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)