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5.3 Periodic Trends (Shannon Newman)
Added by Mary Blackwell, last edited by Shannon Newman on Oct 19, 2007

Ionization Energy

  • Ionization Energy is the energy needed to remove an electron from and in atom in the gas phase.
    • This creates ions and gives evidence of the shell model
    •  Francium (F) has the lowest ionization energy while Helium (He) has the highest.
  • M(g)=M+(g) + e-
  • As more shells are added, electrons become farther away and the force exerted on them by the nucleus decreases. Also, the electrons farther away are repelled by inner electrons in a process called shielding.
  • Core charge: atomic number-number of electrons on the inner shells
  • The ionization energy increases across a period as core charge increases and decreases down a group as the distance from the nucleus increases.
  • The force of attraction between the nucleus and electrons is greater as it goes across the table. This is called an electrostatic force.
    • Electrostatic force=k(Q1xQ2/r2)
    • K=constant, r=distance between charges, q1 and q2 are the charges
  • The most commonly referred to ionization energy is the first, which involves removing one electron but there can also be a second and third ionization energy which entails removing a second and third electron respectively.

Ionization Energies. http://www.webelements.com/webelements/properties/media/tables/intensity/ionization-energy-1.gif


Successive Ionization Energies. http://www.chemistry210.com/notes/pg328a.gif 

 Electron Affinity

  • Electron affinity is the energy needed to add an electron to an atom. Oftentimes, this energy is released in what is called an exothermic process.
    •  Fluorine has the highest and Francium has the lowest. This means that fluorine readily accepts electrons while francium doesn't.
  • X + e- = X- + Energy (electron affinity)
    • Example: 1 mole H + 1 mole e- = H- + 72.22 kJ/mole
  • Electron affinity increases with an increasing core charge. This is because, as the force with which the electrons are pulled into the valence shall of an atom increases, more energy is released when electrons are added because it is easier. There is a more natural force.
  • Some exceptions in which the process is endothermic, or energy is required to add the electrons.


  • Electronegativity is the ability of an atom to attract electrons to itself.
    • Often, this involves attracting electrons in a bond.
    • The electronegativity decreases going down a group because the electrons get farther away and increases going to the right across a period as the core charge increases.
    • Helium, Argon, and Neon, don't have any electronegativity values because they are not normally found in compounds so there is nothing to base the values off of.
    • Flourine is the most electronegative.
    • The electronegativity of an element is not a constant and varies depending on what molecule it's in.

Electronegativity Table. http://www.webelements.com/webelements/properties/media/tables/intensity/electroneg-allred.gif

Atomic Radii

  • The valence shell has no boundary and has an uneven density.
    • Therefore it is necessary to calculate the radius by taking 1/2 the distance between nuclei or diatomic molecules.
      • example: H-H is 74 picometers from nucleus to nucleus so one H atom has a radius of 37 picometers
  • Going down the periodic table, radii become larger as more shells are added.
  • Going across to the right in the periodic table, atomic radii become smaller as the core charge pulls the valence electrons toward the core.
    Atomic Radii. http://www.webelements.com/webelements/properties/media/tables/balls/atomic-radius.gif 

Ionic Radii

  • Cations are smaller and anions are larger than their corresponding atoms.
    • Cations lose electrons and sometimes a shell entirely so become smaller. Also, a greater core charge is pulling in fewer electrons.
    • Anions add electrons which increases shielding so it expands. 
      Ionic Radii in Comparison to Atomic Radii. http://boomeria.org/chemlectures/textass2/table10-9.jpg 

Information gathered from:

Gillespie, Ronald J., et al. Chemistry. Newton, Massachusetts: Allyn and Bacon, Inc., 1986.

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