Brief Summary
This video provides a comprehensive one-shot explanation of the structure of an atom for Class 11 students. It begins with the basics of matter and progresses through the historical discoveries and models of the atom, including Dalton's atomic theory, Thomson's plum pudding model, Rutherford's gold foil experiment, and Bohr's atomic model. It also covers key concepts such as atomic number, mass number, isotopes, isobars, isotones, electromagnetic radiation, black body radiation, the De Broglie hypothesis, the Heisenberg uncertainty principle, and the quantum mechanical model of the atom. Finally, the video explains the Aufbau principle, Hund's rule, and how to write electronic configurations for the first 30 elements, including exceptions like chromium and copper.
- Matter is anything that takes up space and has mass, and it's composed of atoms.
- The structure of the atom has evolved through various scientific theories, culminating in the quantum mechanical model.
- Key concepts include atomic number, mass number, isotopes, isobars, and isotones.
- Quantum numbers (principal, azimuthal, magnetic, and spin) define the state of an electron in an atom.
- Electronic configurations follow the Aufbau principle and Hund's rule, with exceptions for chromium and copper.
Matter
Matter is defined as anything that occupies space and possesses mass. Atoms serve as the fundamental building blocks of matter, similar to how bricks construct a house. Various examples of matter include flowers, fruits, and everyday objects.
Introduction to atom
The term "atom" originates from the Greek word "atomos," signifying indivisible or unbreakable. John Dalton discovered the atom in 1808 and introduced the atomic theory, asserting that atoms are the fundamental units of matter. Different materials exhibit distinct properties due to their composition of different atoms. The video outlines the progression of atomic theories from Dalton's to Thomson's, Rutherford's, Bohr's, and finally, quantum mechanics.
Discovery of electron
Dalton's atomic theory faced limitations as it failed to explain the existence of subatomic particles. The discovery of electrons and protons led to the downfall of Dalton's theory. Michael Faraday initiated the study of electrical discharges in partially evacuated tubes, known as cathode ray discharge tubes. This setup includes a glass tube with a cathode (negative electrode) and an anode (positive electrode) connected to a battery. A vacuum is created within the tube by removing gases to achieve very low pressure. By applying high voltage across the electrodes, a current flows, and particles move from the negative to the positive electrode. The flow of current from the cathode to the anode is verified by creating a hole in the anode and coating the tube behind the anode with zinc sulphide, causing a bright spot when rays strike it.
Charge to mass ratio
Cathode rays originate from the cathode and move towards the anode. These rays are invisible but can be observed using fluorescent materials. In the absence of electric or magnetic fields, these rays travel in a straight line. The behaviour of cathode rays in electric or magnetic fields is similar to that of negatively charged particles. The characteristics of cathode rays do not depend on the material of the electrodes or the nature of the gas in the cathode tube. The charge-to-mass ratio is a physical quantity that relates the mass and electrical charge of a particle. It depends on the magnitude of the charge, the mass of the particle, and the strength of the electric and magnetic fields. The ratio is calculated as 1.75 * 10^11 coulombs per kilogram.
Discovery of protons
Electrons are fundamental constituents of all atoms and carry a negative charge of -1. The mass of an electron is 9.11 * 10^-28 grams, and its exact charge value is -1.6 * 10^-19 coulombs. Protons were discovered through anode rays, with Eugene Goldstein noting positively charged particles in cathode ray experiments in 1886. Atoms are neutral, implying the existence of positive charges. Particles move in the opposite direction of cathode rays and are called canal rays. Unlike cathode rays, the mass of positively charged particles depends on the nature of the gas in the cathode tube, as these are simply positively charged gaseous ions. The charge-to-mass ratio of the particles depends on the gas from which they originate.
Thompson model of an atom
The smallest and lightest positive ion is obtained from hydrogen and is called a proton, carrying a charge of +1. Its mass is 1.67 * 10^-24 grams. Neutrons were discovered in 1932 by James Chadwick, who identified a subatomic particle with no charge but with a mass equal to that of a proton. Neutrons are present in the nucleus of every atom except hydrogen. The mass of an atom is determined by the sum of protons and neutrons in the nucleus. Protons and neutrons combined are called nucleons. J.J. Thomson proposed the plum pudding model, suggesting that atoms are positively charged spheres with negatively charged electrons embedded within, like raisins in pudding. The magnitude of negative and positive charges is equal, making the atom electrically neutral.
Rutherford atomic model
In 1908, Ernest Rutherford conducted the gold foil experiment, firing positively charged alpha particles at a thin gold sheet. The expected result was slight deflection of alpha particles. However, observations included most alpha particles passing straight through the foil, some being deflected at small angles, and very few bouncing back. These observations led to the conclusion that most of the space inside an atom is empty, positive charge occupies minimal space, and there is a mass in the atom which is double the mass of the proton. Rutherford's atomic model led to the discovery of the nucleus, which contains protons and neutrons.
Atomic number & mass number
The movement of electrons was not specified, leading to the expectation of instability. Any particle in a circular orbit undergoes acceleration and radiates energy, causing the revolving particle to fall into the nucleus. Rutherford's model could not explain the stability of an atom, the dual character of electromagnetic radiation, or experimental results regarding atomic spectral emission lines. The atomic number (Z) is the number of protons in an atom. For an element, Z = number of protons (p) = number of electrons (e). The mass number (A) is the sum of neutrons and protons in the nucleus. A = number of protons + number of neutrons.
Isotopes
Isotopes are atoms of the same element with the same atomic number but different mass numbers. Hydrogen has three isotopes: protium, deuterium, and tritium. Isotopes have similar chemical properties but different physical properties.
Isobars
Applications of isotopes include using uranium isotopes as fuel in nuclear reactions, cobalt isotopes in cancer treatment, and iodine isotopes in goitre treatment. Isobars are atoms of different elements with the same mass number but different atomic numbers.
Isotones
Calcium and argon are examples of isobars. Isotones are atoms with different atomic numbers and mass numbers but the same number of neutrons. Calcium and potassium are examples of isotones.
Bohr’s atomic model
In 1913, Niels Bohr proposed improvements to Rutherford's atomic model, placing each electron in a specific energy level. Bohr's model incorporated the dual character of electromagnetic radiation and addressed the instability of Rutherford's model. Electrons occupy stable orbits with special, discrete locations. Atoms have a centre called the nucleus, and electrons revolve only in fixed circular orbits with fixed energy and velocity. The quantization condition states that electrons revolve only in those circular orbits for which their angular momentum is an integral multiple of h/2π. Energy is emitted or absorbed only when electrons jump from a higher energy level to a lower energy level or vice versa.
Line spectrum of hydrogen
The Bohr frequency rule is expressed as hν = E2 - E1. The most stable state of an atom is its ground state or normal state. The spectrum of radiation emitted by a substance that has absorbed energy is called the emission spectrum. The study of emission or absorption spectra is referred to as spectroscopy. The spectrum of visible light is continuous, while the emission spectra of atoms in the gas phase do not show a continuous spread of wavelengths. Instead, they emit light only at specific wavelengths, known as line spectra or atomic spectra. The equation for the line spectrum of hydrogen is 1/λ = RZ^2 (1/n1^2 - 1/n2^2), where R is the Rydberg constant, Z is the atomic number, n1 is the energy level the electron is coming to, and n2 is the energy level the electron is coming from.
Electromagnetic Radiations
When an electric discharge is passed through gaseous hydrogen, the hydrogen molecules dissociate, and the energetically excited hydrogen atoms emit electromagnetic radiation of discrete frequencies. The hydrogen spectrum consists of several series, each named after its discoverer. If n = 1, the series is called the Lyman series, and the radiations are ultraviolet. If n = 2, the series is called the Balmer series, and the radiations are visible. If n = 3, the series is called the Paschen series, and the radiations are infrared. James Maxwell explained the interaction between charged bodies and the behaviour of electric and magnetic fields. When an electrically charged particle moves under acceleration, it produces alternating electric and magnetic fields that are transmitted in the form of waves, known as electromagnetic waves or electromagnetic radiations.
Black body radiations
Electromagnetic waves consist of oscillating electric and magnetic fields that are perpendicular to each other and to the direction of propagation. These waves do not require a medium and can move in a vacuum. Types of electromagnetic radiation include gamma rays, X-rays, ultraviolet rays, infrared rays, and radio waves. In a vacuum, all types of electromagnetic radiation travel at the speed of light. Wavelength (λ) is the distance between two consecutive crests or troughs. Frequency (ν) is the number of waves passing through a point in one second. Velocity (c) is the linear distance travelled by a wave in one second. Amplitude (a) is the height of the crest or depth of the trough. Wave number (ν̄) is the reciprocal of the wavelength. A black body is an ideal body that emits and absorbs radiations of all frequencies uniformly.
De Broglie Hypothesis
The black body spectrum depends only on the temperature of the object and not on its composition. As the temperature of an object increases, it emits more black body energy at all wavelengths. The peak wavelength of the black body spectrum becomes shorter as the temperature increases. The particle nature of electromagnetic radiation is explained by Planck's quantum theory. The photoelectric effect involves the ejection of electrons from a metal surface when radiation strikes it. Albert Einstein explained the photoelectric effect in 1905, proposing that light is made up of tiny particles of energy called photons. Light exhibits a dual nature, behaving as both a wave and a particle. De Broglie proposed that if light, which is a wave, can exhibit particle behaviour, then matter, such as electrons, which are particles, can exhibit wave behaviour.
Heisenberg uncertainty principle
Every moving material particle is associated with a wave. The De Broglie wavelength is given by λ = h/mv, where h is Planck's constant, m is mass, and v is velocity. The wavelength of matter is inversely proportional to the magnitude of its linear momentum. Davisson and Germer performed experiments to explain the wave nature of electrons through electron diffraction, proving De Broglie's hypothesis. The Heisenberg uncertainty principle states that it is impossible to measure or calculate exactly both the position and the momentum of an object. The product of the uncertainty in position (Δx) and the uncertainty in momentum (Δp) is greater than or equal to h/4π.
Quantum mechanical model of an atom
The Heisenberg uncertainty principle is significant in the quantum world, where atoms and subatomic particles have very small masses. The quantum mechanical model of the atom takes into account the dual behaviour of matter. In an atom, electrons reside in shells, which contain subshells, which contain orbitals. An orbital is a 3D space where the probability of finding an electron is maximum. Atomic orbitals are precisely distinguished by quantum numbers. Each electron in an atom is identified by a set of four quantum numbers: the principal quantum number (n), the azimuthal quantum number (l), the magnetic quantum number (m), and the spin quantum number (s).
Pauli’s exclusion principle
The principal quantum number (n) indicates the main shell in which the electron resides and provides an idea about the energy of the shell. The azimuthal quantum number (l) tells us about the number of subshells in any main shell and represents the angular momentum of an electron and the shapes of the shells. The magnetic quantum number (m) tells us about the number of orbitals in a subshell and their orientation. The spin quantum number (s) indicates the direction of the spinning of the electron, either clockwise or anticlockwise. The maximum number of electrons that can be accommodated in a main energy level is 2n^2. Pauli's exclusion principle states that no two electrons in an atom can have identical sets of four quantum numbers. Only two electrons may exist in the same orbital, and these electrons must have opposite spins.
Afbau Principle
The Aufbau principle states that in the ground state of an atom, electrons fill atomic orbitals of the lowest available energy levels before occupying higher energy levels. Orbitals are filled in the increasing order of their energy. The energy is directly proportional to n + l. If two orbitals have equivalent values, the orbital with the lower value of n is said to have the lower energy.
Hund’s Rule
Hund's rule of maximum multiplicity states that for a given electronic configuration, the term with the maximum multiplicity falls lowest in energy. According to this rule, electron pairing in p, d, and f orbitals cannot occur until each orbital of a given subshell contains one electron each in a singly occupied manner.
Electronic Configuration of 30 elements
The video concludes by demonstrating how to write the electronic configurations of the first 30 elements, following the Aufbau principle and Hund's rule. It also mentions that there are two exceptions to the Aufbau principle: chromium and copper. These elements achieve greater stability by having either half-filled or fully filled d orbitals.
