Structure of the Atom Class 9 Notes Science Chapter 4

In the late 19th century, scientists sought to understand the internal structure of the atom and its essential characteristics. A series of experiments and observations over subsequent decades revealed that the atom is not indivisible but contains charged particles and a compact central core. These short notes present the development of atomic models, key definitions and rules, and examples useful for class-level study and exam preparation.

Atom

Charged Particles in Matter

Investigations of static electricity and the conduction behaviour of different substances provided an early clue that atoms contain charged particles. Objects become electrically charged on rubbing, and charged bodies can attract small neutral objects, indicating the presence of movable charges.

Comb and Glass Rod Activity

• Action: A plastic comb is used to comb dry hair; a glass rod is rubbed with a silk cloth.
• Observation: The comb attracts small pieces of paper; the charged glass rod attracts an inflated balloon or small pieces of paper.
• Conclusion: Rubbing transfers charge so that the comb or rod becomes electrically charged. This shows that matter contains mobile charged particles (later identified as electrons and positive charges).

J. J. Thomson and the Discovery of the Electron

Investigations with cathode rays by J. J. Thomson showed that atoms contain small, negatively charged particles called electrons. These experiments suggested that atoms are not indivisible and that negative charges exist within the atom.

Thomson Model of the Atom

Thomson proposed an atomic model often compared to a Christmas pudding or a watermelon with seeds:

• An atom consists of a uniformly distributed positive charge (the pudding or the red part of the watermelon) in which electrons are embedded like currants or seeds.
• The total positive charge balances the total negative charge of the electrons, so the atom is electrically neutral.

Drawback: Thomson's model could not explain the results of later scattering experiments that showed most of the atom is empty space with a very small dense centre.

Thomson's Model of an atom

Rutherford's Gold Foil Experiment and Nuclear Model

Rutherford bombarded a thin gold foil with fast-moving α-particles (helium nuclei with mass ≈ 4 u and charge +2) and observed their scattering. The key observations were:

• Most α-particles passed straight through the foil.
• Some α-particles were deflected by small angles.
• About one in 12000 α-particles was deflected by an angle close to 180° (back-scattered).

Conclusions drawn by Rutherford:

• Most of the atom's volume is empty space, since most α-particles passed through with no deflection.
• The positive charge of the atom is concentrated in a very small region (the nucleus), which contains nearly all the atom's mass.
• Because a few α-particles experienced large-angle deflections, the nucleus must be positively charged and compact.

Scattering of α-particles by a gold foil

Main points of Rutherford's nuclear model:

• The atom has a tiny, dense, positively charged central nucleus containing protons (and later discovered neutrons).
• Electrons revolve around the nucleus in orbits, and the size of the nucleus is much smaller than the overall atom.

Drawbacks of Rutherford's model:

• A charged particle moving in a circular orbit should continuously lose energy by electromagnetic radiation and spiral into the nucleus; Rutherford's model could not explain why atoms are stable.

Bohr's Model of the Atom

Niels Bohr proposed modifications to Rutherford's model to explain atomic stability and the discrete spectral lines of hydrogen. The main points are:

• Electrons move in certain permitted circular orbits called discrete orbits or energy levels.
• While an electron remains in a permitted orbit, it does not emit energy.
• When an electron jumps from a higher permitted orbit to a lower one, it emits radiation whose frequency corresponds to the energy difference between the orbits; absorption occurs for the reverse process.

These orbits or shells are called energy levels, which are shown in Fig. 4.3. 

Bohr's model successfully explained the hydrogen spectrum and the stability of atoms in their ground states, although later quantum mechanics refined these ideas further.

Discovery of Neutrons

• In 1932, James Chadwick discovered a neutral particle inside the nucleus called the neutron.
• The neutron has approximately the same mass as the proton and no electrical charge; it is represented by the symbol n.
• The mass of an atom (its nucleon count) is given by the sum of the numbers of protons and neutrons present in the nucleus.

Distribution of Electrons in Shells (Bohr-Bury Rules)

The arrangement of electrons in different shells (or energy levels) of an atom follows the Bohr-Bury rules:

• The maximum number of electrons that can be present in a shell having principal quantum number n is given by 2n².
• Thus, the maximum electrons in the first (K) shell (n = 1) = 2 × 1² = 2.
• Maximum electrons in the second (L) shell (n = 2) = 2 × 2² = 8.
• Maximum electrons in the third (M) shell (n = 3) = 2 × 3² = 18.
• The outermost shell can hold at most 8 electrons (for the purpose of simple valency rules used at this level).
• Electrons fill inner shells first before occupying outer shells (Aufbau principle in qualitative form).

Structure of the First Eighteen Elements (Hydrogen to Argon)

Using the Bohr-Bury filling rules, the electronic structures of the first eighteen elements (from hydrogen to argon) can be represented by the number of electrons in successive shells (K, L, M, ...). These diagrams help explain chemical properties and valency of these elements.

Valency

• Valency is a measure of an atom's combining capacity and is often determined by the number of electrons in the outermost shell (valence shell).
• Atoms tend to attain a fully filled outermost shell (an octet, where possible). They do so by losing, gaining, or sharing electrons.
• Elements with a completely filled outermost shell show little chemical activity and have valency zero (noble gases).
• Examples:
  • Hydrogen (H), sodium (Na) and potassium (K) have one electron in their outermost shell and commonly lose one electron to attain a stable configuration; their valency is 1 (univalent).
  • Magnesium (Mg) and aluminium (Al) have two and three electrons respectively in their outermost shell, so their valencies are 2 and 3 respectively.
• For elements of Group 1, Group 2 and Group 13, the valency commonly equals the number of valence electrons. For Group 14 the valency is often 4. For Groups 15, 16 and 17, valency is often 8 minus the number of valence electrons.
• It is not always convenient for an element to lose many electrons to attain an octet. For example, fluorine has 7 electrons in its outermost shell; it is much easier to gain 1 electron than to lose 7 electrons, so its valency is 1.

Atomic Number

• The atomic number of an element, denoted by Z, equals the number of protons in the nucleus of an atom.
• For hydrogen, Z = 1; for carbon, Z = 6; and so on.

Mass Number

• The mass number of an atom is the total number of nucleons (protons + neutrons) in the nucleus.
• It is usually denoted by A. For example, carbon-12 has mass number A = 12 because it contains 6 protons and 6 neutrons.
• Atomic notation: the conventional notation for an atom places the mass number as a superscript and the atomic number as a subscript to the left of the chemical symbol, for example:

For example: nitrogen can be represented as shown in the figure below.

Atomic Number and Mass Number

Isotopes

• Isotopes are atoms of the same element that have the same atomic number (same number of protons) but different mass numbers (different number of neutrons).
• Hydrogen has three common isotopes:
• Isotopes have nearly identical chemical properties but different physical properties (such as mass and nuclear stability).
• Chlorine occurs naturally as two isotopes with masses 35 u and 37 u in approximately the ratio 3 : 1. Mass of an atom of any natural element is taken as the average mass of all the naturally occurring atoms of that element. The average atomic mass of chlorine atom can be calculated as under :
  

Applications of Isotopes

• An isotope of uranium (U-235) is used as fuel in nuclear reactors and in nuclear power generation.
• An isotope of cobalt (cobalt-60) is used in radiotherapy for the treatment of cancer.
• An isotope of iodine (iodine-131) is used in the diagnosis and treatment of thyroid disorders such as goitre.

Isobars

• Isobars are atoms of different elements that have the same mass number but different atomic numbers.
• For example, calcium (Ca) has atomic number 20 and argon (Ar) has atomic number 18, but both can have mass number 40 (Ca-40 and Ar-40); they are isobars because their total number of nucleons (protons + neutrons) is equal.