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Periodic trends: Atomic radius

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Intros
Lessons
  1. Chemical "Forces of attraction"
  2. Principles of electrostatic forces.
  3. Using the principles to explain periodic trends.
  4. Trend in melting point across a period.
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Examples
Lessons
  1. Explain trends in atomic radius down a group.
    Explain the trend in atomic radius observed in group 1 of the Periodic Table.
    1. Explain trends in atomic radius across a period.
      Explain the trend in atomic radius observed going across Period 2 of the Periodic Table.
      Topic Notes
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      In this lesson, we will learn:
      • To understand the principles of electrostatic forces and how they are used to explain experimental data.
      • To explain trends in atomic radius down a group using principles of electrostatic forces.
      • To explain trends in atomic radius across a period using principles of electrostatic forces.
      • To understand and explain the trend in melting and boiling points of elements across a period.

      Notes:

      • Chemists have found, through experimenting, some principles of electrostatic forces – forces that exist because charged particles attract or repel each other. The principles are:
        • #1: Oppositely charged particles attract each other, while particles of like charge repel each other.
        • #2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force between a 2+ ion and a 2- ion is stronger than the attractive force between a 1+ ion and a 1- ion).
        • #3: Attractive forces between oppositely charge particles decrease with distance.
        • #4: Repulsive forces between like-charged particles decrease with distance.

      • These principles form a theory that helps explain the trends that chemists see in their experimental data, such as in the change in atomic radius and first ionization energies of the elements

      • Atomic radius measures the distance between the nucleus and the outermost electron(s). There is a clear trend in atomic radius when going down the elements in a group or moving across elements in a period. Using the principles of electrostatic forces, we can explain both trends.

      • Going down a group of elements:
        • Each element further down the group has an extra inner shell of negatively charged electrons between the outermost electrons and the positively charged nucleus.
        • These negative inner electron shells are also attracted to the positive nucleus (see #1 above), and are ‘shielding’ the positive charge of the nucleus from the outermost electron shell. This offsets the extra positive charge from the extra protons in the nucleus.
        • In effect, going down a group, the atomic radius is determined by the number of inner electron shells between the nucleus and the outer electron shell.
        • These extra inner electron shells repel the outer electron shell (see #1) as both are negatively charged. Being too close to the inner electron shells would cause repulsion (see #4). To reduce this, the outer shell is pushed further away from the nucleus due to the repulsion and so it is less attracted to the nucleus (see #3 above). This leads to larger atomic radius going down the group.
      • Going across a period of elements:
        • Each element further across the period has an extra proton in its nucleus, strengthening its positive nuclear charge, and an extra negative electron in its outer shell which is attracted to the nucleus (see #1).
        • This extra positive nuclear charge and extra negative charge of the outer shell electrons leads to a greater force of attraction (see rule #2) and this effect is stronger than the repulsion (see #1) of adding one extra electron to the outer shell of electrons. This causes the outer electrons to be drawn in closer to the nucleus. Because going across a period does not add extra electron shells, there is no extra effect of electron shielding.

      • Melting and boiling points across a period also change across a period for a similar reason to the change in atomic radius
        • From Na through to Al, the elements have a giant metallic structure. This is a giant lattice made of positive metal ions surrounded by an attractive force of delocalized electrons – we call this metallic bonding.
        • Going from left to right, the metal ions of the lattice are increasingly positive (Na+ \enspace → \enspace Mg2+ \enspace → \enspace Al3+) and they each attract more moles of electrons per ion:
          • Na \enspace → \enspace Na+ and one mole of electrons in the lattice
          • Mg \enspace → \enspace Mg2+ and two moles of electrons in the lattice
          • Al \enspace → \enspace Al3+ and three moles of electrons.
        • This creates stronger metallic bonding by principle #2 above – a greater charge difference between positive metal ions and the moles of electrons holding the structure together. This explains the melting/boiling points in Al being substantially higher than Mg, which is higher than Na.