Periodic Trends

  • periodicity: Trends or recurring variations in element properties with increasing atomic numbers

  • electronegativity: The tendency of an atom to attract electrons in a bond to itself

  • ionization energy: Energy required to remove an electron from a gaseous atom or ion/remove from the 1rst electron in the valence electron’s shell

  • Atomic Radius: One half of the distance between the nuclei of 2 identical atoms when they are joined by a covalent bond

  • Nuclear Charge: The pull exerted on a specific electron by the nucleus (# of protons)

  • across a period, atomic radius decreases

    • sheilding is when another electron blocks the attraction of an additional electron to the nucleus

    • effective nuclear charge: the regular amount of protons, or the nuclear charge (z), with sheilding taken into account; zeff=z-sheilding

    • atomic radius decreases because the effective nuclear charge (zeff) increases while shielding remails constant (n is constant), and electrons are pulled closer by increasing atomic number/protons in the nucleus, causing more attraction and a smaller atomic radius

  • down a group, atomic radius increases

    • this happens because there are more energy levels; the principal quantum number (n, vertical numbers) increases as energy levels are added, giving rise to a larger atomic radius

  • Exceptions for atomic radius happen because a new subshell (s, p, d, f) is added to the outermost energy level; instead of decreasing, it increases

    • Mg—>Al

      • [Ne]s2—>[Ne]3s2,3p1

    • Zn—>Ga

      • [Ar] 4s2,3d10—>[Ar] 4s2,3d10, 4p1

  • across a period, ionization energy increases

    • the effective nuclear charge (zeff) goes up with the increasing atomic number, causing more attraction; this makes it more difficult to remove a valence a valence electron, meaning more energy will be required

  • down a group, ionization energy decreases

    • more energy levels are present, meaning that there are more core electrons (electrons close to the center), giving rise to more sheilding; there is therefore less attraction, making it easier to remove a valence electron and causing more blockage; this blockage makes it easier to take an electron away with less energy required

  • exceptions for ionization energy are because when 1 orbital is full or half full, it is more stable and harder for an electron to be removed, making it decrease

    • Be—>B

    • Mg—>Al

    • N→O

    • P—>S

  • Across a period, electronegativity increases

    • this is because there is more effective nuclear charge (zeff), meaning that it is more attractive to valence electrons

  • down a group, electronegativity decreases

    • there are more energy levels, meaning that there are more core electrons and therefore more sheilding; there is more distrance between valence electrons and the nucleusl it is less attractive to the electrons

  • exceptions happen for electronegativity when electrons aren’t necessary because an atom is already stable with a full outer shell

    • noble gases are already stable so they have very little electronegativity

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