Parts of a Wave
Unit 4:
Electron in Atoms
• How are electrons related to
electromagnetic radiation?
• Does the location of an electron
influence the behavior of an atom?
• How do we most accurately describe the
location of electrons around a nucleus?
•Wavelength (l)
•Length of one wave
•Measured in unit of distance (m, nm, etc.)
l
•Frequency(n)
•Number of cycles in one second
•Measured in Hertz (1 Hz = 1 cycle/sec = s-1)
EMR
EMR
• All waves have different l and n
• What is relationship between l and n?
• Electromagnetic Radiation (EMR)
– Energy with wave-like behavior as it travels
through space
– Higher Frequency = Lower Wavelength
• Electromagnetic Spectrum
• Visible Light is only part our eyes can
detect
– Organizes emr according to l and n
– We only see about 1.5% of the spectrum
EMR
• To find n or l of a wave
Atomic Spectra
Ground State
•
–
c = ln
–
Excited State
•
– c = Speed of Light = 3 x 108 m/s
– If you know l or n, you can determine what type
of EM wave it is
Lowest possible energy level
Normal energy level of e-
–
–
–
Higher energy Level
Caused by energy gain of eUnstable and short-lived
state
1
Atomic Spectra
• Emission Process
Atomic Spectra
• Atomic Emission Spectrum
– e- start in ground state
– e- absorb energy from source and jump to
excited state
– e- release energy due to instability
• Return to ground state
• Released energy appears as light
– Shows wavelengths of visible light released
by excited electrons
– Unique for each element
• Intensive property – used to identify
– Viewed with Spectroscope
• Prism that separates colors of light
Energy
• Photoelectric Effect
– Heinrich Hertz – 1887
– Electrons are emitted by
matter when emr is
absorbed
– Short wavelength needed
• Often blue light or UV
– Used in solar cells,
cameras, etc.
Energy
• Photon
– Albert Einstein - 1905
– Particle of radiation
– Zero mass
– Carrying a quantum of
energy
• Radiation is a stream of
photons
Energy
• Quantum
– Max Planck – 1900
– Minimum amount
of energy that can
be gained or lost
by an atom
– Packet of energy
Energy
• To calculate energy of a photon:
E = hn
– h – Planck’s constant = 6.626 x 10-34 J‧s
• For wavelength:
E = hc
l
•Higher E
–Higher n and lower l
2
Energy
• Is energy a wave or a
particle?
– Yes
• Wave-Particle Duality
– Louis de Broglie -1924
– Matter moves like a wave
– Like standing waves of a
stringed instrument
Quantum Mechanical Model
• Quantum Mechanical Model
– AKA: Electron Cloud Model
– Erwin Schrödinger – 1933
– Described changes to system over
time
– Schrödinger equation is solved to
indicate probable regions where
e- is located
Quantum Numbers
• Principal quantum
number (n)
– Energy level of e– Determines size of
area where e- can be
found
– Higher n = larger
area for movement
– Number from 1–7
Bohr Model
• Bohr’s Ring Model
– Electrons travel like
planets around
nucleus
– Electrons move
between energy
levels in atom
– Only works for
hydrogen atoms
Quantum Numbers
• Quantum Numbers
– Describe region where e- should be located
– Regions called Suborbitals
– 4 numbers needed to best describe location of e -
Quantum Numbers
• Angular momentum
quantum number(l)
– Shape of suborbital
– Shapes are bigger for
higher energy levels
3
Suborbitals
• S Suborbitals
Suborbitals
• P Suborbitals
– Spherical
– 1 orbital per energy
level
– 2 e- per energy level
– n>0
– Dumbbell-shaped
– 3 orbitals per
energy level
– 6 e- per energy
level
– n>1
Suborbitals
• D suborbitals
– 5 orbitals per energy
level
– 10 e- per energy level
– n>2
Quantum Numbers
Magnetic quantum
number (mI)
Which orbital the
electron is in
Gives the axis
orientation (X,Y,Z)
Suborbitals
• F Suborbitals
– 7 orbitals per
energy level
– 14 e- per
energy level
– n>3
Quantum Numbers
• Electron spin quantum number (ms)
– Which e- in orbital
– Either +1/2 or -1/2
4
Electrons
• Where exactly are
the electrons?
– Hard to tell
– Electrons are almost
like spinning fans
Electron Configuration
• Electrons fill in an atom in a specific order
• We follow 3 rules to get the correct electron
configuration for each atom:
1. Aufbau Principle
2. Pauli’s Exclusion Principle
3. Hund’s Rule
Electron Configuration
• Pauli’s Exclusion
Principle
– Wolfgang Pauli – 1925
– Only 2 electrons per
orbital
– Electrons in same orbital
must have opposite
spins
– Spin is represented by an
arrow
Electrons
• Uncertainty Principle
– Werner Heisenberg – 1927
– More you know about
position of e-, less you
know about where it’s
going
– Can’t really know all 4
quantum numbers
Electron Configuration
• Aufbau Principle
– As electrons are
added to atom they
arrange themselves in
orbitals
– Fill up in order of
lowest energy (1s) to
the highest energy
Electron Configuration
• Hund’s Rule
– Friedrich Hund – 1925
– Lowest energy
configuration is the
one with maximum
unpaired electrons
– Pair e- in orbital only
when necessary
5
Electron Configuration
• Orbital Diagrams
– Show arrangement of electrons in orbitals
– Max of 2 e- per box
– Draw all boxes in suborbital (even if empty)
– Ex: P – 15 e1s2
2s2
2p6
3s2
3p3
Valence Electrons
• Core electrons
– Inner electrons
• Valence electrons
– The electrons in the outermost energy level
– Determine how atoms bond
– Count electrons by level to determine valence
Noble Gas Configuration
• Noble Gases
– Stable, unreactive gases
• e- config won’t change
– Far right of table – group 18 (He, Ne, etc.)
• Noble Gas Configurations
1. Place symbol of previous noble gas in
brackets
2. Continue configuration after last e- of noble
gas
Valence Electrons
• You can also use periodic table to determine
valence e• Valence e- are same as column of main group
elements (s- and p- block)
• d- and f- block elements always have 2
valence e-
Lewis Dot Diagrams
• Lewis Dot Diagrams
– Indicate arrangement of valence e- in atom
– Draw valence e- as dots around chemical symbol
– Spread out in 4 directions – MAX 8 dots
– Pair only when needed
– Ex: C
• 4 valence e– Ex: O
• 6 valence e-
6