Home Chemistry CHEMISTRY: A Textbook for Grade 10   Chapter 3: Atomic Structure

E-Lecture - Important Points
  • The idea that matter consists of discrete indivisible particles called “atoms” was introduced by Democritus (460-370 BC).
  • Modern chemistry began with Dalton’s atomic theory, which states that all matter is composed of tiny, indivisible particles called atoms.
  • According to the modern atomic theory, an atom consists of a very dense central nucleus containing protons and neutrons, with electrons moving around the nucleus at a relatively large distance from it.
  • Protons are positively charged, neutrons have no charge, and electrons are negatively charged.
  • A proton and a neutron have approximately the same mass; but the mass of an electron is negligible.
  • Rutherford’s atomic model is that of a very small positively charged nucleus and extra-nuclear electrons.
  • The nucleus consists of protons and neutrons and contains practically all the mass of an atom.
  • Atomic masses and relative abundances of the isotopes of an element can be established by mass spectrometry.
  • The atomic number of an element is the number of protons in the nucleus of an atom of the element.
  • In an atom, the number of electrons equals the atomic number hence, an atom is electrically neutral.
  • The mass number is the sum of the number of protons and the number of neutrons in the nucleus of an atom.
  • An atom is represented by the notation, in which AZX is the symbol of an element Z is the atomic number, and A is the mass number.
  • Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
  • The atomic mass of an element is the average mass of its isotopes.
  • Electrons occupy certain energy levels outside the nucleus of an atom; these energy levels are also known as atomic shells.
  • The maximum number of electrons that can occupy each of the atomic shells is specified.
  • Electronic configuration is the arrangement of electrons in atomic energy levels.
  • The n, l and ml quantum numbers define an orbital, but a fourth quantum number is also required to characterize an electron in an orbital - the spin quantum number, ms. This quantum number may have either of two values: +½ or –½.
  • Aufbau Principle is a scheme used to reproduce the electron configurations of the ground states of atoms by successively filling subshells with electrons in a specific order (the building-up order).
  • Hund’s rule states that the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons.
  • Pauli’s Exclusion principle states that no two electrons in an atom can have the same four quantum numbers.
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