Atomic Structure and Chemical Bonding Class 9 ICSE
Atomic structure and chemical bonding class 9 ICSE is Chapter 4 of your Chemistry textbook. It explains how atoms are built and why they join together. You will meet the three subatomic particles, the experiments that revealed them, and the two main kinds of bonds. This walkthrough follows the chapter section by section, with the formulas and worked patterns you need, and you can browse the full set of Class 9 Chemistry chapters on this site alongside it.
Download Atomic Structure and Chemical Bonding Class 9 ICSE Chemistry Chapter 4 PDF
Download the Atomic Structure and Chemical Bonding Class 9 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 26 pages, 14 sections.
What is inside Chapter 4
26 pages · 14 sections · 1 worked examples
| Section | Title |
|---|---|
| 4.2 | DEFINITION OF AN ATOM |
| 4.3 | CONSTITUENTS OF AN ATOM |
| 4.5 | DISCOVERY OF NUCLEUS |
| 4.5.2 | Drawback of Rutherford's atomic model |
| 4.6 | BOHR'S ATOMIC MODEL |
| 4.7 | DISCOVERY OF NEUTRONS |
| 4.8 | ATOM — ITS STRUCTURE |
| 4.9 | ATOMIC NUMBER [Z] |
| 4.10 | DISTRIBUTION OF ELECTRONS IN THE ORBITS — BOHR-BURY SCHEME |
| 4.11 | VALENCE ELECTRONS |
| 4.12 | REASON FOR CHEMICAL ACTIVITY OF AN ATOM |
| 4.13 | ISOTOPES |
| 4.14.1 | Structures of some electrovalent compounds |
| 4.15 | COVALENT (MOLECULAR) BOND |
4.2 Definition of an atom
An atom is the smallest particle of an element that still shows all the properties of that element. It may or may not exist on its own, but it takes part in every chemical reaction.
Think of grinding copper filings into finer and finer powder. Each new powder still behaves like copper, until you reach particles that cannot be split further without losing copper’s properties. Those indivisible particles are atoms of copper.
The definition does not claim atoms are unbreakable. Later sections show atoms are divisible into protons, neutrons and electrons.
4.3 Constituents of an atom
A discharge tube is a glass tube with two metal electrodes, one positive (the anode) and one negative (the cathode). Gases are poor conductors at ordinary pressure. When the pressure drops to about 0.01 mm Hg and a high voltage near 10,000 volts is applied, the gas conducts and rays flow from cathode to anode. These are cathode rays.
Cathode rays travel in straight lines and bend toward the positive field, so they carry a negative charge. Their charge-to-mass ratio is fixed at 1.76 \times 10^{11}\ \text{C/kg}, whatever gas is used. Thomson concluded these are negatively charged particles, called electrons, present in all atoms. The electron is written _{-1}e^{0}, the subscript being the charge and the superscript the mass.
Goldstein used a perforated cathode and found rays moving from anode to cathode, called canal, positive or anode rays. These contain protons, particles of mass 1 amu and unit positive charge, written \mathrm{P}^{1}. Thomson then proposed the plum pudding model: a sphere of uniform positive charge with electrons embedded in it, so the atom stays neutral.
4.5 Discovery of nucleus
In 1911 Rutherford fired alpha particles at a thin gold foil. Alpha particles are doubly charged helium ions, \mathrm{He}^{2+}, and gold was chosen because it can be beaten into an extremely thin layer.
Three observations came out. Most alpha particles passed straight through. Some were slightly deflected. Very few, about one in ten thousand, bounced back.
Each result points to a conclusion. The empty space means the atom is mostly hollow. A small, dense, positively charged nucleus holds nearly all the mass. Electrons circle it in orbits, much as planets circle the Sun. Light nuclei such as lithium behave differently: fast alpha particles may push the light nucleus aside instead of being deflected back.
The model has one drawback. Classical physics says a moving charged particle radiates energy. An orbiting electron should lose energy, spiral into the nucleus, and the atom should collapse. Atoms are stable, so the model cannot be right.
4.6 Bohr’s atomic model
In 1913 Niels Bohr answered the stability problem with three postulates. Electrons move only in fixed circular orbits called shells. While in an orbit, an electron neither gains nor loses energy. An electron jumps to another orbit only when it gains or loses a definite amount of energy.
Each orbit holds a fixed amount of energy, so Bohr called them energy levels. They are labelled K, L, M, N or I, II, III, IV. The K shell is closest to the nucleus and holds the least energy.
Bohr’s fixed orbits explain why the electron does not spiral inward, removing the drawback Rutherford could not solve.
4.7 Discovery of neutrons
The neutron was predicted before it was found. A helium atom has 2 protons, so its mass should be 2 \times 1 = 2 a.m.u. The measured mass is near 4 a.m.u., so another particle of proton mass and no charge must exist.
In 1932 Chadwick bombarded beryllium with alpha particles:
{}_{4}\mathrm{Be}^{9} + {}_{2}\mathrm{He}^{4} \rightarrow {}_{6}\mathrm{C}^{12} + {}_{0}n^{1}
The neutral particle released is the neutron, written _{0}n^{1}. Neutrons are not deflected by electric or magnetic fields. Their mass is about 1 a.m.u., slightly more than a proton’s.
This discovery also corrected Dalton’s atomic theory. Atoms are divisible into subatomic particles, and isotopes show that atoms of one element are not all alike.
4.8 Atom — its structure
An atom has two structural parts: the nucleus and the orbits around it. Protons and neutrons, together called nucleons, sit in the nucleus. Electrons move in imaginary circular paths, the orbits.
| Particle | Symbol | Charge | Mass (a.m.u.) |
|---|---|---|---|
| Electron | _{-1}e^{0} | −1 | 1/1840 |
| Proton | _{1}^{1}H | +1 | 1 |
| Neutron | _{0}^{1}n | 0 | 1 |
The electron’s mass is about 9.1 \times 10^{-28} g, while the proton and neutron each weigh about 1.6 \times 10^{-24} g. Hydrogen is the one exception among elements: its atom has no neutron.
4.9 Atomic number [Z]
The atomic number Z is the number of protons in the nucleus. In a neutral atom it also equals the number of electrons and the positive charge of the nucleus, so Z = p = e.
The mass number A is the total of protons and neutrons, A = n + p. The neutron count follows from n = A - Z. The standard symbol ^{A}_{Z}X puts the mass number as the superscript and the atomic number as the subscript. Mass number is a whole-number approximation of atomic mass.
Worked pattern: an atom with atomic number 20 and mass number 42 has 20 protons, 20 electrons and 42 - 20 = 22 neutrons.
A quick check of common slips:
| Mistake | Correct rule |
|---|---|
| Confusing atomic number with mass number | Z counts protons only; A counts protons plus neutrons |
| Miscounting neutrons | Always A - Z, never Z - A |
| Writing the octet rule for helium | Helium is stable with a duplet of 2 electrons |
4.10 Distribution of electrons in the orbits — Bohr-Bury scheme
Electrons fill shells according to the 2n^{2} rule, where n is the shell number. So K holds 2 \times 1^{2} = 2, L holds 2 \times 2^{2} = 8, M holds 2 \times 3^{2} = 18 and N holds 2 \times 4^{2} = 32.
Three filling rules apply. Inner shells fill before outer ones. The outermost shell never holds more than 8 electrons, even if its capacity is larger. A stable outer shell holds 8 electrons, except helium, which holds 2.
Magnesium, atomic number 12, has 12 protons and 24 - 12 = 12 neutrons. Its 12 electrons fill as 2, 8, 2. Potassium, atomic number 19, is 2, 8, 8, 1, and calcium, atomic number 20, is 2, 8, 8, 2. In elements heavier than argon, the fourth shell starts filling once the third shell has only 8 electrons, even though M can hold 18.
The pattern for elements 1 to 20 runs from hydrogen (1) to calcium (2, 8, 8, 2).
4.11 Valence electrons
The outermost shell is the valence shell, and the electrons in it are valence electrons. They decide an element’s chemical properties, because they are the ones that take part in reactions.
Elements with 1, 2 or 3 valence electrons are metals. They lose electrons to form positive ions, or cations, and act as reducing agents. Sodium shows this:
\mathrm{Na} - \mathrm{e}^{-} \rightarrow \mathrm{Na}^{+},\quad (2, 8, 1) \rightarrow (2, 8)
Elements with 4, 5, 6 or 7 valence electrons are non-metals. They gain electrons to form negative ions, or anions, and act as oxidising agents:
\mathrm{Cl} + \mathrm{e}^{-} \rightarrow \mathrm{Cl}^{-},\quad (2, 8, 7) \rightarrow (2, 8, 8)
Noble gases, with 8 valence electrons, and helium, with 2, are inert. Valency equals the number of valence electrons for 1, 2 or 3, and equals 8 - (valence electrons) for 4, 5, 6 or 7.
One exam trap is applying this valency rule to a shell that is not the outermost one. Valency always refers to the valence shell.
4.12 Reason for chemical activity of an atom
All noble gases except helium have 8 electrons in their outermost shell, an arrangement called an octet. Helium is stable with 2 electrons, a duplet, in its only shell.
Atoms of other elements have incomplete valence shells. They combine to reach the stable configuration of the nearest noble gas, by gaining, losing or sharing electrons. This is the octet rule; for helium and hydrogen the duplet rule applies.
The force that binds atoms together as a molecule is called a chemical bond. Kossel and Lewis put forward this electronic theory of valency in 1916.
Edge case: the octet rule does not apply to helium. Its single shell can hold only 2 electrons, so it follows the duplet rule instead.
4.13 Isotopes
Isotopes are atoms of the same element with the same atomic number but different mass numbers. The difference comes from the neutron count.
Hydrogen has three isotopes: protium (0 neutrons), deuterium (1 neutron) and tritium (2 neutrons), all with atomic number 1. Chlorine has ^{35}_{17}\mathrm{Cl} with 18 neutrons and ^{37}_{17}\mathrm{Cl} with 20.
Isotopes share chemical properties, because their electronic configurations are identical. They differ in physical properties such as density and boiling point. Some, like tritium and carbon-14, are radioactive.
Fractional atomic mass arises because an element’s atomic mass is a weighted average of its isotopes. If a halogen such as bromine occurs equally as mass 79 and mass 81, its average is:
\frac{79 \times 1 + 81 \times 1}{2} = 80
Uses include cobalt-60 in radiotherapy, carbon-14 for dating, uranium-235 as nuclear fuel and iodine-131 for goitre. Isobars are different atoms with the same mass number but different atomic numbers, such as argon-40 and calcium-40.
4.14.1 Structures of some electrovalent compounds
An electrovalent or ionic bond forms when electrons transfer from a metal atom to a non-metal atom. The donor becomes a cation, the acceptor an anion, and their opposite charges attract.
In sodium chloride, sodium (2, 8, 1) loses one electron to become \mathrm{Na}^{+} (2, 8), like neon. Chlorine (2, 8, 7) gains one electron to become \mathrm{Cl}^{-} (2, 8, 8), like argon.
In magnesium chloride, one magnesium atom gives one electron to each of two chlorine atoms. The ratio is 1:2, so the formula is \mathrm{MgCl}_{2}. In calcium oxide, calcium gives 2 electrons to a single oxygen atom, giving \mathrm{CaO}, not \mathrm{Ca}_{2}\mathrm{O}_{2}.
The Lewis or electron dot symbol shows the element symbol surrounded by dots for valence electrons. In the transfer, the electropositive atom is oxidised and the electronegative atom is reduced, together a redox process. Ionic compounds stay stable because the attraction between opposite charges beats the repulsion between like charges.
4.15 Covalent (molecular) bond
A covalent bond forms when two atoms share one or more pairs of electrons, each atom contributing an equal number. Non-metals share rather than transfer electrons, because losing several electrons costs too much energy.
Bonds come in three types. A single bond is one shared pair, written \mathrm{H}-\mathrm{H}. A double bond is two shared pairs, written \mathrm{O}=\mathrm{O}. A triple bond is three shared pairs, written \mathrm{N}\equiv\mathrm{N}.
Covalency is the number of electrons an atom shares. Hydrogen has covalency 1, oxygen 2, nitrogen 3 and carbon 4.
Covalent compounds are non-polar when the shared pair sits equally between identical atoms, as in \mathrm{H}_{2}, \mathrm{Cl}_{2} and \mathrm{O}_{2}. They are polar when the pair is drawn toward one atom, creating fractional charges, as in hydrogen chloride. Hydrogen shares with chlorine for HCl, two hydrogens share with one oxygen for water, nitrogen shares three pairs for ammonia, and carbon shares four pairs for methane.
| Property | Electrovalent compound | Covalent compound |
|---|---|---|
| Bond formed by | Transfer of electrons | Sharing of electrons |
| Typical partners | Metal + non-metal | Non-metal + non-metal |
| Ions present | Cations and anions | No ions |
| Example | NaCl, MgCl₂ | HCl, H₂O |
Across period 3, the electrovalent character of chlorides and oxides decreases from NaCl to SiCl₄.
What to practise from Atomic Structure and Chemical Bonding Class 9 ICSE
Focus your practice on the chapter’s named exercise block and the embedded question sets.
- Exercise 4(C) carries 24 questions covering atomic structure, isotopes, ions, valency and bonding.
- Shorter question sets sit directly after sections 4.3, 4.10, 4.13, 4.14.1 and 4.15. Work through these as you finish each section.
ICSE-style questions on this chapter usually give an element’s atomic number and mass number, then ask for the electronic configuration and valency, or ask you to classify a bond as electrovalent or covalent from the elements involved. The highest-value practice is drawing electron arrangements, writing electronic configurations, and classifying bonds. You can verify the syllabus for this chapter on the official CISCE website, and continue with the next topic in the Class 9 periodic table chapter.
Atomic Structure and Chemical Bonding Class 9 ICSE frequently asked questions
Which topics in atomic structure and chemical bonding class 9 ICSE are most important for the exam?
The core topics are the three subatomic particles, atomic number and mass number, electronic configuration, isotopes, and the octet rule. Valency and the difference between electrovalent and covalent bonds are the most tested ideas. Master neutron counting from A - Z and the 2n^{2} filling rule.
How do I find the number of neutrons from the mass number and atomic number?
Subtract the atomic number from the mass number, n = A - Z. An atom with mass number 42 and atomic number 20 has 42 - 20 = 22 neutrons. The atomic number gives the protons and, in a neutral atom, the electrons too.
Why is the atomic mass of chlorine 35.5 instead of a whole number?
Chlorine’s atomic mass is the weighted average of its isotopes. Natural chlorine holds ^{35}_{17}\mathrm{Cl} and ^{37}_{17}\mathrm{Cl} in about a 3:1 ratio, so the average is (35 \times 3 + 37 \times 1)/4 = 35.5.
What is the difference between an electrovalent bond and a covalent bond in this ICSE chapter?
An electrovalent bond forms when electrons transfer from a metal to a non-metal, making cations and anions, as in NaCl. A covalent bond forms when non-metals share electron pairs, as in HCl and H₂O.
Why does helium need only 2 electrons to be stable in Class 9 ICSE Chemistry?
Helium has only one shell, the K shell, which holds a maximum of 2 electrons by the 2n^{2} rule. With a full duplet, helium is stable and does not follow the octet rule.
Reference: ICSE Class 9 Chemistry textbook material, chapter 4: Atomic Structure and Chemical Bonding — CISCE-aligned.