Atomic Structure | General Awareness for SSC CGL PDF Download

Atoms and Molecules

Atom: The smallest unit of an element (or matter) that participates in a chemical reaction. Atoms consist of electrons, protons, and neutrons.

Molecule: The smallest unit of a substance that does not engage in a chemical reaction. Molecules are composed of two or more atoms, which can be the same or different types.

• Homoatomic Molecule: Consists of the same type of atoms (e.g., O₂, H₂, N₂).
• Heteroatomic Molecule: Consists of different types of atoms (e.g., HCl, N₂O, CO₂).

Fundamental and Subatomic Particles of an Atom

• Fundamental Particles: Electron, proton, and neutron.
• Subatomic Particles: Includes fundamental particles along with mesons, positrons, and neutrinos.

1. Electron:

   • Discovered by Joseph John Thomson.
   • Negatively charged (−1.6 × 10^-19 C) with a mass of 9.11 × 10^-28 g.
   • The charge was determined by Millikan in his oil drop experiment.

2. Proton:

   • Discovered by Ernest Rutherford.
   • Positively charged (+1.6 × 10^-19 C) with a mass of 1.673 × 10^-24 g.

3. Neutron:

   • Discovered by James Chadwick.
   • Neutral particle with a mass of 1.675 × 10^-24 g.

• All elements, except hydrogen, have atoms containing all three particles (protons, neutrons, electrons), including isotopes like deuterium and tritium.

Additional Subatomic Particles

• Positron: Discovered by Anderson in 1932, a positively charged counterpart of the electron.
• Neutrino and Antineutrino: Discovered by Fermi in 1934, particles with small mass and zero charge.
• Meson: Particles with mass between that of an electron and a proton, available in neutral, positively, and negatively charged forms. All mesons are unstable.
• Boson: Subatomic particles, which can be elementary (like photons) or composite (like mesons).
• Baryons: Particles heavier than the proton.
• Quark: Subatomic particles that combine to form hadrons, such as protons and neutrons. The antiparticle of a quark is an antiquark.

Some Important Terms related to Atom

• The atomic number of an element is equal to the number of protons present inside the nucleus of its atom.
  For an isolated atom,
  Atomic number (Z) = Number of protons
  = Number of electrons
• The atomic weight is also called relative atomic mass.
  
• Mass number is the sum of nucleons (protons and neutrons).

  Mass number (A) = Number of protons + Number of neutrons
  * The total number of atoms present in one molecule is called atomicity. For example, the atomicity of nitrogen (N₂) is 2 and of ammonia (NH₃) is 4.
  * One mole of a substance contains 6.023 × 10²³ atoms or molecules. This number is called Avogadro’s number.

Isotopes, Isobars, and Isotones

• Isotopes: Atoms of the same element with the same atomic number but different mass numbers (e.g., C-12 and C-14 for carbon,  , for hydrogen).
• Isobars: Atoms of different elements with the same mass number but different atomic numbers (Ca).
• Isotones: Atoms of different elements with the same number of neutrons but different atomic numbers ).

Atomic Models

To illustrate how fundamental particles are arranged in an atom, several atomic models have been proposed. Some significant models include:

Dalton’s Atomic Theory

Based on the laws of chemical combination, Dalton’s Atomic Theory made several assumptions:

• All matter is made up of indivisible and indestructible atoms.
• All atoms of a given element have identical properties, including identical mass.
• Atoms combine in small whole numbers to form compounds.
• Chemical reactions involve only the combination, separation, or rearrangement of atoms.

Thomson’s Atomic Model

• Each atom consists of a uniformly positively charged sphere with a radius on the order of 10⁻¹⁰ meters, in which the entire mass is uniformly distributed.
• Negatively charged electrons are embedded randomly within this sphere.
• The atom as a whole is neutral.

Limitations of Thomson’s Atomic Model:

• It could not explain the origin of the spectral series of hydrogen and other atoms.
• It failed to explain the large-angle scattering of α-particles.

Rutherford’s Atomic Model

• The entire positive charge and almost all the mass of the atom are concentrated in a very tiny region called the nucleus (about 10⁻¹⁵ meters in size).
• Negatively charged electrons revolve around the nucleus in different orbits.
• The total positive charge on the nucleus equals the total negative charge on the electrons, making the atom overall neutral.
• The existence of the nucleus was proven by Rutherford’s α-particle scattering experiment.
• A nucleus consists of positively charged protons and electrically neutral neutrons. The charge on an electron is 1.67 × 10⁻¹⁹.
• The centripetal force required for an electron to revolve around the nucleus is provided by the electrostatic force of attraction between the electrons and the nucleus.

Limitations of Rutherford’s Atomic Model:

• According to Maxwell’s electromagnetic wave theory, an accelerated charged particle emits energy in the form of electromagnetic waves. Therefore, an electron in its orbital motion under centripetal acceleration would emit energy, causing the radius of its path to decrease gradually, and ultimately, it would fall into the nucleus.

Electromagnetic Radiations

• Electromagnetic radiation is energy transported in the form of waves and is associated with both electric and magnetic fields.
• Various types of electromagnetic radiations exist, and in a vacuum, all types, regardless of their wavelength, travel at the same speed (3.00 × 10⁸ m/s^-1).

Applications of Different Regions:

• Radio Frequency Region (10⁶ Hz): Used for broadcasting.
• Microwave Region (10¹⁰ Hz): Used for radar.
• Infrared Region (10¹³ Hz): Heat radiation.
• Ultraviolet Region (10¹⁶ Hz): Component of the sun’s radiation.
• Visible Light Region (10¹⁵ Hz): Light visible to the human eye.

Photoelectric Effect:

• When certain metals like potassium, rubidium, and cesium are exposed to light of an appropriate frequency, electrons are ejected from their surface. This phenomenon is known as the photoelectric effect.

Planck’s Quantum Theory

• In 1900, Planck proposed a revolutionary theory of radiation known as the quantum theory.
• According to this theory, radiant energy is emitted or absorbed discontinuously in small packets of energy called photons (quanta).
• The amount of energy associated with a quantum of radiation is proportional to the frequency of radiation:
  E = hν (where E is energy, ν is frequency, and h is Planck’s constant).

Bohr’s Model

Based on the quantization of energy, Bohr’s Model includes several postulates:

• Electrons revolve in fixed circular orbits around the nucleus without losing or gaining energy.
• Electrons can only move in orbits where their angular momentum is an integral multiple of $\frac{\mathrm{h/}}{}$2π (quantized).
• Electrons have definite energy characteristic of the orbit in which they are moving. As long as the electron remains in an orbit, it does not lose energy. An electron in n=1 has the lowest possible energy.
• Absorption or emission of energy occurs only through the transition of electrons between lower and higher energy levels.

Heisenberg’s Uncertainty Principle

• According to this principle, "it is impossible to simultaneously measure the position and determine the velocity or momentum of a microscopic particle."
• Heisenberg’s principle is not applicable to macroscopic objects (i.e., large objects).

Shells and Energy Levels

Electrons possess specific energies characteristic of the orbits they occupy, also known as stationary orbits or shells. These orbits are denoted as follows:

• n = 1, 2, 3, 4
• Shells = K, L, M, N

The shell with n = 1 is closest to the nucleus, and electrons in this level have the lowest possible energy due to their proximity to the positively charged nucleus.

Distribution of Electrons into Different Orbits of an Atom

Suggested by Bohr and Burry, the distribution rules are:

• The maximum number of electrons in a shell is given by the formula 2n² (where n = 1, 2, 3, and 4 for K, L, M, and N shells respectively).
• The maximum number of electrons that can be accommodated in the outermost orbit is 8.
• Shells are filled stepwise.

Electronic Configuration

This is the arrangement of electrons in various shells, subshells, and orbitals in an atom. It can be written as 2, 8, 8, 18, 18, 32 or in the form nl^x (where n is the principal quantum number, l is the azimuthal quantum number or subshell, and x is the number of electrons).

For example:

• Number of electrons in n shell = 2n²
• In the second shell: number of electrons = $2\times 2\mathrm{<sup\; style="word-break:\; break-word;\; line-height:\; 1.8;\; color:\; black;">2</sup>}$= 8

Filling of Orbitals in Atoms

The filling of electrons into orbitals follows the Aufbau Principle, Pauli’s Exclusion Principle, and Hund’s Rule of Maximum Multiplicity:

• Aufbau Principle: Electrons enter the orbital of lowest energy first, followed by subsequent electrons in increasing energy order.
• Energy Order: Lower (n + l) value for an orbital means lower energy. If two orbitals have the same (n + l) value, the orbital with a lower n value has lower energy.
• Hund’s Rule of Maximum Multiplicity: Electron pairing does not occur in orbitals of the same energy until each orbital is singly filled with parallel spins.

Quantum Numbers

Each electron in an atom is characterized by a set of three quantum numbers, with a fourth needed to specify the electron's spin:

• Principal Quantum Number (n): Determines the main energy level or shell (K, L, M, N). As n increases, the electron's energy increases.
• Azimuthal Quantum Number (l): Determines the subshell (s, p, d, f) in a given principal energy level.
• Magnetic Quantum Number (m): Provides information about the orbital's orientation.
• Spin Quantum Number (s): Describes the electron's spin orientation (clockwise or anticlockwise).

Pauli’s Exclusion Principle

Proposed by Wolfgang Pauli in 1952, this principle states that no two electrons in an atom can have the same set of four quantum numbers.

Radioactivity

Radioactivity is a nuclear property of some elements where heavy elements disintegrate into lighter elements by emitting radiations, discovered by Henri Becquerel in 1896. Types of radiation include:

• α-rays: Consist of doubly ionized helium ions.
• β-rays: Consist of fast-moving electrons.
• γ-rays: Consist of electromagnetic rays.

Radioactive Emissions:

α-particle emission: Decreases atomic number by 2 and mass number by 4.

β-particle emission: Increases atomic number by 1 and keeps the mass number unchanged.

Soddy Fajan’s Group Displacement Law

This law states:

• Emission of an α-particle results in the formation of a new element two places to the left of the parent element in the periodic table.
  
• Emission of a β-particle results in the formation of a new element one place to the right of the parent element.
  

Emissions and Isotopes:

• Emission of one β-particle results in the formation of an isobar.
• Emission of one α-particle followed by two β-particles results in the formation of an isotope.

Nuclear Isomers and Isodiaphers:

• Nuclear Isomers: Atoms with the same atomic and mass numbers but different radioactive properties (e.g., Uranium with different half-lives).
• Isodiaphers: Atoms with the same difference of neutrons and protons or the same isotopic number (e.g., Uranium-235 and Thorium-231).

Applications of Radioactivity and Radioisotopes

Geological Dating: Uses the radioactive properties of uranium to determine the age of rocks by measuring the ratio of Uranium-239 to Lead-206.

Tracer Technique: Radioactive isotopes are added to study their path by measuring radioactivity, used in medical diagnosis and studying reaction mechanisms.

Medical Treatment:

• Co-60 for treating cancerous tumors.
• P-32 for blood cancer (Leukemia).
• I-131 for thyroid disorders.
• Na-24 for detecting defects in blood circulation.
• Re-59 for detecting anemia.
• Se-79 for detecting pancreatic gland disorders.

• Neutron Activation Analysis: Measures the concentrations of elements in small amounts without destroying the sample.
• Agriculture: P-32 studies phosphorus uptake and mineral salt transportation within plants.

Nuclear Hazards and Safety Measures

• Nuclear radiations can have devastating effects, causing somatic effects (e.g., cancer) and genetic effects (e.g., harmful genetic mutations passed to the next generation).
• In nuclear reactors, radioactive materials are continuously produced, and spent fuel rods become highly radioactive. Proper storage and care are needed to prevent leakage, which can contaminate water bodies and soil, posing risks to human, plant, and animal life.