Acids, Bases and Salts: The pH Scale, Indicators and Everyday Salts

Acids, Bases and Salts: Definition and Everyday Signs

Hydrogen Ions, Hydroxide Ions and Salts: Definitions

An acid is a substance that gives hydrogen ions, H⁺, when it dissolves in water. Sour taste is the everyday sign of one, and lemon, curd, tamarind and vinegar all taste sour because they contain acids. A base is a substance that gives hydroxide ions, OH⁻, in water, and its everyday signs are a bitter taste and a soapy, slippery feel. A salt is what forms when an acid and a base react and cancel each other, and sodium chloride, the salt on the table, is the best-known example.

Acids and bases also reach into environmental science through acid rain and ocean acidification, into agriculture through soil acidity and liming, and into human biology through blood pH, digestion and tooth decay.

Properties of Acids and Bases

Taste, Feel and Litmus Response of Acids and Bases

Acids taste sour and turn blue litmus paper red. Bases taste bitter, feel soapy or slippery, and turn red litmus paper blue. Substances that do neither, such as tap water, sugar solution and salt solution, are called neutral. Tasting is never a safe test in a laboratory, which is why indicators exist.

Taste alone can mislead. Bitter gourd (karela) tastes bitter but is not basic, so a bitter taste does not prove a base.

Bases that dissolve in water are called alkalis. Sodium hydroxide dissolves in water to give Na⁺ and OH⁻ ions, so it is an alkali. Alkalis feel soapy and taste bitter, but they are corrosive, so they must never be tasted or touched.

Arrhenius, Brønsted-Lowry and Lewis Concepts of Acids and Bases

Chemists have defined acids and bases in three steps, each wider than the one before.

• Arrhenius concept: An acid gives H⁺ ions in water and a base gives OH⁻ ions. A bare H⁺ cannot exist alone in water, so it joins a water molecule to form the hydronium ion, H₃O⁺; the two forms are used interchangeably. Svante Arrhenius proposed the definition of a base in 1884, and he received the 1903 Nobel Prize in Chemistry for his theory of electrolytic dissociation. The limit of the concept is that it works only in water and cannot explain why ammonia, which has no hydroxyl group, is a base.
• Brønsted-Lowry concept (1923): An acid is a proton donor and a base is a proton acceptor. Two species that differ by a single proton, such as NH₄⁺ and NH₃, form a conjugate acid-base pair. A strong acid has a weak conjugate base and the reverse.
• Lewis concept (1923): An acid accepts a pair of electrons and a base donates one. Boron trifluoride, BF₃, has no proton, yet it acts as an acid by accepting the lone pair of ammonia. Electron-deficient species such as AlCl₃ are Lewis acids, and H₂O, NH₃ and OH⁻ are Lewis bases.

Each later concept includes the earlier ones. For school-level work and most UPSC questions the Arrhenius picture is enough. Acids make H⁺ (as H₃O⁺) and bases make OH⁻.

Reactions of Acids and Bases: Metals, Carbonates and Oxides

Acids and bases react in a small number of predictable ways, and all of them follow from the ions the acid or base releases.

• Acid with a metal: The metal displaces hydrogen from the acid, giving a salt and hydrogen gas. Zinc granules in dilute sulphuric acid fizz, and when a burning candle is brought near a soap bubble filled with the gas, the gas burns, which shows that it is hydrogen.
• Base with a metal: Zinc also reacts with sodium hydroxide, giving hydrogen and a salt called sodium zincate. Such reactions are not possible with every metal.
• Acid with a metal carbonate or hydrogencarbonate: The products are a salt, carbon dioxide and water. The gas turns lime water milky, because it forms a white precipitate of calcium carbonate. Limestone, chalk and marble are all forms of calcium carbonate, which is why an acid makes limestone and marble fizz.
• Acid with a metal oxide: The products are a salt and water, so metal oxides behave like bases and are called basic oxides. Copper oxide dissolves in dilute hydrochloric acid to give a blue-green copper chloride solution.
• Base with a non-metal oxide: Carbon dioxide reacts with calcium hydroxide to give a salt and water, so non-metal oxides behave like acids and are called acidic oxides.

Acid solutions conduct electricity because their ions carry the current. Glucose and alcohol also contain hydrogen but do not release it as ions, so their solutions do not conduct and they are not acids. Dry hydrogen chloride gas does not change dry litmus paper at all, and only with water does it release H⁺, so acidic behaviour needs water.

Strong and Weak Acids, Alkalis and Safe Dilution

The strength of an acid depends on how many hydrogen ions it releases, not on how much of it is present. At the same concentration, hydrochloric acid releases far more H⁺ than acetic acid does, so hydrochloric acid is a strong acid and acetic acid is a weak one. Strong and weak bases differ in the same way in the OH⁻ ions they produce.

Being organic does not make an acid weak or strong. Formic acid (methanoic acid) and acetic acid are both organic acids, yet formic acid is about ten times stronger, since its ionisation constant is 1.8 × 10⁻⁴ against 1.74 × 10⁻⁵ for acetic acid.

Dissolving an acid or a base in water is strongly exothermic. A concentrated acid must therefore always be added slowly to water, with constant stirring. Water poured into concentrated acid can make the mixture splash out and cause burns, and the glass container may crack from the local heating. Water's large specific heat is what absorbs the heat released when the acid goes into a large amount of water.

Adding water also lowers the concentration of H₃O⁺ or OH⁻ ions per unit volume, a process called dilution.

Acid-Base Indicators

Natural Indicators: Litmus, Turmeric, Red Rose and Red Cabbage

An indicator is a dye, or a mixture of dyes, that shows by a change of colour whether a solution is acidic or basic. Many are natural.

Litmus is a purple dye extracted from lichen. A lichen is a partnership of a fungus and an alga. Litmus is purple when the solution is neutral, red in acid and blue in base, and the colour change happens over roughly the pH range 4.5 to 8.3. It is sold as a solution and as red and blue paper strips.

Turmeric paper behaves differently. Yellow turmeric turns reddish-brown in a base but stays yellow in acidic and neutral solutions, so it can detect a base but cannot tell an acid from a neutral liquid. That is why a curry stain on a white shirt turns reddish-brown when soap is rubbed on it and yellow again when the shirt is rinsed with plenty of water.

Red rose petal extract is red in acid and green in base. Red cabbage, beetroot, jamun, red hibiscus and the petals of Petunia and Geranium also work, because many coloured plant pigments (the anthocyanins) are red in acidic solutions and blue in basic ones. Hydrangea flowers change colour with the soil for a different reason. On acidic soil the plant can take up aluminium, which turns the flowers blue, while on basic soil it cannot, and the flowers stay pink.

Synthetic, Olfactory and Universal Indicators

Laboratories use synthetic indicators because their colour changes are sharp. Phenolphthalein is colourless in acid and pink in base. Methyl orange is red in acid and yellow in base, with orange in between. In an acid-base titration, an acid is added to a base drop by drop until the indicator changes colour at the point of neutralisation; phenolphthalein is often chosen for this work.

Some substances change odour rather than colour, and these are called olfactory indicators. Onion, vanilla essence and clove oil are tested in this way. Strips of cloth kept overnight with finely chopped onion take up its smell, and the smell is checked again after a few drops of a dilute acid go on one strip and a dilute base on the other; a substance whose smell changes in an acidic or a basic medium is an olfactory indicator.

A universal indicator is a mixture of several indicators that shows a different colour at each hydrogen ion concentration. Yellow to red colours show an acid, blue to violet an alkali and green a neutral solution. Paper impregnated with it is the pH paper used in schools. The universal indicator turns "acidic or basic" into a number on the pH scale.

The pH Scale

pH Scale: Meaning, Formula and the 0 to 14 Range

pH is a number that tells how acidic or basic a solution is by its hydrogen ion concentration. The p stands for potenz, German for power. Danish chemist Søren Sørensen introduced the idea in 1909 at the Carlsberg Laboratory. The modern definition is the negative logarithm to base 10 of the hydrogen ion concentration, written pH = −log[H⁺].

On the usual scale, 0 is very acidic, 7 is neutral and 14 is very alkaline. A solution with a pH below 7 is acidic, and a solution with a pH above 7 is basic. The higher the hydronium ion concentration, the lower the pH. Pure water at 25 °C has [H⁺] = 10⁻⁷ M, so its pH is 7. A 0.01 M (10⁻² M) solution of hydrochloric acid has a pH of 2. A sodium hydroxide solution with [OH⁻] = 10⁻⁴ M has [H₃O⁺] = 10⁻¹⁰ M and a pH of 10.

The range 0 to 14 is a convenience, not a physical limit. Very concentrated solutions go beyond it. A saturated sodium hydroxide solution is about 15, and concentrated hydrochloric acid is about −1.

Ionic Product of Water and the Relation of pH and pOH

Water ionises very slightly. One molecule gives up a proton to another, so pure water holds equal, tiny amounts of H₃O⁺ and OH⁻, each being 1.0 × 10⁻⁷ M at 25 °C. Their product is called the ionic product of water, K_w = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴ at 25 °C. Taking negative logarithms gives pH + pOH = 14. That relation is why the common scale ends at 14. In a solution where H⁺ is 10⁻¹⁴ M, OH⁻ is 1 M.

The balance of the two ions decides the nature of a solution. In an acidic solution [H₃O⁺] is greater than [OH⁻]; in a neutral solution the two are equal; in a basic solution [H₃O⁺] is less than [OH⁻]. K_w rises with temperature, so the pH of a neutral solution falls below 7 in hot water; the values given here are for 25 °C.

Two checks make the idea concrete. If solution A has a pH of 6 and solution B a pH of 8, A has the higher hydrogen ion concentration and is acidic while B is basic. Basic solutions do contain H⁺ ions too; they simply contain far more OH⁻ ions, so the solution is basic overall.

Measuring pH and the pH of Common Substances

pH is measured in two ways. Universal indicator solution or pH paper gives a colour that is matched by eye against a printed chart, so it reads only to about half a pH unit across the range 1 to 14. A pH meter measures the pH-dependent electrical potential of the solution through a glass electrode, which is calibrated against buffer solutions of known pH, and it reads to 0.001.

Gastric juice, at pH 1.2, holds more than a million times the hydrogen ions of blood at pH 7.4, because the two differ by more than six pH units. Milk of magnesia, at pH 10, is the mild base used as an antacid.

Why pH Matters for Life and the Environment

pH of Blood, Buffers, Digestion and Tooth Decay

Living organisms survive only within a narrow pH range, and the human body works between 7.0 and 7.8. Blood is slightly alkaline, at about 7.4, and the body holds it there with great precision. The pH of the blood plasma is regulated by chemical buffers, by the respiratory system and by the kidneys, and its normal range is narrow, about 7.35 to 7.45.

A buffer solution resists a change in pH when it is diluted or when a little acid or alkali is added. A mixture of acetic acid and sodium acetate buffers around pH 4.75, and a mixture of ammonium chloride and ammonium hydroxide buffers around pH 9.25. One buffer behind blood pH is the carbonic acid and bicarbonate pair, and when the pH is 7.4, the ratio of carbonic acid to bicarbonate is 1 to 20.

The stomach, by contrast, is meant to be strongly acidic. It produces hydrochloric acid, which helps digest food without harming the stomach. During indigestion the stomach makes too much acid, and the pain and irritation are relieved by antacids, which are bases that neutralise the excess acid. Marketed antacids contain salts of aluminium, calcium, magnesium or sodium, with milk of magnesia (magnesium hydroxide) and baking soda as familiar examples.

Tooth decay starts when the pH of the mouth falls below 5.5. Tooth enamel, made of calcium hydroxyapatite (a crystalline form of calcium phosphate), is the hardest substance in the body. It does not dissolve in water, but acid corrodes it once the mouth pH drops below 5.5. Bacteria in the mouth make acids by breaking down sugar and food particles left after eating, and those acids lower the pH. Toothpastes are generally basic, so cleaning the teeth after eating neutralises the excess acid and slows decay.

Acid Rain: Causes, Effects and the pH Limit

Rain is called acid rain when its pH falls below 5.6. Clean rain is already slightly acidic because of the carbonic acid formed from dissolved carbon dioxide, so the pH line for "acid rain" is set below the pH of clean rain, not at 7.

The current NCERT textbook puts the acid rain limit at pH 5.6, while some other accounts draw it at pH 5.

The chief acid-forming gases are sulphur dioxide and oxides of nitrogen, and they react with water in the air to form sulphuric and nitric acids. They come largely from burning fossil fuels in coal-fired power stations and motor vehicles. Volcanoes and lightning are natural sources, since lightning produces nitric acid. Tall smokestacks carry the gases far from their source, so the acid often falls a considerable distance downwind.

Acid rain does its harm in three places.

• Water bodies: When acid rain flows into rivers it lowers their pH and makes the survival of aquatic life difficult. Below pH 5 most fish eggs do not hatch, and below pH 4 adult fish die.
• Soil and forests: Acid rain washes away nutrients such as calcium and magnesium and releases toxic aluminium from soil.
• Buildings and monuments: It corrodes limestone and marble, and in India the marble of the Taj Mahal has been damaged in this way.

Controls act on the source. Flue-gas desulphurisation in coal power stations removes sulphur dioxide with lime or limestone, and emission control on vehicles cuts oxides of nitrogen.

India shows a different picture. Measurements published in 1984 showed that rain over unpolluted parts of India may have a pH between 6 and 7, and that a fall to acidic values may be restricted to highly industrialised cities such as Bombay (now Mumbai), where the rain pH was 4.5. The reason is that basic dust in the air neutralises the acid, since Indian soils are normally basic and the dust load is high.

Soil pH, Acid Soils and Liming

Soil pH is a master variable of the soil. It controls which forms of plant nutrients are available, and the best pH range for most plants is 5.5 to 7.5. Soils normally fall between pH 3 and 10, and a soil under warm, humid conditions turns acidic with time as water leaches the products of weathering out of it, while soils in dry climates are often neutral or alkaline. Ammonium fertilisers also acidify soil, because they release H⁺ as they are converted to nitrate.

To treat an acidic soil a farmer adds quicklime (calcium oxide), slaked lime (calcium hydroxide) or chalk (calcium carbonate), all bases. A basic soil is treated by adding organic matter such as manure and composted leaves, which release acids.

In India, about 25 million hectares of cultivated land have a pH below 5.5, and these are classed as critically degraded acid lands with low productivity. Liming acid soils raises the efficiency with which crops use fertiliser, saving about 50 per cent of chemical fertilisers, particularly in oilseeds and pulses.

The Soil Health Card scheme records pH. It was introduced on 19 February 2015 at Suratgarh, Rajasthan, and each card reports the soil status on 12 parameters, namely nitrogen, phosphorus, potassium, sulphur, zinc, iron, copper, manganese, boron, pH, electrical conductivity and organic carbon.

Ocean Acidification and Marine Life

Ocean acidification is the ongoing fall in the pH of the ocean as it absorbs carbon dioxide. Dissolved CO₂ forms carbonic acid, which releases hydrogen ions. Between 1950 and 2020 the average pH of the ocean surface is estimated to have fallen from about 8.15 to 8.05. The sea is still alkaline, but the change is large in hydrogen ion terms, since a fall of 0.1 pH unit is a 26 per cent increase in hydrogen ion concentration.

The organisms hit hardest are those that build shells and skeletons of calcium carbonate, such as corals, molluscs and coccolithophores, a phytoplankton. Acid water makes it harder for these organisms to build a shell or skeleton, so calcification is reduced. Pteropod shells dissolve as acidity rises, and brittle stars lose muscle mass when they regrow appendages.

Larvae are affected too. In experiments at pH 7.6, orange clownfish larvae showed no reaction to the smell cues that they normally use to find reefs. Many of these organisms sit at the base of the food web, so the effect can run up to fish and to the one billion people who depend wholly or partly on coral reefs for fishing, tourism and coastal management.

Salts and Their Uses

Neutralisation Reaction and Salt Formation

A neutralisation reaction is the reaction of an acid with a base to give a salt and water, with heat released. In the standard example, sodium hydroxide and hydrochloric acid give sodium chloride and water: NaOH + HCl → NaCl + H₂O. At heart, H⁺ joins OH⁻ to make water.

Neutralisation happens in many everyday settings. A drop of lemon juice in water turns blue litmus red, and lime water added slowly turns it blue again, because the base has used up the acid. A red ant bite leaves formic acid in the skin, and moist baking soda, a base, relieves the pain. Factory waste that is acidic can be neutralised with a base before it is released into a lake, so the fish are not harmed.

Families of Salts and the pH of Salt Solutions

Salts with the same positive or negative part belong to a family. NaCl and Na₂SO₄ are sodium salts, and NaCl and KCl are chloride salts.

A salt solution need not be neutral. The ions of some salts react with water, a process called hydrolysis, and this changes the pH. The result depends on which parent was strong.

The ions of a strong acid or strong base simply become hydrated and do not hydrolyse, which is why the salt of a strong acid with a strong base gives a neutral solution.

Sodium acetate is made from the weak acetic acid and the strong sodium hydroxide. Its acetate ion reacts with water to give acetic acid and OH⁻ ions, and because acetic acid is weak the OH⁻ ions build up and the solution is basic. Ammonium chloride is made from the weak base ammonium hydroxide and the strong hydrochloric acid. Its ammonium ion reacts with water to form NH₄OH and H⁺ ions, and because ammonium hydroxide is a weak base the H⁺ ions build up and the solution is acidic.

Hydrolysis is different from three processes with similar names. Electrolysis uses a direct electric current to drive a reaction, as in the chlor-alkali process; dialysis separates molecules in solution by their different rates of diffusion through a semipermeable membrane; and photolysis breaks molecules down by the absorption of light.

Common Salt and the Chlor-Alkali Process

Common salt is sodium chloride. It is separated from sea water, which contains many dissolved salts, or mined as rock salt, which occurs in large crystals that are often brown from impurities. Beds of rock salt formed where seas of past ages dried up, and the salt is mined like coal.

Common salt is the raw material for sodium hydroxide, baking soda, washing soda and bleaching powder. When electricity is passed through brine, a solution of sodium chloride, it decomposes. The process is called the chlor-alkali process, after chlorine and the alkali sodium hydroxide that it gives. The reaction is 2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂.

Chlorine gas is given off at the anode, hydrogen at the cathode, and sodium hydroxide solution forms near the cathode. All three products are useful, and the chlorine, for instance, goes to make bleaching powder.

The process became the principal source of chlorine during the twentieth century. The older diaphragm and mercury cells were environmentally unfriendly because they used asbestos and mercury, and the newer membrane cell is more energy efficient and avoids those materials.

Baking Soda, Baking Powder and Washing Soda

Baking soda is sodium hydrogencarbonate, NaHCO₃, also called sodium bicarbonate. It is made from sodium chloride, water, carbon dioxide and ammonia. It is a mild, non-corrosive basic salt, and heated during cooking it gives sodium carbonate, water and carbon dioxide, which is why it makes pakoras crisp and speeds cooking.

Its uses follow from its two properties, mild basicity and the carbon dioxide it releases with an acid:

• It is an ingredient in antacids, where being alkaline it neutralises excess acid in the stomach.
• It is used in soda-acid fire extinguishers, where an acid releases carbon dioxide from it.
• It eases a bee sting or ant bite, since the sting leaves an acid.

Baking soda is often confused with baking powder, which is a mixture of baking soda and a mild edible acid.

Washing soda is sodium carbonate decahydrate, Na₂CO₃·10H₂O. Heating baking soda gives sodium carbonate (soda ash), and recrystallising it from water gives washing soda, which is also a basic salt. Washing soda is used in the glass, soap and paper industries. In glass-making, sodium carbonate goes in with quicklime (calcium oxide), which makes the glass more durable. Washing soda is also used to make sodium compounds such as borax, as a household cleaning agent, and to remove the permanent hardness of water.

Bleaching Powder and Water Disinfection

Bleaching powder is made by the action of chlorine on dry slaked lime, Ca(OH)₂, and the chlorine comes from the chlor-alkali process. It is written as Ca(ClO)₂, the formula of calcium hypochlorite, its main active ingredient, although the real composition is quite complex. The reaction is 2Ca(OH)₂ + 2Cl₂ → Ca(ClO)₂ + CaCl₂ + 2H₂O.

Bleaching powder has three main uses.

• Bleaching: It bleaches cotton and linen in the textile industry, wood pulp in paper factories and washed clothes in laundries.
• Oxidising agent: It acts as an oxidising agent in many chemical industries.
• Disinfection: It makes drinking water free from germs, which is the commonest use of calcium hypochlorite.

Water of Crystallisation, Gypsum and Plaster of Paris

Water of crystallisation is the fixed number of water molecules present in one formula unit of a salt. Blue copper sulphate crystals look dry but contain it, with five water molecules per formula unit, written CuSO₄·5H₂O. On heating, the water is removed and the salt turns white; moisten it with a few drops of water and the blue colour comes back. Washing soda, Na₂CO₃·10H₂O, is also a hydrated salt, so a crystal of it is not "wet" even though it has ten water molecules in the formula.

Gypsum is another salt with water of crystallisation, two molecules of it, written CaSO₄·2H₂O. Heating gypsum drives off water and gives calcium sulphate hemihydrate, CaSO₄·½H₂O, which is Plaster of Paris. It is a white powder that sets into a hard solid mass as gypsum again when mixed with water, which is why doctors use it to hold a fractured bone in position and why it must be stored in a moisture-proof container. It is also used for making toys, decorative materials and smooth surfaces. The half in the formula arises because two formula units of CaSO₄ share one water molecule.

The current NCERT textbook gives 100 °C (373 K) for heating gypsum, while industrial kilns work at about 120 to 180 °C.

The name comes from the large gypsum deposit at Montmartre, a hill in the north of Paris, from which the material was first made.

Acids and Bases in Everyday Life

Natural Acids in Food, Stings and Plants

Many foods owe their sour taste to natural acids, and ant and nettle stings carry one too.

Stinging hairs on nettle leaves inject methanoic acid and cause burning pain; a traditional remedy is to rub the spot with the leaf of the dock plant, which often grows beside the nettle.

Sources

• NCERT: Class 10 Science, Chapter 2
• NCERT: Class 7 Curiosity (Science), Chapter 2
• NCERT: Class 11 Chemistry Book 1, Chapter 6
• PIB: Celebrating a Decade of Soil Health Cards
• ICAR: Natural Resource Management, Annual Report 2005‑06
• India Meteorological Department (MAUSAM): Prospects of acid rain over India
• India Meteorological Department (MAUSAM): Natural dust and the acidity of rain over India
• Wikipedia: Acid-base homeostasis
• Wikipedia: Acid rain
• Wikipedia: Antacid
• Wikipedia: Base (chemistry)
• Wikipedia: Calcium hypochlorite
• Wikipedia: Chloralkali process
• Wikipedia: Cloud seeding
• Wikipedia: Coccolithophore
• Wikipedia: Copper(II) sulfate
• Wikipedia: Dialysis (biochemistry)
• Wikipedia: Dialysis (chemistry)
• Wikipedia: Electrolysis
• Wikipedia: Formic acid
• Wikipedia: Glass
• Wikipedia: Litmus
• Wikipedia: Magnesium sulfate
• Wikipedia: Methyl orange
• Wikipedia: Ocean acidification
• Wikipedia: pH
• Wikipedia: pH indicator
• Wikipedia: Phenolphthalein
• Wikipedia: Photodissociation
• Wikipedia: Plaster of Paris
• Wikipedia: Sodium bicarbonate
• Wikipedia: Sodium carbonate
• Wikipedia: Sodium hydroxide
• Wikipedia: Soil pH
• Wikipedia: Universal indicator
• Wikipedia: Water
• UPSC: Civil Services Examination question papers