J. J. Thomson’s “plum pudding” model,
in which electrons are surrounded by a
soup of positive charge to balance the
electron’s negative charge, like
negatively-charged “plums”
surrounded by positively charged
“pudding”.
The Rutherford model of the
atom was an improvement
over previous models, but it
was incomplete.
It did not explain how the atom’s negatively
charged electrons are distributed in the space
surrounding its positively charged nucleus.
In the early twentieth century, a
new atomic model evolved as
a result of investigations into
the absorption and emission of
light by matter.
The studies revealed a relationship between light
and an atom’s electrons.
Before 1900, scientists thought light behaved
solely as a wave. This belief changed when it
was later discovered that light also has particlelike characteristics.
A quick review of these wavelike properties
follows.
Properties of Light
•Visible light is a kind of electromagnetic radiation
(form of energy that exhibits wavelike behavior as
it travels through space). Other examples include
X-rays, ultraviolet and infrared light, microwaves,
and radio waves.
•The electromagnetic spectrum consists of all
types of electromagnetic radiation.
Electromagnetic Spectrum
•All forms of electromagnetic radiation move at a
speed of 3.0 x 108 m/s through air.
•Wavelength – λ (m, cm, or nm) and frequency – ν
(wave/second) are measureable properties of
wave motion. One wave/second is called a hertz
(Hz).
•The relationship between wavelength (λ) and
frequency (ν) is c = λν where c = speed of light.
Wavelength and Frequency
In the early 1900s, scientists conducted two
experiments involving interactions of light and
matter that could not be explained by the wave
theory of light.
One experiment involved a phenomenon
known as the photoelectric effect.
Photoelectric Effect
•The photoelectric effect refers to the emission of
electrons from a metal when light shines on the
metal.
Light may cause electrons to be emitted from
an electrode in a photocell. Long wavelength
light does not have enough energy to cause
the electron to escape, regardless of its
intensity. When light of a shorter wavelength
(higher energy) light strikes the electrode,
electrons are released. The amount of current
produced depends on the intensity of the light
and the energy of the escaping electrons
depends on the wavelength of the light.
•Show video clip.
The wave theory of light predicted that light of
any frequency could supply enough energy to
eject an electron.
Scientists couldn’t explain why the light had to be
of a certain frequency in order for the
photoelectric effect to occur.
The German physicist Max Planck proposed
an explanation for the photoelectric effect.
He proposed that a hot object does not emit
electromagnetic radiation continuously, as
would be expected if the energy emitted were
in the form of waves.
•Max Planck proposed that objects
emit energy in small, specific
amounts called quanta (1900).
•Quantum is the minimum quantity
of energy that can be lost or gained
by an atom.
•The relationship between a quantum of energy and
the frequency of radiation is illustrated by the
following equation:
E = hν
•E is the energy, in joules, of a quantum of radiation, ν is the frequency
in s-1 of the radiation emitted, and h is a physical constant now known
as Planck’s constant.
•Einstein proposed that
electromagnetic radiation has dual
wave-particle nature (1905).
These particles are called photons.
•A photon is a particle of electromagnetic radiation
having zero mass and carrying a quantum of energy.
•In order for an electron to be ejected from a metal
surface, the electron must be struck by a single
photon possessing the minimum energy and
frequency to knock it loose. •Show video clip.
•The energy of a particular photon depends on the frequency of
the radiation.
Ephoton = hν
Hydrogen Atom Line-Emission
Spectrum
•The ground state is the lowest energy level of an
atom. When it has higher potential energy an atom
is in its excited state.
•When an excited atom returns to its ground state it
gives off energy in the form of colored light.
(Example: Neon lights.)
•When doing experiments with hydrogen gas, it was
found that hydrogen atoms emit only specific
frequencies of light.
•The fact that hydrogen atoms emit only specific
frequencies of light indicated that the energy
differences between the atom’s energy states were
fixed.
•This suggested that the electron of a hydrogen
atom exists only in very specific energy states (led to
quantum theory).
Hydrogen’s Line Emission Spectrum
Bohr Model of the Hydrogen Atom
•Niels Bohr proposed a model of the
hydrogen atom that showed that the
electron can circle the nucleus only
in allowed paths (orbits) (1913).
Electrons as Waves
•In the early 1900s, it was found through
experimentation that light could behave as both a
wave and a particle (dual-wave particle of nature).
•In 1924 Louis de Broglie experimented to see if
electrons have a dual-wave particle of nature as
well.
•He found that electrons did.
The Heisenberg Uncertainty Principle
•The idea of electrons having a dual wave-particle
nature troubled scientists. If electrons are both
particles and waves, then where are they in the
atom?
•Heisenberg’s idea involved the detection of
electrons. Electrons are detected by their
interaction with photons. Because photons have
about the same energy as electrons, any attempt to
locate a specific electron with a photon knocks the
electron off its course.
•As a result, there is always a basic uncertainty in
trying to locate an electron.
•The Heisenberg Uncertainty principle states that
is it 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 in atoms as
waves.
•Schrödinger’s wave equation laid the foundation for
the quantum theory (1926).
•The quantum theory describes mathematically the
wave properties of electrons and other very small
particles.
•The theory suggested that electrons do not travel in
neat orbits, as Bohr’s model showed, but in regions
called orbitals.
•An orbital is a three-dimensional region around the
nucleus that indicates the probable location of an
electron.
Atomic Orbitals and Quantum
Numbers
•In order to describe orbitals, scientists use quantum
numbers (specify the properties of atomic orbitals
and the properties of electrons in orbitals).
•The first three quantum numbers indicate the main
energy level, the shape, and orientation of the
orbital.
•The fourth, the spin quantum number, describes the
state of the electron.
•Basically, quantum numbers were devised as a way
to describe where individual electrons are located in
an atom.
Principal Quantum Number
•The principal quantum number, symbolized by n,
indicates the main energy level occupied by the
electron.
•Values of n are positive integers only (e.g., 1, 2, 3…).
•As n increases, the electron’s energy and its average
distance from the nucleus increases.
Principal Quantum Number
n=6
n=5
n=4
Energy
n=3
n=2
n=1
Angular Momentum Quantum
Number
•The angular momentum quantum number,
symbolized by l, indicates the shape of the orbital.
•The values of l allowed are zero and all positive
integers less than or equal to n-1.
•Sublevels in the atoms of the known elements are
s-p-d-f.
Shapes of Orbitals
Magnetic Quantum Number
•The magnetic quantum number, symbolized by m,
indicates the orientation of an orbital around the
nucleus.
•Orbitals contain one or two electrons, never more.
Spin Quantum Number
•The spin quantum number, has only two possible
values (+1/2, -1/2) which indicate the two
fundamental spin states of an electron in an
orbital.
•Electrons in the same orbital must have opposite
spins.
•Possible spins are clockwise or counterclockwise.