Download Chemical Periodicity

Chemical Periodicity
Nuclear Charge
• the total charge of all the protons in the
nucleus
• equals the Atomic Number
• GROUP: INCREASES
• ROW : INCREASES
why?
as the nuclear charge increases, the electrons
are pulled tighter to the nucleus
Electron Shielding…electrons
between the nucleus and the valence electrons
“shield” the valence electrons from the force of
attraction exerted by the positive charge in the
nucleus.
GROUP
 Shielding increases as you go down a group - each
row increases the number of electrons a bond must
“swim” through
ROW
 Shielding effect is constant within a given period
– the elements are in the same row, and therefore must
swim through the same number of electrons
Atomic Radius/Size…distance
between the nucleus and the outer ring
 the size of an atom depends on how far
away the valence/outermost electrons
are from the nucleus
 if the electrons are very close to the
nucleus, then the atom will be very small
 if the electrons are very far from the
nucleus, then the atom will be large
GROUP
• Atomic size increases as you move down a
group (column/family)
WHY? adding principle energy levels
ROW
• Atomic size decreases as you move left to
right across a row (period/series)
• WHY? added protons and electrons are on
the same principal energy level; increase in
nuclear charge pulls electrons closer to the
nucleus, thereby decreasing size
IONIC SIZE…size of the ion formed
when the atom loses or gains electrons
• positive ions (cations) are smaller than the
neutral atom (#p > #e)
• negative ions (anions) are larger than the neutral
atom (#p < #e)
GROUP
 increases
ROW
 decreases
(adding energy levels)
(nuclear charge!)
IONIZATION ENERGY… the
energy required to remove an electron (from a
gaseous atom, resulting in a positive ion)
First Ionization Energy
 energy required to remove the first outermost
electron
Second Ionization Energy
 energy required to remove the second outermost
electron
** the higher the ionization energy, the harder it is to
remove an electron
GROUP
 IE decreases as you move down a group
WHY? electrons further from nucleus more easily
removed
ROW
 IE increases as you move left to right across
a row (for representative elements)
WHY? nuclear charge increases and therefore,
greater attraction between protons in the nucleus
and the valence electrons
ELECTRONEGATIVITY… tendency
for the atom to attract electrons when they are chemically
combined with another element.
GROUP
 Decreases down a group
WHY? due to addition of energy levels and the
shielding effect
ROW
 Increases left to right (for representative
elements)
WHY? due to an increased number of electrons
in the outer ring
*** Fluorine is the MOST electronegative element