Transcript Electron Configurations Chemical Periodicity (Kotz, Ch 8)

Electron Configurations  Chemical Periodicity (Ch 8)
• Electron spin & Pauli exclusion principle
• configurations
• spectroscopic, orbital box notation
• Hund’s rule - electron filling rules
• configurations of ATOMS:
• the basis for chemical valence
• configurations and properties of IONS
• periodic trends in :
• size
• ionization energies
• electron affinities
Na + Cl  NaCl
Mg + O2  MgO
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Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS ()
ORBITALS (m)
Each orbital can be assigned
up to 2 electrons!
WHY ?
. . . Because there is a 4th quantum number,
the electron spin quantum number, ms.
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Electron Spin
Quantum Number, ms
• It can be proved experimentally that the
electron has a spin. This is QUANTIZED.
• The two allowed spin directions are defined by
the magnetic spin quantum number, ms
ms = +1/2 and -1/2 ONLY.
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Electron Spin Quantum Number
MAGNETISM is a macroscopic
result of quantized electron spin
5_magnet.mov
Diamagnetic: NOT attracted to a magnetic field
All electrons are paired
N2
Paramagnetic: attracted to a magnetic field.
Substance has unpaired electrons
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Chemical Periodicity
O2
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Pauli Exclusion Principle
• electrons with the same spin keep as far apart as possible
• electrons of opposite spin may occupy the same
“region of space” (= orbital)
• Consequences:
• No orbital can have more than 2
electrons
• No two electrons in the same atom can
have the same set of 4 quantum
numbers (n, l, ml, ms)
OR
• “Each electron has a unique address.”
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QUANTUM
NUMBERS
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n (shell)
1, 2, 3, 4, ...
 (subshell)
0, 1, 2, ... n - 1
m (orbital)
-  ... 0 ... + 
ms (electron spin)
+1/2, -1/2
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Shells, Subshells, Orbitals

#orbitals
0 s
1
0 s
1
1 p
3
0 s
1
1 p
3
2 d
5
0 s
1
1 p
3
2 d
5
3 f
7
0..(n-1) (2 +1)
n
1
2
3
4
n
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#e2
2
6
2
6
10
2
6
10
14
2*(2 +1)
Chemical Periodicity
Total
2
8
PERIOD
1 (H, He)
2 (Li…Ne)
3 (Na .. Ar)
18
32
2n2
=0
=1
=2
=3
s
p
d
f
etc, for n = 5, 6
7
Element Mnemonic Competition
Hey! Here Lies Ben Brown. Could Not Order Fire. Near
Nancy Margaret Alice Sits Peggy Sucking Clorets. Are
Kids Capable ?
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Chemical Periodicity
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Assigning Electrons to Atoms
• Electrons are assigned to orbitals successively in order
of the energy.
• For H atoms, E = - R(1/n2). E depends only on n.
• For many-electron atoms, orbital energy depends on
both n and .
• E(ns) < E(np) < E(nd) ...
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Assigning Electrons to Subshells
• In H atom all subshells of
same n have same energy.
• In many-electron atom:
a) subshells increase in energy
as value of (n + )
increases.
b) for subshells of same (n +),
subshell with lower n is
lower in energy.
(n + )=
5
(n + )= 4
5_manyelE.mov
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Effective Nuclear Charge
• The difference in SUBSHELL energy
e.g. 2s and 2p subshells
is due to effective nuclear charge, Z*.
Charge felt by 2s
e- of Li atom
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2s e- spends
more time
close to Li3+
nucleus than
the 2p eTherefore
2s is
lower in E
than 3s
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Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by an electron.
• Z* increases across a period owing to incomplete
shielding by inner electrons.
• For VALENCE electrons we estimate Z* as:
Z* = [ Z - (no. of inner electrons) ]
• Charge felt by 2s e- in Li
Be
B
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Z* = 3 - 2 = 1
Z* = 4 - 2 = 2
Z* = 5 - 2 = 3
and so on!
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Photoelectron Spectroscopy - Measuring IE
Photoelectric effect: h + A  A+ + eforms basis for DIRECT determination of IE
Kinetic energy of electron = h - IE
VALENCE
therefore: IE = h - KE(e )
ELECTRONS
Signal
1s
Ne
Inner shell or
CORE ELECTRONS
2p
Ar
1s
309
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2s
2s
50
100
Chemical Periodicity
2p
3s
3p
0
IE (MJ/mol)
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Electron Filling Order
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Chemical Periodicity
(Figure 8.7)
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Writing Atomic Electron Configurations
Two ways of writing configurations.
One is called the spectroscopic notation:
SPECTROSCOPIC NOTATION
for H, atomic number = 1
1
no. of
electrons
1s
value of n
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value of l
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Writing Atomic Electron Configurations (2)
A second way is called the orbital box notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
2
1s
1s
Arrows
depict
electron
spin
One electron has n = 1,  = 0, ml = 0, ms = + 1/2
Other electron has n = 1,  = 0, ml = 0, ms = - 1/2
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Electron Configuration tool - see “toolbox”.
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Beryllium
Lithium
Group 2A
Z=4
1s22s2
Group 1A
Z=3
1s22s1
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
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Boron
Carbon
Z=5
Z=6
1s2 2s2 2p1
1s2 2s2 2p2
3p
3p
3s
3s
2p
2p
2s
2s
Why not  ?
1s
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1s
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Carbon
Z=6
1s2 2s2 2p2
3p
3s
2p
2s
The configuration of C is an
example of HUND’S RULE:
the lowest energy
arrangement of electrons in
a subshell is that with the
MAXIMUM no. of unpaired
electrons
Electrons in a set of orbitals
having the same energy,
are placed singly as long as possible.
1s
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Nitrogen
Oxygen
Z=7
Z=8
1s2 2s2 2p3
1s2 2s2 2p4
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
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Fluorine
Neon
Z=9
Z = 10
1s2 2s2 2p5
1s2 2s2 2p6
3p
3s
3p
3s
2p
2p
2s
1s
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2s
1s
Chemical Periodicity
Note that we have reached
the end of the 2nd period,
. . . and the 2nd shell is full!
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GROUPS and
PERIODS
Sodium
Z = 11
1s2 2s2 2p6 3s1
or “neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Na begins a new period.
Li Na K Rb Cs
All Group 1A elements:
have [core] ns1 configurations. (n = period #)
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Periodic Chemical Properties
REACTIVITY
5_Li.mov
5_Na.mov
5_K.mov
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SIZE
IE (Ionization Energy)
Li
Be
Na
Mg
K
Ca
Rb
Sr
Cs
Ba
Alkalis
Alkaline Earths
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Alkaline Earths
Metals (ns2) - easily oxidized to M2+
- less reactive than alkalis of same period
reactivity: Be < Mg < Ca < Sr < Ba
WHY? - • Size INCREASES as  group
• VALENCE e- are farther from nucleus
• same Z* - Valence e- less tightly held
• Therefore valence e- are easier to remove
Typical reactions / compounds
Oxides: M +1/2O2 (g)  MO (s) CaO (lime) - #5 Ind. Chem
Halides: M + X2 (g)  MX
Carbonates: CaCO3 (limestone)  CaO + CO2
Sulfates:
CaSO4.2H2O (gypsum)  CaSO4. 0.5H2O (plaster-of-paris) + 3/2H2O
RECALL: Solubility rules and PRECIPITATION REACTIONS
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Relationship of Electron Configuration
and Regions of the Periodic Table
s block
d block
p block
f block
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Transition Metals Table 8.4
• Transition metals (e.g. Sc .. Zn in the 4th period)
have the configuration [argon] nsx (n - 1)dy
• also called “d-block” elements.
3d orbitals used for Sc - Zn
Chromium
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Iron
Chemical Periodicity
Copper
27
Ion Configurations
To form cations from elements : remove 1 e- (or more)
from subshell of highest n [or highest (n + )].
P [Ne] 3s2 3p3 - 3e-  P3+ [Ne] 3s2 3p0
3p
3p
3s
3s
2p
2p
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2s
2s
1s
1s
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Ion Configurations (2)
Transition metals ions:
remove ns electrons and then (n - 1)d electrons.
Fe [Ar] 4s2 3d6 loses 2 electrons  Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
4s
3d
E4s ~ E3d - exact energy
of orbitals depend on
whole configuration
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3d
Fe3+
4s
Chemical Periodicity
3d
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Ion Configurations (3)
How do we know the configurations of ions?
From the magnetic properties of ions.
Ions (or atoms) with UNPAIRED ELECTRONS are:
PARAMAGNETIC.
Ions (or atoms) without unpaired electrons are:
DIAMAGNETIC.
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General Periodic Trends
• Atomic and ionic radii : SIZE
• Ionization energy :
E(A+) - E(A)
• Electron affinity :
E(A-) - E(A)
Higher Z*.
Electrons held
more tightly.
Larger orbitals.
Electrons held less
tightly.
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Atomic Size INCREASES
down a Group
• Size goes UP on going down a GROUP
• Because electrons are added further from the
nucleus, there is less attraction.
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Atomic Size DECREASES across a period
Size goes DOWN on going across a PERIOD.
Size decreases due to increase in Z*.
Each added electron feels a greater and greater
+ve charge.
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Atomic Radii
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Trends in Atomic Size
(Figure 8.10)
Radius (pm)
250
K
1st transition
series
3rd period
200
Na
2nd period
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
30
35
40
Atomic Number
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Sizes of Transition Elements(Figure 8.11)
• 3d subshell is inside the 4s subshell.
• 4s electrons feel a more or less constant Z*.
• Sizes stay about the same and chemistries are similar!
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Ion Sizes - CATIONS
Does the size go up or down when an atom loses an
electron to form a cation?
+
Li, 152 pm
3 e-, 3 p
Forming
a cation
Li+, 60 pm
2 e-, 3 p
• CATIONS are SMALLER than the parent atoms.
• The electron/proton attraction goes UP so size DECREASES.
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Ion Sizes - ANIONS
Does the size go up or down when gaining an
electron to form an anion?
F, 64 pm
9 e-, 9 p
Forming
an anion
F-, 136 pm
10 e-, 9 p
• ANIONS are LARGER than the parent atoms.
• electron/proton attraction goes DOWN so size INCREASES.
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Trends in Ion Sizes
CATIONS
ANIONS
(59 pm)
(207 pm)
Trends in relative ion sizes are the same as atom sizes.
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Oxidation-Reduction Reactions
• Why do metals lose electrons in
their reactions?
• Why does Mg form Mg2+ ions
and not Mg3+?
• Why do nonmetals take on
electrons?
- related to IE and EA
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Ionization Energy (IE)
Mg (g) atom
[Ne]2s
735 kJ  Mg+ (g) + e-
[Ne]2s1
Mg+ (g) + 1451 kJ  Mg2+ (g) + e-
[Ne]2s0
Mg (g) +
Mg3+
Mg2+ (g) + 7733 kJ  Mg3+ (g) + e- [He]2s22p5
Mg2+
• Energy ‘cost’ is very high to remove an
INNER SHELL e- (shell of n < nVALENCE).
• This is why oxidation. no. = Group no.
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Chemical Periodicity
Mg+
Mg
41
Trends in First Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
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3
Li
5
7
9
11
Na
13
15
17
19
21
K
Chemical Periodicity
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25
27
29
31
33
35
Atomic Number
42
Trends in Ionization Energy (2)
•
•
•
•
IE increases across a period because Z* increases.
Metals lose electrons more easily than nonmetals.
Metals are good reducing agents.
Nonmetals lose electrons with difficulty.
• IE decreases down a group
• Because size increases, reducing
ability generally increases down the
periodic table.
• E.g. reactions of Li, Na, K
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2nd IE / 1st IE
2nd IE:
A+  A++ + e-
Li
Na
K
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Electron Affinity (EA)
• A few elements GAIN electrons to form anions.
• Electron affinity is the energy released when an
atom gains an electron.
A(g) + e-  A-(g)
E.A. = DE = E(A-) - E(A)
• If E(A-) < E(A) then the anion is more stable
and there is an exothermic reaction
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Trends in Electron Affinity (Table 8.5, Figure 8.14)
Atom EA (kJ)
• Affinity for electron increases B
C
across a period
(EA becomes more negative). N
O
-27
-122
0
-141
F
-328
F
Cl
Br
I
-328
-349
-325
-295
• Affinity decreases down a
group
(EA becomes less negative).
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SUMMARY
• Electron spin: diamagnetism vs. paramagnetism
• Pauli exclusion principle - allowable quantum numbers
• configurations
• spectroscopic notation
• orbital box notation
• Hund’s rule - electron filling rules
• configurations of ATOMS: the basis for chemical valence
• period 2 ; groups
• transition metals
• configurations and properties of IONS
• periodic trends in :
• size
• ionization energies
• electron affinities
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