter; CBSE Science curriculum.

The observations here are for reasoning, not unsupervised experiments. Never identify an unknown substance by tasting, touching or directly smelling it. Do not mix household cleaners or handle acids, alkalis, flames or electrical apparatus for these activities. Use the virtual evidence cards with this guide; practical chemistry belongs in a properly supervised laboratory.

Acids

In the aqueous model used here, an acid increases the concentration of hydrogen ions in water. We write these as H⁺(aq); a more explicit representation is the hydronium ion, H₃O⁺. The proton is associated with water molecules, rather than floating as an isolated bare particle in the solution.

For hydrogen chloride dissolving in water:
HCl + H₂O → H₃O⁺ + Cl⁻.

Hydrochloric acid is its aqueous solution. The familiar acidic properties arise in that watery environment. Dry hydrogen chloride does not turn dry blue litmus red; moisture allows the acidic response. This does not make the gas harmless.

Containing hydrogen is not enough to make a substance an acid in this model. Glucose and ethanol contain hydrogen but do not supply appreciable H₃O⁺ to water as hydrochloric acid does. Their ordinary aqueous solutions are poor conductors compared with acid solutions of similar suitable concentrations.

Vinegar contains ethanoic, or acetic, acid; curd contains lactic acid; citrus fruits contain citric acid; tamarind contains tartaric acid. These examples connect formulas with familiar materials, but foods contain mixtures and their pH is not one fixed number for every sample. “Natural” and “acidic” are not safety classifications.

Bases

A base can neutralise an acid. The water-soluble bases called alkalis increase hydroxide-ion concentration in water. Sodium hydroxide dissolves to give sodium ions and hydroxide ions, OH⁻:

NaOH(s) → Na⁺(aq) + OH⁻(aq), with water as the solvent.

OH⁻ is hydroxide, not a hydrogen ion. One formula unit of magnesium hydroxide contains two hydroxide groups; the small amount that dissolves is represented by Mg(OH)₂(s) ⇌ Mg²⁺(aq) + 2OH⁻(aq). The coefficient 2 preserves both atoms and charge.

Not every base contains OH in its formula. Ammonia reacts with water to form ammonium and hydroxide ions: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. This is a useful boundary on the simple “bases supply OH⁻” description. You do not need advanced acid–base theories to use the chapter’s examples. OpenStax explains the aqueous particle model and its extension.

Acidic and alkaline solutions conduct because they contain mobile ions. Both positive and negative ions contribute. A salt solution can also conduct, so conductivity alone does not prove that a solution is acidic. Pure water has a very low concentration of ions and is a poor conductor; it is not perfectly devoid of ions.

A solution lights a conductivity tester. Has the test proved it is an acid?

No. It indicates mobile charged particles under the test conditions. An alkali or dissolved salt can also conduct. Use a suitable indicator or pH measurement to investigate acidity. Do not build or test this circuit with household chemicals.

Indicators: evidence, not identity labels

An indicator changes in a characteristic way over a pH range. It can support a conclusion about a solution, but cannot identify the exact chemical on its own.

• Litmus: an acidic solution turns blue litmus red; an alkaline solution turns red litmus blue. Red paper staying red alone does not distinguish an acid from a neutral solution.
• Turmeric: normally yellow; becomes reddish-brown in a sufficiently alkaline medium. It is less useful for distinguishing an acid from a neutral sample.
• Phenolphthalein: colourless in acidic and neutral solutions, and pink in its alkaline indicator range. Colourless does not automatically mean acidic.
• Methyl orange: red at the acidic end of its transition range, orange during transition and yellow above it. Its yellow colour is not exclusive to alkaline solutions.
• Universal indicator: a mixture whose colour is compared with its supplied scale to estimate pH. Colour charts are approximate and depend on the indicator.

Litmus dyes come from lichens. A lichen association includes a fungus and photosynthetic algae or cyanobacteria; describing every lichen simply as “fungus plus bacteria” is incomplete. OpenStax’s lichen account explains the partners.

Olfactory indicators such as onion or vanilla show an odour-related change in particular acidic or basic conditions. They are examples of another observation, not a reason to sniff an unknown chemical. The virtual activity uses reported indicator results instead.

An unknown solution leaves red litmus unchanged and phenolphthalein colourless. Is it definitely an acid?

No. A neutral solution can give both results. A blue-litmus test or an appropriate pH measurement would supply additional evidence. Two observations that are compatible with several explanations do not uniquely identify one.

Reaction patterns: explain the products

A word equation is a useful prediction, but check that it applies to the reactants and conditions. A balanced formula equation then checks conservation of atoms. The symbols (s), (l), (g) and (aq) mean solid, liquid, gas and dissolved in water.

Reactive metal with a suitable dilute acid

Many reactive metals react with dilute hydrochloric or sulphuric acid to form a salt and hydrogen. For example:
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g).

Zinc sulphate remains dissolved in this example; hydrogen is the gas. A supervised laboratory hydrogen test gives a characteristic pop on ignition. Bubbles by themselves do not identify a gas, and this is not an activity to try at home. Not every metal reacts this way, and the same product rule does not apply to every acid.

Some metals also react with strong alkalis. NCERT gives the simplified sodium-zincate representation Zn + 2NaOH → Na₂ZnO₂ + H₂ for zinc reacting with warmed sodium hydroxide. This is a specific reaction, not “every base plus every metal”. The formula is the chapter’s simplified salt representation; detailed aqueous zinc complexes are beyond this guide.

Carbonates and hydrogencarbonates with acids

Here the gas is carbon dioxide. The products are a salt, water and CO₂:

Na₂CO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g).
NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g).

Notice why the acid coefficient changes. Carbonate, CO₃²⁻, needs two H⁺ to become water and carbon dioxide; hydrogencarbonate, HCO₃⁻, already carries one hydrogen and needs one more. Sodium and chloride balance the salt alongside this change.

In the laboratory, passing CO₂ into limewater can produce a white calcium carbonate precipitate:
Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l).

The cloudiness is a new solid, not the gas itself turning white. With excess CO₂, the solid can dissolve as soluble calcium hydrogencarbonate forms:
CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq).

Why is “carbonate plus acid makes hydrogen gas” the wrong pattern?

The carbonate provides carbon and oxygen that leave as CO₂. In Na₂CO₃ + 2HCl, both acid hydrogens enter H₂O. The reactive-metal pattern cannot be transferred to a carbonate merely because both reactions may bubble.

Neutralisation

For an acid reacting with a hydroxide base, the chapter’s general pattern is salt plus water:
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l).

At the particle level, H⁺(aq) + OH⁻(aq) → H₂O(l), or explicitly H₃O⁺ + OH⁻ → 2H₂O. Sodium and chloride ions remain in solution; they have not vanished. Neutralisation uses reacting amounts, not equal volumes by default.

If acid remains after the base is used up, the mixture can still be acidic. If excess base remains, it can be alkaline. Even a mixture with no excess of the original acid/base need not have pH 7 when the resulting salt reacts with water. “A neutralisation reaction happened” and “the final solution is neutral” are different statements.

In a virtual trial, 15 mL of one NaOH solution requires 10 mL of one HCl solution. How much of the same acid is needed for 45 mL of the same base?

Three times as much base solution contains three times the reacting amount, because its concentration is unchanged. It needs 3 × 10 = 30 mL acid. Equal volumes were not the condition; fixed concentration and the reaction ratio were.

Oxides connect with the same pattern

Copper(II) oxide behaves as a basic oxide:
CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l).
The black solid dissolves, and copper(II) chloride gives the solution a blue-green appearance under the chapter’s conditions.

Carbon dioxide behaves as an acidic oxide, as shown by its reaction with calcium hydroxide above. Many metal oxides are basic and many non-metal oxides acidic, but these are trends, not universal definitions of everything containing metal or non-metal. Some oxides react with both acids and bases; others do neither readily.

pH, dilution and strength: three different ideas

At about 25°C, neutral water has equal concentrations of H₃O⁺ and OH⁻ and pH 7. An acidic aqueous solution has more H₃O⁺ than OH⁻ and pH below 7; an alkaline solution has more OH⁻ than H₃O⁺ and pH above 7. Alkaline solutions still contain hydrogen ions. They are less abundant than hydroxide ions.

The familiar school scale from 0 to 14 covers many dilute aqueous examples. It is not a universal limit, and the pH of neutrality changes with temperature. For the comparisons in this guide, assume dilute aqueous solutions near 25°C unless stated otherwise. OpenStax’s pH explanation supplies these limits.

Lower pH means higher hydronium-ion concentration. The scale is not an ordinary ruler: a difference of one pH unit corresponds approximately to a tenfold hydronium concentration difference in these dilute solutions. Thus pH 3 has about 100 times the H₃O⁺ concentration of pH 5. This comparison is an extension that does not require you to calculate logarithms.

• Concentration: how much solute is present in a given volume of solution.
• Strength: how extensively an acid ionises or a base reacts/dissociates to supply the relevant ions in water. HCl is a strong acid; ethanoic acid a weak acid.
• pH: a measure of the resulting solution’s acidity, influenced by both identity and concentration.

A strong acid can be dilute. A weak acid can be concentrated. At equal suitable concentrations, HCl supplies more H₃O⁺ than ethanoic acid because it ionises much more extensively. Comparing two unlabelled pH readings alone cannot tell you which acid is intrinsically stronger. OpenStax distinguishes acid strength from concentration.

Dilution spreads the solution through more water

Diluting an ordinary acid solution reduces H₃O⁺ concentration and moves its pH towards neutrality. Diluting an alkali reduces OH⁻ concentration and also moves its pH towards neutrality. Dilution alone does not turn an acid into an alkali. For weak acids, additional ionisation can occur during dilution, so do not assume the fraction ionised stays unchanged.

Mixing concentrated acids with water, or dissolving alkalis such as sodium hydroxide, can release substantial heat. The laboratory rule is to add acid slowly to water under proper supervision, with suitable protection and stirring. Pouring water into concentrated acid can cause intense local heating and splashing. This is a reason for the rule, not a home procedure.

Solution A has pH 4 and solution B pH 6. What can you conclude, and what remains unknown?

At the same temperature in the dilute model, both are acidic and A has about 100 times B’s hydronium concentration. You cannot identify the chemicals, their total acid concentrations or intrinsic strengths from those two pH readings alone.

Why pH matters in living systems

Different biological locations need different conditions. The stomach uses acid and enzymes during digestion; this does not mean the whole body has the stomach’s pH, or that one pH range describes every tissue and fluid. NIDDK explains the stomach’s digestive role. Some antacid ingredients react with acid, which is a chemistry application rather than a reason to self-prescribe household chemicals.

Acids made by mouth bacteria from sugars can remove minerals from tooth enamel. NCERT uses approximately pH 5.5 to illustrate the risk of enamel demineralisation; actual conditions and mineral availability matter. Fluoride, saliva and cleaning also matter, so cavity prevention cannot be reduced to “all toothpaste is basic”. NIDCR describes the balance between mineral loss and repair.

Plants have suitable soil-pH ranges that vary with species. A soil result is evidence to interpret, not a reason to add lime automatically. Acid deposition can affect soils and aquatic habitats. Unpolluted rain is already mildly acidic, near pH 5.6 because of dissolved CO₂; sulphur and nitrogen oxides can make precipitation substantially more acidic. The effect on a lake also depends on its buffering surroundings. US EPA explains acid rain.

The textbook uses animal and plant stings as familiar acid examples. Do not extend a simple neutralisation model into a treatment rule: a sting’s effects are not explained by pH alone. NHS guidance specifically advises against bicarbonate-of-soda home remedies for bites and stings. The chemistry model has a boundary.

Salts: a family name does not fix pH

A salt contains positive and negative ions in an electrically neutral ratio. Sodium chloride, NaCl, and sodium sulphate, Na₂SO₄, share Na⁺ and belong to the sodium family. NaCl and KCl share chloride and belong to the chloride family. The subscript 2 in Na₂SO₄ balances the sulphate ion’s 2− charge with two 1+ sodium ions.

When discussing a salt’s pH, we mean the pH of its aqueous solution, not a pH number attached to a dry crystal.

• A typical salt of a strong acid and strong base, such as NaCl, gives a near-neutral solution in pure water.
• A salt such as NH₄Cl, associated with a strong acid and weak base, gives an acidic solution.
• A salt such as sodium ethanoate, associated with a weak acid and strong base, gives an alkaline solution.

The reason is whether its ions react with water to change the balance of H₃O⁺ and OH⁻. You need not calculate salt hydrolysis equilibria here, but the explanation prevents the false rule “all salts are neutral”. OpenStax explains the ion–water reactions.

NaCl and Na₂CO₃ are both sodium salts. Must their solutions have the same pH?

No. Their anions differ. Chloride has negligible acid–base effect in this model, while carbonate reacts with water to produce an alkaline solution. A shared sodium ion is not the whole chemical story.

Common salt as an industrial starting point

Sodium chloride is obtained from saline water and rock-salt deposits. Its importance goes beyond seasoning: the sodium and chlorine it supplies enter several manufacturing routes. The process map with this guide separates products from uses, so you can explain how the named chemicals connect.

Sodium hydroxide: the chlor-alkali process

Industrial electrolysis of brine, an aqueous sodium chloride solution, produces sodium hydroxide, chlorine and hydrogen:
2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + Cl₂(g) + H₂(g).

Chlorine forms at the anode, hydrogen at the cathode, and sodium hydroxide solution is obtained from the cathode side in a suitable separated cell. Product separation is part of the industrial design. The name joins chlorine with alkali.

NaOH is used in soap and paper manufacture and other industrial processes; chlorine is used in water treatment and chemical manufacture; hydrogen can be a feedstock for ammonia. This equation describes an industrial process, not a safe saltwater experiment.

Bleaching powder: read the current formula carefully

Chlorine reacts with dry slaked lime, Ca(OH)₂. The 2026–27 NCERT chapter represents bleaching powder using Ca(ClO)₂, calcium hypochlorite, while noting that its actual composition is complex. It gives:

2Ca(OH)₂ + 2Cl₂ → Ca(ClO)₂ + CaCl₂ + 2H₂O.

Older notes commonly use the simplified formula CaOCl₂. These expressions should not be silently swapped: Ca(ClO)₂ contains two oxygen atoms per formula unit, while CaOCl₂ contains one. Use the representation and balanced equation in your prescribed edition, and recognise that commercial bleaching powder is not adequately described as one perfectly pure simple compound. The current-source wording is visible on NCERT page 30.

Bleaching powder is used for bleaching textiles and pulp, as an oxidising agent and in controlled disinfection. A disinfectant’s use requires the correct product and dose; the school equation is not drinking-water treatment or cleaner-mixing guidance.

Baking soda and washing soda: one name change, different chemistry

Baking soda is sodium hydrogencarbonate, NaHCO₃, also called sodium bicarbonate. NCERT represents its manufacture using:

NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃.

On heating:
2NaHCO₃ → Na₂CO₃ + H₂O + CO₂.
The carbon dioxide is useful in raising food mixtures. Baking powder is a mixture containing baking soda and an edible acid, often with a drying/stabilising ingredient; it is not another name for pure baking soda. NCERT uses tartaric acid as the acid example.

The gas-producing acid reaction can be written without an undefined “sodium salt” on one side:
HCO₃⁻ + H⁺ → CO₂ + H₂O.
The accompanying ions form the appropriate salt in the complete equation. Baking soda also appears in some antacids and soda-acid fire-extinguisher designs. These are uses to explain, not permission to make a medicine or pressurised extinguisher.

Washing soda is sodium carbonate decahydrate, Na₂CO₃·10H₂O. Crystallising sodium carbonate with water gives the hydrated salt, represented by Na₂CO₃ + 10H₂O → Na₂CO₃·10H₂O. The solid incorporates water in a fixed formula ratio.

Sodium carbonate is used in glass, soap and paper industries and to make other sodium compounds such as borax. Its carbonate ions can remove permanent hardness by precipitating calcium or magnesium compounds; it is also used in cleaning. Baking soda and washing soda have different formulas and properties and should not be treated as interchangeable products.

Check the heating equation: why must baking soda have coefficient 2?

The sodium carbonate product contains two sodium atoms, so two NaHCO₃ units are required. The totals then balance: Na₂, H₂, C₂ and O₆ on each side. One carbon enters Na₂CO₃ and the other CO₂; both hydrogens enter water.

Water of crystallisation and Plaster of Paris

A crystal can look dry while containing water as part of its structure. Water of crystallisation is the fixed ratio of water molecules associated with each formula unit in a crystalline hydrate. It is different from an arbitrary film of liquid on the outside.

Hydrated copper(II) sulphate is CuSO₄·5H₂O. Under appropriate laboratory heating it loses its water of crystallisation and changes from blue towards the white anhydrous form. Rehydration restores the hydrated blue material. This is a teacher-supervised observation; colour change alone in an unknown substance does not identify copper sulphate.

Gypsum is CaSO₄·2H₂O. Controlled heating, described by NCERT at 373 K, converts it to calcium sulphate hemihydrate, Plaster of Paris, CaSO₄·½H₂O:

CaSO₄·2H₂O → CaSO₄·½H₂O + 1½H₂O.

The fraction records a ratio, not a water molecule cut in half. For two CaSO₄ units in the hemihydrate there is one H₂O molecule. Rewrite the change with whole-number ratios if that makes it easier to check:
2(CaSO₄·2H₂O) → 2(CaSO₄·½H₂O) + 3H₂O.

Adding water allows the reverse hydration and a hard gypsum mass forms:
CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O.

This setting process explains plaster’s uses in moulds, decoration and professionally applied casts. It also explains why the powder needs moisture-resistant storage. Setting can release heat, so do not make casts around a body part or carry out unsupervised mixing.

Two formula units of gypsum become two formula units of the hemihydrate. How many water molecules are released in the formula ratio?

The gypsum side has 2 × 2 = 4 water molecules. The hemihydrate retains 2 × ½ = 1. Therefore 3 are released. This is a whole-number atom-balanced description of the fractional formula.

Discover it: become an evidence detective

Use the virtual cards accompanying this guide. First infer what indicator evidence allows you to say. Then compare pH readings and ask what is still unknown. Finally follow sodium chloride through a product map and balance the two different bleaching-powder representations without pretending they are the same formula.

Make your own fourth evidence card on paper: choose an observation that is consistent with both an acidic and a neutral solution. Ask a friend what further evidence would distinguish them. An effective investigation asks for the next useful test, rather than guessing a chemical name from one clue.

Checks worth taking into the next chapter

• Match observations with ions and a balanced reaction; bubbles or one indicator colour do not identify a substance.
• H⁺/H₃O⁺ and OH⁻ are different ions. Balance charge as well as atoms.
• Strength, concentration and pH answer different questions.
• Neutralisation depends on reacting amounts; its occurrence does not guarantee final pH 7.
• A salt solution can be acidic, alkaline or neutral because its ions interact differently with water.
• Recognise formulas: NaHCO₃, Na₂CO₃·10H₂O, CaSO₄·2H₂O and CaSO₄·½H₂O describe different materials.
• Use the current prescribed-edition bleaching-powder representation explicitly, and keep chemical learning separate from household treatment or mixing advice.

Sources

The coverage and current-edition formulas were checked against the complete NCERT Chapter 2, 2026–27 reprint and CBSE Science curriculum. Explanations, evidence cards and numerical checks are original; textbook laboratory descriptions have not been turned into home instructions.

Additional primary-publisher checks: OpenStax on aqueous acids and bases, pH, strength and salts in water. The health and environmental examples link directly to NIDDK, NIDCR, NHS and US EPA guidance above.