Hsslive.net provided Plus One Chemistry notes for students in their higher secondary years in two languages English Medium & Malayalam Medium. Topics- “Structure of Atom” that are usually covered in the first year of chemistry at the higher secondary level. valuable for both Kerala Syllabus and CBSE students with knowledge cultivated over two decades of teaching experience.
Structure of Atom
An atom is the smallest unit of matter that retains the identity and properties of an element. An atom comprises three subatomic particles: electrons, protons, and neutrons. Electrons are negatively charged particles that move around the nucleus, which is the atom’s central core. Protons are positively charged particles that are present in the nucleus. Neutrons are neutral particles that are also present in the nucleus. The number of protons in an atom determines its atomic number and its identity as an element. The number of neutrons in an atom determines its mass number and its isotopic variation. The number of electrons in an atom determines its charge and chemical behaviours.
The structure of an atom has been a subject of interest and investigation for many scientists since the nineteenth century. Various models and theories have been proposed to explain the nature and behaviour of atoms and their interactions with other atoms and radiation. Some of the important models and theories that we will learn in this chapter are:
- Thomson’s model of an atom
- Rutherford’s model of the atom
- Bohr’s model of atom
- Quantum mechanical model of atom
- Dual nature of matter and radiation
- Photoelectric effect
- Atomic spectra
- De Broglie’s hypothesis
- Heisenberg’s uncertainty principle
- Quantum numbers and orbitals
- Aufbau principle, Pauli exclusion principle, and Hund’s rule
- Electronic configurations of atoms
Thomson’s Model of Atom
Thomson’s model of atoms was proposed by J.J. Thomson in 1898, based on his discovery of electrons using cathode ray tubes. He proposed that an atom is a sphere of positive charge in which electrons are embedded like seeds in a watermelon. He also calculated an electron’s charge-to-mass ratio (e/m) as -1.76 x 108 C g-1. This model could explain the electrical neutrality of atoms, but it failed to explain the results of later experiments, such as the scattering of alpha particles by thin metal foils.
! Thomson’s model)
Rutherford’s Model of Atom
Rutherford’s model of the atom was proposed by Ernest Rutherford in 1911, based on his famous alpha-particle scattering experiment. He bombarded a thin gold foil with a beam of alpha particles (positively charged helium nuclei) and observed their deflection using a fluorescent screen. He found that most alpha particles passed through the foil without deflection, small angles deflected some, and large angles deflected a few or even bounced back. He concluded that an atom consists of a tiny, dense, positively charged nucleus at the Centre, surrounded by a cloud of electrons at a large distance from the nucleus. He also calculated the size of the nucleus as about 10-15 m, which is much smaller than the size of the atom (about 10-10 m). This model could explain the stability and electrical neutrality of atoms, but it failed to explain the origin and nature of atomic spectra.
! Rutherford’s experiment)
Bohr’s Model of Atom
Bohr’s model of atoms was proposed by Niels Bohr in 1913, based on his study of hydrogen atomic spectra. He modified Rutherford’s model by introducing the concept of quantization of energy levels for electrons in an atom. He proposed that electrons revolve around the nucleus in fixed circular orbits, called shells or energy levels, with definite energies. He also proposed that electrons can jump from one orbit to another by absorbing or emitting a quantum of energy (photon) equal to the difference between their energies. He derived an expression for the radius and energy of an electron in an orbit as:
r_n = n^2 h^2 / (4π^2 m k e^2)
E_n = – m k e^4 / (2 h^2 n^2)
where n is the principal quantum number (1, 2, 3,…), h is Planck’s constant (6.626 x 10-34 J s), m is the mass of electron (9.109 x 10-31 kg), k is Coulomb’s constant (8.987 x 109 N m^2 C^-2), e is the charge on electron (1.602 x 10-19 C).
This model explained the origin and nature of hydrogen atomic spectra. Still, it failed to explain the spectra of other atoms, the fine structure of spectral lines, and the dual nature of matter and radiation.
! Bohr’s model)
Quantum Mechanical Model of Atom
The Quantum mechanical model of the atom is the most accepted and advanced model of the atom, based on the principles of quantum mechanics and wave mechanics. It was developed by various scientists, such as Louis de Broglie, Erwin Schrödinger, Werner Heisenberg, Max Born, and others, in the 1920s and 1930s. It considers an electron as a wave-particle duality, meaning it has both wave-like and particle-like properties. It describes an electron in terms of a wave function (ψ), which is a mathematical function that gives the probability of finding an electron in a given region of space. It also describes an electron in terms of four quantum numbers (n, l, m_l, m_s), which specify an atomic orbital’s size, shape, orientation, and spin. An atomic orbital is a three-dimensional region around the nucleus where the probability of finding an electron is maximum. The quantum mechanical model can explain the spectra of all atoms, the fine structure of spectral lines, and the dual nature of matter and radiation.
!Quantum mechanical model)
Dual Nature of Matter and Radiation
The dual nature of matter and radiation is the concept that matter and radiation have wave-like and particle-like properties. It was proposed by Louis de Broglie in 1924, based on his hypothesis that if radiation can behave like particles (photons), then particles can also behave like waves. He derived an expression for the wavelength (λ) associated with a particle of mass (m) and velocity (v) as:
λ = h / (m v)
where h is Planck’s constant.
This concept was experimentally verified by Davisson and Germer in 1927, who observed the diffraction pattern of electrons scattered by a crystal. This concept is also supported by Heisenberg’s uncertainty principle, which states that it is impossible to measure simultaneously both the position and momentum (or velocity) of a particle with absolute accuracy. The product of the uncertainties in position (Δx) and momentum (Δp) is always equal to or greater than h/4π.
Δx Δp ≥ h/4π
This implies that the more precisely we know the position of a particle, the less precisely we know its momentum, and vice versa.
Photoelectric Effect
The photoelectric effect is the emission of electrons from a metal surface when light of a suitable frequency falls on it. It was observed by Hertz in 1887 and explained by Einstein in 1905. It shows that light behaves like particles (photons) with energy proportional to their frequency (ν). The energy (E) of a photon is given by:
E = h ν
where h is Planck’s constant.
The photoelectric effect can be explained by assuming that when a photon strikes a metal surface, it transfers its energy to an electron in the metal. Suppose the energy of the photon is greater than or equal to the minimum energy required to remove an electron from the metal (called work function, φ). In that case, the electron is ejected from the metal surface with some kinetic energy (K). The kinetic energy of the ejected electron is given by:
K = h ν – φ
This equation is called Einstein’s photoelectric equation. It shows that the kinetic energy of the ejected electron depends on the frequency of the incident light and not on its intensity. The intensity only affects the number of electrons ejected per unit time.
The photoelectric effect has many applications, such as photovoltaic cells, photocells, photomultipliers, etc.
!Photoelectric effect)
Atomic Spectra
Atomic spectra are the spectra of electromagnetic radiation emitted or absorbed by atoms when they undergo transitions between different energy levels. Atomic spectra can be classified into two types: continuous spectra and line spectra. Continuous spectra contain all wavelengths or frequencies within a given range, such as the spectrum of white light or black body radiation. Line spectra contain only certain discrete wavelengths or frequencies corresponding to specific transitions between energy levels, such as the spectrum of hydrogen or sodium atoms.
Atomic spectra can identify elements, determine their atomic structure, measure their physical properties, etc.
!Atomic spectra)
De Broglie’s Hypothesis
De Broglie’s hypothesis is that if radiation can behave like particles (photons), then particles can also behave like waves. It was proposed by Louis de Broglie in 1924 based on his analogy between light and matter. He derived an expression for the wavelength (λ) associated with a particle of mass (m) and velocity (v) as:
λ = h / (m v)
where m is the particle’s mass, v is the particle’s velocity, and h is Planck’s constant. This hypothesis was experimentally verified by Davisson and Germer in 1927, who observed the diffraction pattern of electrons scattered by a crystal. This hypothesis implies that matter has both wave-like and particle-like properties and that the wavelength of matter depends on its momentum.
De Broglie’s hypothesis was a breakthrough in the development of quantum mechanics, as it provided a basis for the wave equation of Schrödinger, which describes the behaviour of electrons in atoms. De Broglie’s hypothesis also explains the origin of atomic spectra, as it suggests that electrons can only occupy those orbits in an atom with a circumference equal to an integral multiple of the wavelength of the electron. This means that only certain discrete energy levels are possible for electrons in an atom. They emit or absorb photons with specific frequencies when they jump from one level to another.
De Broglie’s hypothesis also has many applications in various fields of science and technology, such as electron microscopy, electron diffraction, electron beam lithography, etc.
Quantum numbers and orbitals:
Quantum numbers specify the size, shape, orientation, and spin of an atomic orbital. An atomic orbital is a three-dimensional region around the nucleus where the probability of finding an electron is maximum. There are four quantum numbers: principal (n), angular (l), magnetic (m_l), and spin (m_s). The principal quantum number (n) determines the energy level and the distance of the orbital from the nucleus. It can have values from 1 to infinity. The angular quantum number (l) determines the shape or sublevel of the orbital. It can have values from 0 to n-1. The magnetic quantum number (m_l) determines the orientation or direction of the orbital in space. It can have values from -l to +l. The spin quantum number (m_s) determines the spin or direction of rotation of the electron in the orbital. It can have values of -1/2 or +1/2. The different shapes of orbitals are denoted by letters: s (spherical), p (dumbbell), d (double dumbbell), and f (complex).
Aufbau principle, Pauli exclusion principle, and Hund’s rule:
These rules govern the filling of electrons in orbitals in an atom. The Aufbau principle states that electrons are filled in orbitals in order of increasing energy, starting from the lowest energy orbital. The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers, which means that each orbital can hold a maximum of two electrons with opposite spins. Hund’s rule states that when there are more than one orbital of equal energy (degenerate orbitals), electrons are filled singly with parallel spins before pairing up.
Electronic configurations of atoms:
Electronic configurations are the arrangements of electrons in orbitals in an atom. They are written using the notation n lx, where n is the principal quantum number, l is the letter denoting the shape of the orbital, and x is the number of electrons in that orbital. For example, the electronic configuration of hydrogen (H) is 1s1, which means that it has one electron in the 1s orbital. The electronic configurations of atoms follow the Aufbau principle, Pauli exclusion principle, and Hund’s rule, as well as some exceptions due to stability and symmetry factors.
This summary helps you to understand the structure of atoms better. If you have any questions or doubts, please feel free to ask me. I’m here to help. 😊