The Arrangement of Electrons
in Atoms
CH. 4 MODERN CHEMISTRY
A New Atomic Model
 Cathode Ray Tube Summary
 As of 1911, Rutherford’s Gold Foil Experiment had
shown scientists the nucleus of an atom was very
small, dense, and positively charged.
 Why was Rutherford’s model of the atom incomplete?
 Rutherford’s
model did not explain how electrons
were distributed around the nucleus.
Just assumed electrons located around nucleus
Wave-like Properties of Light
 Over the next decade, scientists began to discover the
dual nature of electrons - particles and waves.
 Electromagnetic radiation is a form of energy that
exhibits wavelike behavior as it travels through space.
 Together, all the forms of electromagnetic radiation
form the electromagnetic spectrum.
 Visible
spectrum: visible portion of spectrum
 But there are many other forms of energy
Electromagnetic Spectrum
Properties of Light
 Wavelength (λ) :
distance between
corresponding points on
adjacent waves
 Frequency (ν):
number of waves that
pass a given point per
unit time time (usually 1
sec)
Properties of Light
 Frequency and wavelength are mathematically related
to each other:
c = λν
c
: speed of light (m/s)
 λ : wavelength (m)
 ν : frequency of wave (s−1 or 1/s).
Photoelectric Effect
Particle Description of Light
 Electromagnetic
radiation strikes the
surface of the metal
 Electrons are ejected
from the metal causing
an electric current
Photoelectric Effect
 The photoelectric effect refers to the emission of
electrons from a metal when light shines on the metal.
 Not
all forms of electromagnetic radiation were
“powerful” enough to eject electrons from the surface
 A minimum amount of energy was needed
 A quantum of energy - the minimum quantity of
energy that can be lost or gained by an atom
Photoelectric Effect
 German physicist Max Planck proposed the following
relationship:
E = hν
E
: quantum of energy (J – joules)
 ν : frequency of the radiation emitted; s−1
 h : fundamental constant called Planck’s constant;
h = 6.626 × 10−34 J• s
Particle Description of Light
 photon : particle of electromagnetic radiation having
zero mass and carrying a quantum of energy.
 The energy of a particular photon depends on the
frequency of the radiation.
Ephoton = hν
The Hydrogen Line Emission Spectrum
The Hydrogen Line Emission Spectrum
 When investigators passed electric current through a
vacuum tube containing hydrogen gas at low pressure,
they observed the emission of a characteristic pinkish
glow.
 When a narrow beam of the emitted light was shined
through a prism, it was separated into four specific
colors of the visible spectrum.
 The four bands of light were part of what is known as
hydrogen’s line-emission spectrum.
The Hydrogen Line Emission Spectrum
 What had scientists expected to observe?
A
continuous range of frequency (all bands of ER)
The Hydrogen Line Emission Spectrum
 The lowest energy state of an atom is its ground
state.
 A state in which an atom has a higher potential energy
than it has in its ground state is an excited state.
Bohr Model of the Atom
 By 1922, Niels Bohr proposed a hydrogen-atom model
that linked the atom’s electron to photon emission.
 electrons can circle the nucleus only in allowed paths,
called orbits.
The energy of the electron is higher when it is in an
orbit farther from the nucleus.

Bohr Model of the Atom
 When electrons move from an excited state to a ground
state or to a lower energy state, a photon of energy is
emitted from an atom
Process called emission (high to low orbit)
absorption (low to high orbit): energy is added
to an atom to move an electron from lower energy level
to higher energy level

3) The energy of a photon can be directly calculated if
you know the frequency of radiation being emitted
Photo Absorption & Emission
Section 4.2
QUANTUM MODEL OF THE
ATOM
Electrons as Waves
 French scientist Louis de Broglie suggested that
electrons be considered waves confined to the space
around an atomic nucleus.
 electron waves could exist only at specific
frequencies.
 It followed that if each electron has a specific
frequency, it also has a corresponding energy.
 Recall equation E = hv
Heisenberg Uncertainty Principle
 German physicist Werner Heisenberg proposed that
any attempt to locate a specific electron with a photon
knocks the electron off its course.
 The Heisenberg uncertainty principle: it is
impossible to determine simultaneously both the
position and velocity of an electron or any other
particle.
The Schrödinger Wave Equation
 In 1926, Austrian physicist Erwin Schrödinger
developed an equation that treated electrons as waves
in an atom.
 Heisenberg & Schrödinger ideas laid the foundation
for modern quantum theory.
 Quantum theory describes mathematically the wave
properties of electrons and other very small particles.
The Schrödinger Wave Equation
 Electrons do not travel around the nucleus in neat
orbits, as Bohr had postulated, but rather in certain
regions called orbitals.
 An orbital: three-dimensional region around the
nucleus - indicates the probable location of an
electron.
 Quantum numbers specify the properties of atomic
orbitals and the properties (location & spin) of
electrons in orbitals.
4 quantum numbers
How should we look at the atom?
 Models of the Atom – An Atomic Rant