Acids, Bases and Salts

1. What This Chapter Covers

From class VII you already know that acids are sour and turn blue litmus red, and that bases are soapy to touch and turn red litmus blue. This chapter asks what is actually happening underneath that.

It is the longest chapter in the volume — 9 periods in July, textbook pages 70 to 122 — and it is built almost entirely out of sixteen activities. The theory is assembled from what those activities show, so the activities are worth following in order rather than skipping to the conclusions.

The chapter's own opening question is a practical one: a person suffering from acidity takes an antacid tonic or tablet — what chemical reaction takes place? Section 3.6.2 answers it, but only after eight sections of groundwork.

Indicators, before anything else

Many natural materials contain dye molecules that are themselves weak acids or bases, and these serve as acid-base indicators: litmus, extract of red cabbage, turmeric solution, and extracts of the coloured petals of some flowers. Alongside them are synthetic indicators such as methyl orange and phenolphthalein.

Litmus itself is a dye extracted from lichen, a plant belonging to the division Thallophyta. In a neutral solution its colour is purple. The coloured petals of Hydrangea, Petunia and Geranium also work as indicators.

2. Chemical Properties of Acids and Bases (Textbook 3.1)

Activity 1 — the indicator table

Collect dilute solutions of hydrochloric acid, sulphuric acid, nitric acid, acetic acid, sodium hydroxide, calcium hydroxide, magnesium hydroxide, ammonium hydroxide and potassium hydroxide.

Take four watch glasses and put one drop of the first solution in each. Then dip blue litmus in the first, red litmus in the second, add a drop of methyl orange to the third, and a drop of phenolphthalein to the fourth. Repeat for every solution and record the colour changes in Table 1.

The point of running all nine substances through all four indicators is that the pattern, not any single result, identifies the sample as acidic or basic.

Activity 2 — olfactory indicators

Some substances change their odour in acidic or basic media. These are olfactory indicators.

Put finely chopped onions in a plastic bag with some clean cloth, tie it tightly, and leave it overnight in the fridge. Check the odour of the cloth strips. Put a few drops of dilute HCl on one strip and dilute NaOH on another, rinse both separately with distilled water, and check the odour again.

Repeat the test with clove oil and vanilla essence, adding a drop of each to test tubes of dilute HCl and dilute NaOH. From the observations, decide which of the three can serve as olfactory indicators.

The section closes with a question worth keeping: why are pickles and sour substances not stored in brass and copper vessels?

3. Reaction With Metals (Textbook 3.1.1)

Lab activity — acid with metal

Materials. Test tube, delivery tube, glass trough, candle, soap water, dilute HCl, zinc granules, cork.

Procedure. Take about 10 ml of dilute HCl in a test tube and add a few zinc granules. Watch the surface of the granules. Pass the evolved gas through soap water, and bring a burning candle near the gas-filled bubble.

The gas burns with a pop sound, which identifies it as hydrogen:

2HCl(aq) + Zn(s) -> ZnCl₂(aq) + H₂(g)

Repeating with H₂SO₄ and HNO₃ gives the same result, so the general rule is:

Acid + Metal -> Salt + Hydrogen

Activity 3 — base with metal

Place a few zinc granules in a test tube, add 10 ml of sodium hydroxide solution, and warm the contents. The gas evolved is again hydrogen, and the salt formed is sodium zincate:

2NaOH(aq) + Zn(s) -> Na₂ZnO₂(aq) + H₂(g)

So hydrogen is evolved when a base reacts with a metal too. This is a result students often miss, because it breaks the neat expectation that only acids attack metals.

4. Reaction With Carbonates and Hydrogen Carbonates (Textbook 3.1.2)

Activity 4

Label two test tubes A and B. Put about 0.5 g of sodium carbonate in A and 0.5 g of sodium hydrogen carbonate in B. Add about 2 ml of dilute HCl to both, and pass the gas produced through lime water:

Na₂CO₃(s) + 2HCl(aq) -> 2NaCl(aq) + H₂O(l) + CO₂(g)

NaHCO₃(s) + HCl(aq) -> NaCl(aq) + H₂O(l) + CO₂(g)

The gas turns lime water milky, because of a white precipitate of calcium carbonate:

Ca(OH)₂(aq) + CO₂(g) -> CaCO₃(s) + H₂O(l)

On passing excess carbon dioxide, the milkiness disappears again, since the precipitate redissolves:

CaCO₃(s) + H₂O(l) + CO₂(g) -> Ca(HCO₃)₂(aq)

The generalisations follow:

metal carbonate + acid -> salt + carbon dioxide + water

metal hydrogen carbonate + acid -> salt + carbon dioxide + water

5. Neutralisation (Textbook 3.2)

Activity 5

Take about 2 ml of dilute NaOH in a test tube and add one drop of phenolphthalein. The solution turns pink. Now add dilute HCl drop by drop — the pink colour disappears, because the NaOH has been completely reacted with the HCl and the effect of the base is nullified.

Add one or two drops of NaOH again and the pink colour reappears, because the solution has become basic once more.

NaOH(aq) + HCl(aq) -> NaCl(aq) + H₂O(l)

The reaction of an acid with a base to give a salt and water is a neutralisation reaction: Base + Acid -> Salt + Water

Activity 6 — acids with metal oxides (3.2.1)

Take a small amount of copper oxide in a beaker and slowly add dilute hydrochloric acid while stirring. The black copper oxide dissolves and the solution turns bluish-green, because copper(II) chloride has formed:

CuO(s) + 2HCl(aq) -> CuCl₂(aq) + H₂O(l)

Metal oxide + Acid -> Salt + Water

Comparing activities 5 and 6, both give a salt and water. So metal oxides are basic in nature, like metal hydroxides.

Bases with non-metal oxides (3.2.2)

The reaction of carbon dioxide with lime water from activity 4 has the same shape: a base reacting to give a salt and water. Therefore carbon dioxide, a non-metal oxide, is acidic in nature — and in general all non-metal oxides are acidic.

6. What Do All Acids Have in Common? (Textbook 3.3)

Acids generate hydrogen gas with metals, so hydrogen looks like the common element. But is every hydrogen-containing compound an acid?

Activity 7 — the conductivity test

Prepare solutions of glucose, alcohol, hydrochloric acid and sulphuric acid. Connect two graphite rods in a 100 ml beaker to a 230 volt AC supply through a bulb, as in figure 3. Pour in dilute HCl and switch on.

The bulb glows in the acid solutions but not in glucose or alcohol. A glowing bulb means current is flowing, which means ions are present and moving.

The positive ion in HCl solution is H⁺. So acids produce hydrogen ions in solution, and these are responsible for acidic properties. Glucose and alcohol contain hydrogen but produce no H⁺ ions, which is why they are not acids.

Repeating the activity with alkalis such as sodium hydroxide and calcium hydroxide also makes the bulb glow, because bases give the negative hydroxide ion, OH⁻.

7. Do Acids Produce Ions Only in Water? (Textbook 3.4)

Activity 8 — dry gas versus wet gas

Take about 1.0 g of solid NaCl in a clean, dry test tube and add some concentrated sulphuric acid:

2NaCl(s) + H₂SO₄(l) -> 2HCl(g) + Na₂SO₄(s)

Test the evolved gas with dry blue litmus paper and then with wet blue litmus paper.

The dry paper does not change colour. The wet paper turns red. So dry HCl gas is not an acid; the aqueous solution is. Dissociation into ions does not occur without water.

HCl + H₂O -> H₃O⁺ + Cl⁻

Hydrogen ions cannot exist as bare ions. They associate with water molecules, each H⁺ attached by four to six of them, and are represented as the hydronium ion:

H⁺ + H₂O -> H₃O⁺

The book adds two practical notes. In very humid weather, pass the gas through a guard tube containing calcium chloride to dry it. And always use a test tube holder with concentrated solutions — working with bare hands is dangerous.

Bases in water

Dissolving a base in water produces hydroxide ions:

NaOH(s) -> Na⁺(aq) + OH⁻(aq)

KOH(s) -> K⁺(aq) + OH⁻(aq)

Mg(OH)₂(s) -> Mg²⁺(aq) + 2OH⁻(aq)

Bases that are soluble in water are called alkalis. Not all bases dissolve; Be(OH)₂ is only slightly soluble. So every alkali is a base, but not every base is an alkali.

8. Mixing With Water (Textbook 3.5)

Activity 9 and the safety rule

Take 10 ml of water in a test tube, add a few drops of concentrated H₂SO₄, swirl slowly, and touch the bottom of the tube. It is hot — the process is exothermic. Repeating with sodium hydroxide pellets gives the same result.

This produces the most important safety rule in the chapter:

The acid must always be added slowly to water, with constant stirring. If water is added to a concentrated acid, the heat generated may cause the mixture to splash out and cause burns, and the glass container may break from excessive local heating.

Mixing an acid or base with water decreases the concentration of H₃O⁺ or OH⁻ ions per unit volume. This is dilution, and the acid or base is said to be diluted.

9. Strength of Acids and Bases (Textbook 3.6)

Activity 10

Fill beaker A with dilute acetic acid and beaker B with dilute hydrochloric acid of the same concentration, and pass current through each as in activity 7.

The bulb glows brightly in HCl but only dimly in acetic acid. More ions means more H₃O⁺, so HCl is a strong acid and acetic acid is a weak acid.

The distinction matters: strength is about the degree of dissociation at a given concentration, not about how much acid you have.

The pH scale (3.6.1)

A scale for measuring hydrogen ion concentration is the pH scale. The 'p' stands for Potenz, the German word for power.

• pH of a neutral solution is 7.
• Values less than 7 are acidic.
• Values above 7 are basic — as pH rises from 7 to 14, H₃O⁺ concentration falls and OH⁻ concentration rises.

The scale runs 0 to 14 because the concentration of H⁺ in most solutions falls between 1 M, which is pH 0, and 10⁻¹⁴ M, which is pH 14. Sorensen introduced the pH concept to avoid writing the negative powers of hydrogen ion concentration, so it is properly restricted to solutions below 1 molar.

The pH scale as the chapter uses it. The two thresholds worth memorising are 5.5 for tooth decay and 5.6 for acid rain — close together, but about different things.

Activity 11 — the universal indicator

The universal indicator is a mixture of several indicators, showing different colours at different hydrogen ion concentrations.

Test the pH of seventeen listed substances with pH paper and record the colour, the approximate pH and the nature of each. The list is deliberately mixed — laboratory reagents such as HCl, CH₃COOH, NH₄Cl, CH₃COONa, NaHCO₃, Na₂CO₃ and NaOH sit alongside distilled water, lemon juice, carrot juice, coffee, tomato juice, tap water, banana juice, a colourless aerated drink, and saliva before and after a meal.

That last pair is the one to notice. It sets up the tooth-decay discussion directly.

10. Why pH Matters in Everyday Life (Textbook 3.6.2)

Plants and animals are pH sensitive. Living organisms survive only within a narrow range of pH. When the pH of rain water is less than 5.6 it is called acid rain, and when acid rain flows into rivers it lowers the river's pH, making survival of aquatic life difficult.

Tooth decay. Decay starts when the pH of the mouth falls below 5.5. Tooth enamel, made of calcium phosphate, is the hardest substance in the body and does not dissolve in water — but it is corroded below pH 5.5. Bacteria in the mouth produce acids by degrading sugar and food particles left behind. Cleaning the mouth after eating is the best prevention, and toothpastes, being generally basic, neutralise the excess acid.

Digestion. The stomach produces hydrochloric acid, which helps digestion without harming the stomach. During indigestion it produces too much, causing pain and irritation. Antacids — bases such as magnesium hydroxide, milk of magnesia — neutralise the excess. This answers the question the chapter opened with.

Soil pH. Plants need a specific pH range. In Activity 13, put about 2 g of soil in a test tube, add 5 ml of water, shake, filter, and test the filtrate with universal indicator paper. The follow-up question is when a farmer would treat his fields with quicklime or calcium carbonate.

Chemical warfare in nature (3.6.3)

A bee sting leaves an acid that causes pain and irritation; a mild base such as baking soda on the stung area gives relief.

The stinging hairs on nettle leaves inject methanoic acid, also called formic acid, causing a burning pain. The traditional remedy is to rub the area with the leaf of the dock plant, which is basic in nature and often grows beside the nettle in the wild.

11. Salts (Textbook 3.7)

Salts are ionic compounds produced by the neutralisation of an acid with a base, and they are electrically neutral.

Families and pH (3.7.1, 3.7.2)

Salts sharing the same positive or negative radical belong to a family. NaCl and Na₂SO₄ are both sodium salts; NaCl and KCl are both chloride salts, so one salt can belong to two families.

Activity 15 asks you to dissolve various salts and test them. The pattern that emerges:

Common salt (3.7.3, 3.7.4)

Sodium chloride is the most common salt, also called table salt. Sea water contains many dissolved salts with sodium chloride predominant. Solid deposits also exist, often brown from impurities, called rock salt — formed when seas of bygone ages dried up, and mined like coal.

The chapter pauses on history here. The British tax on salt, levied on poor and rich alike, helped unite people in the freedom struggle — Mahatma Gandhi's Dandi March and the salt satyagraha.

Common salt is the raw material for four further chemicals, and the next four sections take them in turn.

Sodium hydroxide — the chloro-alkali process (3.7.5)

Passing electricity through brine, an aqueous solution of sodium chloride, decomposes it:

2NaCl(aq) + 2H₂O(l) -> 2NaOH(aq) + Cl₂(g) + H₂(g)

It is called the chloro-alkali process after its products: chloro for chlorine, alkali for sodium hydroxide. Chlorine is given off at the anode, hydrogen at the cathode, and sodium hydroxide forms near the cathode. All three products are useful — hydrogen for fuels, margarine and ammonia for fertilisers; chlorine for water treatment, PVC, disinfectants and pesticides; sodium hydroxide for degreasing metals, soaps and detergents, paper making and artificial fibres.

Bleaching powder (3.7.6, 3.7.7)

Chlorine from the chloro-alkali process acts on dry slaked lime to give bleaching powder, represented as CaOCl₂ although its actual composition is quite complex:

Ca(OH)₂ + Cl₂ -> CaOCl₂ + H₂O

Its four listed uses are bleaching cotton and linen in textiles, wood pulp in paper and washed clothes in laundry; as an oxidising agent in chemical industries; for disinfecting drinking water; and as a reagent in preparing chloroform.

Baking soda (3.7.8)

Sodium hydrogen carbonate, NaHCO₃, prepared as:

NaCl + H₂O + CO₂ + NH₃ -> NH₄Cl + NaHCO₃

It is a mild, non-corrosive base. On heating during cooking:

2NaHCO₃ -> Na₂CO₃ + H₂O + CO₂

Its uses: baking powder, which is mainly NaHCO₃ together with Ca(H₂PO₄)₂ and starch, where the CO₂ produced rises through the dough as bubbles and makes cake and bread smooth and spongy; as an ingredient in antacids; as soda-acid in fire extinguishers; and as a mild antiseptic.

Washing soda (3.7.9)

Heating baking soda gives sodium carbonate, and recrystallisation of it gives washing soda:

Na₂CO₃ + 10H₂O -> Na₂CO₃·10H₂O

It is a basic salt, used in glass, soap and paper industries, in manufacturing sodium compounds such as borax, as a domestic cleaning agent, and for removing permanent hardness of water.

12. Water of Crystallisation (Textbook 3.8)

Activity 16

Heat a few crystals of copper sulphate in a dry test tube. The blue colour fades to white, and water droplets appear on the sides of the tube. Add two or three drops of water to the heated sample and the blue colour returns.

Water of crystallisation is the fixed number of water molecules present in one formula unit of a salt.

Copper sulphate has five, so hydrated copper sulphate is CuSO₄·5H₂O. This also answers the earlier question about Na₂CO₃·10H₂O — the ten water molecules are chemically bound, so the crystals are not wet.

Plaster of Paris (3.8.1)

Gypsum is another hydrated salt, CaSO₄·2H₂O. Heating it carefully at 373 K drives off part of its water to give calcium sulphate hemihydrate:

CaSO₄·½H₂O, which is Plaster of Paris — the white powder doctors use to support fractured bones. On mixing with water it sets into a hard solid mass, because gypsum re-forms:

CaSO₄·½H₂O + 1½H₂O -> CaSO₄·2H₂O

The half water molecule looks impossible, and the book addresses it directly: it is written that way because two formula units of CaSO₄ share one molecule of water.

Plaster of Paris is also used for making toys, decorative materials and for smoothing surfaces. It follows that it must be stored in a moisture-proof container, since any water will set it.

Key words from the chapter

Indicator, acid, base, red litmus, blue litmus, phenolphthalein, methyl orange, salts, neutralisation, guard tube, hydronium ion, alkali, strong acid, strong base, universal indicator, pH scale, antacids, tooth decay, family of salts, common salt, bleaching powder, baking soda, washing soda, hydrated salt, water of crystallisation, Plaster of Paris.

13. Summary

Acid-base indicators are dyes, or mixtures of dyes, used to detect acids and bases. Besides litmus, methyl orange and phenolphthalein, the chapter introduces olfactory indicators such as onion, clove oil and vanilla essence, which change odour rather than colour.

The characteristic reactions are: acid with metal giving salt and hydrogen; base with metal also giving hydrogen, as with zinc and NaOH forming sodium zincate; metal carbonates and hydrogen carbonates with acid giving salt, carbon dioxide and water; acid with metal oxide giving salt and water, showing metal oxides to be basic; and base with non-metal oxide doing the same, showing non-metal oxides to be acidic. An acid and a base together give salt and water — neutralisation.

The conductivity activity shows acidity is due to H⁺, existing as the hydronium ion H₃O⁺, and basicity to OH⁻. Glucose and alcohol contain hydrogen but produce no ions, so they are not acids. The dry-versus-wet litmus test proves that water is required for dissociation: dry HCl gas is not an acid. Bases soluble in water are alkalis.

Dissolving an acid or base in water is highly exothermic, so acid is always added to water and never the reverse. Strength depends on the concentration of H₃O⁺ or OH⁻ produced, measured on the pH scale from 0 to 14, where 7 is neutral. The thresholds that matter in daily life are 5.5 for tooth decay and 5.6 for acid rain, along with the role of antacids in the stomach and pH testing of soil.

Salts are ionic compounds from neutralisation, grouped into families and classified as acidic, basic or neutral by the strengths of their parent acid and base. From common salt come sodium hydroxide by the chloro-alkali process, bleaching powder, baking soda and washing soda. Finally, water of crystallisation is the fixed number of water molecules in one formula unit of a salt, illustrated by CuSO₄·5H₂O and by gypsum, whose partial dehydration at 373 K gives Plaster of Paris.