Atomic Models: From Dalton to Quantum Mechanics
Atomic Models: A Historical Perspective
Dalton’s Atomic Model (1808)
Dalton proposed that matter is composed of indivisible particles called atoms. All atoms of a given element are identical, and atoms of different elements have different properties. Atoms combine in simple whole-number ratios to form compounds.
Thomson’s Atomic Model (1904)
Thomson described the atom as a solid sphere with a positively charged interior, in which negatively charged electrons are embedded, similar to raisins in a pudding. This model explained the formation of ions: an atom that gains or loses electrons becomes an ion. Losing electrons results in a positive ion (cation), while gaining electrons results in a negative ion (anion).
Rutherford’s Atomic Model (1912)
Rutherford discovered that atoms are mostly empty space, with a tiny, dense, positively charged nucleus at the center. Electrons orbit the nucleus at great distances and high speeds.
Bohr’s Atomic Model (1913)
Bohr proposed that electrons revolve around the nucleus in specific energy levels. Each energy level can hold a specific number of electrons. When an electron changes energy levels, it absorbs or emits energy.
Modern Atomic Model
The modern atomic model consists of a nucleus and an electron cloud. The nucleus contains protons (positively charged particles with a mass of approximately 1.67 x 10-27 kg) and neutrons (uncharged particles with a mass similar to protons). The electron cloud contains electrons (negatively charged particles much smaller than protons and neutrons). Electrons are found in different energy levels and occupy orbitals, which are regions around the nucleus where the probability of finding an electron is high. Energy levels are numbered from 1 to 7, with increasing energy. Orbitals are regions around the nucleus where the probability of finding an electron is high.
Quantum Mechanical Model (1924-1927)
In 1924, it was concluded that particles exhibit wave-like properties. All moving particles are associated with a wave. In 1927, considering the wave-particle duality of electrons, Heisenberg proposed the uncertainty principle, which states that it is impossible to know both the position and momentum of an electron simultaneously. The more accurately one variable is known, the less accurately the other is known.
Based on the uncertainty principle, it is postulated that electrons exist in regions of high probability called energy levels or layers.
Energy Levels
Energy levels are identified using two notations:
- Letters: K, L, M, N, etc.
- Numbers: 1, 2, 3, 4, etc.
These numbers are known as the principal quantum number. The maximum number of electrons in a given energy level is given by the formula: 2n2, where n is the energy level.
Sublevels
Not all electrons in a given energy level have the same energy. Each main energy level has one or more sublevels. Sublevels are designated by the letters s, p, d, and f. A sublevel is specified by the main energy level number followed by the corresponding letter (e.g., 2s). The maximum number of electrons in each sublevel is:
- s = 2 electrons
- p = 6 electrons
- d = 10 electrons
- f = 14 electrons
Sublevels are also symbolized by a secondary quantum number (l), with values ranging from 0 to n-1.
Orbitals
Within each sublevel, there are orbitals, which are regions where electrons are most likely to be found. Each orbital can hold a maximum of two electrons. The number of orbitals in each sublevel is:
- s = 1 orbital
- p = 3 orbitals
- d = 5 orbitals
- f = 7 orbitals
Electrons are individually located within an orbital using a magnetic quantum number (m), which describes the orientation of the orbital in space. The values of m range from -l to +l.
Electron Spin
Electrons possess a property called spin, which is analogous to the Earth’s rotation. This property is described by the spin quantum number (s), which can have values of +1/2 or -1/2. Two electrons in the same orbital must have opposite spins.
Electron Configuration
Electron configuration describes the location of electrons in the orbitals of a given atom.
Pauli Exclusion Principle
The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers.
Aufbau Principle (Minimum Energy Principle)
Electrons in an atom’s ground state occupy the lowest energy levels and sublevels available.
Hund’s Rule (Maximum Multiplicity Principle)
Hund’s rule states that when filling a sublevel, electrons will individually occupy each orbital with the same spin before pairing up in the same orbital.