the Atom
Human Body
House
Room
Organ
Atoms
Si
Wall
Tissue
Brick
Cell
Cell
Silicates
Protein
Think It Over
Everything that you see or observe, or feel around is matter. You have learnt that matter consists of tiny particles called atoms. Take a closer look at the picture given above. What do you observe? Do you notice that both living beings, like us, and non-living things, like a house, are ultimately composed of atoms? These atoms are so tiny that they cannot be seen with the naked eye.
Grade 8CuriosityChapter 7
• Are atoms the smallest indivisible particles? • Why do electrons not fall into the nucleus even though they are attracted to protons in it? • Why did scientists keep modifying atomic models?
You may be wondering — Is an atom truly the smallest unit of matter, or can it be divided even further?
Scientists, too, have been exploring whether atoms are divisible. If so, what are their constituents and how are these arranged? Let us examine how the concept of atoms emerged and how it has been evolved since then.
141Journey Inside the Atom
8.1 Rediscovering the Roots of Atomic Theory
Let us embark on a journey that takes us back more than 2,000 years, to the intellectual landscapes of ancient India and ancient Greece. In these distant yet remarkably parallel civilisations, profound thinkers, such as Acharya Kanada in India, and Leucippus and Democritus in Greece, pondered over the same fundamental question that has continued to inspire human inquiry across centuries — What is everything made up of?
Acharya Kanada suggested that if matter (dravya) is divided repeatedly, you will reach a stage where you would encounter the smallest particles that can no longer be divided. He called these particles parmanus. His ideas are recorded in the Sanskrit text Vaisesika Sutras. A parmanu is infinitely small and cannot be perceived by the senses. Combinations of these forms dyads (groups of two parmanus) and triads (groups of three parmanus), and so on. It is out of these combinations that the whole of the material universe, including the bodies of living beings, is created. However, this description does not specify the proportions in which parmanus combine to form different substances.
The Greek philosophers Leucippus and Democritus also proposed a similar idea. They called these indivisible particles atomos (in Greek, atomos means indivisible).
You must remember that the concept of ‘atom’ originated as an imaginary idea rather than from experimental observations.
Many centuries later, in 1808, John Dalton proposed his atomic theory. It was based on scientific experiments of that time. He proposed that all matter is composed of indivisible particles called atoms. That is, the atoms are the fundamental building blocks of matter that cannot be broken down into smaller parts. Dalton’s atomic theory was the first scientific description of how matter is made. It became the starting point for the current understanding of atomic structure.
You may be wondering how Dalton’s idea evolved into a model for the structure of the atom.
Following Dalton’s theory, scientists were curious to know the answers of the questions:
y
What are atoms made up of ? • What would atoms look like if we could see them? • What makes the atoms of one element different from the atoms of
another element?
8.2 A Short Historical Journey Through Atomic Models
More than a hundred years ago, scientists tried to imagine what atoms might look like by proposing simple models. As new experiments were performed and new evidence came to light, these models were changed and improved. Although we now know that the early models were not fully correct, they are still important because they show how science moves forward — one step at a time, driven by curiosity, questioning, and experimentation.
Until the late 19th century, atoms were thought to be the smallest, indivisible units of matter. However, scientists discovered that certain elements emit invisible energy and particles called radiation,
a phenomenon known as radioactivity. This showed that atoms must be composed of smaller particles, proving that they were not indivisible as previously believed.
Note
Atoms do not show any colour. The colours depicted in the diagrams are for illustrative purposes only.
In 1897, J. J. Thomson studied the conduction of electric current through gases at a very low pressure. He used a glass tube with two electrodes and applied a high voltage. He observed rays moving from the cathode (negative electrode) to the anode (positive electrode) (Fig. 8.1). These were called cathode rays. By studying these cathode rays in electric and magnetic fields, he concluded that they are streams of negatively charged particles, with a much smaller mass than atoms. These particles, later called electrons, were emitted from atoms, indicating that atoms are composed of smaller subatomic components.
–+
Meet a Scientist
High voltage power
supply
J. J. Thomson’s most significant discovery was that of the electron, the first subatomic particle to be identified and a part of every atom. He received the Nobel Prize in Physics in 1906 for his studies of the electrical conductivity of gases. This research led him to discover electrons. As the head of the famous Cavendish Laboratory in Cambridge, he guided and inspired many scientists, including Ernest Rutherford.
Cathode
Anode
(–)
(+)
Vacuum pump
Fig. 8.1: A line diagram of cathode ray tube
It was found that the nature of cathode rays was independent of the material of the cathode and the gas filled in the cathode ray tube. It showed that electrons are a fundamental component of all atoms, present in every element. The charge of an electron (–1.602 × 10–19 C) is taken as –1 as a matter of convention and convenience.
8.2.1 Thomson’s model of an atom
When J. J. Thomson discovered tiny negatively charged particles called electrons, he faced a puzzle — atoms are neutral, so where is the positive charge present? To solve this, Thomson proposed the atom to be a sphere of positive charge with electrons distributed throughout it (Fig. 8.2). This model was compared to a pudding with plums embedded in it, called the plum pudding model. A more familiar picture would be that of a watermelon (Fig. 8.3), where the red pulp represents the positively charged matter, and the seeds represent electrons distributed throughout the atom. This simple picture, though later replaced, was the first genuine attempt to describe how the atom’s positive and negative charges stay balanced.
Electron
Positive
sphere
an atom
143Journey Inside the Atom
Pause and Ponder
1. Suppose you made up your own ‘atom’, as Thomson described, using clay for the
positive charge and small beads for the electrons spread through it. What will happen if:
(i) the positive charge on the clay is lesser than the total negative charge of the beads?(ii) by mistake, the clay itself carries a bit of negative charge? Would your model still
represent a neutral atom?2. Could an orange or a lemon, which also contain seeds inside soft pulp, be a good
comparison? In what ways does it match Thomson’s idea and where does it fall short?3. Why did Thomson conclude that electrons are present in all atoms?
8.2.2 Testing Thomson’s model: The gold foil experiment
In 1911, Geiger and Marsden, working under Ernest Rutherford, tested Thomson’s model of the atom through what became famous as the gold foil experiment. They aimed a narrow beam of alpha particles at an extremely thin sheet of gold foil. Alpha (symbol α) particles are tiny, positively charged particles emitted from certain radioactive elements. Later in this chapter, you will learn that an alpha particle is actually a nucleus of a helium atom containing two protons and two neutrons. According to Thomson’s model, the positive charge in the atom was spread out evenly. So they expected the alpha particles to pass straight through the gold foil or be deflected only slightly. But to their surprise, while most particles passed through undeflected, some were sharply deflected (Fig. 8.4), and a few even bounced back. This deflection from the straight path is called scattering. Hence, the gold foil experiment is also called an α-ray scattering experiment.
Thin gold foil
Beam of α-particles
Fig. 8.4: Schematic view of the gold foil experiment
T homson’s model failed to explain the results of the gold foil experiment, particularly the deflection of some α-particles through large angles and that most of the α-particles passed undeflected.
A. Rutherford’s model of an atom From the gold foil experiment, Rutherford concluded that the positive charge of an atom is not spread throughout but remains concentrated in an extremely small region called the nucleus. He proposed that:
y
Most of an atom is empty space, as most α-particles passed through the
gold foil without any deflection. • The nucleus is dense, contains all the positive charge and most of the
mass of an atom. • The electrons revolve around the nucleus, somewhat like planets
orbiting the Sun. Hence, this model is called the planetary model of the atom (Fig. 8.5).Rutherford found that the nucleus is extremely small — about 105 (one lakh) times smaller than the atom. He calculated that the diameter of an atom is ≈ 10−10 m, and the diameter of the nucleus is ≈ 10–15 m. You can imagine that if an atom were of the size of a cricket ground (about 100 m across), the nucleus would be just a tiny black pepper grain (a few mm) at the centre!
suggested by Rutherford
Ready to Go Beyond
Can you calculate how many atoms would be needed to make a sheet of paper that is 0.1 mm thick, like the one in your textbook?If the diameter of one atom is about 10–10 m, and the sheet is 0.1 mm (10–4 m) thick,number of atoms ≈ (10–4 m) ÷ (10–10 m) = 106.
That is about one million atoms need to be stacked together! When these atoms come together, they form the solid materials we see and touch. What feels like a simple sheet is actually a vast assembly of tiny building blocks, all neatly arranged. It’s amazing how the unseen atomic world shapes everything we see and use every day!
Think as a Scientist
Observe Fig. 8.4 of the gold foil experiment. Predict the observations you would expect if the gold foil in the experiment were made thicker. Also, draw a simple diagram to show the observations you expect.
Hint: Compare thin foil vs thick foil. How does the thickness affect the chances of hitting a nucleus?
Pause and Ponder
4. What do you think would happen if α-particles were replaced with negatively
charged particles in Rutherford’s gold foil experiment?
5. Rutherford found that a few α-particles bounced back sharply. How does this
single surprising result completely rule out Thomson’s ʻplum pudding modelʼ of the atom?6. If you could ask Rutherford one question about his work, what would it be?
145Journey Inside the Atom
Rutherford’s atomic model was better than Thomson’s atomic model in explaining the results of the gold foil experiment. However, it had a limitation. It could not explain the stability of the atom.
Meet a Scientist
Born in New Zealand, Ernest Rutherford moved to Cambridge to work with J. J. Thomson and later became known as the Father of Nuclear Physics. He discovered the atomic nucleus and explained how some elements naturally break down, for which he won the 1908 Nobel Prize in Chemistry. In 1911, he proposed the nuclear model of the atom. His portrait now appears on New Zealand’s $100 banknote.
B. Limitations of Rutherford’s model While the idea of a central nucleus was a major step forward, Rutherford’s model could not explain why atoms are stable.
As you have learnt in Chapter 4, Describing Motion Around Us, a particle moving in a circular path is constantly changing direction, which means it is accelerating. If a negatively charged electron keeps accelerating around the nucleus, it should lose energy. Losing energy would make it spiral inward and eventually fall into the positively charged nucleus (Fig. 8.6). If that really
happened, atoms would collapse and would not exist! But in reality, atoms are stable that is why matter around us stays intact. This meant that Rutherford’s model was not completely correct, and a new explanation was needed to describe how electrons remain in motion without collapsing into the nucleus.
But before we come to that, you need to know about another subatomic particle, the proton, discovered and named by Rutherford.
Fig. 8.6: Spiral path followed by a charged particle on losing energy
C. Discovery of the proton Rutherford showed that the nucleus carries positive charge which comes from the particles called protons. Protons are much heavier than electrons and possess a charge equal and opposite to that of electrons.
For an atom to be electrically neutral, the number of protons must be equal to the number of electrons. For example, a helium atom has 2 protons and 2 electrons, while a sodium atom has 11 protons and 11 electrons. Since the total positive charge equals the total negative charge, the helium or the sodium atoms are also electrically neutral. Similarly, all the atoms are electrically neutral.
Pause and Ponder
7. Assertion (A): Rutherford concluded that most of the mass of an atom is concentrated in
a small region at the centre called the nucleus.Reason (R): According to Thomson’s model, electrons are embedded in a uniformly distributed positive charge sphere.Choose the correct option:
(i) Both A and R are true, and R is the correct explanation of A.(ii) Both A and R are true, but R is not the correct explanation of A.(iii) A is true, but R is false.
(iv) A is false, but R is true.
8.2.3 Bohr’s model of the atom
To explain why atoms are stable, Niels Bohr proposed a new model of the atom in 1913. According to Bohr:
L-shell (n=2)M-shell (n=3)N-shell (n=4)
K-shell (n=1)
Nucleus
y
Electrons do not move randomly around the
nucleus but follow fixed circular paths called stationary states, orbits, or shells. In each shell, an electron has a definite amount of energy, so these shells are also called energy levels. • These shells are represented by the letters K, L,
M, N, ... or by the numbers n = 1, 2, 3, 4, ... (Fig. 8.7). • Electrons can revolve only in these allowed shells, and not in between
atom
them. While moving in a fixed shell, an electron does not lose energy. • The first energy level K (n = 1) is the one closest to the nucleus (Fig. 8.7)
Meet a Scientist
and has the least energy. • The energy of these levels increases as we move away from the
nucleus. That is, the energy of an electron in the L-shell (n = 2) is more than that of an electron in the K-shell (n = 1). The farther away a shell is from the nucleus, the higher is its energy. • An electron can move to another shell by absorbing or releasing a
fixed amount of energy equal to the difference between the energies of the two levels. • Each shell can hold only a certain number of electrons.
Niels Bohr was a professor of physics at Copenhagen University, Denmark. He was curious about how atoms exist because the old models could not explain why electrons stay around the nucleus without collapsing. His explanation of the atomic structure provided more clarity. Niels Bohr received the Nobel Prize in 1922, for his work on the structure of the atom.
You may be wondering, how does Bohr’s model explain stability? In Bohr’s model, electrons also move in circular paths around the positively charged nucleus, as in Rutherford’s model. But why don’t they lose energy while doing so? Bohr addressed this by introducing the concept of stationary states as a postulate. In a stationary state, the energy of an electron remains constant, even though it is in motion around the nucleus. Bohr’s model could explain many experimental observations and marked a major step in understanding the atomic structure.
Threads of Curiosity
Why are Bohr’s shells called K, L, M, N... and not, A, B, C, D? The naming came from early X‑ray experiments by the physicist, Charles Barkla, who called the first observed X-ray line K. He didn’t start naming from A to leave room for possible discovery of a series earlier than the K series, although none were ever found. Bohr adopted the same notation for atomic shells.
Later, even Bohr’s model was found to have limitations, and yet
another model, the quantum mechanical model, was proposed. You will learn about it in higher grades.
NextLevel
Up
147Journey Inside the Atom
8.3 What Components Contribute to the Mass of an Atom?
Rutherford’s model showed that most of the mass of an atom is concentrated in its nucleus. Electrons that revolve around the nucleus are so light that their mass can be ignored.
What if ...
an atom had no empty space? How would this have affected the size of various objects?
However, something puzzling appeared early in the 20th century. For instance, a hydrogen atom has one proton, whereas a helium atom has two protons, yet the mass of a helium atom is about four times that of a hydrogen atom, not double. This led scientists to wonder whether besides protons, is there something else in the nucleus adding mass without affecting its charge?
Ready to Go Beyond
The discovery of the neutron opened a new era in atomic physics. Neutrons, being uncharged, can easily penetrate nuclei, leading to breakthroughs, such as the creation of artificial radioactive elements and the splitting of uranium atoms. This gave birth to the ʻatomic ageʼ, allowing the development of both nuclear power and nuclear weapons.
8.3.1 Discovery of the Neutron
In 1932, this problem was solved by James Chadwick (a student of Ernest Rutherford). He discovered a new subatomic particle with a mass nearly equal to that of a proton but no electrical charge. This neutral particle was named as neutron and is usually represented by the symbol ‘n’. The neutrons are found in the nucleus of all atoms except hydrogen. Thus, the mass of an atom comes mainly from its protons and neutrons packed tightly in the nucleus. This also explains why atoms are heavier than the mass of their total number of protons. Table 8.1 shows the symbols and relative charges of subatomic particles.
Meet a Scientist
S.No.Subatomic particleSymbolRelative charge
1.Electrone ––1
James Chadwick, working under Rutherford at the famous Cavendish Laboratory at the University of Cambridge, solved a key puzzle in 1932, when he discovered the neutron. This breakthrough explained atomic mass and earned him the Nobel Prize in Physics in 1935. The neutron immediately transformed research, enabling scientists to probe nuclear secrets, and sparked a chain of discoveries in understanding and harnessing atomic energy.
2.Proton p++1
3.Neutronn00
Threads of Curiosity
Lighter atoms often have an equal number of protons and neutrons (like carbon with six each or oxygen with eight each). However, as atoms get heavier, their nuclei have many more neutrons than protons. Iron has 26 protons and 30 neutrons, and by the time we reach uranium, the nucleus has 92 protons and 146 neutrons.
You may wonder that since all the protons with like charges are squished together in a nucleus why do they not push each other away?
Think of it this way — every proton inside the nucleus repels every other proton because they all carry positive charge. Neutrons being neutral help reduce this repulsion by intervening and increasing the distance between protons, and also by strengthening the force, called the nuclear force, that binds all particles together. So, heavier atoms need many more neutrons to hold everything in the nucleus tightly bound.
India's Scientific Contributions
The Bhabha Atomic Research Centre (BARC, Fig. 8.8), Mumbai, leads advanced neutron-scattering experiments using reactors, such as Dhruva. This has revealed key insights into materials like superconductors, battery electrodes, and drug molecules, helping develop better medicines, energy storage, and industrial alloys right here in India.
Research and explore more about Dhruva!
By 1869, scientists knew about 69 elements, most of which were found naturally on the Earth. Today, we know about 118 unique chemical elements. Some of these are artificially made, and the search for even more continues.
Now, let us explore how elements began to be represented for the sake of simplicity.
8.4 Symbols of Elements
John Dalton realised the need for a standard way to represent elements and compounds to make the study of chemistry easier. In 1803, he introduced the first pictorial symbols to represent the known elements. The Fig. 8.9 shows some of the symbols.
HydrogenPhosphorusCarbonOxygen
CopperSulphurLeadIron
GoldPlatinaSilverMercury
Fig. 8.9: Symbols of some elements given by Dalton
In 1813, Berzelius suggested that the symbols of elements should be derived from their Latin names. Thus, came alphabetic chemical symbols. Nowadays, the International Union of Pure and Applied Chemistry (IUPAC), an international scientific organisation, approves the names and symbols of elements. Some of the norms of writing these symbols are:
y
Many symbols are the first letter or the first two letters of the name of
the element. • The first letter of a symbol is always written as a capital letter
(uppercase) while the second letter (if there is one) is written as a small letter (lowercase). For example, hydrogen, H; aluminium, Al (not AL); cobalt, Co (not CO), etc. • Symbols of some elements are formed from the first letter of the name
and a letter other than the second letter in the name. For example, chlorine, Cl; zinc, Zn, etc.
149Journey Inside the Atom
y
Symbols of some elements come from their Latin, Greek, or German
names rather than English. For example, the symbol for iron is Fe (from Latin ferrum), for mercury is Hg (from Greek hydrargyros), and for tungsten is W (from German wolfram). • Names and symbols of some commonly used elements are given in
Name of the elementSymbolName of the elementSymbolName of the elementSymbol
AluminiumAlCopper (Cuprum)CuNitrogenN
ArgonArFluorineFOxygenO
BariumBaGold (Aurum)AuPotassium (Kalium)K
BoronBHydrogenHSiliconSi
BromineBrIodineISilver (Argentum)Ag
CalciumCaIron (Ferrum)FeSodium (Natrium) Na
CarbonCLead (Plumbum)PbSulfurS
ChlorineClMagnesiumMgUraniumU
CobaltCoNeonNeZincZn
Scientists use these symbols instead of full names because they are internationally recognised and allow scientists worldwide to communicate clearly, regardless of language barriers.
Pause and Ponder
8. Imagine you are a scientist who has discovered a new element. Name
this element after yourself and justify that the symbol you have chosen follows the IUPAC rules.9. What problems could arise if every scientist used different symbols for
the same element?
8.5 Atomic Number
You have learnt that the atoms of an element are all alike but different from the atoms of other elements. They differ in the number of electrons and protons in them. The number of protons in the nucleus of an atom of an element is known as its atomic number. It is designated by the symbol Z, this number determines the identity of an element and its chemical behaviour.
Since the atom as a whole is neutral, the number of protons in it equals the number of electrons orbiting the nucleus. For example, hydrogen has one proton and one electron, so its atomic number is 1. Helium, with atomic number 2, has 2 protons and 2 electrons.
A piece of lithiummetal
Can you now say that elements with different atomic numbers are distinct from each other, and the atomic number uniquely identifies an element? Observe Fig. 8.10. How many neutrons and protons are present in a lithium atom, and what is its atomic number?
Lithium atoms
8.6 Mass Number
As stated earlier, helium has two protons, but its mass is about four times that of a proton. The total number of protons and neutrons present in the nucleus of an atom is called its mass number, and is denoted by A. The protons and neutrons present in the nucleus are called nucleons.Mass number = Number of protons + Number of neutrons
NucleusLithium
atom
Electron
Neutrons and protons
Since, the mass of a neutron is roughly equal to that of a proton, the mass of a helium atom can be accounted by two protons and two neutrons. A few examples are given in Table 8.3.
in the nucleus
ElementProtons (p+) Neutrons (n0)Mass number (A)
Hydrogen101
Helium224
Lithium347
Pause and Ponder
The electron, in comparison, has almost negligible mass, and hence, can be ignored in calculations.
10. An atom with an atomic
In the standard notation for an atom, the symbol, the atomic number (Z) and mass number (A) of the element are written as —
number of 26 has 56 nucleons. Find out its number of electrons, protons and neutrons. 11. The nucleus of an atom
Symbol of elementMass Number
contains 20 protons. If its mass number is 41, find the number of neutrons in it.12. An atom has 18 neutrons
Atomic Number
For example, the symbol for carbon is C, its atomic number is 6, and its mass number is 12. In notation, it would be written as —
and an atomic number of 17. What is its mass number?13. An atom 23A has 11 electrons. Find the number of neutrons in it.
6 C
So far, you know that protons and neutrons are located in the nucleus of an atom, while electrons move around the nucleus. Now, let us explore how electrons are arranged around the nucleus.
151Journey Inside the Atom
8.7 How Are Electrons Distributed in Different Energy Levels?
Bohr and Bury suggested the following rules:
y
The maximum number of electrons present in a shell is given by the
formula 2n², where ‘n’ is the number of the shell (Fig. 8.7). Hence, K-shell (n = 1) 2×1² = 2, can accommodate 2 electrons, L-shell (n = 2) 2×2² = 8 electrons, and M-shell (n = 3) 2×3² = 18 electrons. • The maximum number of electrons that can be accommodated in the
outermost shell is 8 (the first shell can accommodate a maximum of two electrons). • Electrons are filled in these shells in a stepwise manner, starting from
the one closest to the nucleus and moving outward, i.e., in the order K, L, M, N, … The L-shell will be filled only after the K-shell is complete, and so on.To understand the placement of electrons in atoms, let us begin with hydrogen (atomic number 1). The only electron in it has to be in the K-shell. The electron distribution for hydrogen (H) is depicted in Fig. 8.11. Helium contains two protons in its nucleus and two electrons. In which way will the two electrons be arranged in its atomic shell?
8.7.1 Building up atomsLet us create two-dimensional (2-D) atomic structures for various elements by adding one electron to the appropriate energy level each time atomic number is increased by 1. The distribution of electrons among various shells is known as the electronic configuration of the atom (Fig. 8.11). Table 8.4 lists the symbols, atomic numbers, number of protons, number of neutrons, number of electrons, and the electron distribution in the shells of the first eighteen elements.
NaMgAlSiPSClAr
Fig. 8.11: Schematic atomic structure of the first eighteen elements showing how the electrons
are filled in the K, L and M shells
electrons, and the electronic distribution of atoms of the first eighteen elements
Distribution
Name of the element SymbolAtomic numberNumber of protonsNumber of neutronsNumber of electrons
KLMN
HydrogenH11-11---
HeliumHe22222---
LithiumLi334321--
BerylliumBe445422--
BoronB556523--
CarbonC666624--
NitrogenN777725--
OxygenO888826--
FluorineF9910927--
NeonNe1010101028--
SodiumNa11111211281-
MagnesiumMg12121212282-
AluminiumAl13131413283-
SiliconSi14141414284-
PhosphorusP15151615285-
SulfurS16161616286-
ChlorineCl17171817287-
ArgonAr18182218288-
Pause and Ponder
14. Identify the number of electrons in the outermost shell of the following elements:
(i) 12
6 C
(ii) 19
9 F
(iii) 28
14 Si
15. Write the electronic configuration of the elements having atomic numbers 12,
16 and 18.
16. Solve this riddle: I am an atom with a mass number of 23 and 11 protons. I
am a soft metal and react vigorously with water. Who am I and how many neutrons do I have? You can also create one such riddle.
153Journey Inside the Atom
8.8 Combining Capacity of an Atom: Valency
You have learnt that atoms of the same or different elements can combine to form molecules. The number of atoms of hydrogen or chlorine with which one atom of an element can combine to form a compound is called its combining capacity. It is expressed in terms of hydrogen and chlorine because both possess a combining capacity of one. For example, in H2O (water), oxygen combines with two hydrogen atoms, so the combining capacity of oxygen is two. In NH3 (ammonia) and MgCl2 (magnesium chloride), what will be the combining capacities of nitrogen and magnesium respectively? These examples show how the combining capacity of elements is determined by the number of hydrogen or chlorine atoms they combine with.
Grade 8CuriosityChapter 7
Let us learn how the combining capacity of an atom is decided by its electronic configuration.
The outermost shell containing electrons of an atom is known as its valence shell. The electrons present in it are known as valence electrons.
If the outermost shell of an atom has 8 electrons, it is called an octet. It has been observed that elements with complete octet of electrons (8 electrons), or 2 electrons in the case of helium in their valence shell are largely unreactive and more stable. On the other hand, atoms with incomplete valence shells are usually more reactive. Such elements lose, gain, or share electrons to complete their octet.
The number of electrons gained, lost, or shared to complete the octet is called the valency of the element. Generally, if the element has fewer than four electrons in its valence shell, it tends to lose electrons to complete its octet and become stable. On the other hand, if the number of valence electrons is more than four, it tends to gain electrons to complete its octet. For example, the electronic configuration of sodium is 2, 8, 1. It can get an octet by losing one electron. Therefore, its valency or combining capacity is 1. On the other hand, oxygen has an electronic configuration of 2, 6. It has six valence electrons. Therefore, it can gain two electrons to attain an octet, its valency is 2. The electronic configuration of carbon is 2, 4. Carbon has four valence electrons and cannot easily gain or lose them. Therefore, it can share four electrons with other atoms to complete its octet. Thus, the valency of carbon is 4. Some compounds appear to violate the usual valency rule, about which you will learn in higher grades.
NextLevel
Up
Can you predict what happens to the atoms that already have eight electrons in their outermost shell (except the elements with one shell only, where only two electrons are possible)? Will they still try to lose or gain electrons?
Examine Table 8.4. Add one more column to it, and write down the common valency of each element.
You have learnt that all the atoms of an element have the same number of electrons and protons, which is equal to its atomic number. Can you say the same regarding the number of neutrons too? Scientists have observed that atoms of the same element can have different numbers of neutrons. What effect does this difference have on the properties of the atom? Let us try to find out!
8.9 A Deeper Look into Atomic Structure
8.9.1 Isotopes
Dalton proposed that all atoms of an element are identical and have the same mass. But scientists later discovered that there were atoms of the same element that could have the same number of protons (atomic number, Z) yet could have different numbers of neutrons, and thus, different mass numbers (A = p+ + n0). These ʻtwin atomsʼ with the same atomic number but different mass numbers are called isotopes.
Let us take hydrogen as an example. Naturally occurring hydrogen is a mixture of three different isotopes: 11 H (protium, ~99.98%), 21 H(deuterium, ~0.015%), and
31 H (tritium, in traces) (Fig. 8.12). All of these contain one proton each, whereas deuterium contains one neutron, and tritium contains two neutrons. Can you guess how many electrons each of these isotopes have?
Fig. 8.12: Schematic representation of isotopes of hydrogen ProtiumDeuteriumTritium
The chemical properties of isotopes are similar. Do you know why? It is so because they have the same number of electrons and the same electronic configuration. As you have learnt, chemical properties depend mainly on the number of valence electrons. All the isotopes will have the same chemical properties although they differ in their physical properties, for example, boiling and melting points.
Ready to Go Beyond
Atoms are too tiny to be weighed in kilograms or grams like everyday objects. Just as it is easier to weigh a grain of wheat in milligrams rather than kilograms, scientists use a special unit called the unified atomic mass unit (u) to measure the mass of atoms. This is because a kilogram is too large for such small particles. Earlier, atomic mass was expressed in atomic mass unit (abbreviated as ‘amu’).
126 C,
136 C, and
146 C(Fig. 8.13). Each of these has six protons and six electrons.
Similarly, carbon has three isotopes:
6 C is the most abundant isotope in nature. You may note that these three isotopes differ in terms of the number of neutrons in them.
Carbon – 12Carbon – 13Carbon – 14
Fig. 8.13: Schematic representation of isotopes of carbon
Let us learn about the essential uses of some isotopes in our daily lives.
155Journey Inside the Atom
Bridging Science and Society
Some isotopes have special properties that are useful in various fields. Some of the applications of isotopes are:
92 U, an isotope of uranium, is used as a fuel in a nuclear reactor to generate electricity in a nuclear power plant (Fig. 8.14).
6027 Co, a radioactive isotope of cobalt, is used in radiation treatment for cancer.
53 I , an isotope of iodine, is used to treat goitre and thyroid cancer.
6 C , an isotope of carbon, is used in archaeology and geology to determine the age of ancient fossils and artefacts.
A. Average atomic mass
Chlorine occurs in nature in two isotopic forms. One isotope has a mass of 35 u, and the other has a mass of 37 u. They occur in the ratio 3:1. This raises an interesting question — should we consider the mass of a chlorine atom to be 35 u or 37 u?
If the atomic mass of a natural element is taken as the average mass of all its naturally occurring isotopes, it would be the simple arithmetic mean of the masses of its isotopes, calculated without accounting for the relative abundances of each isotope. For chlorine, with isotopes ³⁵ Cl and ³⁷ Cl, the simple average would be —
3537Average atomic mass = 2+
= 36 u
However, this does not accurately reflect nature, since isotopes do not occur in equal ratios. The more common isotope is ³⁵ Cl, which constitutes about 75% while the other isotope, ³⁷ Cl makes up about 25%. The accurate average atomic mass can be calculated by considering their natural abundances. This is called a weighted average atomic mass, which can be calculated by multiplying the mass of each isotope by its percent relative abundance and then adding the two values, as explained below.
Mathematically,
75253537100100=×+×
10537u44=+
142 u4=
35.5 u=
Meet a Scientist
This does not mean that any one atom of chlorine has a fractional mass of 35.5 u. It means that if you take, say 1 million chlorine atoms, they will contain 7.5 lakh 3517 Cl and 2.5 lakh 3717 Cl atoms with a weighted average atomic mass of 35.5 u.
Homi Jehangir Bhabha was an Indian physicist. He is known as the father of the Indian nuclear programme. He made a pioneering contribution to the development of atomic energy in India. He established key institutions like the Tata Institute of Fundamental Research (TIFR) and the Bhabha Atomic Research Centre (BARC), for peaceful uses of atomic energy to generate electricity, support agriculture, and advanced medical treatments.
This comparison shows that the simple average ignores abundance, whereas the weighted average accurately reflects the mass of the element as it occurs in nature.
Pause and Ponder
17. Two different atoms have 11 protons each, but one has
12 neutrons, and the other has 13 neutrons. How do their atomic numbers and mass numbers compare? Are they the same element or different elements?18. If a bromine atom is available in the form of, say two
7935 Br (49.7%) and
8135 Br (50.3%), calculate the average atomic mass of the bromine atom.
isotopes,
You have learnt that the two atoms that have the same atomic number but different mass numbers are called isotopes. What if they have the same mass number but different atomic numbers? What are such atoms called? Let us find out!
Ready to Go Beyond
Electron microscopes, such as Scanning Tunnelling Microscopes (STMs) and Transmission Electron Microscopes (TEMs) can produce images of materials with atomic-level details. STMs mainly study surfaces (Fig. 8.15), while TEMs reveal how atoms are arranged inside very thin samples.
8.9.2 Isobars
Let us consider three elements — calcium (atomic number 20), potassium (atomic number 19), and argon (atomic number 18). These elements have different numbers of protons, yet each has a mass number of 40. This shows that the total number of nucleons in their atoms is the same, even though they are different elements. When atoms of different elements have the same mass number, but different atomic numbers, they are called isobars.
As we end our journey into the structure of the atom, it is important to know that the story does not end here. The exploration of the structure of the atom is still being discovered (Fig. 8.16). Later, scientists discovered that even Bohr’s model was not entirely correct. Electrons do not follow well-defined paths like the fixed Bohr orbits. Today, we understand that they exist as ʻelectron cloudsʼ around the nucleus. We can predict regions where they are most
Fig. 8.15: STM image showing individual atoms on a surface
likely to be, not exactly where they are. You will learn all these details in higher grades. The journey of exploring the mysteries of the atom is far from over, and exciting discoveries still lie ahead!
NextLevel
Up
157Journey Inside the Atom
Dalton’s Atomic Model
Thomson’s Atomic Model
Rutherford’s Atomic Model
Bohr’s Atomic Model
Modern Atomic Model
Still being discovered
Atom as indivisible
Positive and negative charges
The nuclear
Energy levels
Quantum mechanical
model
particle
model
Fig. 8.16: Journey of the development of atomic models
At a Glance
y
Atoms are the building blocks of matter. • J. J. Thomson proposed that in an atom, electrons are embedded in
a positively charged sphere. • Rutherford’s model described the atom as mostly empty space,
with a dense, positively charged nucleus at its centre and electrons orbiting it. • Niels Bohr’s model proposed that electrons move in fixed energy
levels (shells) around the nucleus. • The shells of an atom are named as K, L, M, N, and so on. • James Chadwick discovered the presence of neutrons in the atom. • The three subatomic particles of an atom are electrons, protons,
and neutrons. • If the outermost shell of an atom has an octet of electrons (or
two electrons in case of helium), the atom is stable and largely unreactive. • Valency is the combining capacity of an atom. It is equal to the
number of electrons which can be gained, lost or shared by an atom to achieve a stable configuration. • The atomic number of an element is equal to the number of
protons in its nucleus. • The mass number of an atom is equal to the total number of
nucleons (protons and neutrons) in its nucleus. • Isotopes are atoms of same element that have the same atomic
number but different mass numbers. • The average atomic mass of an element is calculated based on
the relative abundance of its isotopes in nature. • Isobars are atoms of different elements with the same mass
number but different atomic numbers.