Comprehensive theory, key formulas, diagrams, and memory aids for Structure of the Atom.
In the previous chapter, we learned that Dalton’s atomic theory suggested that the atom was indivisible and indestructible. However, the discovery of two fundamental particles (electrons and protons) inside the atom led to the failure of this aspect of Dalton’s theory. It became necessary to know how electrons and protons are arranged within an atom. This chapter explores the journey of scientific discoveries that revealed the intricate structure of the atom.
The fact that atoms are divisible was supported by the discovery of charged particles. * Electrons: J.J. Thomson discovered the electron ($e^-$), a negatively charged subatomic particle, through cathode ray tube experiments. An electron has a negligible mass and a charge of $-1$. * Protons: E. Goldstein discovered the presence of new radiations in a gas discharge called canal rays. These rays were positively charged and led to the discovery of the proton ($p^+$). A proton has a mass approximately 2000 times that of an electron and a charge of $+1$.
By 1900, it was known that the atom was not a simple, indivisible particle but contained at least one subatomic particle—the electron.
Since atoms were neutral overall, it was clear that they must contain equal amounts of positive and negative charge. How were these arranged? Several models were proposed.
J.J. Thomson proposed that an atom is similar to a Christmas pudding or a watermelon. * He proposed that the atom consists of a positively charged sphere and the electrons are embedded in it (like seeds in a watermelon). * The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral. While this explained electrical neutrality, it failed to explain the results of experiments carried out by other scientists.
Ernest Rutherford was interested in knowing how the electrons are arranged. He conducted the famous $\alpha$-particle scattering experiment. He bombarded a very thin gold foil with fast-moving alpha ($\alpha$) particles (doubly-charged helium ions, $He^{2+}$). Observations: 1. Most of the fast-moving $\alpha$-particles passed straight through the gold foil. 2. Some of the $\alpha$-particles were deflected by the foil by small angles. 3. Surprisingly, one out of every 12,000 particles appeared to rebound (bounce straight back).
Conclusions: 1. Most of the space inside the atom is empty because most $\alpha$-particles passed without getting deflected. 2. The positive charge of the atom occupies very little space because very few particles were deflected from their path. 3. A very small fraction of $\alpha$-particles rebounded, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom (which he called the nucleus).
Rutherford's Nuclear Model: * There is a positively charged center in an atom called the nucleus. Nearly all the mass resides here. * The electrons revolve around the nucleus in circular paths. * The size of the nucleus is very small as compared to the size of the atom.
Drawback: Any particle in a circular orbit would undergo acceleration and radiate energy. Thus, the revolving electron would lose energy and finally fall into the nucleus, making the atom highly unstable. But we know atoms are stable.
To overcome Rutherford's objections, Niels Bohr put forward the following postulates: 1. Only certain special orbits known as discrete orbits of electrons are allowed inside the atom. 2. While revolving in discrete orbits, the electrons do not radiate energy. These orbits or shells are called energy levels. They are represented by letters K, L, M, N... or the numbers $n=1, 2, 3, 4\dots$
In 1932, J. Chadwick discovered another subatomic particle which had no charge and a mass nearly equal to that of a proton. It was eventually named the neutron ($n$). Neutrons are present in the nucleus of all atoms, except hydrogen. Thus, the mass of an atom is given by the sum of the masses of protons and neutrons present in the nucleus.
The distribution of electrons into different orbits of an atom is governed by the following rules: 1. The maximum number of electrons present in a shell is given by the formula $2n^2$, where '$n$' is the orbit number (1, 2, 3...). * K shell ($n=1$): $2(1)^2 = 2$ electrons max. * L shell ($n=2$): $2(2)^2 = 8$ electrons max. * M shell ($n=3$): $2(3)^2 = 18$ electrons max. 2. The maximum number of electrons that can be accommodated in the outermost orbit is 8 (Octet rule). 3. Electrons are not accommodated in a given shell unless the inner shells are filled. (Step-wise filling).
The electrons present in the outermost shell of an atom are known as valence electrons. Atoms of elements having a completely filled outermost shell show little chemical activity (e.g., noble gases like Helium and Neon). Their combining capacity or valency is zero. Other elements try to achieve this stable octet (8 electrons in the outermost shell) by reacting with other atoms. They can gain, lose, or share electrons. The number of electrons gained, lost, or shared is the valency of the element. * Sodium (Atomic number 11, config: 2, 8, 1) loses 1 electron. Valency = 1. * Oxygen (Atomic number 8, config: 2, 6) gains 2 electrons to reach 8. Valency = 2.
An atom is represented as: $$^A_Z \text{Symbol}$$ For example, $^{14}_7 N$ means Nitrogen has atomic number 7 and mass number 14.
The atom is a complex system consisting of a dense, positively charged nucleus (containing protons and neutrons) surrounded by negatively charged electrons in discrete energy shells. From Thomson's "plum pudding" to Rutherford's nucleus and Bohr's orbits, scientific models have evolved to explain the stability and behavior of matter. Concepts like valency, isotopes, and atomic number allow us to predict how atoms will interact to form the diverse array of chemicals in our universe.