Download lecture 10

What’s coming up???
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Oct 25
Oct 27
Oct 29
Nov 1
Nov 3,5
Nov 8,10
Nov 12
The atmosphere, part 1
Midterm … No lecture
The atmosphere, part 2
Light, blackbodies, Bohr
Postulates of QM, p-in-a-box
Hydrogen and multi – e atoms
Multi-electron atoms
Ch. 8
Ch. 8
Ch. 9
Ch. 9
Ch. 9
Ch.9,10
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Nov 15
Periodic properties
Ch. 10
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Nov 17
Nov 19,22
Nov 24
Nov 26
Nov 29
Dec 1
Dec 2
Periodic properties
Valence-bond; Lewis structures
Hybrid orbitals; VSEPR
VSEPR
MO theory
MO theory
Review for exam
Ch. 10
Ch. 11
Ch. 11, 12
Ch. 12
Ch. 12
Ch. 12
SCREENING AND PENETRATION
PENETRATION
Is to get close to the nucleus
SCREENING
Is to block the view of other
electrons of the nucleus
LOOK AT RADIAL PROBABILITY DISTRIBUTIONS
In a multi-electron atom
Zeff(s) > Zeff(p) > Zeff(d)
ENERGY LEVELS REVERSE!!
E4s < E3d
En = − A ×
Z
2
eff
2
n
0.25
S-orbital penetrates closer to nucleus
0.2
0.15
0.1
0.05
0
0
5
10
15
20
25
30
Closer proximity to nucleus
Æ Higher effective nuclear
charge, Zeff
En = − A ×
Z
2
eff
2
n
The result is that 4s lies below 3d
5s
4d
4p
3d
4s
3p
E
3s
2p
2s
1s
MAGNETIC PROPERTIES
Atoms, molecules or ions with all electrons
paired are diamagnetic.
Diamagnetic materials are weakly repelled by
magnetic fields
Example:
NEON
1s
Ne
2s
1s22s22p6
ALL ELECTRONS PAIRED
2p
Spin quantum number
Electrons are influenced by a magnetic field as though they were spinning charges.
They are not really, but we think of them as having “spin up” or “spin down” levels.
MAGNETIC PROPERTIES
Atoms, molecules or ions with at least one
unpaired electron are paramagnetic.
Paramagnetic materials are strongly attracted
to magnetic fields.
Nitrogen atom is PARAMAGNETIC
1s
N
1s22s22p3
2s
2p
• Which of these is paramagnetic?
– Na
Ca
– K+
O2−
PERIODIC TRENDS
We will look at some trends in experimentally determined
properties such as:
ATOMIC RADII
IONIC RADII
IONIZATION ENERGIES
ELECTRON AFFINITIES
TRENDS IN ATOMIC RADII
Atomic radii increase
down the group
ATOMIC RADIUS
ATOMIC RADIUS
FOR GROUPS 1,2 and 13 to 18
Orbitals bigger as n increases
Radial probability density extends further.
TRENDS IN ATOMIC RADII
Atomic radii decrease
across a period
ATOMIC RADIUS
ATOMIC RADIUS
ATOMIC RADIUS
FOR GROUPS 1,2 and 13 to 18
SCREENING AND PENETRATION
ATOMIC RADIUS DECREASES ACROSS PERIOD!
WHY?
AS WE ADD ELECTRONS
The electrons go into same shell and do not
screen the nucleus from each other
As nuclear charge goes up, so does Zeff
Therefore radial probability pulled in….
En = − A ×
Z
2
eff
2
n
Zeff increases .. no screening
.. radius decreases
0.12
0.1
0.08
0.06
0.04
0.02
0
0
5
10
15
20
25
30
Zeff INCREASES
1
1 1s
18
13 14 15 16 17 1s
2
RADIUS DECREASES
2
2s
3
3s
4
4s
3d
4p
5
5s
4d
5p
6
6s
La
5d
6p
7
7s
Ac
6d
2p
3p
DOWN GROUP...
4f
5f
2
2
2s
3
3s
4
4s
5
5s
6
6s
La
7
7s
Ac
18
Zeff DECREASES
1 1s
RADIUS INCREASES
1
13 14 15 16 17 1s
2p
3p
3d
4p
4d
5p
5d
6p
6d
4f
5f
TRENDS IN IONIC RADII
IONIC RADIUS
Same reasons as for atomic radii…….
Cations and anions…...
IONIC RADIUS
IONIC RADIUS
FOR GROUPS 1,2 and 13 to 18
Cations are smaller than their corresponding
neutral atoms.
Na is 186 pm and Na+ is 95 pm
One less electron electrons pulled in by nuclear charge
Anions are larger than their corresponding neutral
atoms.
F is 64 pm and F- is 133 pm
same nuclear charge and repulsion among
electrons increases radius
O < O– < O2–
QUESTIONS…...
EXAMPLES
Which is bigger?
Na or Rb
Rb …. higher n, bigger orbitals
K or Ca
K …. poorer screening for Ca
Ca or Ca2+
Br or Br-
What about
Ca …. bigger than cation
Br …. smaller than anion
ISOELECTRONIC SPECIES?
QUESTION
The species F-, Na+,Mg2+ have relative sizes in
which order?
22s22p6
ALL
1s
+
2+
1
F < Na <Mg
ALL are isoelectronic
2
F-> Na+>Mg2+
3
Na+>Mg2+> F-
4
Na+=Mg2+= F-
5
Mg2+> Na+>F-
QUESTION
The species F-, Na+,Mg2+ have relative sizes in
which order ?
1
F-< Na+<Mg2+
2
F-> Na+>Mg2+
Check numbers
3
Na+>Mg2+> F-
4
Na+=Mg2+= F-
5
Mg2+> Na+>F-
Na+ is 95 pm
Mg2+ is 66 pm
F- is 133 pm
ELECTRON AFFINITY
the energy change associated with the addition
of an electron to a gaseous atom.
X + e– → X–
F(g) + e– → F–(g)
∆E= -328 kJmol-1
TWO DEFINITIONS!
EA = ∆E the electron affinity is negative if ∆E<0.
EA = -∆E the electron affinity is positive if ∆E<0.
In general no clear cut trends…….
MORE NEGATIVE
ELECTRON AFFINITY
ELECTRON AFFINITY
MORE NEGATIVE
MORE NEGATIVE
ELECTRON AFFINITY
TRENDS IN EA
ELECTRON AFFINITIES
400
Cl
F
-ELECTRON AFFINITY
300
S
200
C
H
100
Si
O
Li
Na
B
0
-100
0
1
2
He
3
4
Be
Al
5
6
7
8
9
10
11
N
12
13
15
16
17
18
Mg
Ne
-200
ATOMIC NUMBER
What is special about … He, Ne and Ar?
Be, N and Mg?
14
P
Ar
19
• He, Ne, Ar: rare gases … adding electron
to filled shell … must increase n
• Be, Mg: alkali earths … adding electron to
filled sub-shell … must increase l
• N: Hund’s rule stability… adding electron
to ½-filled degenerate p-shell
• Q: what would you predict for… Cr, Fe?
TRENDS IN IONIZATION
ENERGIES
M(g) → M+(g) + eThe ionization energy of gaseous atoms
of the elements have been measured….
And we find…….
First Ionization
energies decrease
down the group
IONIZATION ENERGY
Zeff DECREASES
IONIZATION ENERGY
TRENDS IN FIRST IE
Zeff INCREASES UP THE GROUP
Electrons closer to nucleus more tightly held
TRENDS IN FIRST IE
Zeff INCREASES
IONIZATION ENERGY
Greater effective nuclear charge across period
Poor shielding by electrons added
n
IONIZATION ENERGY
I . E. = A ×
Z
2
eff
2
TRENDS IN FIRST IE
IONIZATION ENERGY(kJ/mol)
2500
n=1
n=2
2000
n=3
1500
n=4
1000
500
0
0
11
2
2
155
166
133 144
GROUP NUMBER
177
8
18
9
Closed shells most stable
The noble gases have the highest ionization energy!