Transcript No Slide Title

Chapter 5
Periodicity & Atomic Structure
1
The Periodic Table 01
• The periodic table is the most important
organizing principle in chemistry.
• Chemical and physical properties of elements in
the same group are similar.
• All chemical and physical properties vary in a
periodic manner, hence the name periodic table.
2
Development of the Periodic
Table
Mendeleev’s Periodic Table (1871)
Until the discovery of the proton, the elements were
typically organized by increasing atomic weight.
The modern organization is by increasing atomic number.
3
The Periodic Table 04
4
Ar = 39.95 amu, would appear on the right of K = 39.10 amu
Development of Modern Periodic
Table
• In 1913 Moseley discovered that when
elements were irradiated with high energy
radiation y emitted X-ray. He used the
frequency of the emitted radiation to
calculate the atomic number.
√ υ = a( Z-b)
The Periodic Table 03
6
What is light made of ?
1) Made of Waves? Waves interfere,
2) Made of particles?
3) Made of both?
4) What are electron’s made of?
8
Electromagnetic Radiation 02
9
Electromagnetic Radiation 01
• Frequency (, Greek nu): Number of peaks that
pass a given point per unit time.
• Wavelength (, Greek lambda): Distance from
one wave peak to the next.
• Amplitude: Height measured from the center of
the wave. The square of the amplitude gives
intensity.
10
Electromagnetic Radiation 03
11
Electromagnetic Radiation 04
• Speed of a wave is the wavelength (in meters)
multiplied by its frequency (in reciprocal seconds,
(s–1) ).
–
–
Wavelength x Frequency = Speed
 (m)
x  (s–1) = c (m/s)
12
Learning Check
• The red light in a laser pointer comes from a diode
laser that has a wavelength of about 630 nm. What
is the frequency of the light?
C = 2.9979 x 108 m.s–1
13
A photon has a frequency of 6.0 x 104 Hz (s-1). Convert
this frequency into wavelength (nm). Does this frequency
fall in the visible region?
c = 2.9979 x 108 m/s–1
x=c
 = c/
 = 2.9979 x 108 m/s / 6.0 x 104 Hz
 = 5.0 x 103 m
 = 5.0 x 1012 nm
Radio wave
14
Atomic Spectra 01
• Atomic spectra:
Result from excited
atoms emitting light.
• Line spectra: Result
from electron
transitions between
specific energy
levels.
15
Maxwell (1873), proposed that visible light consists of
electromagnetic waves.
Electromagnetic
radiation is the emission
and transmission of energy
in the form of
electromagnetic waves.
c = 2.9979 x 108 m/s–1
Speed of light (c) in vacuum = 3.00 x 108 m/s
All electromagnetic radiation
x=c
16
7.1
Line Emission Spectrum of Hydrogen Atoms
17
7.3
18
“Photoelectric Effect”
h
Depending on metal used only
lights of certain minimum frequency
could cause ejection of electron
KE e-
Light must be composed of particles
called photon
Planks and Einstein
Light has both:
1. wave nature
2. particle nature
19
Particle like Properties of
Electromagnetic Energy
20
Planck in 1900
Energy (light) is emitted or
absorbed in discrete units
(quantum).
E=hx
Planck’s constant (h)
h = 6.63 x 10-34 J•s
21
“Photoelectric Effect”
22
“Photoelectric Effect”
Solved by Einstein in 1905
Light has both:
1. wave nature
2. particle nature
h
KE e-
Photon is a “particle” of light
23
Bohr’s Model of
the Atom (1913)
What happens in a hydrogen lamp?
Line Emission Spectrum of Hydrogen Atoms
1. e- can only have specific
(quantized) energy
values
2. light is emitted as emoves from one energy
level to a lower energy
level
En = -RH (
1
n2
)
n (principal quantum number) = 1,2,3,…
RH (constant) = 2.18 x 10-18J
24
E = h
E = h
25
Hydrogen Atomic Spectra
Line Emission Spectrum of Hydrogen Atoms
26
ni = 3
ni = 3
ni = 2
nf = 2
Ephoton = DE = Ef - Ei
1
Ef = -RH ( 2
nf
1
Ei = -RH ( 2
ni
1
DE = RH( 2
ni
)
)
1
n2f
)
RH (constant) = 2.18 x 10-18J
nnf f==11
m = initial n = final
1
1
1
=R 2
E = h x C/λ
m
n2

27
R = 1.097 X 10-2 nm-1
Why is e- energy quantized?
De Broglie (1924) reasoned
that e- is both particle and
wave.
 = h/mv
v = velocity of em = mass of eh = 6.63 x 10-34 J•s
28
What is the de Broglie wavelength (in nm)
associated with a 2.5 g Ping-Pong ball
traveling at 15.6 m/s?
 = h/mv
h in J•s
m(mass) in kg
V( velocity) in (m/s), J = Kg.m2/S2)
= 6.63 x 10-34 J.S ((Kg.m2/S2)/J) / ((2.5 x 10-3 Kg)
x (15.6 m/s))
 = 1.7 x 10-32 m = 1.7 x 10-23 nm
29
Uncertainty Principle
W. Heisenberg
1901-1976
Problem of defining nature of
electrons in atoms solved
by W. Heisenberg.
Cannot simultaneously define
the position and momentum
(= m•v) of an electron.
We define e- energy exactly
but accept limitation that we
do not know exact position.
30
http://www.youtube.com/watch?v=DfPeprQ7oGc
Schrodinger Wave Equation
In 1926 Schrodinger wrote an equation that
described both the particle and wave nature of the eWave function (Y, psi) describes:
1. energy of e- with a given Y
2. probability of finding e- in a volume of space
Schrodinger’s equation can only be solved exactly
for the hydrogen atom. Must approximate its
solution for multi-electron systems.
31
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
Existence (and energy) of electron in atom is described
by its unique wave function Y.
Pauli exclusion principle - no two electrons in an atom
can have the same four quantum numbers.
Each seat is uniquely identified (E, R12, S8)
Each seat can hold only one individual at a
time
32
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
principal quantum number n
n = 1, 2, 3, 4, ….
distance of e- from the nucleus
n=1
n=2
n=3
33
Electron Radial Distribution 01
34
Electron Radial Distribution 02
• s Orbital Shapes:
35
1s Orbital
36
2s Orbital
37
3s Orbital
38
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
39
QUANTUM NUMBERS
Y = fn(n, l, ml, ms)
The shape, size, and energy of each orbital is a
function of 3 quantum numbers:
n (major)
---> shell
l (angular) ---> subshell
ml (magnetic) ---> designates an orbital
within a subshell
40
QUANTUM NUMBERS
Symbol
Values
Description
n (major)
1, 2, 3, .. Orbital size
and energy
l (angular)
0, 1, 2, .. n-1
Orbital shape and energy
(subshell)
ml (magnetic) -l..0..+l
Orbital orientation
# of orbitals in subshell = 2 l + 1
41
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
42
l = 0 (s orbitals)
l = 1 (p orbitals)
43
7.6
Electron Radial Distribution 03
• p Orbital Shapes:
44
l = 2 (d orbitals)
45
7.6
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
magnetic quantum number ml
for a given value of l
ml = -l, …., 0, …. +l
if l = 1 (p orbital), ml = -1, 0, or 1
if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2
orientation of the orbital in space
46
7.6
ml = -1
ml = -2
ml = 0
ml = -1
ml = 0
ml = 1
ml = 1
ml =
2
47
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
spin quantum number ms
ms = +½ or -½
ms = +½
ms = -½
48
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
Existence (and energy) of electron in atom is described
by its unique wave function Y.
Pauli exclusion principle - no two electrons in an atom
can have the same four quantum numbers.
Each seat is uniquely identified (E, R12, S8)
Each seat can hold only one individual at a
time
49
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
50
7.6
f-Orbitals
•
51
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
Shell – electrons with the same value of n
Subshell – electrons with the same values of n and l
Orbital – electrons with the same values of n, l, and ml
How many electrons can an orbital hold?
If n, l, and ml are fixed, then ms = ½ or - ½
Y = (n, l, ml, ½) or Y = (n, l, ml, -½)
An orbital can hold 2 electrons
52
How many 2p orbitals are there in an atom?
n=2
If l = 1, then ml = -1, 0, or +1
2p
3 orbitals
l=1
How many electrons can be placed in the 3d
subshell?
n=3
3d
l=2
If l = 2, then ml = -2, -1, 0, +1, or +2
5 orbitals which can hold a total of 10 e53
Effective Nuclear Charge
01
• Electron shielding leads
to energy differences
among orbitals within a
shell.
• Net nuclear charge felt
by an electron is called
the effective nuclear
charge (Zeff).
54
Effective Nuclear Charge
02
• Zeff is lower than actual
nuclear charge.
• Zeff increases
toward nucleus
• Energy of electron
1)The higher the main shell the
more the energy of electron.
2)Subshell also contribute to the energy of
electron:
ns > np > nd > nf
• This explains certain periodic changes
observed.
55
Effective Nuclear Charge
03
56
Energy of orbitals in a multi-electron atom
Energy depends on n and l
n=3 l = 2
n=3 l = 0
n=2 l = 0
n=1 l = 0
n=3 l = 1
n=2 l = 1
57
“Fill up” electrons in lowest energy orbitals (Aufbau principle)
??
Be
Li
B5
C
3
64electrons
electrons
22s
222s
22p
12 1
BBe
Li1s1s
1s
2s
H
He12electron
electrons
He
H 1s
1s12
58
The most stable arrangement of electrons
in subshells is the one with the greatest
number of parallel spins (Hund’s rule).
Ne97
C
N
O
F
6
810
electrons
electrons
electrons
22s
222p
22p
5
246
3
Ne
C
N
O
F 1s
1s222s
59
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
60
< 4f< 5d< 6p< 7s< 5f< 6d< 7p
Electron Configuration of Atoms
•
Rules of Aufbau Principle:
1. Lower energy orbitals fill first.
2. Each orbital holds
two electrons; each
with different ms.
3. Half-fill degenerate
orbitals before pairing
electrons.
61
What is the electron configuration of Mg?
Mg 12 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s2
2 + 2 + 6 + 2 = 12 electrons
Abbreviated as [Ne]3s2
What are the possible quantum numbers for the
last (outermost) electron in Cl?
Cl 17 electrons
1s22s22p63s23p5
1s < 2s < 2p < 3s < 3p < 4s
2 + 2 + 6 + 2 + 5 = 17 electrons
Last electron added to 3p orbital
n=3
l=1
ml = -1, 0, or +1
ms = ½ or -½
62
4f
5f
63
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Electron Configuration and the Periodic Table
Electron Configuration of Atoms
06
64
65
Using periodic table write Noble gas notation
for the following elements:
a)S
b)Fe
[Ne]3s23p4
[Ar] 4s23d6
c)Se
[Ar] 4s23d104p4
d)Gd
[Xe]6s24f75d1
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
66
< 4f< 5d< 6p< 7s< 5f< 6d< 7p
67
Electron Configuration of Atoms
08
• Anomalous Electron Configurations: Result from
unusual stability of half-filled & full-filled subshells.
•
Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5
•
Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10
68
Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic
field
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired
electrons.
69
* See page 261 of your book for molecular structure of N2 vs O2
Molecular Orbital Theory: Other Diatomic
Molecules
Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
71
Periodic Properties 01
72
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
(Sigma)
Zeff  Z – number of inner or core electrons
Z
Core
Na
11
10
1
186
Mg
12
10
2
160
~ Zeff
Radius/pm
Al
13
10
3
143
Si
14
10
4
132
Within a Period
as Zeff increases
radius decreases
73
Size of the atoms
74
(pm)
75