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Introduction to Sub-Atomic Particles
Matter is composed of fundamental building blocks known as atoms. For a long time, atoms were considered indivisible, but scientific discoveries through the late nineteenth and early twentieth centuries revealed that atoms are themselves made up of smaller entities called sub-atomic particles. These particles are the electron, the proton, and the neutron.
The electron was discovered by J.J. Thomson through cathode ray tube experiments, identifying a particle with a negative charge. Protons were subsequently identified, carrying a positive charge equal in magnitude to the negative charge of an electron. Neutrons, discovered by James Chadwick, possess no electric charge and have a mass nearly equal to that of a proton.
Each sub-atomic particle possesses distinct properties regarding its mass and electrical charge:
- Electron: Carries a relative electric charge of negative one (-1) and has an extremely small mass, considered negligible when calculating the total mass of an atom.
- Proton: Carries a relative electric charge of positive one (+1) and has a relative mass of one unified atomic mass unit.
- Neutron: Carries zero electric charge (neutral) and has a relative mass of one unified atomic mass unit.
In a neutral atom, the number of negatively charged electrons equals the number of positively charged protons, ensuring that the net electrical charge of the entire atom is zero. These particles are not distributed uniformly throughout the volume of the atom; instead, they are segregated into distinct regions, forming the complex internal architecture of the atomic structure.
Models of the Atom and the Rutherford Discovery
Understanding how these sub-atomic particles are arranged required a series of theoretical models proposed by various scientists. J.J. Thomson first proposed the plum pudding model, suggesting that an atom consisted of a positively charged sphere with negatively charged electrons embedded within it, much like dried fruit in a Christmas pudding.
This model was tested and disproved by Ernest Rutherford and his coworkers through the famous alpha-particle scattering experiment. In this experiment, fast-moving alpha particles, which are doubly charged helium ions, were directed at a very thin gold foil sheet. Based on Thomson's model, the alpha particles were expected to pass straight through the foil with minor deflections because the positive charge was thought to be spread out evenly.
However, the results revealed three distinct observations:
- Most of the fast-moving alpha particles passed straight through the gold foil without any deflection, indicating that the interior of an atom is mostly empty space.
- A small fraction of the alpha particles was deflected by small angles, showing that the positive charge of the atom occupies a very concentrated region.
- A very few alpha particles rebounded completely, retracing their path, which demonstrated that all the positive charge and mass of the atom are concentrated in an extremely small volume at the centre.
From these observations, Rutherford proposed the Rutherford model of the atom, also known as the nuclear model. This model states that:
- There is a positively charged centre in an atom called the nucleus, where nearly all the mass of the atom is concentrated.
- Electrons revolve around the nucleus in circular paths.
- The size of the nucleus is extremely small compared to the total size of the atom.
Atomic Number, Mass Number, and Electronic Configuration
To distinguish one element from another, scientists use quantitative measures based on the particles inside the nucleus. The atomic number of an element, denoted by the symbol Z, is defined as the total number of protons present in the nucleus of an atom. Since atoms are electrically neutral, the atomic number also equals the number of electrons in a neutral atom. For example, hydrogen has an atomic number of 1, meaning every neutral hydrogen atom has one proton and one electron.
The total mass of an atom is concentrated primarily in its nucleus due to the presence of protons and neutrons. These two particles are collectively referred to as nucleons. The mass number, denoted by the symbol A, is defined as the total number of protons and neutrons present in the nucleus of an atom.
Mass Number (A) = Number of protons (Z) + Number of neutrons (n)
For instance, a carbon atom has 6 protons and 6 neutrons, giving it an atomic number of 6 and a mass number of 12. If the atomic number and mass number of an element are known, the number of neutrons can be calculated easily by subtracting the atomic number from the mass number.
Regarding the arrangement of electrons outside the nucleus, the Bohr-Bury scheme dictates how electrons are distributed into different energy levels or shells, designated by the letters K, L, M, N, and so on. The maximum number of electrons accommodated in any shell is given by the formula 2n squared, where n is the orbit number or energy level index. The outermost shell of an atom is called the valence shell. Electrons cannot accommodate in a given shell unless the inner shells are filled completely, following a step-by-step filling order.
Valency
The chemical reactivity of an element is determined by the number of electrons residing in its outermost shell. Atoms tend to attain a fully stable configuration of eight electrons in their outermost shell, known as an octet, or two electrons in the case of the innermost shell (K shell) if it is the only shell. This drive for stability dictates the valency of an element.
Valency is defined as the combining capacity of an atom, representing the number of electrons gained, lost, or shared to achieve a stable electronic configuration. For elements with 1, 2, 3, or 4 valence electrons, the valency is generally equal to the number of valence electrons. For instance, sodium has an electronic configuration of 2, 8, 1; its valence electron is 1, so its valency is 1.
For elements with 5, 6, or 7 valence electrons in their outermost shell, the valency is calculated by subtracting the number of valence electrons from eight. For example, chlorine has an electronic configuration of 2, 8, 7. It has 7 valence electrons, so its valency is 8 minus 7, which equals 1. Elements that already possess a complete octet, such as helium, neon, and argon, have a valency of zero and are chemically inert.
Isotopes and Isobars
While all atoms of a specific element share the same atomic number, they do not always possess the exact same mass. Atoms of the same element that have the same atomic number but different mass numbers are called isotopes. This difference in mass number occurs because isotopes have varying numbers of neutrons in their nuclei.
A classic example is hydrogen, which exists as three distinct isotopes: protium (one proton, zero neutrons), deuterium (one proton, one neutron), and tritium (one proton, two neutrons). Another prominent example is carbon, which exists as carbon-12 and carbon-14. Because chemical properties are dictated primarily by the number of electrons and protons, isotopes of the same element exhibit identical chemical properties, though their physical properties may vary slightly due to mass differences.
In contrast, isobars are atoms of different chemical elements that have different atomic numbers but the same mass number. Because they belong to different elements, isobars have different numbers of protons and different numbers of electrons, yet their total sum of protons and neutrons remains identical. An example of a pair of isobars is calcium (atomic number 20, mass number 40) and argon (atomic number 18, mass number 40). Both possess a mass number of 40, despite having completely different chemical characteristics and elemental identities.
