Structure of the Atom

Structure of the Atom - Class 9 Science

  • Atoms
  • Subatomic Particles
  • The Structure of an Atom
  • Thomson's Plum Pudding Model
  • Rutherford's Nuclear Model of an Atom
  • Bohr's Model of an Atom
  • Arrangement of Electrons in an Atom
  • Valency
  • Atomic and Mass Numbers
  • Isotopes, Isobars and Isotones
  • Solved Questions on the Structure of the Atom
  • Atoms

    An atom is the basic unit of matter, representing the smallest particle of an element that retains the chemical properties of that element. Everything around us is made up of atoms, from the air we breathe to the water we drink and the metals we use. Atoms are incredibly small, on the order of picometers (10-12 meters) in size.

    Subatomic Particles

    Atoms are composed of three primary subatomic particles: protons, neutrons, and electrons.


    1. Protons are positively charged particles found within the nucleus of an atom. Each proton carries a fundamental positive charge of approximately +1.6 x 10-19 coulombs.
    2. The number of protons in an atom determines its atomic number, which defines the element itself. For example, an atom with 1 proton is hydrogen, while an atom with 6 protons is carbon.
    3. Protons contribute to the mass of the atom and are held together in the nucleus by a strong nuclear force.


    1. Neutrons are electrically neutral particles also located within the nucleus of an atom.
    2. They have a similar mass to protons, approximately 1.67 x 10-27 kg.
    3. Neutrons play a crucial role in maintaining the stability of the nucleus by counteracting the repulsive forces between positively charged protons. Their presence or absence can affect an atom's stability and isotopic properties.


    1. Electrons are negatively charged particles that orbit the nucleus in energy levels or electron shells.
    2. They have a much smaller mass compared to protons and neutrons, approximately 9.1 x 10-31 kg.
    3. Electrons are organised into different energy levels, with the innermost levels being lower in energy and closer to the nucleus. The outermost level, known as the valence shell, contains electrons that are involved in chemical bonding and interactions.

    Comparing the Properties of Electrons Protons and Neutrons






    J.J. Thomson

    E. Goldstein

    J. Chadwick


    -1.6 x 10-19 C

    +1.6 x 10-19 C



    9.1 x 10-31 kg

    1.672 x 10-27 kg

    1.67 x 10-27 kg


    Inside nucleus

    Inside nucleus

    Outside nucleus





    Role in Atom

    Determines chemical properties

    Determines the atom's identity

    Adds to the mass of the nucleus

    Unit Charge




    Structure of Atom

    The discovery of subatomic particles like electrons and protons challenged the notion of indivisibility proposed by Dalton.
    Scientists realised that they needed to develop new atomic models that could explain the distribution of these subatomic particles within the atom and how they contributed to an atom's overall properties.

    This need for new models ultimately led to the development of various atomic models, starting with J.J. Thomson's Plum Pudding Model, then progressing to Ernest Rutherford's Nuclear Model, Niels Bohr's Planetary Model, and eventually the modern Quantum Mechanical Model. Each of these models aimed to better explain the structure and behaviour of atoms based on the new understanding of subatomic particles.

     Explore more about Metals and Non-Metals

    Thomson's Plum Pudding Model

    J.J. Thomson's model of the atom, often referred to as the "plum pudding" or "raisin pudding" model, was proposed as an analogy to help explain the distribution of charges within an atom. According to this model:

    a) Positive Sphere: The atom is visualised as a sphere filled with a positively charged substance. This positive charge is thought to be spread uniformly throughout the atom.

    b) Electron Embedding: Electrons, which carry a negative charge, are embedded within the positively charged sphere. They are distributed in such a way that they are somewhat evenly dispersed throughout the atom.

    c) Neutral Atom: According to Thomson's model, the negative charge of the electrons was exactly balanced by the positive charge of the surrounding sphere. As a result, the atom was considered electrically neutral as the positive and negative charges cancelled each other out.

    d) Mass Distribution: Thomson's model assumed that the mass of the atom, as well as the positive charge, was uniformly distributed. This was a representation of the understanding at that time.

    Thomson's Atomic Model - Labelled Diagram of Plum Pudding Model

    Limitations of Thomson's Model

    a) Charge Distribution: The model assumed a uniform distribution of positive charge and electrons, which couldn't explain experimental results involving alpha particles interacting with atoms.

    b) Stability: It didn't explain how negatively charged electrons in circular orbits could remain stable without emitting energy and collapsing into the nucleus.

    c) Nature of Positive Charge: It didn't address how the positive charge was confined to the centre of the atom without repelling electrons.

    d) Absence of Neutrons: The model didn't account for the existence of neutrons, which play a crucial role in atomic stability.

    e) Atomic Mass Variation: The model didn't explain the varying atomic masses observed among different elements.

    These limitations prompted the development of more accurate atomic models to better explain the complexities of atomic structure and behaviour.

     Explore more about Periodic Classification of Elements

    Rutherford's Nuclear Model of an Atom

    Rutherford's Nuclear Model of an Atom, also known as the Rutherford Model or Planetary Model, was a significant advancement in our understanding of atomic structure.
    Proposed by the New Zealand-born physicist Ernest Rutherford in 1911, this model introduced the concept of a central nucleus within the atom, which contains the majority of the atom's mass and positive charge.

    Key Points about Rutherford's Nuclear Model:

    1. Experimental Setup: Rutherford and his colleagues conducted an experiment in which they directed fast-moving alpha particles (α-particles) at a thin sheet of gold foil. They expected the α-particles to pass through the foil with minor deflections, as per the prevailing Thomson's Plum Pudding Model.

    2. Observations

    1. Most Particles Passed Through: The majority of the α-particles indeed passed through the gold foil with little or no deflection. This was expected based on Thomson's model, which suggested that an atom's positive charge was uniformly distributed.
    2. Some Deflections: However, a small fraction of α-particles experienced significant deflections at various angles, and very few were even deflected back toward the source.

    3. Inferences

    1. Most Space is Empty: The fact that most α-particles passed through the foil without deflection indicated that atoms contain a lot of empty space, disproving the idea of a uniformly distributed positive charge.
    2. Nucleus: The deflection of α-particles by significant angles and some being deflected backwards led Rutherford to propose that an atom's positive charge and mass are concentrated in a tiny, dense nucleus at the centre of the atom. This nucleus contains protons, positively charged subatomic particles, and almost all of the atom's mass.
    3. Electrons in Orbit: Rutherford's model suggested that electrons, which are negatively charged subatomic particles, orbit around the nucleus at a distance. He compared this arrangement to the planets orbiting around the sun, hence the term "Planetary Model."
    Rutherfords Gold Foil Experiment - Scattering of Alpha Particles

    4. Significance and Limitations: Rutherford's Nuclear Model was groundbreaking because it fundamentally changed our understanding of atomic structure by introducing the concept of a nucleus and the idea of empty space within atoms. However, it also had limitations:

    1. Stability of Atoms: Rutherford's model couldn't explain why electrons, carrying a negative charge, don't spiral into the positively charged nucleus due to electromagnetic attraction between opposite charges.
    2. Orbital Motion: Classical physics predicted that electrons in orbit should continuously emit radiation and lose energy, causing them to spiral into the nucleus. But this contradicted observations of stable atoms.

    Despite its limitations, Rutherford's model paved the way for further developments in atomic theory, notably Niels Bohr's model, which addressed the issues of stability and quantised energy levels of electrons in orbits.

    Bohr's Model of an Atom

    Bohr's Model of an Atom, proposed by Niels Bohr in 1913, was a significant improvement over Rutherford's Nuclear Model. It introduced the concept of quantised energy levels for electrons and helped explain some of the shortcomings of Rutherford's model.

    Key Postulates of Bohr's Model

    a) Electron Orbits: Electrons orbit the nucleus in well-defined and quantized energy levels or shells. These shells are labelled as K, L, M, N, etc. Electrons in the inner shells possess lower energy levels compared to those in the outer shells.

    b) Quantized Energy: Electrons can only exist in specific energy levels, and the energy of an electron in a particular shell is fixed and quantized. This prevents the electron from spiralling into the nucleus due to radiation emission.

    c) Stable Orbits: Electrons do not emit radiant energy while in stable orbits. This was a crucial departure from classical physics, where an accelerated charged particle was expected to lose energy in the form of radiation.

    d) Energy Transitions: When an electron absorbs energy, it moves to a higher energy level (excited state). Conversely, when it releases energy, it transitions to a lower energy level (ground state). The energy emitted or absorbed is exactly the difference in energy between the two levels.

    e) Radiation Emission: Electrons in excited states can spontaneously transition to lower energy levels while emitting energy in the form of photons. The emitted energy corresponds to specific wavelengths, resulting in spectral lines.

    Bohrs Atomic Model - Labelled Diagram


    While Bohr's Model successfully explained many aspects of atomic spectra and stability, it had limitations:

    a) Limited to Hydrogen-like Atoms: Bohr's model was initially developed for hydrogen-like atoms (atoms with one electron, like hydrogen and singly ionized helium). It faced challenges when applied to multi-electron atoms due to their more complex electron-electron interactions.

    b) Failures in Complex Atoms: Bohr's model failed to predict the electron configuration and spectral properties of atoms beyond hydrogen-like atoms. It couldn't explain the complexities of multi-electron systems, including the chemical behaviour and spectral lines of heavier elements.

    Bohr's model relied on classical physics principles, which didn't fully capture the behaviour of electrons in atomic dimensions.

    Arrangement of Electrons in an Atom

    The arrangement of electrons in an atom is governed by certain rules and principles. Here are the key points regarding the arrangement of electrons in an atom:

    1. Maximum Electron Capacity of Energy Levels

    The maximum number of electrons that can be accommodated in a particular energy level (shell) is given by the formula 2n2, where n represents the energy level. Each energy level is represented by a whole number starting from 1, known as the principal quantum number.

    For example:

    • The first energy level (n = 1) can hold a maximum of 2 electrons.
    • The second energy level (n = 2) can hold a maximum of 8 electrons.
    • The third energy level (n = 3) can hold a maximum of 18 electrons, and so on.

    2. Maximum Number of Electrons in Outermost Orbit

    The outermost energy level of an atom is referred to as the valence shell. This shell has the most influence on the chemical behaviour of the atom. The maximum number of electrons that can be accommodated in the outermost orbit (valence shell) is 8.

    3. Filling of Inner Shells before Outer Shells

    Electrons are arranged in an atom in such a way that the inner energy levels are filled before the outer ones. This principle is known as the "Aufbau principle." It means that electrons will occupy the lowest available energy levels before moving to higher energy levels.


    1. Valence Electrons: Valence electrons are the electrons located in the outermost energy level (valence shell) of an atom. These electrons are most involved in chemical reactions and bonding with other atoms.

    2. Combining Capacity and Valency: Valency represents the combined capacity of an atom. It indicates how many other atoms an atom can bond with to achieve a stable electron configuration. The valency is determined by the number of valence electrons.

    3. Valency and Octet Rule: The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration with 8 electrons in their valence shell. In the case of elements with only the first energy level (K shell), the rule is called the duplet rule, and they aim to have 2 electrons in their valence shell. The number of electrons an atom gains or loses to complete the octet (or duplet) determines its valency.

    4. Valency of Elements with Full Outer Shells: Elements with completely filled outermost energy levels (valence shells) have a valency of zero. This is because they already have a stable electron configuration, so they don't need to gain or lose electrons to achieve stability.

    Overall, understanding the valency of elements is crucial in predicting how they will interact with other elements to form compounds and molecules. It helps explain the formation of chemical bonds and the variety of chemical reactions that occur in nature.

    Atomic and Mass Number

    1. Atomic Number (Z)

    The atomic number of an element is a fundamental property that uniquely defines an element. It is denoted by the symbol "Z." Atomic number represents the number of protons present in the nucleus of an atom of that element. In a neutral atom, the number of protons is equal to the number of electrons, ensuring that the atom is electrically balanced.

    Key points about Atomic Number:

    1. Each element on the periodic table has a distinct atomic number.
    2. Atomic number determines an element's chemical identity and its position in the periodic table.
    3. Elements are arranged in increasing order of atomic number in the periodic table.
    4. For example, hydrogen (H) has an atomic number of 1, meaning it has one proton in its nucleus. Carbon (C) has an atomic number of 6, indicating it has six protons.

    2. Mass Number (A)

    The mass number of an atom is the sum of the number of protons and the number of neutrons in its nucleus. It is represented by the symbol "A." Since protons and neutrons have approximately equal mass, the mass number is a good approximation of the total number of nucleons (protons and neutrons) in an atom.

    Key points about Mass Number:

    1. Mass number is a whole number because it's the sum of protons (whole numbers) and neutrons (also whole numbers).
    2. Mass number is used to calculate the atomic mass or atomic weight of an element, which is the weighted average of the masses of its isotopes.
    3. For example, carbon-12 (12C) has a mass number of 12, indicating it has 6 protons and 6 neutrons. Carbon-14 (14C) has a mass number of 14, with 6 protons and 8 neutrons.

    Isotopes, Isobars and Isotones


    1. Isotopes are different forms of the same chemical element that have the same number of protons (which defines the element) but different numbers of neutrons in their atomic nuclei.
    2. Because the number of protons remains constant, isotopes of the same element share the same chemical properties and occupy the same position on the periodic table. However, due to the varying number of neutrons, isotopes can have different atomic masses. This leads to slight differences in their physical properties, such as their stability and radioactive behaviour.
    3. Example: Carbon has three isotopes - Carbon-12 (6 protons and 6 neutrons), Carbon-13 (6 protons and 7 neutrons), and Carbon-14 (6 protons and 8 neutrons).


    1. Isobars are atoms of different chemical elements that have the same total number of nucleons (protons + neutrons) in their nuclei, resulting in the same atomic mass number.
    2. While the atomic numbers of these elements differ, their atomic masses are equal due to the balance between protons and neutrons. Isobars have distinct chemical properties and positions on the periodic table because the number of protons determines an element's chemical behaviour. In other words, isobars have the same overall nuclear mass but belong to different elements.
    3. Example: Potassium-40 (19 protons and 21 neutrons) and Calcium-40 (20 protons and 20 neutrons) are isobars.


    1. Isotones are atoms of different elements that have the same number of neutrons in their atomic nuclei. As a result, isotones share similar nuclear properties.
    2. However, their atomic numbers and chemical properties are different due to varying numbers of protons. Isotones have different atomic masses as well since atomic mass depends on both protons and neutrons.
    3. Example: Chlorine-37 (17 protons and 20 neutrons), Argon-38 (18 protons and 20 neutrons), Potassium-39 (19 protons and 20 neutrons), and Calcium-40 (20 protons and 20 neutrons) are all isotones with 20 neutrons.

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