Revision: Atomic Structure and Chemical Bonding Chemistry (English Medium) ICSE Class 9 CISCE

Reduction is defined as the phenomenon in which an atom gains an electron to form a negatively charged ion called an anion.
During the formation of ionic bond one atom undergoes oxidation while another atom undergoes reduction.

Oxidation is the loss of electrons during a reaction by a molecule, atom or ion. In terms of electron transfer, oxidation is defined as the phenomenon in which an atom loses an electron to form a positively charged cation.
During the formation of ionic bond one atom undergoes oxidation while another atom undergoes reduction.

A chemical bond may be defined as the force of attraction between any two atoms in a molecule to maintain stability.

or

A chemical bond may be defined as the force of attraction between any two atoms, in a molecule, to maintain stability.

Element is a substance which cannot be split up into two or more simple substances by usual chemical methods of applying heat, light or electric energy; for example, hydrogen, oxygen and chlorine.

Molecule: A molecule can be defined as the smallest unit of an element or a compound which exhibits all the properties of that element or compound and has an independent existence. They are divisible into atoms.

Compounds are pure substances composed of two or more elements in definite proportion by mass and has properties, entirely different from those of its constituents elements.
Compound, are made up of different types of atoms combined chemically.

Elements: An element is defined as a pure substance made up of only one kind of atoms that cannot be converted into anything simpler than itself by any physical or chemical process. 

“Mixtures can be defined as. a kind of matter which is formed by mixing two or more pure substances (elements and compounds) in any proportion, such that they do not undergo any chemical change and retain their individual properties. Therefore they are impure substances.

The number of atoms in a molecule of an element is called its atomicity. 

Element is a substance which cannot be broken further into simpler substances and has a definite set of properties. Elements are made up of only one kind of atoms.

Atom: An atom is the smallest indivisible unit of an element which exhibits all the properties of that element and may or may not have an independent existence. An atom is the smallest indivisible unit of an element which exhibits all the properties of that element and may or may not have an independent existence. 

Formula: Formula is a short way of representing the molecule of an element or a compound

Relative atomic mass— Relative atomic mass is the mass of an atom of an element as a multiple of the standard atomic mass unit.

The relative atomic mass of an element is the ratio between the average mass of its isotopes to 1/12th part of the mass of a carbon – 12 atoms. It is denoted as A_r.

Relative atomic mass = `" Average mass of the isotopes of the element"/(1"/"12^{"th"}" of the mass of one Carbon- 12 atom")`

Compound: A compound is a pure substance that is formed when the atoms of two or more elements combine chemically in definite proportions.

Ex: H20, NaCl.

Non-Metal: Non-metal is an element that doesn’t have the characteristics of metal including, (i.e.) ability to conduct heat or electricity luster or flexibility.

Ex. Carbon Iodine, Sulphur.

Mass number— Mass number is the sum of the number of protons and neutrons present in the nucleus of an atom. It is denoted by A.

An atom which becomes charged by losing or gaining electrons is called an ion.

Atom: An atom is the smallest indivisible unit of an
OR
Atom is the smallest unit of matter.

Molecule : Molecule is the smallest unit of a compound (or an element) which always has an independent existance.

Covalent bond— When atoms of different non-metals neither donate nor accept electrons and hence no ions are formed, such a bond is called covalent bond.

Metal:  A chemical element that is an effective conductor of electricity and heat can be defined as a metal.

Ex.: Copper, Iron, Silver, etc.

Metalloid: Metalloid is a chemical element that exhibits some properties of metals and some of non-metals. Metalloids are generally semi-conductors.

Ex.: Silicon. Arsenic, Antimony and Boron.

An atom is the smallest particle of a chemical element that retains its chemical properties.

Chemical bond— A chemical bond is the binding force between two or more atoms of a molecule.

Element: It is a substance that cannot be broken down into simpler substance by chemical means

Ex.: Oxygen, Hydrogen, Gold & Helium.

An atom is the smallest particle of an element which retains its chemical identity in all physical and chemical changes.

Radicals : A radical is an atom of an element or a group of atoms of different elements that behaves as a single unit with a positive or negative charge on it.

An Atom: Smallest particle of an element that can exist and have properties of an element.

The protons and neutrons collectively are known as nucleons.

Atomic number refers to the number of protons present in an atom. It is denoted by Z. Example: An atom of oxygen contains 8 proton Therefore its atomic number is 8.

Mass number refers to the sum of the number of protons and neutrons present in the nucleus of an atom and denoted by A Mass number = Number of protons + Number of neutrons.

The outermost shell of an atom is known as its valence shell.

The number of protons in the nucleus is known as the atomic number of the element and is denoted by Z.

The number of protons in the nucleus of an atom, which is characteristic of a chemical element and determines its place in the periodic table. Atomic number is also equal to the number of electrons in an atom.

The mass number of an atom is equal to the total number of nucleons (i.e., the sum of the number of protons and the number of neutrons) in its nucleus.

The atomic number of an atom is equal to the number of protons in its nucleus (which is same as the number of electrons in a neutral atom).

The total number of neutrons and protons in the nucleus is called the mass number of the element and is denoted by A.

Relative atomic mass is defined as the ratio of the average atomic mass to the unified atomic mass unit.

Relative atomic mass (A_r) = `"Average mass of the atom"/"Unified atomic mass"`

The mass of a single atom of an element is called the atomic mass.

The valency of an element is determined by the number of electrons present in the outermost shell of its atoms, that is, the valence electrons.

When the properties of elements in a period or a group of the modern periodic table are compared, certain regularity is observed in their variations. It is called the periodic trends in the modern periodic table.

The atoms of the same element, having same atomic number Z, but different mass number A, are called isotopes.

OR

Atoms having the same atomic number (Z) but different mass numbers (A).

The number of electrons that an atom of an element loses or gains to form a electrovalent bond is called its electrovalency.

The chemical compounds formed as a result of the transfer of electrons from one atom of an element to one atom of another element are called ionic (or electrovalent) compounds.

An ion is a charged particle which is formed due to the gain or the loss of one or more electrons by an atom.

A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation.

A non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion.

A metallic element, whose one atom readily loses electron(s) to form a positively charged ion, is an electropositive element.

A non-metallic element, whose atom readily accepts electron(s) to form a negatively charged ion, is an electronegative element.

The cation and the anion being oppositely charged attract each other and form a chemical bond. Since this chemical bond formation is due to the electrostatic force of attraction between a cation and an anion, it is called an electrovalent (or an ionic) bond.

An ion is a charged particle which is formed due to the gain or the loss of one or more electrons by an atom.

A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation.

A non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion.

A metallic element, whose one atom readily loses electron(s) to form a positively charged ion, is an electropositive element.

A non-metallic element, whose atom readily accepts electron(s) to form a negatively charged ion, is an electronegative element.

The cation and the anion being oppositely charged attract each other and form a chemical bond. Since this chemical bond formation is due to the electrostatic force of attraction between a cation and an anion, it is called an electrovalent (or an ionic) bond.

The number of electrons that an atom of an element loses or gains to form a electrovalent bond is called its electrovalency.

The chemical compounds formed as a result of the transfer of electrons from one atom of an element to one atom of another element are called ionic (or electrovalent) compounds.

The chemical bond formed between two combining atoms by mutual sharing of one or more pairs of electrons is called a covalent bond.

The molecule formed due to the sharing of electrons (covalent bond) is called a covalent molecule.

The chemical bond that is formed between two combining atoms by mutual sharing of one or more pairs of electrons is called a covalent (or a molecular) bond, and the compound formed due to this bond is called a covalent compound.

The bond formed between two atoms by sharing a pair of electrons, provided entirely by one of the combining atoms but shared by both, is called a coordinate bond. 

The covalency of an atom is the number of its electrons taking part in the formation of shared pairs.

\[L=mvr=\frac{nh}{2\pi},\quad n=1,2,3\ldots\]

\[h\nu=E_2-E_1=\frac{hc}{\lambda}\]

The average atomic mass accounts for the different isotopes of an element and their natural abundances.

\[M_{\mathrm{avg.}}=\frac{M_{1}\times r_{1}+M_{2}\times r_{2}+M_{3}\times r_{3}}{r_{1}+r_{2}+r_{3}}\]

where M₁, M₂, M₃ are atomic masses of isotopes and r₁, r₂, r₃ are their relative abundances.

Bohr's First Postulate:
An atom consists of a small, massive central core called the nucleus, around which planetary electrons revolve. The centripetal force required for their rotation is provided by the electrostatic attraction between the electrons and the nucleus.

Bohr's Second Postulate (Quantum Condition):
The electrons are permitted to circulate only in those orbits in which the angular momentum of an electron is an integral multiple of \[\frac{h}{2\pi}\]; h being Planck's constant.

Bohr's Third Postulate:
While revolving in the permissible orbits, an electron does not radiate energy. These non-radiating orbits are called stationary orbits.

Bohr's Fourth Postulate:
An atom can emit or absorb radiation in the form of discrete energy photons only when an electron jumps from a higher to a lower orbit or from a lower to a higher orbit, respectively.

• The word atom comes from the Greek word atomos meaning uncuttable or indivisible.
• Generally, the size of an atom is about 10⁻¹⁰ m (this distance is the average distance between the nucleus and the outermost shell carrying electrons).
• Atomic number (Z) = Number of protons = Number of electrons (in neutral atom).
• Atomic mass = Number of neutrons + Number of protons.
• Atomic mass = Equivalent mass × Valency = 6.4 × Specific heat (cal).

Dalton's atomic theory laid the foundation of modern chemistry with four core postulates:

1. All matter is made up of extremely small particles called atoms.
2. Atoms of the same element are identical to each other in mass and properties; atoms of different elements differ.
3. Atoms can neither be created nor destroyed — they are indestructible.
4. Atoms combine in fixed, simple whole-number ratios to form compound atoms (molecules).

Note: Modern discoveries have refined some postulates (e.g., isotopes show atoms of the same element can differ in mass), but the core framework remains foundational.

• Discovered by J. J. Thomson (1897) using the cathode ray tube experiment
• Charge = –1.6 × 10⁻¹⁹ C
• Mass = 9.109 × 10⁻³¹ kg (very small)
• Cathode rays travel in straight lines
• Deflected by electric and magnetic fields (proves negative charge)

• Discovered by E. Goldstein using a discharge tube (canal rays)
• Charge = +1.6 × 10⁻¹⁹ C
• Mass = 1.673 × 10⁻²⁷ kg
• Present in the nucleus of an atom
• Determines the atomic number (Z) of an element

• Discovered by James Chadwick (1932)
• Charge = 0 (neutral)
• Mass = 1.675 × 10⁻²⁷ kg (almost equal to proton)
• Present in the nucleus along with protons
• Responsible for isotopes and atomic mass

Reaction: \[_4^9Be+_2^4He\to_6^{12}C+_0^1n\]

• Proposed by J. J. Thomson in 1904 after the discovery of electrons.
• The atom is a uniform sphere of positive charge.
• Electrons are embedded within this sphere.
• The positive charge is spread evenly throughout the atom.
• Total positive charge = total negative charge, so the atom is neutral.
• The model explained the presence of electrons in atoms.
• It did not include a nucleus in the atom.
• It failed to explain Rutherford’s results from the gold foil experiment.

• Proposed by Ernest Rutherford in 1911 based on the gold foil (α-particle scattering) experiment.
• Most α-particles passed straight through, showing that the atom is mostly empty space.
• Some α-particles were deflected, indicating the presence of a positively charged centre.
• Very few α-particles were deflected at large angles or bounced back, proving a dense nucleus.
• All the positive charge and most of the mass are concentrated in a tiny nucleus (~10⁻¹⁵ m).
• Electrons revolve around the nucleus in circular orbits.
• The electrostatic force of attraction between nucleus and electrons keeps them in orbit.
• Limitation: Could not explain stability of atom and line spectra of hydrogen.

• Bohr modified Rutherford's model - electrons move in fixed orbital shells, each with fixed energy levels.
• The centripetal force for electron revolution is provided by electrostatic attraction between the electron and the nucleus.
• An electron does not radiate energy while revolving in a stationary orbit.
• Energy is emitted or absorbed only during electron transitions between orbits.
• Limitations of Bohr's Model:
• Fails to explain the Zeeman Effect (effect of high magnetic fields on atomic spectra).
• Contradicts the Heisenberg Uncertainty Principle.
• Unable to explain the spectra of larger/multi-electron atoms.

• The structure of an atom and its nucleus was developed from the discovery of electrons by J.J. Thomson and alpha particle scattering experiments by Rutherford.
• An atom consists of electrons, protons, and neutrons, with protons and neutrons in the nucleus and electrons revolving in stationary orbits.
• The maximum number of electrons in a shell is given by 2n², and the shells are named K, L, M, N, O, P, and Q.

Isotopes are atoms of the same element that have the same atomic number but different mass numbers (different number of neutrons).

Same in isotopes:

• Atomic number (Z)
• Number of protons and electrons
• Electronic configuration
• Position in periodic table
• Chemical properties (nearly identical)

Different in isotopes:

• Mass number (A)
• Number of neutrons
• Physical properties

Examples: \[_1H^1and_1H^2\]

An ionic bond is formed by the complete transfer of one or more electrons from an electropositive atom to an electronegative atom, resulting in oppositely charged ions that attract each other.

Key conditions for ionic bond formation:

• One atom must have low ionisation enthalpy (easily loses electron) — typically a metal

• The other must have high electron affinity (easily gains electron) — typically a non-metal

• Large difference in electronegativity between the two atoms

Example: Na + Cl → Na⁺ + Cl⁻ → NaCl

• Sodium (2,8,1) loses 1 electron → Na⁺ (2,8)

• Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8)

Ionic solids are crystalline structures containing cations and anions held together by strong electrostatic ionic bonds.

An ionic bond is formed by the complete transfer of one or more electrons from an electropositive atom to an electronegative atom, resulting in oppositely charged ions that attract each other.

Key conditions for ionic bond formation:

• One atom must have low ionisation enthalpy (easily loses electron) — typically a metal

• The other must have high electron affinity (easily gains electron) — typically a non-metal

• Large difference in electronegativity between the two atoms

Example: Na + Cl → Na⁺ + Cl⁻ → NaCl

• Sodium (2,8,1) loses 1 electron → Na⁺ (2,8)

• Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8)

Ionic solids are crystalline structures containing cations and anions held together by strong electrostatic ionic bonds.

• Carbon forms covalent bonds by sharing electrons to achieve a noble gas configuration.
• Covalent bonds can be single, double, or triple, as seen in molecules like H₂, O₂, and N₂.
• Covalent compounds have low melting and boiling points and are poor conductors of electricity.
• Carbon has allotropes such as diamond, graphite, and fullerene (C₆₀), each with different physical properties.

• A single covalent bond involves sharing one pair of electrons; seen in H₂, Cl₂, CH₄, NH₃, and H₂O.
• Double bond shares two pairs of electrons (e.g., O₂, CO₂); triple bond shares three pairs (e.g., N₂, C₂H₂).
• Molecules like ethene (C₂H₄) and ethyne (C₂H₂) have combinations of single and double/triple bonds.

• Atoms involved in covalent bonding must be non‑metals.
• High electronegativity in both atoms favors covalent bond formation.
• High electron affinity in both atoms helps them attract shared electrons.
• High ionization energy makes atoms less likely to lose electrons, supporting electron sharing instead.
• The electronegativity difference between the two atoms should be zero or very small for a covalent bond to form.