The atom is the tiniest building block of matter with the characteristics of a chemical element. Atoms do not exist in isolation; rather, they create ions and molecules, which then amalgamate in significant quantities to construct the matter we perceive, experience and touch. The majority of an atom’s volume is void. The remaining part includes a core with a positive charge made up of protons and neutrons encircled by a cluster of negatively charged electrons. The atomic nucleus is compact and heavy in comparison to electrons, which are the least weighty charged particles known. An electric force drives electrons towards a positive charge, while in an atom, these forces glue the electrons to the nucleus.
The atom’s diverse attributes cannot be completely depicted by any single illustration due to the principles of quantum mechanics. Consequently, physicists must employ multiple images of the atom to clarify its different properties. The behaviour of electrons in an atom can be described in two ways: as particles orbiting the nucleus and as wave patterns frozen in position around the nucleus. These wave patterns are called orbitals, and they help to explain how individual electrons are distributed in the atom. The properties of an atom’s orbitals strongly impact its behaviour, and its chemical characteristics are defined by sets of orbitals called shells.
The notion of an undividable particle has been present since ancient times across various cultures, positing that matter is composed of minuscule, indivisible particles. With the progression of cognitive abilities, our perceptions and philosophy towards the nature of matter and energy also have advanced. Kaṇāda, an Indian philosopher, hypothesised that paramanu (atom) is an indestructible particle of matter that cannot be divided, as it represents a state that cannot be measured. The term “atom” originates from the Greek word “atomos”, meaning “uncuttable”, and was originally a concept rooted in philosophy rather than science.
Unlike outdated concepts, modern atomic theory was developed upon new principles. During the early 19th century, John Dalton, a scientist, observed that chemical elements combined in discrete units of weight. He coined the term “atom” to describe these units, believing them to be the fundamental building blocks of matter. Although it was later found that Dalton’s atoms were not truly indivisible, the term remained in use for over a century. Now, let’s examine the crucial principles of the atomic model that were developed through modern physical science.
During the early 1800s, John Dalton, an English chemist, examined experimental data collected by himself and other scientists and identified a pattern that is currently recognised as the “law of multiple proportions. He observed that the proportion of a specific chemical element within a compound varied in weight by ratios of small integers. This pattern implied that chemical elements combined with one another through fundamental units of weight, which Dalton named “atoms”.
Among various historical, scientific models of the atom, the plum pudding model, first introduced by J. J. Thomson in 1904, endeavoured to account for two known properties of atoms at that time: the negative charge of electrons and the absence of a net electric charge in atoms, despite containing both positive and negative charges. Thomson proposed this model prior to the discovery of the atomic nucleus. In the plum pudding model, electrons are embedded in a positively charged “pudding,” forming negatively charged “plums”.
For several years, it has been established that atoms consist of subatomic particles with a negative charge. These particles were originally referred to as “corpuscles” by Thomson, but they became better known as “electrons,” a term coined by G.J. Stoney in 1891 to describe the “fundamental unit quantity of electricity.” For many years, it has also been understood that atoms are electrically neutral. Thomson’s belief was that atoms contained a positive charge that could balance out the negative charge of electrons. This proposed model was published in the famous English science journal of the time, the Philosophical Magazine, in March 1904.
Ernest Rutherford, a physicist from New Zealand, proposed a model of the structure of atoms in 1911, which is commonly referred to as Rutherford’s model, the nuclear atom or the planetary model of the atom. J.J. Thomson’s plum pudding model was unable to account for certain experimental findings related to the atomic structure of elements. The atomic model proposed by Rutherford depicts a nucleus, which is a small, dense, positively charged core containing most of the atom’s mass. Electrons, the light, negatively charged components, orbit the nucleus at a distance similar to planets revolving around the sun.
Rutherford formulated the atomic structure of elements based on the aforementioned observations and conclusions. The nucleus of an atom contains the majority of its mass and a positive charge, which is concentrated in a very small volume. Rutherford named the region of an atom containing a positive charge the nucleus. In his model, he proposed that the negatively charged electrons orbit the nucleus. Rutherford’s model posited that electrons orbit the nucleus in circular paths at high speeds, which he referred to as orbits. Due to the strong electrostatic force of attraction between the negatively charged electrons and the positively charged nucleus, the latter being a densely concentrated mass of particles, they remain held together.
Niels Bohr, a physicist from Denmark, proposed a model of the structure of atoms, particularly that of hydrogen, in 1913, which is known as the Bohr model. The Bohr model of the atom was a significant deviation from classical depictions and was the first to integrate quantum theory. It served as a forerunner to purely quantum-mechanical models. The Bohr model and its successors characterise the characteristics of atomic electrons in relation to a set of permissible values. Electrons in atoms only absorb or emit radiation when they make abrupt transitions between these permitted, stationary states. In 1914, Gustav Hertz and James Franck, both physicists from Germany, provided direct experimental proof of the presence of these discrete states. The Bohr radius is a physical constant that represents the most probable distance between the nucleus and the electron in the ground state of a hydrogen atom.
The primary accomplishment of the model was its ability to account for the spectral emission lines of atomic hydrogen through the Rydberg formula. Although the Rydberg formula had been established experimentally, it lacked a theoretical foundation until the advent of the Bohr model. The Bohr model not only clarified the basis of the Rydberg formula’s structure but also validated the essential physical constants that constitute the formula’s experimental outcomes.
When compared to the valence shell atomic model, the Bohr model is comparatively rudimentary as a model of the hydrogen atom. The Bohr model can be viewed as an outdated scientific theory since it can be derived as a first-order approximation of the hydrogen atom utilising the more comprehensive and precise framework of quantum mechanics. Despite being an outdated scientific theory, the Bohr model is still frequently taught to students as an introductory concept to quantum mechanics due to its simplicity and its accurate predictions for some systems before transitioning to the more intricate and precise valence shell atom model.
Schrödinger’s wave equation and its solution form the basis of quantum mechanics, which introduces the concepts of shells, subshells and orbitals. The likelihood of finding an electron at a specific point within an atom is directly proportional to the square of the electron’s wave function, denoted as |ψ|2. Approximate methods were used to overcome the difficulty of putting Schrödinger’s equation to multiple-electron atoms, as the wave equation cannot be solved exactly for such atoms.
In quantum mechanical models of atoms, the complex shapes of orbits (also called electron clouds) are used to represent the volumes of space where electrons are most likely to be found. This model is probabilistic in nature rather than deterministic and describes the complex shape of orbitals (also known as electron clouds) where electrons are likely to be found. The energy of an electron in an orbital is quantised, meaning it can only have certain allowed values. The solution to the Schrodinger wave equation gives the quantised energy of the electron, which arises from the wave-like properties of the electron.
Due to Heisenberg’s Uncertainty Principle, the exact position and momentum of an electron cannot be known simultaneously. Hence, we can only determine the probability of locating an electron in a specific orbital. The wave properties (ψ) that describe the behaviour of electrons in an atom are referred to as atomic orbitals. Atomic orbitals are occupied by electrons whose positions are described by wave functions. Since an electron can have many wave functions, it can occupy multiple atomic orbitals. Each atomic orbital has a specific shape and energy associated with the wave function.
Quantum mechanics can extract information about the electrons in an atom from the orbital wave properties of the atom, which contain all the relevant information about the electrons. Quantum theory postulates that electrons can only occupy a discrete set of quantised energy levels: all orbits have fixed atomic radii respective to their energy levels. Quantum theory also postulates that the Pauli Exclusion Principle holds, which implies that no two electrons within a system can occupy the same energy state simultaneously and that all energy levels are occupied in ascending order, starting from the lowest level.