The Nuclear Atom

Atomic Structure

    Rutherford's Scattering Experiments:

  • Ernest Rutherford fired alpha particles at a sheet of gold foil
  • The experiment resulted in some particles reflecting back
  • Rutherford concluded that an atom is mostly empty space occupied by negative charged electrons, while having a positive and dense nucleus

    Basic Atom Structure:

  • Protons (+ charge) are located in the nucleus, and have a relative mass of 1
  • Neutrons (0 charge) are located in the nucleus, and have a relative mass of 1
  • Electrons (- charge) are orbiting around the nucleus, and have a relative mass of 1/2000

Nuclear Symbol Notation

    Notation Components:

  • Atomic Number (Z): the number of protons in an atom’s nucleus
  • Nucleon/Mass Number (A): the sum of the number of protons and the number of neutrons in an atom’s nucleus
  • Nuclear Symbol (X): represents the identity of an element, using its chemical symbol
  • Notation: AZX
    • Example: Iron with 26 protons and 28 neutrons → 5426Fe

    Ions:

  • Formed when neutral atoms either lose or gain electrons, creating positive/negative ions
  • A positive ion's charge is shown in nuclear symbol notation by (# of protons - # of electrons)+
    • Example: 3919K2+ → 19 protons, 20 neutrons, 17 electrons
  • A negative ion's charge is shown in nuclear symbol notation by (# of electrons - # of protons)-
    • Example: 3919K2- → 19 protons, 20 neutrons, 21 electrons

Isotopes

    Definition:

  • Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons
  • Because isotopes only differ by the number of neutrons, their masses are different but their charges are the same
  • Thus, their chemical properties remain the same, but their physical properties can vary

    Example:

  • Hydrogen has three common isotopes:
    • Hydrogen-1 (protium): 1 proton, 0 neutrons
    • Hydrogen-2 (deuterium): 1 proton, 1 neutron
    • Hydrogen-3 (tritium): 1 proton, 2 neutrons

    Relative Atomic Mass (Ar):

  • The weighted average mass of an atom compared to 1/12 the mass of an atom of carbon-12
  • Formula: Ar = (Isotope1 Mass * Isotope1 % Abundance + Isotope2 Mass * Isotope2 % Abundance ...) / 100%
  • Example: Magnesium has three isotopes:
    • Mg-24 (78.99%), Mg-25 (10.00%), Mg-26 (11.01%)
    • Relative atomic mass = (24 * 78.99% + 25 * 10.00% + 26 * 11.01%) / 100% = 24.32

Mass Spectrometry

    Purpose:

  • Determines isotope abundance by measuring mass-to-charge ratio (m/z) of ions
  • Allows for the calculation of relative atomic masses from experimental data

    Main Stages:

  • Vaporisation: the sample is heated/vaporized to produce gaseous molecules
  • Ionisation: the molecules are bombarded with high-energy electrons, knocking off their own electrons, where only ions with a single positive charge are formed
  • Acceleration: these charged particles are then accelerated through slits in parallel plates under the influence of an electric field
  • Deflection: the paths of the ions are bent by a magnetic field, where lighter ions deflect more than heavier ones
  • Detection: the gaseous ions are detected with a sensor, and a signal is sent to a recorder to produce a mass spectrum

    Diatomic Elements in Mass Spectra:

  • Elements like Cl2, O2, and H2 exist as diatomic molecules
  • Different isotope combinations produce multiple peaks (e.g. 35Cl2, 37Cl2, 35Cl37Cl)
  • Keep in mind the different possibilities for relative atomic masses and abundances involved with diatomic elements