Transcript Periodic Trends Notes

Trends of the Periodic Table
Periodic Law
• When elements are arranged in order of
increasing atomic #, elements with similar
chemical and physical properties appear at
regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
Properties of Atoms
• Atomic Radius
– size of atom
• Ionization Energy
© 1998 LOGAL
• Energy required to remove an e- from a neutral
atom
• Electronegativity
• Energy required to remove
an e- from a neutral atom
© 1998 LOGAL
Shielding Effect
• Attraction between the positive nucleus and negative electrons
• All about what happens when electrons and energy levels are added
to an atom as you go through the PT
• Outer electrons feel a reduction of attractive force or “love” from
nucleus due to being farther away from nucleus
– Inner electrons shield the attractive force or “love”
• The same charge of electrons force inner to push away outer electrons
Periodic Trend,
1. Shielding effect increases down a group.
2. Shielding effect remains constant across a period.
Atomic Radius
• Atomic Radius = ½ the distance between
two identical bonded atoms
Atomic Radius
• What is Atomic Radius?
-- Distance from the nucleus
to the outermost level of
electrons
-- It’s the size of the atom
• What trend do you see as
you go across the period?
-- Atomic radius decreases
• Down the group?
-- Atomic radius increases
• WHY???
Atomic Radius
Explaining the Trend
• As you go across a period, the number of protons
increases (e- increase too, but on the same energy
level). More p+ can pull in e- closer, decreasing the
radius.
More attractions = SMALL atomic radius
• As you go down a group, e- are added to new energy
levels. Each level is further from the nucleus, which
increases the radius.
More energy levels = LARGE atomic radius
Atomic Radius
• Atomic Radius
Decreases to the Right and Up
Decreases
Decreases
1
2
3
4
5
6
7
Ionization Energy
• What is ionization energy?
-- The energy needed to
remove an electron
• What trend do you see
across the period?
-- Ionization Energy
increases
• Down the group?
-- Ionization Energy
decreases
• WHY???
Ionization Energy
Explaining the Trend
• As you go across a period, electrons are held more
closely because the atomic radius decreases. It takes
more energy to remove an electron, so the ionization
energy increases.
Small radius (more attractions) = HIGH ionization energy
• As you go down a group, electrons are further from
nucleus because the atomic radius increases. It takes
less energy to remove an electron, so the ionization
energy decreases.
Large radius (less attractions) = LOW ionization energy
Ionization Energy
• First Ionization Energy
Increases UP and to the RIGHT
Increases
Increases
1
2
3
4
5
6
7
Electron Affinity
• Most atoms can attract e- to form negatively charged ions
• The energy change that occurs when an e- is added to a
gaseous atom or ion.
• The ease with which an atom gains an e-.
Electron Affinity
• Electron affinity increases up a group
• decreases the atomic radius taking the electrons closer to
the nucleus’ positive attraction.
• decreasing shielding effect increases the effective
positive nuclear charge (+) as additional shells are added
and e- are held on tighter.
• Electron affinity increases across a period
– atomic radius decreases
– effective positive nuclear charge increases steadily
and the e- are drawn closer to the nucleus making it
easier to add e- to unfilled sublevels.
Electronegativity
• What is electronegativity?
-- An atom’s ability to
attract electrons
(An atom’s love of e-)
• What trend do you see as
you go across the period?
-- Electronegativity
increases
• Down the group?
-- Electronegativity
decreases
• WHY???
Electronegativity
Explaining the Trend
• As you go across a period, electrons are held more
closely because the atomic radius decreases. It is
easier to attract electrons, so electronegativity
increases.
Small radius (more attractions) = HIGH electronegativity
• As you move down a group, electrons are further
away from nucleus because the atomic radius
increases. It is harder to attract electrons, so
electronegativity decreases.
Large radius (less attractions) = LOW electronegativity
Electronegativity
• The measure of the ability of an atom to attract e- closer to itself
when forming a chemical bond with another atom.
Decreases
Increases
1
2
3
4
5
6
7
Examples
• Which atom has the larger radius?
Be or Ba
 Ca or Br
Examples
• Which atom has the higher ionization energy?
 N or Bi
Ba or Ne
Examples
• Which element has the higher electronegativity?
Cl or F
Be or Ca