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Trends of the Periodic Table
Review!• Periodic Table was first organized by…
– Dmitri Mendeleev in the mid 1800’s– Mendeleev organized the elements by chemical
reaction in rows, then by atomic mass in columns
• Henry Moseley then took Mendeleev’s table, kept the chemical reactivities together, but placed them in columns instead. He also ordered the elements by increasing atomic number in rows.
• When Moseley did this, all the periodic trends just fell into place.
• Remember: columns = groups/families, rows = periods
Periodic Trends
Electrons
• Electrons do not freely float in space
• Orbit around nucleus: Electron shells
• Each shell corresponds to an amount of energy.
Valence Electrons• The valence electrons are the outermost electrons of an atom.
• The valence electrons determine the chemical properties
• Number of valence electrons equals the column number in the “A” columns
• Elements with the same number of valence electrons are very similar chemically
– Alkali metals in Group 1A – 1 valence electron
Li, Na, K, Rb, Cs
– Halogens in Group 7A – 7 valence electrons
• F, Cl, Br, I
Atomic Radius• What is Atomic Radii?• Distance from the
nucleus to the outermost level of e- (aka the valence shell)
• What trend do you see as you go across (left to right) the period?
• Atomic radius decreases
• Down the group?• Atomic Radius
increases• WHY???
Explaining the Trend
• As you go L to R, the atomic radius decreases because as you go L to R, the amount of attraction between p+ and e- increase.
More attractions = smaller atomic radius• As you go down a column, atomic radius increases
because the e- are farther away from the nucleus. There are weaker attractions.
Weaker attractions = larger atomic radius
Electronegativity• What is Electro-
negativity?• An atom’s Luuuvvv
for electrons!• The tendency to
attract another atom’s electrons
• What trend do you see as you go across the period?
• Electronegativity increases!
• Down the group?• Electronegativity
decreases!• WHY???
Explaining the Trend
• As you go L to R, electronegativity increases because of the increase in protons. The more protons, the more able it will be to attract other atom’s electrons.
More attractions (small radius) = large electronegativity
• As you move down a column, electronegativity decreases because of the increase in number electron an atoms already has. This means the atom will be less able to attract another atom’s electrons.
• Less attractions (large radius) = small electronegativity
Ionization Energy• What is Ionization
Energy?• The energy needed to
remove an electron• What trend do you see
as you go across the period?
• Ionization E increases• Down the Group?• Ionization E decreases• WHY???
Explaining the Trend
• As you go L to R, the ionization energy increases because of the increase in the number of protons. The more protons, the more energy that is needed to remove an electron.
More attractions (small radius) = large ionization energy• As you go down a column, the ionization energy decreases
because of the decrease in attractions. – Due to electron shielding– More electrons, leads to outer electrons less tightly held.
• The less attractions, the lower the energy that is needed to remove an electron.
Less attractions (large radius) = small ionization energy
Ionization Energy
• Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.– First ionization energy is that energy required
to remove first electron.– Second ionization energy is that energy
required to remove second electron, etc.
Ionization Energy
• It requires more energy to remove each successive electron.
• When all valence electrons have been removed, the ionization energy takes a quantum leap.
Electron Affinity• What is Electron
Affinity?• The energy needed to
add an electron• As you go across the
period electron affinity increases .
• Electron affinity decreases down the family
• WHY???
Explaining the trend• As you go L to R, the electron affinity increases because of
the increase in the number of protons. The more protons, the greater the attraction the protons have for electrons.
More attractions (small radius) = large electron affinity• As you go down a family, the electron affinity decreases
because of the decrease in attractions. – Due to electron shielding– More electrons, leads to outer electrons less tightly
held.• The less attractions, the lower the electron affinity Less attractions (large radius) = small electron affinity
Homework
• Worksheet(s)