Chemistry Unit 2: the
nd
2
half!
Electrons and their Properties
I. Why isn’t the electron pulled into
the nucleus?
A. Remember the nucleus is positive and
the electron is negative.
B. However, there is a special relationship
between light (electromagnetic radiation)
and the electrons of an atom
II. Properties of Light
A. made of many particles
B. described as a wave
C. Electromagnetic Radiation
1. Includes: x rays, ultraviolet, infrared,
microwaves, radio waves, visible light
2. Constant speed: 3.0 x 108 m/s
D. Wavelength (λ)
1. distance between 2 waves
2. unit varies depending on type of EM ray
3. List the types of EM radiation in order from
longest wavelength to shortest wavelength.
E. Frequency (v)
1. number of waves passing a point in a given
amount of time
2. Unit: waves per second
1 wave/second = 1 Hz (Hertz)
3. Arrange EM radiation types form highest
frequency to lowest frequency
F. Mathematical Relationship
c = λv
c : speed of light
λ : wavelength
v : frequency
III. Photoelectric Effect
A. Emission of electrons from a metal
when light shines on the metal
B. Wave model could NOT explain this
1. light below a minimum frequency would not
knock off any electrons
C. Max Planck
1. Objects emit energy in small specific
amounts know as quanta
2. Quantum – minimum amount of energy that
can be lost or gained by an atom
3. Mathematical Relationship
Equantum = hv
E: energy (Unit is Joules)
v: frequency (Hz)
h: Planck’s Constant
• h = 6.626 x 10-34 J • s
D. Einstein
1. Light has dual nature
2. wave and particle-like properties
3. made of photons
1. photons have zero mass and a quantum of energy
4. Mathematical relationship
Ephoton = hv
IV. Atomic Emission Spectrum
A. When energy passes through a gas at
low pressure  the potential energy of
some of the atoms increases
B. Ground State  Excited State
1. Ground state: lowest energy state of an atom
2. Excited state: atom has higher potential
energy than ground state
C. Excited State  Ground State
1. gives off energy gained
2. produces colors
3. Example: Neon signs
D. Atomic Emission Spectra
1. Set of frequencies emitted by the element
2. Unique to each element
Ex. Strontium – Red
Hydrogen – Pink
3. Known as a line-emission spectrum
Hydrogen line emission spectrum
V. Bohr’s Model & Atomic Emission
Spectrum
A. Bohr proposed that electrons can circle
the nucleus in fixed paths known as orbits
1. Electron in orbit = fixed energy
2. Therefore, the lowest energy state is the orbit
closest to the nucleus
a. energy increases with each successive orbit away
from the nucle
B. Compare Bohr’s model to the rungs of a
ladder. (Think about potential energy)
C. Bohr’s model can be used to explain
atomic emission lines
1. in orbit, energy is neither gained nor lost
2. Gain energy  move to higher orbit (Excited
state)
3. Electron drops down  photon emitted
Ephoton = higher orbit energy – lower orbit energy
D. Bohr’s Model only worked for the lineemission spectrum of hydrogen
1. Did not fully account chemical behavior of
atoms
2. Electrons DO NOT move in fixed circular
orbits
VI. Quantum Model of the Atom
A. Louis de Broglie
1. Electrons could behave as waves at specific
frequencies
2. Electrons could be bent and diffracted like
light waves
3. Electron waves can also interfere with one
another
a. when waves overlap
b. reduces energy in some areas, increases it in
others
Light Diffraction
Electron Diffraction
4. De Broglie’s equation for electron
behavior:
λ = h__
mv
B. Heisenberg Uncertainty Principle
1. Dealt with detection of electrons
2. Photons have about the same energy as
electrons
a. When photons are used to locate electrons, they
knock them off.
b. There is always uncertainty when trying to locate
electrons
3. Heisenberg Uncertainty Principle states:
it is impossible to determine
simultaneously both the position and
velocity of an electron or any other particle
a. fundamental to the foundation of scientists
understanding of light and matter
C. Schrodinger Wave Equation
1. developed equation that treated electrons as
waves
2. Along with Heisenberg Principle, laid the
foundation of Quantum Theory
3. Quantum Theory: describes mathematically
the wave properties of electrons and other very
small particles
a. Proposed probable locations of electrons
b. Exist in regions known as orbitals (3-dimensional
region around the nucleus that indicates the probable
location of an electron)
VII. Atomic Orbitals & Quantum
Numbers
A. Quantum Numbers
1. Specify the properties of atomic orbitals and
the properties of electrons in the orbitals
2. Indicate the following for an orbital:
Main energy level
The shape
The orientation
3. Types of Quantum Numbers
Principal Quantum Number (n)
Angular Momentum Quantum Number (l)
Magnetic Quantum Number (m)
Spin Quantum Number (+1/2 or -1/2)
a.Principal Quantum Number (n)
1. main energy level
2. n = a positive integer (1, 2, 3, 4, etc)
3. Often referred to as shells
b. Angular Quantum Numbers (l)
1. Exist at each Main energy level AFTER the
first one
2. Known as sublevels – orbitals of different
shapes
3. Indicates the shape of the orbital
s, p, d, f
4. Each orbital is represented by a principal
quantum number and the letter of the sublevel
c. Magnetic Quantum Numbers (m)
1. Indicates orientation around the nucleus
2. s sublevel has only one orientation, but p, d,
and f have multiple possibilities.
3 orientations for the p sublevel
5 orientations for the d sublevel
d. Spin Quantum Number (+1/2 or -1/2)
1. electrons spin on an internal axis
2. single orbital holds 2 electrons and each has
an opposite spin
Pauli Exclusion Principle
VIII. Electron Configuration
A. method of indicating the arrangement of
electrons around a nucleus
B. Consists of the following:
1. A number indicates energy level (n)
2. A letter indicates type of orbital: s, p, d, f (l)
3. A superscript indicating the number of
electrons in the orbital
Example: 1s2 (This is Helium)
C. Rules for Electron Configuration
1. Aufbau Principle
Shows the order in which electrons occupy
orbitals
States: an electron occupies the lowest-energy
orbital that can receive it
2. Pauli Exclusion Principle
Importance of the spin quantum number is
reflected here
States: no two electrons in the same atom can
have the same 4 quantum numbers
How can an orbital hold two electrons then?
OPPOSITE SPINS!!!!
3. Hund’s Rule
Placing unpaired electrons in separate orbitals
in the same sublevel
States: orbitals of equal energy are each
occupied by a one electron before any orbital is
occupied by a second electron, and all
electrons in singly occupied orbitals must have
the same spin.
D. Methods
1. There are three methods for writing
electron notation
A. Orbital Notation – Shows spin direction of
each electron
B. Electron-Configuration Notation – Uses the
superscript with the orbital letter to represent
number of electrons in each
C. Noble Gas Notation – shorthand method
used to represent larger elements
Noble gases are in Group 18!
E. How do we do this? (STEPS)
1. Determine the total number of electrons
to be represented
2. Use the Aufbau process to fill the
orbitals with electrons
3. The sum of the superscripts should
equal the total number of electrons
F. Practice (Assume atoms are
neutral)
Hydrogen
Orbital Notation:
EC Notation:
Helium
Orbital Notation:
EC Notation
Boron
Orbital Notation:
EC Notation:
G. Noble Gas Notation
1. Noble Gases – Group 18 Elements
Helium, neon, argon, krypton, xenon, and radon
2. Method of simplifying the electron
configuration of larger elements
H. Practice Problems
Write the complete electron configuration
for Iron (Fe) and the noble-gas notation.
Write both the electron configuration and
noble gas notation for an atom of Barium.