Download WAVE NATURE OF ELECTRONS IN ATOMS

WAVE NATURE OF
ELECTRONS IN ATOMS
Electrons have wave properties
Can not specify exact location of a wave,
or of an electron
Electrons occupy orbitals in atoms
size, shape, names
Are quantized
QUANTUM NUMBERS
(1) Principal quantum number (n)
values: 1, 2, 3, …
average distance from nucleus: size
energy of electrons in orbital
(2) Azimuthal quantum number
(A)
values: 0, 1, 2, …, (n–1)
shape of electron orbital
(3) Magnetic quantum number
values: −A, …, −1, 0, 1, …, A
orientation of orbital in space
(4) Spin quantum number
two values: +½ and –½
(ms)
(mA)
Symbols used for A quantum numbers
A value
0
1
2
3
subshell name
s
p
d
f
no. of
electrons
n=1
n=2
n=2
n=3
A=0
A=0
A=1
A=0
1s
2s
2p
3s
(mA = 0)
(mA = 0)
(mA = –1,0,1)
(mA = 0)
Table 6.2
n2 = number of states
= number of orbitals in shell n
2
2
6
2
s orbitals
First s orbital: 1s
n=1
A=0
mA = 0
Second s orbital: 2s
n=2
A=0
mA = 0
SHAPES OF ORBITALS
s orbitals
electron density
or probability
Ψ2
(1s)
0
r
at nucleus, r = 0
2s is larger than 1s
Size of s orbital increases as n increases
Shape: spherical symmetry
p and d orbitals
First p orbitals: 2p
n=2
same size & shape
A=1
mA = –1, 0, 1 }
3 different orientations
First d orbitals: 3d
n=3
A=2
mA = –2, –1, 0, 1, 2
5 orbitals
5 orientations
SHAPES OF ORBITALS
p orbitals
2 lobes with node between
2p
n=2
A=1
mA = –1, 0, 1
3 orbitals
Because n same
& A same, they have
same size & shape
They differ in
orientation
p orbitals are directional
SHAPES OF ORBITALS
d orbitals
3d
n=3
A=2
mA = –2, –1, 0, 1, 2
5 orbitals with different orientations
SHAPES OF ORBITALS
sphere
1s
dumbbell
2pz
3d yz
2px
2py
3d xy
3d xz
clover-leaf
& friend
3d x 2 -y2
3d z2
REVIEW
ORBITALS – region of space with size,
shape, characteristic energy
Name
s
p
d
f
Number Shape
1
spherical
3
dumbbell
5
5 shapes
7
------
QUANTUM NUMBERS
n
A
mA
principal
size
azimuthal shape
magnetic orientation
FOURTH QUANTUM NUMBER
Electron has magnetic moment,
as if were spinning
Experimental observations confirm
ms = ½ or –½
Electrons have 4 quantum numbers
n A mA ms
defines orbital
Since ms has only 2 values
Therefore, max of 2 electrons per orbital
Pauli Exclusion Principle
STERN-GERLACH EXP.
Silver (Ag) atoms have one unpaired
electron.
A beam of Ag atoms splits according to
sign of electron spin in magnetic field
PAULI EXCLUSION PRINCIPLE
No two electrons in an atom can have
the same four quantum numbers
( n A mA ms )
Electrons in the same orbital have the
same values for the first three quantum
numbers ( n A mA )
ms can have only two values: ½ or –½
Therefore, an orbital can hold only two
electrons, and they must have opposite
spins
Subshell
s (A = 0)
p (A = 1)
d (A = 2)
f (A = 3)
No. of Orbitals Max. No. of e–
1
2
3
6
5
10
7
14
ENERGIES OF ORBITALS
1
One electron cases: E ∝
n2
not
dependent
on A or mA
Two or more electrons:
E does depend on n and A (but not mA)
Therefore: E2s ≠ E2p
E3s ≠ E3p ≠ E3d
Value of n determines shell
Same n and A means same subshell
…and….
same subshell means same energy
ORBITAL FILLING
SEQUENCE
1
1s
2s
3s
4s
5s
6s
3s
2p
3p
4p
5p
6p
2
3
4
5
3d
4d 4f
5d 5f
6d
6
3p
3d
increasing energy
Orbital size AND shape effect energy
FORCES ACTING ON
ORBITAL ELECTRONS
Electrons in outer orbitals see nucleus
and also the inner electrons
Shielding, Screening
s and p orbitals have different shapes
Therefore, they experience shielding
differently
s orbital has density at the nucleus
p orbital does not
s electrons see more of nuclear charge, Z
have lower energy, are more stable
p electrons see less of Z
have higher energy, are less stable
ELECTRON CONFIGURATIONS
Orbitals are filled by electrons in
sequence determined by energy
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
1s
1s2 ← filled shell
1s22s ← new row of PT
1s22s2
1s22s22p
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6 ← filled shell
1s22s22p63s ← new row of PT
1s22s22p63s2
HUND’S RULE
When electrons are filling orbitals of
equal energy, they go singly into orbitals
before starting to double up. Electrons in
partially-filled orbitals have the same spin.
Carbon
1s2
2s2
2pz
6 electrons
2p2
2px
2py
e- repel each other, go into different orbitals
Each orbital unique region of space.
ELECTRON CONFIGURATIONS
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
1s22s22p6 ← filled shell
(Ne)3s ← new row of PT
(Ne)3s2
(Ne)3s23p
(Ne)3s23p2
(Ne)3s23p3
(Ne)3s23p4
(Ne)3s23p5
(Ne)3s23p6 ← subshell filled
(Ar)4s ← new row of PT
(Ar)4s2
(Ar)4s23d
(Ar)4s23d2
transition
(Ar)4s23d3
metals
(Ar)4s13d5
(Ar)4s23d5
(Ar)4s23d6
half or full d orbital
2
7
(Ar)4s 3d
more stable than
2
8
(Ar)4s 3d
filled s orbital
1
10
(Ar)4s 3d
(Ar)4s23d10
PERIODIC TABLE
APPRECIATION
What first comes to mind when
looking at the Periodic Table
of the Elements?
Marvel at the order and symmetry
of the building blocks of our world.
We are going to use what we have
learned about atomic structure to
understand that order.
ELECTRON CONFIGURATIONS
AND THE PERIODIC TABLE
Electron configurations relate to the
elements’ location in the periodic table
Example:
Group 1A
Li [He]2s
Na [Ne]3s
K [Ar]4s
Rb [Kr]5s
Cs [Xe]6s
ELECTRON CONFIGURATIONS
AND THE PERIODIC TABLE
Alkali metals have ns1 outer shell
Li, Na, K, Rb, ...
Halogens have np5 outer shell
F, Cl, Br, I, ...
Noble gases have filled outer shell np6
Ne, Ar, Kr, ...
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Valence electrons determine chemistry
Valence electrons are those outside
the noble-gas core
ION ELECTRON
CONFIGURATIONS
Metal atoms lose electrons to form
cations with charge equal to the
group number
Example: Mg (2A) Æ Mg2+: [Ne]
Nonmetals gain electrons to form
anions with charge equal to the
group number minus 8.
Example: O (6A) Æ O2-: [Ne]
Transition metals lose s electrons
before d electrons
EXAMPLE
Fe
[Ar]3d64s2
Fe2+
[Ar]3d6
Fe2+ + 2e–
[Ar]3d6
Fe3+ + e–
[Ar]3d5
ISOELECTRONIC
SERIES
Isoelectronic: same no. of electrons
EXAMPLES
O–2 F– Ne Na+ Mg2+ Al3+
10 electrons each: 1s22s22p6 = [Ne]
S2– Cl– Ar K+ Ca2+
18 electrons each: [Ne]3s23p6 = [Ar]
Chapter 6 review
6.25
Chapter 6 review
6.34
Chapter 6 review
6.60
TRENDS IN ATOMIC PROPERTIES:
THE PERIODIC TABLE
9 Electron configurations determine
organization of the periodic table
9 Next… properties of elements and
their periodic behavior
9 Elemental properties determined by:
– size (n) and shape (l) of orbitals
– atomic number (nuclear charge)
Atomic sizes
Ionization energies
Electron affinities
ATOMIC SIZE
Size of atom increases going down a group
Why?
As we go down a group, n increases.
As n increases, orbital radius increases.
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Size of atom decreases going from left
to right along a period
Why?
Going across increases no. of protons
and the nuclear charge
Added outer electrons shielded ineffectively
Effective nuclear charge increases, so the
electrons are drawn closer
ORBITAL SIZE INCREASES
WITH n
2s
1s
1s
2s
3s
3s
ORBITAL SIZE DECREASES
WITH INCREASING Zeff
Zeff = Z - S
S = core electron screening charge
S similar for elements in same period
Zeff increases with Z in same period
Na: [He]3s1
3s eNe core (10 e-)
2p
Z = 11+
2s
lower shells
smaller distance
screening
same shell
similar distance
little screening
PERIODIC TREND
IN ATOMIC RADII
Fig. 7.6
IONIC SIZE
Periodic trends same as for atoms
Cation smaller than related atom
Na+
Na
97 pm
154 pm
why?
Na: [Ne]3s1
Anion larger than related atom
Cl–
Cl
181 pm
99 pm
Cl-: [Ne]2s22p6
why?
ATOMIC SIZE AND
ISOELECTRONIC
SERIES
Isoelectronic: same no. of electrons
EXAMPLES
O–2 F– Ne Na+ Mg2+ Al3+
10 electrons each: 1s22s22p6 = [Ne]
nuclear charge increases →
size decreases →
Ca2+ K+ Ar Cl– S2–
what trends in nuclear charge and
atomic or ionic size?