Chapter 4
Spectroscopy &
Arrangement of Electrons(e-)
Properties of Light
• Electromagnetic Radiation- form of
energy w/ wavelike properties as it
travels through space
• Electromagnetic Spectrum- classifies as
electromagnetic radiation based on
wavelength( ) and frequency(v)
The Electromagnetic Spectrum
Visible Light
Violet
Blue
Green Yellow Orange
400nm 450nm 500nm 550nm 600nm 700nm
Gamma Rays X-Rays UV-Rays
v
Red
1019 Hz
1017 Hz
1015 Hz
Infrared Microwave Radio
1014 Hz
109 Hz
Long-Wave
108-106 Hz 105 Hz
Short Wavelength
Long Wavelength
High Frequency
Low Frequency
• Frequency(v)- # of waves that pass a given
point in a specific time
• Wavelength( )- the distance between
corresponding points on a wave
(ex: peak to peak)
Speed of Light(c)
• Speed of Light = wavelength x frequency
•
c= v
• Speed of Light is a Constant:
c = 3.0 x 108 m/s
• Frequency and Wavelength are Inversely Proportional
:
v
Photoelectric Effect
• When light particles(photons) collide with a
metal, the photons knock electrons (e-) loose
• These electrons move toward the positive terminal
creating an electric current(electricity)
Light as Particles
• Light has wave and particle properties
(Max Planck)
• A Quantum of Energy- the minimum quantity if
energy that can be lost or gained by an atom
E = hv E = energy of a photon (J)
h = Planck’s constant (6.626 x 10-34JS)
v = frequency (Hz)
• Photons- particles of light carrying q quantum of
energy
Dual Wave-Particle Nature of Light
• Light has both the properties of waves and
of particles
– Wave Properties: light can be bent as it passes
through objects
– Particle Properties: photons have mass and
exert force on other objects
Emission Spectra
• Emission spectra are fingerprints of an atom
• Every atom gives off different colors from
the visible light spectrum when they
release absorbed energy
Energy States of Atoms/Electrons
• Ground State- lowest energy level of an
electron within an atom
• Excited State- a higher energy level within
an atom that an electron may exist in
– Energy must be absorbed for an electron to go
from ground to excited state
– Energy is given off as visible light when an
atom returns to ground state
– Every atom gives off a unique spectrum based
on the movement of it’s electrons
Energy of a Photon
• Ephoton = Efinal - Einitial
E2
Excited State
E2 – E1 = Ephoton = hv
E1
Ground State
Bohr Model of the Hydrogen Atom
• Bohr’s model indicates that as atoms absorb
energy their electrons move to higher
energy levels
• When the absorbed energy is given off as
visible light the electrons return to their
ground state
Bohr Model
Quantum Model of the Atom
• The Bohr model was more accurate than
previous models but was only completely
accurate for Hydrogen, other elements did
not behave exactly as Bohr predicted
• The Quantum model was later developed
based on work of many scientists including
Schrodinger, Heisenberg, & Einstein
Quantum Model of the Atom
Electrons as Waves
• Louis deBroglie proved that electrons had
wave properties by showing that electrons
could produce interference patterns like
sound and light waves
• Passing electrons through a crystal also
caused the stream of electrons to bend like
light waves do
Interference Patterns
Heisenburg Uncertainty Principle
• This principle states that it is impossible to
determine simultaneously both the position
and velocity of an electron
• This theory led to the concept of the
electron cloud
Quantum Theory
• This theory describes mathematically the
wave properties of electrons based on
probability
– Electrons are not in set energy levels
– Electrons are in 3D orbits around the nucleus
called orbitals
Orbitals
• Orbitals are 3D regions around the nucleus
of an atom that indicate the probable
location of an electron
Quantum #’s
• Quantum #’s are used to specify the
properties and location of electrons in
orbitals around the nucleus
• There are 4 quantum #’s, each is
represented by a letter : n, l, m, & s
Principle Quantum #(n)
• The principle quantum number indicates the
main energy level of an electron and it’s
distance from the nucleus
• n=1 : ground state, e- close to nucleus
• n=7 : excited state, e- further from nucleus
* As the n value increases so does the energy
of the e-, and the distance from the nucleus
Angular Momentum Quantum #(l)
• The angular momentum quantum # indicates the
shape of an orbital
• Each l value has a corresponding shape
l = 0 : s-shape, sphere orbital around nucleus
- an s-orbital can hold up to 2 electrons
l = 1 : p-shape, 2 lobes on either side of nucleus
- an atom can have 3 p-orbitals, one in each
plane(x,y,z)
- a p-orbital can hold up to
6 electrons
Angular Momentum Quantum #(l)
• l = 2 : d-shape, 4 lobes “clover” around
nucleus
– An atom can have up to 5 d-orbitals
– A d-orbital can hold up to 10 electrons
Angular Momentum Quantum #(l)
• l = 3 : f-shape, 8 lobes around nucleus
– An f-orbital can hold up to 16 electrons
Magnetic Quantum #(m)
• The magnetic quantum # indicates the
orientation of an orbital around the nucleus
Spin Quantum #(s)
• The spin quantum number indicates the
direction that an electron is spinning, either
clockwise or counterclockwise
• s = +1/2 : clockwise spin
• s = -1/2 : counterclockwise spin
* 2 electrons in the same orbital must have
opposite spins
Quantum #’s
n – principle quantum #
distance from nucleus / main energy level
l – angular momentum quantum #
shape of orbital
m – magnetic quantum #
orientation of orbitals around nucleus
s – spin quantum #
direction of e- spin around nucleus
Pauli Exclusion Principle
• The Pauli Exclusion principle states that no
2 electrons in the same atom can have the
same combination of 4 quantum #’s
• This means that no 2 electrons could be in
the same place at the same time
Electron Configuration Notation
• 1s22s22p63s23p64s23d104p65s24d105p66s2…
Principle Quantum #
or Energy Level
Angular Momentum Quantum #
or Orbital Shape
• Ex: Carbon = 6 electrons
• C = 1s22s22p2
• Ex: Sodium = 11 electrons
• Na = 1s22s22p63s1
# of Electrons in Orbital
Noble Gas Notation(Shortcut)
• To eliminate repetitive electron configuration for
elements with large #’s of electrons the symbol of
a Nobel Gas can be substituted for a portion of the
electron configuration
• Ex: K = 1s22s22p63s23p64s1
[Ar]4s1
• Ex: Zn = 1s22s22p63s23p64s23d10
[Ar]4s23d10
Orbital Notation
• Orbital notation uses boxes and arrows to
indicate electrons in orbital by energy level
• Aufbau Principle- an electron will occupy
the lowest possible energy level that can
hold it
• Hund’s Rule- orbitals of equal energy will
each receive one electron before they
receive a second electron
Order of Atomic Sublevels
• Orbital Notation:
• P:
• O:
* Each p-orbital gets 1 ebefore it gets a 2nd e-
Examples
• N:
• Si :
• Fe :
• Mg :
e
Configurations
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
Flame Test Lab
Compound
NaCl
NaNO3
Sr(NO3)2
Ca(NO3)2
Color Description
Unknown
A
B
C
Ba(NO3)2
KNO3
D
Cu(NO3)2
E
CuSO4
F
LiNO3
Color
Compound
Description
Calculations
Compound
Flame
Color
λ (nm)
ν (Hz)
LiNO3
Ba(NO3)2
Ca(NO3)2
General Equations:
Speed of Light: c
c = 3.0 x 108 m/s
=λν
Energy: E
=hν
h = 6.626 x 10-34 Js
E (J)
END OF CHAPTER 4
NOTES !!!