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Periodic Trends and the Periodicity of Atomic Properties

Periodic Trends
  • Non-metals are on the right side of the periodic table, and metals are on the left
  • Down a period of metals, they become more reactive
  • With non-metals, going down a period, the non-metal becomes less reactive
Alkali Metals
  • Lithium, sodium, potassium
  • Patterns:
    • Reactivity, density and softness increases down the group
    • Melting and boiling points decrease down the group
    • Reactivity increases as more electron shells means there is a weaker attraction between nucleus and the valence electron, so it is lost more easily
Halogens

Transition Metals
  • They:
    • have high melting points (except mercury)
    • are malleable and ductile
    • are good conductors of heat and electricity (silver is the best)
    • have high density
    • have no trend in reactivity
    • can be used as catalysts
    • form colored compounds
    • can form complex ions as they have variable valencies
Noble Gases

Periodicity of Atomic Properties
  • An atomic property exhibits periodicity in the Periodic Table if its variation across the period or down the group occurs at regular intervals
Atomic Radius
  • Atomic radius is half the distance between two bonded atoms in a molecule. It may be taken as the distance between the center of the nucleus and the outermost shell of electrons
  • Factors that determine the atomic radius of an element include nuclear charge and screening effects
  • The nuclear charge acts as a positive center which pulls the negatively charged shells of electrons toward the nucleus. This tends to decrease the size of the atom.
  • The screening effect is the mutual repulsion between the negatively charged shells of electrons. This tends to increase the size of the atom and its effect increases as the number of shells increase
  • Nuclear charge effect and screening effect are therefore two opposing forces whose net effect will determine the atomic radius of an element
  • Atomic radius decreases from left to right across a period because nuclear charge effect (which increases progressively) outweighs screening effect which remains virtually constant because the number of shells in a period is constant
  • Atomic radius increases down the group because although the nuclear charge and screening effect increase down the group, the increase in nuclear charge is not enough to counteract the screening effect
  • The atomic radius of a metal is always bigger than the ionic radius of its corresponding positive ion. When a metal loses electron(s), its screening effect decreases while its nuclear charge remains unchanged.
  • The atomic radius of a negatively charged ion is always larger than its corresponding neutral atom. The extra electron(s) gained by a non-metal causes the screening effect to increase while the nuclear charge remains unchanged
Ionization Potential
  • Ionization potential (or Ionization energy) is the energy required to remove the most loosely held electron from an isolated gaseous atom or ion
  • First ionization potential is the energy required to remove the most loosely held electron from an isolated gaseous atom
  • Ionization potential increases from left to right across a period the nuclear charge increases progressively from left to right while the screening effect, which is proportional to the number of shells, remains constant. These factors combine to make the removal of the most loosely held electron difficult as the atoms get smaller.
  • Ionization potential but decreases from top to bottom within a group because screening effect outweighs he nuclear charge effect as the atoms get bigger in size
  • In each period, the noble gas at the end of the period has the highest ionization energy.
  • The trend of ionization potential within the period is not smooth. Discontinuities occur for beryllium and boron, as well as nitrogen and oxygen in the second period
  • The ionization potential of Be is higher than the ionization potential of B because the energy required to remove an unpaired electron in the 2p orbital of B is less than that required to remove one of the two paired electrons in the 2s orbital of Be.
  • The ionization potential of nitrogen is higher than that of oxygen because the half-filled 2p orbital in nitrogen is more stable than the 2p orbital in oxygen which is neither fully-filled nor half-filled
  • Second Ionization potential is defined as the energy required to remove the most loosely held electron from an isolated unipositive gaseous ion
  • Second ionization potential is always higher than the first ionization potential because it is more difficult to remove an electron from the positive ion which has a smaller size than its corresponding neutral atom
Electron Affinity
  • Electron affinity is defined as the energy released when an electron is added to an isolated gaseous atom to form a uninegative gaseous ion
  • Electron affinity increases from left to right across a period and decreases from top to bottom within a group
  • A lot of heat is liberated and the electron affinity is consequently high for small sized atoms such as F, O and Cl which can hold on tightly to the added electron
  • Contrary to expectation, the electron affinity of fluorine is smaller than that of chlorine due to its small size as there are already 7 electrons crowded into its outermost shell, an additional electron will experience a greater repulsion from these electrons than the attraction from the positively charged nucleus
  • The electron affinity of a noble gas is positive because energy is required to force an electron into the already fully-filled shell
  • Second electron affinity is the energy change associated with the addition of an electron to a uninegative gaseous ion. It is always positive because energy is required to overcome the repulsion between an anion and the electron to be added
Electronegativity
  • Electronegativity is defined as the ability of an atom to attract shared pair of electrons to itself when covalently bonded
  • Electronegativity increases from left to right across a period but decreases from top to bottom within a group
  • Fluorine is the most electronegative atom in the Periodic Table
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