Transcript first ionization energy. - Greenwich Public Schools

Section 14.2
Periodic Trends
OBJECTIVES:
 Interpret group trends in atomic
radii, ionic radii, ionization energies,
and electronegativities.
 Interpret period trends in atomic
radii, ionic radii, ionization energies,
and electronegativities.
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Trends in Atomic Size
First problem: Where do you start
measuring from?
 The electron cloud doesn’t have a
definite edge.
 They get around this by measuring
more than 1 atom at a time.

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Atomic Size
}
Radius
Atomic Radius = half the distance between
two nuclei of atoms


in the solid state (by X-ray diffraction)
or of a diatomic molecule.
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Atomic Size
Influenced by three factors:
1.
Energy Level
•
Higher energy level is further away.
Charge on nucleus
2.
•
More positive charge pulls electrons in
closer.
Shielding effect
3.
•
The inner electrons shield the outer
electrons from the nuclear
charge/attraction.
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Shielding
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The electron on the outermost
energy level has to look through all
the other energy levels to see the
nucleus.
Second electron has same
shielding, if it is in the same
period
Shielding Increases down a Group,
and is Constant across a Period.
Shielding Across a Group is
Constant, but the EFFECTIVE
NUCLEAR CHARGE INCREASES.
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Group Trends in Atomic Size
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As we go down a group...
each atom has another
energy level,
so the atoms get bigger.
Shielding increases as
well, so the nucleus has
less of a hold on e-…
distance is longer.
The Increased size of
the Energy Levels down a
group outweighs the
increased nuclear charge
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H
Li
Na
K
Rb
Periodic Trends in Atomic Size
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As you go across a period, the radius
gets smaller.
Electrons are in same energy level.
More nuclear charge.
Shielding is constant … not an issue.
Outermost electrons are closer.
Na
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Mg
Al
Si
P
S Cl Ar
Atomic Radius
Overall
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Atomic Radius (nm)
Overall Periodic
Trend for
Atomic Radii
Rb
K
Na
Xe
Li
Kr
Ar
Ne
H
He
10
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Atomic Number
Trends in Ionization Energy
The amount of energy required to
completely remove an electron from
a gaseous atom.
 Removing one electron makes a 1+
ion.
 The energy required to remove the
first electron is called the first
ionization energy.

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Ionization Energy
The second IE is the energy
required to remove the second
electron.
 Always greater than first IE.
 The third IE is the energy required
to remove a third electron.
 Greater than 1st or 2nd IE.

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Ionization Energy Table
Where are the Group Effects?
Symbol
H
He
1A Li
2A Be
3A B
C
N
O
F
Ne
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First
Second
Third
1312
2371
520
900
800
1086
1402
1314
1681
2080
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
Can we See the Effect of
Nuclear Charge in this Table?
Symbol
First
Second
Third
H
He
Li
Be
B
C
N
O
F
Ne
1312
2371
520
900
800
1086
1402
1314
1681
2080
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
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What Affects the IE

The greater the nuclear charge, the
greater the IE.
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
Larger positive nucleus has a greater
attraction for the electrons, so the IE
increases.
Greater distance from nucleus
decreases IE

Electrons are further away from the
attractive nucleus, and are easier to
remove.
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What Affects the IE


Filled and half-filled orbitals have lower
energy, so the removal of an electron to
achieve this ½ filled orbital requires
unusually low IE.
Shielding effect

As Shielding increases, it is easier to “pluck” the
outer electron, so the IE would decrease.
Let’s look at Group & Period Trends for IE
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H
Li
Group trends on IE
1
3
Na
11
K
19
Rb
37
Cs
55
As you go down a group,
first IE decreases
because...
 The electron is further
away.
 More shielding.

Fr
87
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Periodic trends on IE
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Across the representative elements, the
atoms are in the same period & have the same
energy level.
Same shielding.
But, increasing nuclear charge holds e-’s
tighter.
So IE generally increases from left to right.
Exceptions at full and 1/2 full orbitals.
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Summarizing 1st Ionization
Energy
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First Ionization energy
He
First Ionization Energy
He has a greater
IE than H.
 same shielding
 greater nuclear
charge

H
Atomic number
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First Ionization Energy
First Ionization energy
He
 Li
H
Li
has lower IE
than H
 more shielding
 further away
 these outweigh
greater nuclear
charge
Atomic number
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First Ionization Energy
First Ionization energy
He
 Be
H
Be
has higher IE
than Li
 same shielding
 greater nuclear
charge
Li
Atomic number
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First Ionization Energy
First Ionization energy
He
H
B has lower IE than
Be
 same shielding
 greater nuclear
charge
 By removing an
electron we make
the s orbital halffilled

Be
B
Li
Atomic number
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First Ionization Energy
First Ionization energy
He
H
C
Be
B
Li
Atomic number
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First Ionization Energy
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
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First Ionization Energy
First Ionization energy
He

N
H
C O
Be
B
Li
Breaks the
pattern,
because
removing an
electron leaves
1/2 filled p
orbital
Atomic number
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First Ionization Energy
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
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First Ionization Energy
Ne
First Ionization energy
He
N F
H
C O
Be
B
Li
Ne has a lower
IE than He
 Both are full,
 Ne has more
shielding
 Greater
distance

Atomic number
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First Ionization Energy
Ne
First Ionization energy
He
N F
H
C O
Be
B
Li
 Na
has a lower
IE than Li
 Both are s1
 Na has more
shielding
 Greater
distance
Na
Atomic number
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He
Ne
First Ionization Energy
First Ionization energy
Ar
Kr
Li
Na
K
Atomic number
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What’s the Driving Force?
Full Energy Levels require lots of
energy to remove their electrons.
 Noble Gases have full orbitals.
 Atoms behave in ways to achieve
noble gas configuration.

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2nd Ionization Energy
For elements that reach a filled or
half-filled orbital by removing 2
electrons, 2nd IE is lower than
expected.
 True for s2
 Alkaline earth metals form 2+ ions.

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3rd Ionization Energy
Using the same logic s2p1 atoms
have a low 3rd IE.
 Atoms in the aluminum family form
3+ ions.
 2nd IE and 3rd IE are always
higher than 1st IE!!!

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Trends in Electron Affinity
What is Electron Affinity?
It’s the energy change
associated with adding an
electron to a gaseous atom.
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Trends in Electron Affinity
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It’s easiest to add an electron to
Group 7A.
It gets them to a full energy level, or
completes the OCTET.
Increase from left to right: atoms
become smaller, with greater nuclear
charge.
Decrease as we go down a group.
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Electron Affinity in the Periodic
Table
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Electron Affinity in 3-D
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Trends in Ionic Size
Cations form by losing electrons.
 Cations are smaller that the atom
they come from.
 Metals form cations.
 Cations of representative elements
have noble gas configuration.

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Trends in Ionic Size
Anions form by gaining electrons.
 Anions are bigger that the atom
they come from.
 Nonmetals form anions.
 Anions of representative elements
have noble gas configuration.

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Revisiting Configuration of Ions
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Ions always have noble gas
configuration.
Na is: 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6
Same configuration as Neon.
Metals form ions with the
configuration of the noble gas before
them - they lose electrons.
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Revisiting Configuration of Ions
Non-metals form ions by gaining
electrons to achieve noble gas
configuration.
 They end up with the configuration
of the noble gas after them.

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Group Trends
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Going down a Group,
you are adding energy
levels
Ions get bigger as you
go down.
Li1+
Na1+
K1+
Rb1+
Cs1+
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Periodic Trends
Across the period, nuclear charge
increases so they get smaller.
 Energy level changes between
anions and cations.

B3+
Li1+
Be2+
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C4+
N3-
O2-
F1-
Size of Isoelectronic ions
Iso- means the same
 Iso electronic ions have the same #
of electrons
 Al3+ Mg2+ Na1+ Ne F1- O2- and N3- all
have 10 electrons
 all have the configuration:
1s22s22p6

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Size of Isoelectronic ions
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Positive ions that have more protons
would be smaller.
Increase in size from most positive to
most negative
Al3+
Na1+
Mg2+
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Ne
F1-
2O
N3-
Electronegativity
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The tendency for an atom to attract
electrons to itself when it is chemically
combined with another element.
How fair is the sharing?
Big electronegativity means it pulls the
electron towards it.
Atoms with large negative electron
affinity have larger electronegativity.
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Electronegativity Group Trend
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The further down a group, the farther
the electron is away, and the more
electrons an atom has.
Pull/Attraction of the positive nucleus is
lessened due to increased distance and
Shielding.
Electronegativity decreases.
More willing to share.
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Electronegativity Period Trend
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As you move across a Period, there are
the same number of energy levels, the
same shielding, however …
Pull/Attraction of the positive nucleus on
other’s electrons increases as the
nucleus gets larger
Electronegativity Increases.
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Electronegativity Periodic Trend
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Metals are at the left of the table.
They let their electrons go easily
Low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away from others
High electronegativity.
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Electronegativity in 3-D
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Can We Possibly Summarize all of
this Stuff???
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Can We Possibly Summarize all of
this Stuff???
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Can We Possibly Summarize all of
this Stuff???
Shielding Increases
Shielding is constant
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