Atomic Structure and Periodicity

Atomic Structure and Periodicity

1. Trends on a Periodic Table

   • Electronegativity
     • Electronegativity is a measure of the attraction between a nucleus and a bonded pair of electrons
       • This attraction that an atom has for its bonded electrons depends on:
         • nuclear charge (i.e. the number of protons)
         • distance between the nucleus and bonded electrons (i.e. the atomic radius)
         • shielding effect of inner shells of electrons
     • Down a group
       • The bonded electrons are in an energy level further from the nucleus, with increased shielding from inner shells
         • This results in a decreased electronegativity, despite the increase in nuclear charge
     • Across a period (left to right)
       • The bonded electrons are in the same energy level but there is increased positive charge on the nucleus (more protons) so the attraction increases, as does the electronegativity
   • Ionisation Energy
     • The first ionisation energy, IE, is the minimum energy needed to remove one electron from each atom in a mole of atoms in the gaseous state
       • Measured in kJ mol-1
       • The process is endothermic because energy is required to overcome the electrostatic attraction between the nucleus and the valence electrons, and hence, to remove the electron
     • Down a group
       • There is a decrease in ionisation energies
         • The electron lost is further from the nucleus (by an extra shell of electrons) than the element above and there is more shielding by the increased number of inner shells of electrons (core electrons)
           • This results in a decrease in the electrostatic attraction between the nucleus and valence electron, and less energy is needed to remove the electron
     • Across a period (left to right)
       • Ionisation energies increase
         • As atomic number increases, there are more protons in the nucleus and electrons are being added to the same level with no additional shielding
           • As a result, there is a stronger electrostatic attraction between the nucleus and valence electrons, and more energy is needed to remove the electron
   • Atomic Radius
     • The actual sizes of atoms are calculated from the average distances between neighbouring atoms
     • Across a period (left to right)
       • The atomic number increases (because the number of protons increases)
       • Electrons are being added to the same energy level (with no added shielding from inner electron shells)
       • The atoms are smaller
         • The charge on the nucleus increases so the electrostatic attraction between nucleus and valence electrons increases, the valence electrons are pulled closer to the nucleus and the atomic radius decreases
     • Down a group
       • The positive charge of the nuclei increases
       • Valence electrons are in energy levels further from the nucleus, with increased shielding by inner shells of electrons
       • The electrostatic attraction between nucleus and valence electrons decreases and the radius increases

2. Electron Configuration

   • In the ground state of any atom, electrons always fill orbitals with the lowest energy - this is known as the Aufbau principle
   • Two atoms in the first row of transition elements have unusual electron arrangements
     • Chronium is [Ar]3d5 4s1
     • Copper is [Ar]3d10 4s1
     • These atoms gain stability by promoting an electron from the 4s orbital to 3d orbital
       • The half-filled 3d subshell in Cr, and filled subshell in Cu, make these atoms more stable
   • Electron configuration of ions
     • Ions and atoms are said to be isoelectronic when they have the same electron configuration
     • Transition metals lose electrons from the 4s orbitals before 3d

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