A Soap Molecule Has One End That Loves Water and One That Hates It

Grease does not dissolve in water. Pour oil into a glass of water and it floats in a separate layer, however hard you shake it — which is why rinsing a greasy plate under a tap achieves nothing at all.

Add soap and the grease comes away. That is not because soap dissolves grease in the ordinary sense, and it is not because soap dissolves in oil. It is because a soap molecule has two ends with opposite tastes.

- A long hydrocarbon tail, which is repelled by water and attracted to oil — the hydrophobic end
- A short ionic head, which is attracted to water — the hydrophilic end

No single substance could do the job; a two-ended molecule can. The tail buries itself in the grease and the head stays out in the water, so the grease ends up wrapped in a coat of soap molecules with their water-loving heads facing outwards — and a parcel like that can be rinsed away.

That structure is built from the two compounds this part of the chapter studies: a long-chain carboxylic acid provides the tail and the acid group, and sodium hydroxide converts the acid group into the ionic head. So ethanol, ethanoic acid and soap are one story, and the first two are the ingredients of the third.

This page covers the third part of the CBSE Class 10 Science chapter on carbon and its compounds: the properties of ethanol, the properties of ethanoic acid, the cleansing action of soap, and the comparison of soaps with detergents.

**Ethanol is a liquid at room temperature, mixes freely with water, and reacts as an alcohol through its $-\text{OH}$ group.

Reaction with sodium.** Ethanol reacts with sodium metal to give hydrogen and sodium ethoxide:

$$2\text{Na}+2{\text{C}}_2{\text{H}}_5\text{OH}\to 2{\text{C}}_2{\text{H}}_5\text{ONa}+{\text{H}}_2$$

Check the counts: sodium $2$ and $2$; carbon $4$ and $4$; hydrogen $12$ on the left and $10+2=12$ on the right; oxygen $2$ and $2$.

This reaction is a test and a comparison at once. Sodium in water reacts violently; sodium in ethanol reacts much more gently, but it still releases hydrogen. **So ethanol's $-\text{OH}$ behaves a little like the $-\text{OH}$ of water, just far less reactively — and that similarity is why ethanol is written as an alcohol rather than as a hydrocarbon.

Dehydration to ethene.** Heated with excess hot concentrated sulphuric acid, ethanol loses a water molecule and becomes an unsaturated compound:

$${\text{C}}_2{\text{H}}_5\text{OH}\underset{443\text{ K}}{\overset{\text{hot conc. }{\text{H}}_2{\text{SO}}_4}{\to }}{\text{C}}_2{\text{H}}_4+{\text{H}}_2\text{O}$$

The sulphuric acid is a dehydrating agent, not a catalyst in the ordinary sense — its job is to remove water, and naming that role is the mark in this question. The product, ethene, is unsaturated, so it will now decolourise bromine water while the ethanol it came from would not.

Uses of ethanol.

- As a solvent in medicines such as tincture of iodine, in cough syrups and in many tonics, because it dissolves substances that water cannot
- In the manufacture of other chemicals and as a fuel additive, since it burns cleanly
- Ethanol obtained by fermentation of sugar cane is added to petrol in some places, which is why it is also called a biofuel

Two cautions that the syllabus states. Ethanol is a depressant and consuming it is harmful, and methanol is far worse — it is poisonous even in small quantities and can cause blindness. To stop industrial alcohol being drunk, it is made unfit for consumption by adding methanol along with a dye; the product is called denatured alcohol, and the colour warns you what it is.

The distinction worth fixing. Ethanol and methanol differ by a single $-{\text{CH}}_2-$, so they are consecutive members of the same homologous series with almost identical chemistry — and completely different effects on the body. Chemical similarity does not mean biological similarity, which is a useful thing to remember whenever a series is described as very similar throughout.

Ethanoic acid behaves as a typical carboxylic acid: it turns blue litmus red, reacts with bases, releases carbon dioxide from carbonates, and forms esters with alcohols.

A five to eight per cent solution of ethanoic acid in water is the vinegar used in cooking, and pure ethanoic acid is called glacial acetic acid because it freezes into an ice-like solid in cold weather, at about $290$ K.

Reaction with a base — neutralisation.

$${\text{CH}}_3\text{COOH}+\text{NaOH}\to {\text{CH}}_3\text{COONa}+{\text{H}}_2\text{O}$$

The salt formed is sodium ethanoate, also called sodium acetate.

Reaction with a carbonate or a hydrogencarbonate — carbon dioxide is released.

$$2{\text{CH}}_3\text{COOH}+{\text{Na}}_2{\text{CO}}_3\to 2{\text{CH}}_3\text{COONa}+{\text{H}}_2\text{O}+{\text{CO}}_2$$
$${\text{CH}}_3\text{COOH}+{\text{NaHCO}}_3\to {\text{CH}}_3\text{COONa}+{\text{H}}_2\text{O}+{\text{CO}}_2$$

The brisk fizzing is the test, and the gas turns lime water milky. This is the reaction that distinguishes a carboxylic acid from an alcohol — ethanol does nothing with baking soda, while ethanoic acid effervesces at once.

Esterification — the reaction with an alcohol.

$${\text{CH}}_3\text{COOH}+{\text{C}}_2{\text{H}}_5\text{OH}\stackrel{\text{acid catalyst}}{\to }{\text{CH}}_3{\text{COOC}}_2{\text{H}}_5+{\text{H}}_2\text{O}$$

The product is an ester, and esters have sweet fruity smells — they are used in perfumes and as flavouring agents. A small amount of concentrated sulphuric acid acts as the catalyst here, which is a different job from the dehydration of the previous section, where it removed water from a single molecule.

Running esterification backwards makes soap. Heat an ester with sodium hydroxide and it splits, giving the alcohol back and the sodium salt of the acid:

$$\text{ester}+\text{NaOH}\to \text{alcohol}+\text{sodium salt of the acid}$$

That reaction is called saponification, from the Latin for soap, and it is how soap is actually manufactured: the esters in a vegetable oil or an animal fat are heated with sodium hydroxide, and the sodium salts of their long-chain acids are the soap. The alcohol released is glycerol, sold separately.

So the chapter's three topics are one chain of reactions. An alcohol plus an acid gives an ester; an ester plus a base gives a soap. Learning esterification and saponification as forward and reverse of the same equation halves the work and explains why the soap section follows immediately.

The hydrophobic tails bury themselves in the grease while the hydrophilic heads face outwards into the water, forming a ball called a micelle with the dirt trapped inside.

Soap is the sodium or potassium salt of a long-chain carboxylic acid. Its molecule has two distinct parts:

- The long hydrocarbon chain, which is hydrophobic — water-repelling and grease-attracting
- The **ionic $-{\text{COO}}^{-}{\text{Na}}^{+}$ end, which is hydrophilic — water-attracting

What happens when you wash a greasy cloth, step by step.

- Soap dissolves in the water, and the molecules cannot settle comfortably: their heads want water and their tails do not
- The tails escape from the water by burying themselves in any grease they find on the cloth
- Many soap molecules surround one droplet of grease, tails inward and heads outward. That spherical cluster is a micelle
- The outside of the micelle is now covered in water-loving ionic heads, so the whole parcel — grease and all — stays suspended in the water
- Agitating the cloth detaches the micelles from the fibres, and rinsing carries them away

Why the rubbing matters. The micelle has trapped the dirt, but it is still sitting on the cloth. Mechanical agitation is what removes it**, which is why clothes are rubbed, beaten or churned in a machine rather than merely soaked. A question asking why do we have to agitate clothes while washing wants exactly that answer.

Why the grease stays away instead of settling back. All the micelles carry the same negative charge on their outer surface, so they repel one another and cannot join up into a big drop that would sink back onto the cloth. The charge keeps the dirt dispersed, and that dispersion is called an emulsion.

The boundary case that tests understanding. Soap does not form micelles in ethanol, because soap dissolves in ethanol completely — the tails are comfortable there, so they have no reason to hide. Micelles form only in water, which is why a cleaning action that works in a bucket does not work in a solvent.

One visible consequence. Soapy water is cloudy even though soap dissolves, and that cloudiness is the micelles scattering light. Micelles are large enough to scatter light but small enough to stay suspended, which is the definition of a colloid — and it is why soap solution does not clear on standing while a salt solution does.

Hard water contains dissolved calcium and magnesium ions, and soap reacts with them to form an insoluble scum instead of a lather.

The calcium and magnesium salts of long-chain carboxylic acids are insoluble, so instead of dissolving and forming micelles, the soap is thrown out of solution as a sticky curd:

$$\text{soap}+\text{calcium or magnesium ions}\to \text{insoluble scum}$$

Three consequences, all of them familiar in a hard-water area.

- No lather forms until all the calcium and magnesium has been used up, so a great deal of soap is wasted
- The scum sticks to the cloth and to the bucket, leaving a grey deposit and a ring
- The cleaning is poor, because the soap that should have been making micelles has been destroyed

A detergent avoids all three. Detergents are ammonium or sulphonate salts of long-chain hydrocarbons instead of carboxylate salts. Their structure is the same in principle — a long hydrophobic tail and an ionic hydrophilic head, so they form micelles in exactly the same way — but their calcium and magnesium salts are soluble.

So a detergent lathers even in hard water, and that is the whole reason it exists. Detergents are used in shampoos, in washing powders and in products for cleaning clothes, and they work in cold water as well.

Worked comparison — the same cloth, two cleansers, hard water.

- With soap: little lather, a scum on the surface, a grey deposit on the cloth, and much of the soap wasted
- With detergent: full lather, no scum, and the dirt removed

The one disadvantage of detergents, which the syllabus asks about. Some detergents have branched hydrocarbon chains that microorganisms cannot break down, so they are not biodegradable. They persist in rivers and lakes, causing foaming and water pollution. Soap is biodegradable; a badly chosen detergent is not — which is why detergents with straight chains are preferred.

A practical test for hard water that follows from all this. Shake a little soap solution with the water and see whether it lathers. Plenty of lather means soft water; scum with little lather means hard water — a test using nothing but the chemistry of this section.

And one link back to the salts chapter. Permanent hardness is removed by adding washing soda, sodium carbonate, which precipitates the calcium and magnesium as insoluble carbonates and leaves the water soft. That is why washing soda appears in both chapters — it is the reagent that lets soap work where it otherwise could not.