Structure of Atom

The One Thing

Atomic structure is best explained by the quantum mechanical model: a small, dense nucleus is surrounded by electrons occupying quantised orbitals rather than definite classical paths. Experimental evidence progressively replaced earlier models and established the roles of quantum numbers, wave–particle duality, and electronic configuration rules.

Reactions, Processes and Experiments

Key Terms

• Atom: The smallest unit of an element that retains the chemical properties of that element.
• Electron: A negatively charged subatomic particle with very small mass, approximately 9.109 x 10^-31 kg.
• Proton: A positively charged subatomic particle present in the nucleus, with charge equal in magnitude to the electron's charge.
• Neutron: An electrically neutral subatomic particle present in the nucleus, with mass approximately equal to that of a proton.
• Atomic number: The number of protons in the nucleus, represented by Z; it identifies an element.
• Mass number: The total number of protons and neutrons in an atom, represented by A.
• Isotopes: Atoms of the same element having the same atomic number but different mass numbers.
• Thomson model: An early model describing the atom as a positively charged sphere with electrons embedded in it.
• Rutherford nuclear model: A model proposing that an atom has a tiny, dense, positively charged nucleus containing most of its mass, with electrons around it.
• Bohr model: A model in which electrons move in certain permitted circular orbits with fixed energies and do not continuously lose energy.
• Electromagnetic radiation: Energy transmitted through space as oscillating electric and magnetic fields.
• Wavelength: The distance between two successive crests or troughs of a wave, represented by lambda.
• Frequency: The number of waves passing a point per second, represented by nu and measured in hertz.
• Photon: A small packet or quantum of electromagnetic energy.
• Line spectrum: A spectrum containing separate lines of specific wavelengths produced when atoms emit or absorb particular energies.
• Quantum theory: The idea that energy is absorbed or emitted in discrete packets rather than in a continuous manner.
• de Broglie hypothesis: The proposal that moving particles such as electrons have wave-like properties.
• Heisenberg uncertainty principle: It is impossible to determine simultaneously and exactly both the position and momentum of an electron.
• Orbital: A region of space around the nucleus where the probability of finding an electron is high.
• Quantum numbers: Four numbers that describe the energy, size, shape, orientation, and spin state of an electron.
• Principal quantum number: The quantum number n, which identifies the main shell and indicates the approximate energy and size of an orbital.
• Azimuthal quantum number: The quantum number l, which identifies the subshell and shape of an orbital; its values range from 0 to n - 1.
• Magnetic quantum number: The quantum number m_l, which describes the orientation of an orbital and has values from -l to +l.
• Spin quantum number: The quantum number m_s, which describes electron spin and has values of +1/2 or -1/2.
• s, p, d, and f subshells: Subshells corresponding to l values 0, 1, 2, and 3, containing 1, 3, 5, and 7 orbitals respectively.
• Electronic configuration: The arrangement of electrons in shells, subshells, and orbitals.
• Aufbau principle: Electrons occupy orbitals in order of increasing energy.
• Pauli exclusion principle: No two electrons in an atom can have the same set of all four quantum numbers.
• Hund's rule: Electrons occupy degenerate orbitals singly with parallel spins before pairing begins.

Easily Confused

• Atomic number and mass number: Atomic number, Z, is the number of protons; mass number, A, is the total number of protons and neutrons.
• Isotopes and ions: Isotopes differ in neutron number and mass number, whereas ions differ in electron number and therefore carry a charge.
• Thomson model and Rutherford nuclear model: Thomson proposed electrons embedded in a positively charged sphere; Rutherford proposed a small, dense, positively charged nucleus surrounded by electrons.
• Rutherford model and Bohr model: Rutherford established the nuclear structure but did not explain stability or line spectra; Bohr introduced permitted orbits with fixed energies.
• Bohr model and quantum mechanical model: Bohr described electrons in fixed circular orbits, whereas the quantum mechanical model describes probability distributions in orbitals.
• Wavelength and frequency: Wavelength is the distance between successive wave crests or troughs; frequency is the number of waves passing a point per second.
• Photon and line spectrum: A photon is a packet of electromagnetic energy; a line spectrum consists of separate wavelengths produced by particular photon energies.
• Shell, subshell and orbital: A shell is identified by n, a subshell by l, and an orbital by its spatial orientation and probability distribution.
• Azimuthal and magnetic quantum numbers: The azimuthal quantum number, l, identifies subshell shape; the magnetic quantum number, m_l, identifies orbital orientation.
• Aufbau principle, Pauli exclusion principle and Hund's rule: Aufbau determines orbital filling order, Pauli limits each orbital to two opposite-spin electrons, and Hund's rule determines filling among degenerate orbitals.
• Emission and absorption: Emission occurs when an electron moves to a lower energy level; absorption occurs when it moves to a higher energy level.
• Orbit and orbital: An orbit is a permitted circular path in the Bohr model; an orbital is a region of high probability in the quantum mechanical model.

What Gets Asked

• Identifying subatomic particles and calculating atomic quantities: Questions may require use of $Z=$ number of protons, neutral-atom electron number $=Z$, $A=$ number of protons + number of neutrons, and number of neutrons $=A-Z$. A common mark-losing error is confusing atomic number with mass number.
• Explaining the development of atomic models: Questions may ask why Thomson's model was replaced by Rutherford's model, or why Rutherford's model was replaced by Bohr's model and subsequently by the quantum mechanical model. The specific limitations must be stated: Rutherford's model could not explain atomic stability or line spectra, while Bohr's model was not adequate for multi-electron atoms.
• Interpreting Rutherford's alpha-particle scattering experiment: Answers must connect the observations to the conclusions that most of the atom is empty space and that positive charge and most mass are concentrated in a small nucleus.
• Using electromagnetic and photon equations: Questions may require $c=lambdanu$, where c is approximately $3.00x10^{8}m{s}^{-1}$, and $E=hnu=hc/lambda$, where h is approximately $6.626x10^{-34}Js$. Marks may be lost by confusing wavelength with frequency or omitting units.
• Explaining hydrogen spectra and Bohr quantisation: Questions may involve $mvr=nh/(2pi)$, $E_n=-2.18x10^{-18}{Z}^2/n^{2}J$ per atom, $r_n=0.529n^2/Z$ angstrom, or identification of the Balmer series as transitions ending at $n=2$. The key distinction is whether energy is emitted during a downward transition or absorbed during an upward transition.
• Writing or evaluating electronic configurations: Questions may test the energy order $1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d,7p$, together with Aufbau, Pauli exclusion and Hund's rule. A frequent error is pairing electrons in degenerate orbitals before placing them singly with parallel spins.