Revision: Atomic Structure Chemistry ICSE ICSE Class 8 CISCE

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.

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.

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 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.

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

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

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 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 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.

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

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

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.

• 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.

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

• 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: \[_1H^1and_1H^2\]