Transcript periodic_trends

The Periodic Table
The how and why
Atomic Size
 First
problem where do you start
measuring.
 The electron cloud doesn’t have a
definite edge.
 Chemists 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 a diatomic molecule (homo-nuclear
molecule).
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Trends in Atomic Size
Influenced
by two factors.
Energy Level
Higher energy level is further away.
The effective charge from the
nucleus
The greater the nuclear charge
reaching the valence electrons the
closer these electrons are pulled in.
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Group trends
 As
we go down a group
 Each atom has another
energy level, so the
atoms get bigger.
 There are more levels in
the kernel and therefore
greater shielding of
valence electrons
(weaker attraction).
H
Li
Na
K
Rb
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Periodic Trends






As you go across a period the radius gets smaller.
Valence electrons are in the same energy level.
Greater nuclear charge gets to the valence electrons –
same shielding e- but greater nuclear charge.
Outermost electrons are closer.
Row 3 elements all have 10 shielding electrons;
+11
(nuclear charges)
+18
Na
Mg
Al
Si
P
S Cl Ar
6
Rb
K
Atomic Radius (nm)
Overall
Na
Li
Kr
Ar
Ne
H
10
Atomic Number
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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 is called the first
ionization energy.
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Ionization Energy
 The
second ionization energy 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 of 2nd IE.
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Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11749
14840
3569
4619
4577
5301
6045
6276
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Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
5247
7297
1757
2430
2352
2857
3391
3375
3963
Third
11749
14840
3569
4619
4577
5301
6045
6276
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What determines IE
 The
greater the effective nuclear
charge the greater IE.
 Less distance from nucleus
increases IE
 Filled and half filled sublevels have
lower energy, so removing them
raises the IE.
 Shielding: effective blocking of
nuclear charge weakens the
attraction of valence electrons.
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Shielding
 The
electron on the
outside energy level
has to look through
all the other energy
levels to see the
nucleus
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Shielding
 The
electron on the
outside energy level
has to look through
all the other energy
levels to see the
nucleus.
 A second electron
has the same
shielding.
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Group trends
As
you go down a group first IE
decreases because
The electron is further away.
More effective shielding.
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Periodic trends
 All
the atoms in the same period have
the same energy level.
 Same shielding.
 Increasing nuclear charge
 So IE generally increases from left to
right.
 Exceptions at full and 1/2 fill sublevels.
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First Ionization energy
He
 He
has a greater IE than H.
 no shielding
 greater nuclear charge
H
Atomic number
17
First Ionization energy
He  Li has lower IE than H
more shielding
 further away (outweighs greater
nuclear charge)

H
Li
Atomic number
18
First Ionization energy
He
Be has higher IE than Li
 same shielding
 greater nuclear charge
 removing an electron
from a full s sublevel

H
Be
Li
Atomic number
19
First Ionization energy
He
B has lower IE than Be
 same shielding
 greater nuclear charge
 removing an electron from a
Be
partially filled p sublevel

H
B
Li
Atomic number
20
First Ionization energy
He
H
C
Be
B
Li
Atomic number
21
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
22
First Ionization energy
He
N
H
C O
Be
 Breaks
B
Li
the pattern
because in N the electron
gets removed from a ½
filled p sublevel
Atomic number
23
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
24
First Ionization energy
He
Ne
 Ne
N F
H
C O
Be
B
has a lower IE
than He
 Both have full
valence shells
 Ne has more
shielding
 Greater distance
Li
Atomic number
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Ne
First Ionization energy
He

N F
Na has a lower
IE than Li
Both are s1
 Na has more
shielding
 Greater distance

H
C O
Be
B
Li
Na
Atomic number
26
Atomic number
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First Ionization energy
Driving Force
 Full
Energy Levels are very low
energy.
 Noble Gases have full energy levels.
 Atoms behave in ways to achieve
noble gas configuration (become
isoelectronic).
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nd
2
Ionization Energy
 For
elements that reach a filled or
half filled sublevel by removing 2
electrons 2nd IE is lower than
expected.
 True for s2
 Alkaline earth metals form +2 ions.
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3rd IE
the same logic s2p1 atoms have
an low 3rd IE.
 Atoms in the aluminum family form +3
ions.
 2nd IE and 3rd IE are always higher than
1st IE!!!
 Using
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Electron Affinity
The energy change associated with
adding an electron to a gaseous atom.
 Easiest to add to group 17 (7A).
 Gets them to full energy level becomes
stable – releases a large amount of
energy.
 Increase from left to right atoms –metals
are losers so to gain an electron will only
bring about a small increase in stability (if
any) – small energy change.
 Decrease as we go down a group.

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Ionic Size
 Cations
form by losing electrons.
 Cations are smaller than the atom
they come from.
 Metals form cations.
 Cations of representative elements
have noble gas configuration.
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Ionic size
 Anions
form by gaining electrons.
 Anions are bigger than the atom they
come from.
 Nonmetals form anions.
 Anions of representative elements
have noble gas configuration.
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Configuration of Ions
 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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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
 Adding
energy level
 Ions get bigger as
you go down.
Li+1
Na+1
K+1
Rb+1
Cs+1
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Periodic Trends
 Across
the period nuclear charge
increases so they get smaller.
 Energy level changes between
anions and cations.
Li+1
B+3
Be+2
N-3
O-2
F-1
C+4
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Radii vs atomic #
atomic radii
ionic radii
51
33
0.21
19
52
37
15
34 35
38
16 17
20
Radii (nonometers)
7
0.16
11
49
37
8
9
12
13
3
31 32 33
14
0.11
15
11
4
5
16
17
34 35
38
20
49
6
7
0.06
50 51 52
19
8
3
9
50
12
31
32
13
14
4
5
6
0.01
3
4
5
6
7
8
9 11 12 13 14 15 16 17 19 20 31 32 33 34 35 37 38 49 50 51 52
Atomic #
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Size of Isoelectronic ions
 Iso
- same
 Iso electronic ions have the same #
of electrons (same configuration)
 Al+3
Mg+2 Na+1 (Ne) F-1 O-2 and N-3
 all have 10 electrons
 all have the configuration 1s12s22p6
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Size of Isoelectronic ions
 Positive
ions have more protons than
electrons so they are smaller (strong
attraction for electrons).
Al+3
Na+1
Ne
F-1
O-2
N-3
Mg+2
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Electronegativity
Electronegativity
 The
tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
 How fair it shares.
 Big electronegativity means it pulls
the electron toward it.
 Atoms with large negative electron
affinity have larger electronegativity.
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Group Trend
 The
further down a group the farther
the electron is away and the more
electrons an atom has – the lower the
electronegativity value.
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Periodic Trend
 Metals
are at the left end.
 They let their electrons go easily
 Low electronegativity
 At the right end are the nonmetals.
 They want more electrons.
 Try to take them away.
 High electronegativity.
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Ionization energy, electronegativity
Electron affinity INCREASE
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Atomic size decreases,
shielding constant
Atomic
size
increases
Shielding
increases
Ionic size decreases
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