Acids, Bases and Salts

In Class 7 you watched litmus flip from blue to red in lemon juice, and turmeric flip to red in soap solution, and you learned the rule: acids turn blue litmus red, bases turn red litmus blue. But a rule is not an explanation. Litmus paper cannot tell you why it changes colour, only that it does. This chapter asks the question Class 7 left open: what is actually happening, at a level too small to see, that makes an acid an acid and a base a base? The answer turns out to be a single, tiny, electrically charged particle, and once you understand it, almost everything else in this chapter, from the fizz of an antacid tablet to why a farmer adds lime to a field, falls into place as a consequence of that one idea.

Set up a simple circuit: two nails dipped into a beaker of liquid, wired through a battery, a switch and a bulb. Pour in dilute hydrochloric acid and switch on: the bulb glows. Try dilute sulphuric acid instead, and it glows again. But pour in a glucose solution, or an alcohol solution, both of which also contain hydrogen atoms, and the bulb stays dark. A glowing bulb means an electric current is flowing, and current can only flow through a liquid if charged particles, ions, are free to carry it. Acids conduct; glucose and alcohol do not, even though both contain hydrogen. The difference is what kind of hydrogen: in HCl, in HNO3, in H2SO4, in every acid, the hydrogen is present as a positively charged hydrogen ion, H+, and it is specifically this ion, free and mobile in solution, that carries the current and is responsible for acidic behaviour. Bases turn out to have their own mirror-image particle: repeat the same circuit test with sodium hydroxide or calcium hydroxide solution, and the bulb glows there too, this time because the base has released negatively charged hydroxide ions, OH-, into the water.

If free H+ ions are what makes an acid acidic, then dry hydrogen chloride gas, with no water anywhere near it, should behave exactly like hydrochloric acid. It does not: pass dry HCl gas over dry litmus paper and absolutely nothing happens, the paper stays whichever colour it started. Only once that same gas is bubbled into water does the litmus paper turn red. The gas is not acidic by itself; something about mixing with water switches its acidity on. What happens is that a water molecule pulls the H+ off the HCl molecule, and the two combine as a hydronium ion, H3O+: HCl + H2O -> H3O+ + Cl-. This is why an acid's H+ is more precisely written H+(aq), or H3O+, never a bare, standalone H+; the ion cannot exist floating free, it always needs water to carry it. Bases dissolve into their ions the same water-dependent way, for example NaOH(s) dissolves in water to give Na+(aq) and OH-(aq). A base that dissolves in water like this is specifically called an alkali; not every base does dissolve, so not every base is an alkali.

Drop a few pieces of zinc granules into dilute sulphuric acid and bubbles form immediately on the metal's surface. Collect that gas and hold a burning candle near it: it burns with a characteristic pop. That pop is the standard test for hydrogen gas, and the reaction is Zn + H2SO4 -> ZnSO4 + H2. This is not unique to zinc and sulphuric acid; in general, acid + metal -> salt + hydrogen. A second, equally distinctive reaction happens when an acid meets a carbonate: add dilute hydrochloric acid to sodium carbonate (or sodium hydrogencarbonate, baking soda) and the mixture fizzes visibly. Pass that gas through limewater, and the limewater turns milky. That milkiness is the standard test for carbon dioxide, and the reaction is metal carbonate + acid -> salt + water + carbon dioxide, for example Na2CO3 + 2HCl -> 2NaCl + H2O + CO2. Two acids, two completely different partners, and two completely different, easily testable gases, hydrogen with a pop, carbon dioxide turning limewater milky, both signalling that a real chemical reaction, not just a colour change, has taken place.

Add dilute hydrochloric acid to black copper oxide powder, stirring, and the mixture turns a distinct blue-green as it dissolves. That colour is copper(II) chloride, formed by Metal oxide + Acid -> Salt + Water, in this case CuO + 2HCl -> CuCl2 + H2O. Because metal oxides react with acids the same way a base does, chemists classify metal oxides as basic oxides. Now recall passing carbon dioxide through limewater: Ca(OH)2 + CO2 -> CaCO3 + H2O, a base reacting with a non-metal oxide the same way it would react with an acid. That symmetry is exactly why non-metal oxides, like carbon dioxide, are classified as acidic oxides. Put together, this gives a clean way to sort the two families of elements just by watching how their oxides behave: if an oxide acts basic (reacting with acids), the element is a metal; if it acts acidic (reacting with bases), the element is a non-metal.

Add a few drops of phenolphthalein to sodium hydroxide solution and it turns pink, phenolphthalein's colour in a base. Add dilute hydrochloric acid to that pink solution, drop by drop: the pink fades to colourless. Add a few more drops of sodium hydroxide, and the pink reappears. This is the same neutralisation idea from Class 7, acid and base cancelling out, but now it can be written far more precisely. The base supplied OH-(aq) ions and the acid supplied H+(aq) ions; the moment they meet, they combine directly into water, H+(aq) + OH-(aq) -> H2O(l). That single ionic equation is the real engine behind every acid-base neutralisation there is, regardless of which particular acid or which particular base is involved: NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l) is just one example of the general pattern Base + Acid -> Salt + Water, all of them driven by the same H+ + OH- -> H2O reaction underneath.

Class 7 could only sort a liquid into acidic, basic or neutral; it had no way to say how acidic. A universal indicator solves that: unlike litmus's single colour flip, it shifts smoothly through a whole spectrum of colours depending on exactly how many H+ ions are present, and that spectrum is mapped onto a numbered scale called pH, running from 0 to 14. The word pH comes from the German potenz, meaning power or strength; the higher the concentration of H+(aq) ions in a solution, the lower its pH. A pH below 7 is acidic, exactly 7 is neutral, and above 7 is basic, with strong acids sitting near 0 and strong alkalis near 14. Strength itself has a precise meaning too: acids that release more H+ ions per molecule are strong acids, and those that release fewer are weak acids, regardless of how concentrated the solution is; strength is about how completely a substance ionises, concentration is about how much of it is dissolved, and the two are genuinely different properties.

The human body runs within a narrow pH window, roughly 7.0 to 7.8, and drifting outside it makes normal metabolism impossible. Rain is naturally a little acidic, but once its pH falls below about 5.6 it is classed as acid rain, and acid rain draining into rivers can lower the water's pH enough to threaten fish and other aquatic life. Soil has its own preferred pH range depending on what is grown in it, which is why farmers periodically test soil pH before deciding whether it needs correcting. Inside your own stomach, hydrochloric acid helps digest food, but overeating can push the stomach into producing too much acid, causing the pain of indigestion; the standard remedy is a mild base, an antacid such as milk of magnesia (magnesium hydroxide), which neutralises the excess acid. Tooth enamel, the hardest substance in the body, starts to corrode once the pH inside the mouth drops below about 5.5, which happens when bacteria in the mouth break down leftover sugar into acid; brushing with a mildly basic toothpaste after eating helps neutralise that acid before it can do damage. Even insect and plant defences run on the same chemistry: a bee sting injects venom that is acidic, so a mild base can ease it, while nettle leaves sting with methanoic acid, traditionally treated by rubbing the sap of the dock plant, itself mildly basic, on the affected patch.

It is tempting to assume every salt, being the product of an acid and a base cancelling out, must itself be neutral. Testing actual salt solutions with pH paper shows this is false. A salt made from a strong acid and a strong base, such as sodium chloride from HCl and NaOH, is indeed neutral, pH 7. But a salt made from a strong acid and a weak base turns out acidic, pH below 7, and a salt made from a strong base and a weak acid turns out basic, pH above 7. The parent acid and base do not simply vanish without a trace; something of their original strength carries over into the salt they form, and that leftover imbalance is exactly what shifts the salt's own pH away from neutral.

Ordinary sodium chloride, common salt, is not just a seasoning; it is a raw material for a whole branch of industrial chemistry. Pass electricity through a concentrated solution of sodium chloride in water (called brine), and it splits into three separately useful products at once: sodium hydroxide forms near the cathode, chlorine gas is given off at the anode, and hydrogen gas is released too. Because the two main products are chlorine and alkali (sodium hydroxide), this is called the chlor-alkali process. Take the chlorine produced this way and react it with dry slaked lime, and the result is bleaching powder, used to bleach textiles and paper pulp and to disinfect drinking water. Common salt is also the starting material for baking soda (sodium hydrogencarbonate, NaHCO3), used in cooking, in antacids, and in soda-acid fire extinguishers, and for washing soda (Na2CO3.10H2O), used in the glass, soap and paper industries and to soften hard water. Common salt has carried real historical weight in India too: it was central enough to daily life and to colonial-era taxation that Mahatma Gandhi's 1930 Dandi March, a defining act of the freedom struggle, was organised specifically around the right to make salt.

Blue copper sulphate crystals look, and feel, completely dry, yet heat them gently in a dry test tube and two things happen at once: water droplets appear on the cool part of the tube, and the crystals themselves turn white. Add a couple of drops of water back onto that white powder, and the blue returns. The crystals were never truly dry: each formula unit of copper sulphate has exactly five water molecules locked into its very structure, written as CuSO4.5H2O, and heating drives that water out, leaving anhydrous (water-free) white copper sulphate behind, which turns blue again the moment it is rehydrated. Water held this way, as a fixed part of a crystal's structure rather than just dampness on its surface, is called water of crystallisation, and copper sulphate is not unique: gypsum, CaSO4.2H2O, holds two such water molecules. Heat gypsum to about 373 K and it loses some of that water to become CaSO4.½H2O, calcium sulphate hemihydrate, better known as Plaster of Paris, the same white powder doctors use to make supportive casts for fractured bones. Mix Plaster of Paris with water and it sets hard, because the reaction simply runs in reverse, turning it back into gypsum: CaSO4.½H2O + 1½H2O -> CaSO4.2H2O.

This chapter answered Class 7's open question, why litmus changes colour, with a single particle, H+, and its counterpart OH-. But a few things here have only been described in words, not measured: strong and weak acids were told apart qualitatively, and pH was defined as a scale without yet being calculated from anything. Class 11 completes that picture. There, strength becomes a genuine number, the ionization constant Ka for an acid and Kb for a base, pH becomes something you calculate from an actual H+ concentration rather than just read off a colour chart, and you meet buffer solutions, mixtures cleverly designed to resist changing pH even when acid or base is added to them, the same real chemistry your own blood relies on to keep its pH inside that narrow safe range mentioned earlier in this chapter.