Journey Inside the Atom: Class 9 Science · Chapter 8

Sharpen a pencil and rub the grey dust between your fingers. It smears into a layer finer than any powder you could grind. How far could you go on splitting it? In Class 8 you met the answer that chemists give: everything around you, from the air in a room to the cells of your hand, is built from atoms, far too small to see even with a good microscope.

This chapter opens up the atom itself. We follow the experiments that showed it has parts, the models scientists drew to explain each new result, and the rules that tell us how many particles an atom holds and where they sit. By the end you will be able to look at a label such as ${\phantom{X}}_{\phantom{11}}^{\phantom{23}}{\phantom{X}}_{\phantom{2}11}^{\phantom{2}23}\mathrm{N}\mathrm{a}$ and say what is inside that atom and how it is likely to combine with others.11Atoms have no colour of their own. The colours in our drawings are only a code to tell the parts apart.

8.1 Rediscovering the Roots of Atomic Theory

The question what is everything made of? is far older than any laboratory. More than two thousand years ago, thinkers in India and in Greece tackled it by reasoning alone, and they reached a strikingly similar answer.

Imagine breaking a lump of clay in half, then breaking a half in half, and so on. Acharya Kanada reasoned that this cannot go on forever: the breaking must stop at a particle without parts, the parmanu (परमाणु). In the Vaisheshika Sutra, the text that carries his teaching, no eye or touch can ever sense a single parmanu. Kanada pictured parmanus pairing up and then grouping further, so that a stone, a river and a living body are all built from such groups. One thing his scheme left open was the recipe: it does not say how many parmanus of each kind make a given substance.

Far away in Greece, Leucippus and his pupil Democritus arrived at a similar picture. They named the last, uncuttable bit of matter atomos, Greek for 'that which cannot be cut', which gave us our word atom. Both traditions argued their case with logic alone; nobody yet had a way to test it (Fig. 8.1).

The idea became science in 1808, when the English teacher John Dalton used it to explain measured facts. Chemists had found that elements always combine in fixed proportions by mass. Dalton proposed that each element is made of its own kind of atom, that atoms can be neither split, created nor destroyed, and that compounds form when atoms join in simple whole-number ratios. His was the first theory of atoms built on experimental evidence, and every later model grew out of it.

Dalton's theory raised as many questions as it answered. Does an atom have an inside? Could we ever see one, even indirectly? And if every element has its own atoms, how exactly does a gold atom differ from an oxygen atom? The rest of this chapter takes these up one at a time.

8.2 A Short Historical Journey Through Atomic Models

A scientific model is a simplified picture that explains what we observe and predicts what we have not yet tried. When a new experiment disagrees with it, the model is changed or replaced. The atom is a fine example: four models in about a hundred years, each one fixing what the one before it could not explain.

The first crack in the 'uncuttable' atom appeared at the end of the 1800s. Some elements, such as uranium and radium, were found to give off rays and particles all by themselves, a property called radioactivity. If particles can come out of an atom, the atom must have parts.

By the 1890s physicists had a curious device: a sealed glass tube, emptied of nearly all its air, with a metal plate at each end. Connect the plates to a supply of several thousand volts and the far end of the tube glows. Something invisible is streaming out of the negative plate, the cathode, towards the positive one, the anode; these streams were named cathode rays. In 1897 J. J. Thomson at Cambridge set out to find what they were. He steered the beam with charged plates and with magnets and measured how far it swerved (Fig. 8.2).

The beam bent towards the positive plate, so it had to be a stream of negatively charged particles. His measurements showed that each particle is far lighter than even a hydrogen atom, the lightest atom of all; we now know it is about 1840 times lighter. And the result did not depend on the metal of the cathode or on which gas was left in the tube. A particle that every material gives out must be part of every atom. These particles are the electrons (इलेक्ट्रॉन).

An electron carries a charge of $-1.602\times {10}^{-19}\mathrm{C}$. That number is awkward to write again and again, so chemists simply call it a charge of $-1$ and measure other charges in the same unit.

8.2.1 Thomson's model of an atom

Electrons raised an obvious question. Atoms carry no overall charge, so whatever balances the electrons' negative charge must be in there too. Since no light positive particle had turned up, Thomson imagined the positive charge smeared evenly through the whole atom, like a ball of positive 'jelly', with electrons stuck in it here and there (Fig. 8.3).

Because the electrons sit in the positive ball like dried fruit in an English Christmas pudding, it became known as the plum pudding model. A besan laddoo studded with raisins gives the same picture: laddoo for the positive charge, raisins for the electrons, the two charges exactly cancelling. The model did not last long, but it set the pattern every later model followed: say where each kind of charge sits, and show how the charges balance.

8.2.2 Testing Thomson's model: The gold foil experiment

A good model has to survive tests. Around 1910, Hans Geiger and Ernest Marsden, working in Ernest Rutherford's laboratory in Manchester, put Thomson's model through one. Their 'bullets' were alpha particles ($\alpha$-particles, अल्फ़ा कण): fast, positively charged particles shot out by radioactive elements such as radium. Each one is about 7300 times heavier than an electron and carries a charge of $+2$; it is, in fact, a bare helium nucleus.

They aimed a narrow beam of $\alpha$-particles at a sheet of gold beaten so thin that it was only a thousand or so atoms thick. Around the foil was a screen coated with zinc sulphide ($\mathrm{Z}\mathrm{n}\mathrm{S}$), which gives a tiny flash of light wherever an $\alpha$-particle strikes it. Sitting in a darkened room, the experimenters counted the flashes at different angles through a microscope (Fig. 8.4).

What should happen if Thomson was right? His positive charge is spread thinly through each gold atom, so no part of the atom can push hard on a heavy, fast $\alpha$-particle. Every particle should go almost straight through, swerving by a degree or so at most.

Most particles did pass straight through. But a few were turned through large angles, and about one in 8000 came back towards the source. Rutherford said later that this was almost as surprising as firing a heavy shell at a sheet of tissue paper and having it bounce back at you. Physicists call this kind of swerving scattering, which is why the work is often called the $\alpha$-particle scattering experiment.

Thomson's model could not produce those large deflections. A thin, spread-out positive charge is far too weak to turn back a particle as heavy and fast as an $\alpha$-particle. Something much more concentrated had to be inside the atom.

Rutherford's model of an atom

In 1911 Rutherford reasoned from each observation to a feature of the atom:

• Most particles went straight through, so most of an atom is empty space.
• A few were turned through large angles, so the atom's positive charge is packed into a tiny region that pushes hard on any $\alpha$-particle that comes very close. He called it the nucleus (नाभिक).
• Very few bounced straight back, so this tiny nucleus also holds nearly all of the atom's mass. Only a compact, heavy object can reverse a heavy, fast particle.
• Far out around this nucleus move the electrons, the way planets circle the Sun. Rutherford's picture is therefore known as the nuclear model or, informally, the planetary model.

How small is the nucleus? Atoms are about ${10}^{-10}\mathrm{m}$ across and nuclei about ${10}^{-15}\mathrm{m}$, so a nucleus is roughly a hundred thousand (${10}^5$) times narrower than its atom. Picture the nucleus as a 2 cm glass marble, a kancha, lying in the middle of a huge field. The atom's outermost electrons would then be about a kilometre away on every side, so the atom would be a ball some 2 km across, almost all of it empty.

The sim shows why only a close approach matters: each time the distance halves, the push grows four times. Because the nucleus is so tiny, only the rare particle aimed almost straight at one feels a strong enough push to turn back.

Limitations of Rutherford's model

The nuclear model had a hidden flaw. Physicists knew, from experiments with radio aerials, that a charged particle radiates energy whenever it speeds up, slows down or turns. An electron going round a nucleus is turning all the time; Chapter 4 showed that motion in a circle is accelerated motion. So the electron should leak energy as light, sink closer to the nucleus with every turn and crash into it in less than a billionth of a second (Fig. 8.7).

Every atom in the universe would then have collapsed long ago. Yet the iron in a century-old bridge and the gold in an ancient coin are still there, unchanged. Something in Rutherford's picture had to give.

Discovery of the proton

Meanwhile Rutherford was probing the nucleus itself. In 1919 he fired $\alpha$-particles into nitrogen gas and found that they knocked out positively charged particles identical to the nucleus of a hydrogen atom. He concluded that every nucleus contains such particles, and in 1920 he named them protons (प्रोटॉन).

A proton carries a charge of $+1$, exactly equal and opposite to an electron's, but it is about 1836 times heavier. A whole atom is neutral, so the number of protons in an atom equals the number of electrons. A lithium atom has 3 protons and 3 electrons; a chlorine atom has 17 of each. In both, the positive and negative charges cancel exactly.

8.2.3 Bohr's model of the atom

In 1913 the Danish physicist Niels Bohr kept Rutherford's nucleus but changed the rules for the electrons. He did not prove his new rules. He proposed them as postulates, statements accepted because what they predict matches experiments.

• Only certain orbits are allowed. An electron can circle the nucleus only along particular paths, the shells (कोश), also called orbits or stationary states, and never in the space between them.
• Each shell has a fixed energy, so shells are also called energy levels. Going outwards they are named K, L, M, N, … or numbered $n=1,2,3,4,\dots$; K is the lowest in energy and each shell further out is higher.
• No energy is lost inside a shell. An electron circling in its shell keeps exactly the same energy, however long it goes round.
• Changing shells takes an exact amount of energy. To climb to a higher shell an electron must absorb exactly the energy gap between the two; dropping back, it releases that same amount.
• Shells have limited room. Each shell has space for only so many electrons.

Think of a staircase. A ball can rest on any step but never in the air between two steps, and moving it up or down always takes a whole step's worth of energy. Bohr's shells work like those steps. An electron in its shell has a fixed energy that does not drain away, so it does not spiral inwards, and the atom stays stable. Bohr's rules also explained why glowing hydrogen gives out light of only a few sharp colours, and that success made physicists take his strange rules seriously.

8.3 What Components Contribute to the Mass of an Atom?

Electrons are so light that together they make up less than a tenth of a percent of an atom's mass, so for most purposes we can ignore them. The mass must be in the nucleus. If a nucleus held only protons, an atom's mass would simply be its number of protons times the mass of one proton. Hydrogen fits: one proton, mass 1 unit. But helium, with 2 protons, weighs about 4 units, not 2 (Fig. 8.9). Heavier atoms fall even further short: a gold nucleus has 79 protons yet weighs about 197 units. Something heavy but uncharged had to be in the nucleus as well.

8.3.1 Discovery of the Neutron

The missing particle was tracked down in 1932. Experimenters in Germany and France had noticed a strange, very penetrating radiation that could knock protons out of paraffin wax. James Chadwick, Rutherford's long-time colleague at Cambridge, showed that it was a stream of particles about as heavy as protons but carrying no charge: neutrons (न्यूट्रॉन), symbol $\mathrm{n}$. Having no charge, they had slipped past every earlier test that relied on electric forces.

With neutrons in the nucleus, helium's mass makes sense: 2 protons plus 2 neutrons give about 4 units. Ordinary hydrogen is the one exception, a lone proton with no neutron. For every atom, nearly all the mass sits in the nucleus. Table 8.1 sums up the three particles.

Count the particles in a few nuclei and a pattern appears. Nitrogen-14 has 7 protons and 7 neutrons; magnesium-24 has 12 and 12. Silver-107, though, has 47 protons but 60 neutrons, and gold-197 has 79 protons and 118 neutrons. The heavier the nucleus, the further the neutrons pull ahead of the protons; the box below explains why.

8.4 Symbols of Elements

When Dmitri Mendeleev drew up his famous table of elements in 1869, he had a little over sixty to arrange. Today 118 are known, most of them found in nature and a few made only in laboratories. Writing their names in full would make chemistry slow, and the names differ from language to language. So every element has a short symbol that chemists everywhere recognise.

Dalton was the first to give elements symbols, in the early 1800s: small circles with dots, lines or letters inside. They were hard to print and hard to remember. In 1813 the Swedish chemist Jöns Jacob Berzelius suggested using one or two letters of each element's name instead, often of its Latin name. Today an international body, IUPAC (International Union of Pure and Applied Chemistry), approves every new name and symbol. Symbols follow a few simple rules (Fig. 8.10).

• Many symbols are the first letter of the English name: hydrogen H, carbon C, nitrogen N.
• When first letters clash, a second letter is added: helium He, calcium Ca, aluminium Al.
• Sometimes the second letter is a later letter of the name: magnesium Mg, manganese Mn, chromium Cr.
• Several come from Latin names (Table 8.2), and tungsten's symbol W comes from its German name, wolfram.
• The first letter is always a capital and the second always small. Co is the element cobalt, but $\mathrm{C}\mathrm{O}$ is carbon monoxide, a compound of carbon and oxygen.

8.5 Atomic Number

What makes a gold atom gold and not oxygen? It is not the electrons, which an atom can gain or lose. Nor is it the neutrons, as you will see in Section 8.9.1. It is the number of protons in the nucleus. This number is the element's atomic number (परमाणु संख्या), written $Z$.

A neutral atom has as many electrons as protons, so $Z$ also counts its electrons: six for carbon, eight for oxygen, twenty for calcium. Add one more proton to a carbon nucleus and it is no longer carbon at all: it becomes nitrogen, with $Z=7$ (Fig. 8.11). The atomic number works like an element's roll number. It tells you which element you have, and through the electrons it fixes, it decides how the element behaves in reactions.

8.6 Mass Number

Protons and neutrons together are called nucleons, since both live in the nucleus. Each has a mass of about 1 unit and the electrons hardly count, so the total number of nucleons tells us roughly how heavy an atom is. That total is the atom's mass number (द्रव्यमान संख्या), written $A$.

Turn it round and you get the number of neutrons: $N=A-Z$. Table 8.3 applies this to three atoms.

Chemists pack all of this into one compact label. The mass number goes at the top left of the symbol and the atomic number at the bottom left. Aluminium, with 13 protons and 14 neutrons, is written ${\phantom{X}}_{\phantom{13}}^{\phantom{27}}{\phantom{X}}_{\phantom{2}13}^{\phantom{2}27}\mathrm{A}\mathrm{l}$ (Fig. 8.12). Often the atomic number is left out, since the symbol already fixes it, and we write ${\phantom{X}}_{\phantom{}}^{\phantom{27}}{\phantom{X}}_{\phantom{2}}^{\phantom{2}27}\mathrm{A}\mathrm{l}$ or simply aluminium-27.

8.7 How Are Electrons Distributed in Different Energy Levels?

Bohr's model puts electrons in shells, but how many go into each one? In the early 1920s Bohr and the chemist Charles Bury proposed three simple rules that work well for the first twenty elements:

1. A shell numbered $n$ can hold at most $2n^2$ electrons.
2. The outermost shell of an atom never holds more than 8 electrons. (When the K shell is the only shell, it holds at most 2.)
3. Shells fill in order from the inside out: K, then L, then M. An outer shell receives electrons only when the shells inside it have taken all they may.

Try the rules on hydrogen first. Its single electron goes into the K shell. Helium's two electrons fill the K shell exactly. Where must the third electron of lithium go?

8.7.1 Building up atoms

Now build the elements in order of atomic number, like stacking bricks. Going from one element to the next adds a proton to the nucleus (plus a neutron or two) and an electron, which drops into the innermost shell with a free place. The resulting share-out of electrons among the shells is the atom's electronic configuration (इलेक्ट्रॉनिक विन्यास). We write it as the number of electrons in each shell, from K outwards: lithium is 2, 1 and sodium is 2, 8, 1 (Fig. 8.13).

Table 8.5 lists the configurations of the first twenty elements. Notice how each new shell starts with a single electron (lithium, sodium, potassium), and how helium, neon and argon end with a full outermost shell.

8.8 Combining Capacity of an Atom: Valency

Atoms join to form molecules, but not in just any numbers. One carbon atom holds four hydrogen atoms in methane ($\mathrm{C}\mathrm{H}{\phantom{X}}_4$), one sulphur atom holds two in hydrogen sulphide ($\mathrm{H}{\phantom{X}}_2\mathrm{S}$), and one aluminium atom holds three chlorine atoms in aluminium chloride ($\mathrm{A}\mathrm{l}\mathrm{C}\mathrm{l}{\phantom{X}}_3$). Hydrogen and chlorine each link to just one partner, so they make a handy yardstick: we measure an element's combining capacity by counting how many hydrogen or chlorine atoms a single atom of it holds on to. It is 4 for carbon, 2 for sulphur and 3 for aluminium.

Why these numbers? Start with the loners of chemistry. Helium, neon and argon, the noble gases, hardly form compounds at all, which is why argon can safely fill electric bulbs and neon can glow inside signboards. Their configurations are 2; 2, 8; and 2, 8, 8. In each, the outermost occupied shell, the valence shell, is completely full: 8 electrons, called an octet, or 2 for helium. The electrons in an atom's valence shell are its valence electrons, and they are the ones that take part in reactions.

Every other atom has gaps in its valence shell, and it reacts in whatever way brings it to a filled shell like a noble gas: by handing electrons over, taking them in, or sharing them with a partner. How many electrons it has to move or share for this is its valency (संयोजकता), the same number as its combining capacity. (For hydrogen, a filled K shell means 2 electrons.)

A simple rule predicts the valency from the configuration:

• 1, 2 or 3 valence electrons: the atom loses them, so the valency equals the number of valence electrons. Magnesium, 2, 8, 2, loses 2: valency 2.
• 5, 6 or 7 valence electrons: the atom gains enough to reach 8, so the valency is 8 minus the number of valence electrons. Fluorine, 2, 7, gains 1: valency 1.
• 4 valence electrons: losing or gaining four is too hard, so the atom shares them. Carbon, 2, 4, shares 4: valency 4.
• A full outermost shell (helium, neon, argon): nothing to lose or gain, so the valency is 0.

A few compounds seem to break these simple rules; you will meet them in higher classes.

8.9 A Deeper Look into Atomic Structure

One of Dalton's ideas was that every atom of an element is a perfect copy of every other, mass included. Now that you know what is inside an atom, you can test it. The proton count fixes the element, so all atoms of an element share the same $Z$. Must they also share the same number of neutrons?

8.9.1 Isotopes

It turns out that it need not. Atoms of the same element can carry different numbers of neutrons, and so have different mass numbers. Atoms like these, which share an atomic number but not a mass number, are called isotopes (समस्थानिक) of the element.

Hydrogen shows this best. Its nucleus can be a lone proton (protium, ${\phantom{X}}_{\phantom{1}}^{\phantom{1}}{\phantom{X}}_{\phantom{2}1}^{\phantom{2}1}\mathrm{H}$), a proton with one neutron (deuterium, ${\phantom{X}}_{\phantom{1}}^{\phantom{2}}{\phantom{X}}_{\phantom{2}1}^{\phantom{2}2}\mathrm{H}$) or a proton with two neutrons (tritium, ${\phantom{X}}_{\phantom{1}}^{\phantom{3}}{\phantom{X}}_{\phantom{2}1}^{\phantom{2}3}\mathrm{H}$, which is radioactive). Each has a single electron (Fig. 8.16). In nature protium wins easily, at about 99.98 %: only one hydrogen atom in roughly 6400 is deuterium, and tritium exists only in traces.

Water made with deuterium, called heavy water ($\mathrm{D}{\phantom{X}}_2\mathrm{O}$), looks just like ordinary water. Yet it is about 11 % denser, freezes at 3.8 °C and boils at 101.4 °C, because its molecules are heavier. In reactions, though, deuterium behaves like ordinary hydrogen: chemistry is decided by electrons, and isotopes have identical electron arrangements. So isotopes share their chemical properties but differ in physical ones such as density and boiling point. Many of India's nuclear power plants use heavy water to slow down the neutrons in their reactors.

Carbon is another example: carbon-12 makes up about 99 % of it, carbon-13 about 1 %, and radioactive carbon-14 only a trace. All three have 6 protons; their neutrons number 6, 7 and 8.

Average atomic mass

Most elements in nature are mixtures of isotopes. Boron, for example, is about 20 % boron-10 and 80 % boron-11. So what single number should we give as the mass of a boron atom?

One idea is the simple average of the two masses: $(10+11)÷2=10.5\mathrm{u}$. But that treats the two isotopes as equally common, and they are not. Picture a jar of 10 boron atoms in the natural ratio, 2 of mass 10 u and 8 of mass 11 u (Fig. 8.17). Their total mass is $2\times 10+8\times 11=108\mathrm{u}$. Shared among 10 atoms, that is 10.8 u per atom. This weighted average, in which each mass counts according to how common it is, is the element's average atomic mass.

No single boron atom has a mass of 10.8 u: each one is either 10 u or 11 u. The value 10.8 u describes the average atom in any ordinary sample of boron, which is why it is the number printed in the periodic table.

8.9.2 Isobars

Now turn the idea around. Can two different elements have atoms with equal mass numbers? Carbon-14 has 6 protons and 8 neutrons; nitrogen-14 has 7 protons and 7 neutrons. Both contain 14 nucleons, so both have $A=14$, yet they are different elements because their atomic numbers differ (Fig. 8.18). Pairs like this, from different elements but with equal mass numbers, are called isobars (समभारिक). Iron-58 and nickel-58 are another such pair.

Bohr's neat circles are themselves a simplification. Modern physics says that an electron has no definite path at all; we can only work out how likely it is to be found in each region around the nucleus. Drawn as a picture, this becomes a fuzzy electron cloud, densest where the electron is most likely to be. That model grew out of work in the 1920s, and you will study it in higher classes.