Download Waves - Walton High

Waves
What do light waves have
to do with chemistry?
Now What?
• 3 sub-atomic particles
• Protons and Neutrons in
nucleus
• Electrons?
• How are they arranged in space?
• Why weren’t the negatively charged
electrons pulled into the positively
charged nucleus?
Light?
• Certain elements emit visible
light when heated in a flame
• Analysis showed that an
element’s chemical behavior is
related to the arrangement of
electrons in its atoms
Electromagnetic
Radiation
• Light is just one form of
electromagnetic radiation
• Electromagnetic radiation
travels in the form of waves
• All waves have amplitude,
wavelength, frequency and
speed
Amplitude
• The height of
the wave
measured from
the origin (or
base line) to its
crest or trough
Wavelength
• Symbol is λ
• The distance
between two
crests or two
troughs
Frequency
• Its symbol is 
• The number of wave cycles that
pass a given point in one
second
• Measured in hertz (Hz)
(cycles/sec)
• c = λ
Frequency vs
Wavelength
• Long wavelengths have low
frequencies
• Short wavelengths have high
frequencies
Speed of Light
• Is it really a constant?
• It equals 3.00 x 108 m/s
• Symbol is c
Electromagnetic Spectrum
Planck’s Theory
• Max Planck (1900) - accurately predicted
how the spectrum of radiation emitted or
absorbed by an object changes with its
temperature
• proposed that there is a restriction on
the amounts of energy that an object
emits or absorbs (quantum)
Equantum= h 
E = amount of energy
h = Planck's constant (6.6262 x 10-34 J*s)
 = frequency of the radiation
Photoelectric Effect
• Albert Einstein (1905) - used
Planck's equation to explain that
energy has mass (E = mc2)
• Wave Model could not explain when
electrons are emitted from a metal’s
surface when light of a certain
frequency shines on it
• light consists of quanta of energy
that behave like tiny particles of
light (photons)
Photoelectric Effect
• the mass of each photon is
dependent on its wavelength
• Concerns the frequency (color) of
the incident light
• Einstein proposed that
electromagnetic radiation has both
wavelike and particlelike natures
Photoelectric Effect
• Einstein proposes that a photons
energy depends on its frequency:
Ephoton= h 
E = amount of energy
h = Planck’s constant
 = frequency
• Also proposes that the energy of the
light photon must possess
(minimally) the energy required to
free an electron from an atom of the
metal
What does this mean?
Matter can emit or
absorb energy
only in wholenumber
multiples of h :
1 h, 2 h, etc.
Louis de Broglie
• Louis de Broglie (1923) – used Planck’s
and Einstein’s work to describe matter
as having wavelike behavior
• proposed that electrons (or any other
particle) have a characteristic
wavelength dependent on their mass
and velocity
• λ=h
m
λ = wavelength
 = frequency
h = Planck’s constant
m = mass of particle
This principle is used in
electron microscopes,
where electrons streams
are diffracted the same
way light waves are
diffracted by lenses.
A picture taken with an
electron scanning
Microscope.
Photons
• Particles of light are called photons
• No mass
• Each photon carries with it a
specific value of energy (quantun)
based on its frequency
• High frequency photons carry large
amounts of energy
Line Spectra
• A spectrum composed of just
lines that are characteristic of
an element
• Created when an element is
vaporized in a flame, or when
electricity is passed through the
element’s gaseous form
When the light passes
through a slit and
then a prism,
a line spectrum is formed.
Photograph of a line
spectrum from helium
Several examples
of line spectra
Atomic Emission
Spectra
• A line spectrum is also called
an atomic emission spectrum
• It’s formed by the energy given
off, in the form of light, when an
electron moves from its excited
state to its ground state
Bohr’s Model of the
Atom
• Used the dual wave-particle model of
light
• Proposed the hydrogen atom had only
certain allowable energy states
• Ground State – The lowest allowable energy
level (n=1)
• Excited State – When an electron gains energy
and jumps to a higher quantum level
• Quantum Number (n) – The number
assigned to each electron orbit
• Model explained hydrogen’s spectrum
but no other element
How are electrons arranged
around an atom?
• Bohr determined that electrons
gained or lost energy in quanta, or
specific amounts. These he labeled
as quantum numbers, or n.
• The n numbers refer to the energy
levels, which are the same as the
period numbers on the periodic
table.
Quantum Numbers
• Each electron can be described
with its own individual set of
quantum numbers
• Acts like an address for each
electron in an atom
• Set of 4 different numbers
Principle Quantum
Number
• (n) – gives the energy level the
electron is located in
• integral value (whole numbers)
Angular Momentum
Quantum Number
• (l) – describes the sublevel
location of the electron
• l values: 0 – s; 1 – p; 2 – d; 3 – f
• a.k.a: azimuthal quantum number
Magnetic Quantum
Number
• (ml) – describes the exact orbital
location in the given sublevel
0
s
-1, 0,+1
-2,-1, 0, +1,+2
-3, -2,-1, 0,+1,+2,+3
p
d
f
Electron Spin
Quantum Number
• (ms) – describes the spin of the
electron
• “up” = +½; “down” = - ½
Electron Cloud Pictures:
l, along with n and the third quantum
number, m, is responsible for determining the
shape of an electron's probability cloud.
n=1, l=0, m=0
1S1 or 1S2
n=3, l=2, m=1
3d4
n=3, l=2, m=2
3d5
n=4, l=2, m=2
4d6
What are sublevels?
• Each energy level is divided into
sublevels
• There are four different
sublevels: s, p, d and f
You can tell which sublevel the
electrons of an element are in
by looking at the element’s location in the
Periodic Table.
S = pink
p = blue
d = green
f = yellow
Probability and Orbitals
• Orbital: region around the
nucleus where an electron is
likely to be found
• Have characteristic shapes, sizes
and energies
• Each can only hold 2 electrons
with opposite spins
How do orbitals relate
to energy?
• All principle energy levels are
designated with a quantum number,
n
• These numbers are the same as the
period numbers on the Periodic
Table
• Each principle energy level is
divided into sublevels: s,p,d,f
• Each sublevel has an odd number of
orbitals: s has 1, p has 3, d has 5,
and f has 7
Energy Level
Sublevels
Orbitals
# of
Electrons
1
s
1
2
2
s
p
1
3
2
6
3
s
p
d
1
3
5
2
6
10
4
s
p
d
f
1
3
5
7
2
6
10
14
What do the orbitals
look like?
Orbital in the s sublevel
Orbitals in the p sublevel
Orbitals in the d sublevel
Orbitals in the f sublevel
Each one of these orbitals
holds only two electrons.
The d sublevel holds 10
electrons total, and the f
sublevel holds 14 electrons
total.
Electrons are Weird
• Heisenberg Uncertainty Principle:
• It is impossible to know both the
location and velocity of an electron
at the same time
• An electron can only be located
when a photon strikes it. The
collision causes it to change
direction and velocity.
What is electron
configuration?
• Each electron has a characteristic
energy and a characteristic spin
• Electron Configuration - Each
electron is located in a specific
orbital with a specific quantum
number
• 3 rules to determine where each
electron is located and what spin it
has
What are the rules for
assigning configurations?
• Aufbau Principle: electrons are
added to an atom one at a time
starting with the lowest energy
orbital
• 1s will get the first 2 electrons (lowest
energy)
• 2s will get the second 2 electrons (next
higher energy)
• 2p will get the next 6 electrons (next
higher energy)
What are the next rules for
assigning configurations?
• Pauli Exclusion Principle: Each
orbital can hold up to two electrons;
must have opposite spins
• Hund’s Rule: When electrons occupy
orbitals of equal energy, one
electron enters each orbital until all
the orbitals contain one electron
with spins that are parallel
What are electron
spins?
• Electrons have their own spins,
represented by up and down
arrows
• They can spin clockwise or
counterclockwise
• Each orbital holds two electrons
and they have to be spinning in
opposite directions
How do I represent this?
• Electron spins and the order in
which they fill orbitals can be
represented using an orbital
diagram
• Orbital diagrams use arrows
and lines or squares to
represent electrons in their
orbitals
What are some
examples?
• What is the orbital diagram for
boron?
• Its electron configuration is
1s22s22p1
• Its orbital diagram is:
1s
2s
2p
What is sublevel
notation?
• A type of electron configuration
where we describe the
arrangement of electrons in an
atom
• The order in which the levels fill
is shown on the next slide and
on your handout
Electron Filling Order
Long form of the periodic table
Color code represents in a general way
Filling of s orbitals
Filling of p orbitals
Filling of d orbitals
Filling of f orbitals
Long Form of Periodic Table
1
H
3
1s
2s
3s
2
2p
4
5
Li
Be
B
11
12
13
Na
Mg
4s
5s
Al
3d
4d
19
20
21
22
23
24
K
Ca
Sc
Ti
V
Cr
37
38
39
40
41
42
Rb
Sr
Y
6s
7s
55
56
57
Cs
Ba
La
87
88
89
Fr
Ra
Ac
5d
6d
58
59
60
61
62
63
Ce
Pr
Nd
Pm
Sm
Eu
90
91
92
93
94
95
Th
Pa
U
Np
Pu
Am
Copyright (c) 1997 Tanner McCarron http://www.tannerm.com
Zr
Nb
Mo
64
65
66
67
68
69
70
71
72
73
74
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
Hf
Ta
W
96
97
98
99
100
101
102
103
Cm
Bk
Cf
Es
Fm
Md
No
Lr
4f
5f
3d
4d
25
26
27
28
29
30
31
Mn
Fe
Co
Ni
Cu
Zn
Ga
43
44
45
46
47
48
49
Tc
He
6
7
8
9
10
C
N
O
F
Ne
14
3p
4p
5p
15
16
17
18
Si
P
S
Cl
Ar
32
33
34
35
36
Ge
As
Se
Br
Kr
50
51
52
53
54
Sn
Ru
Rh
Pd
Ag
Cd
In
Sb
Te
I
Xe
75
76
77
78
79
80
81
82
83
84
85
86
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
5d
6p
What are some
examples?
• Hydrogen: has 1 electron
• From the Periodic Table, it is in
row 1, so n = 1
• It has a 1s orbital
• Its configuration is 1s1
Another example
• Helium has 2 electrons
• According to the P. Table, it is
in the 1st row, so n = 1
• It has an s orbital
• Its configuration is 1s2
Another example
• Lithium has 3 electrons
• The first two electrons are the
same as in He, so the first term
is 1s2
• It is in the 2nd row, so n=2
• Its configuration is 1s2 2s1
Noble-Gas Notation
• Sub-level notation for an element using
the noble gas in the previous period
(Core e-) and the electron
configuration for the energy level
being filled (Valence e-)
• Core e-: Go up one row and over to the
Noble Gas.
• Valence e-: On the next row, fill in the # of
e- in each sublevel
Examples
• Sodium: [Ne] 3s1 or
1s22s22p6 3s1
• Chlorine: [Ne] 3s23p5 or 1s22s22p6 3s23p5
The Exceptions
• Cr and Cu – Should fill 4s then 3d,
then 4p
• Cr [Ar] 3d5 4s1
• Cu [Ar] 3d10 4s1
• Quantum Mechanics says there is a special
stability associated with half-filled and
filled sets of equivalent orbitals
• Lowers the 3d orbitals below the 4s orbitals
so that [Ar] 3d5 4s1 is lower in energy than
[Ar] 3d4 4s2
What are valence
electrons?
• Electrons in the outermost energy
level (the n #)
• Can be determined by counting the
number of s- and p-electrons since
the most recent noble gas
• These electrons are primarily
responsible for the reactivity's of the
elements in the s- and p-sections of
the Periodic Table
What are some
examples?
• How many valence electrons
does chlorine have?
• How many valence electrons
does sodium have?
• How many valence electrons
does aluminum have?
Lewis Dot Structures
• Element Symbol surrounded by
dots representing the valence
electrons
• For example:
• H.
.Be.
..
.N.
.
..