Pre-IB/Pre-AP CHEMISTRY
Chapter 4 – Arrangement of Electrons in Atoms
Section 1 Objectives
 Be able to define: electromagnetic
radiation, electromagnetic spectrum,
wavelength, amplitude, frequency,
photoelectric effect, quantum(pl. quanta),
photon, ground state, excited state, line
emission spectrum, continuous spectrum,
energy level.
Section 1 Objectives
 Be able to explain the mathematical
relationship between speed, wavelength,
and frequency of a wave.
 Be able to describe what is meant by the
wave-particle duality of light.
 Be able to discuss how the photoelectric
effect and the line emission spectrum of
hydrogen lead to the development of the
atomic model.
Section 1 Objectives
 Be able to describe the Bohr model of the
atom.
Wave
 A wave is a method
of transferring
energy. This
transfer does not
require matter as a
medium.
Wave
 Some waves travel
through matter
(sound, water
waves, etc.).
Wave
 Some waves do not
require matter and
can travel through
empty space (light).
Wave Properties
 Waves can be described by their
wavelength, amplitude, and frequency.
Wavelength
 A crest is the highest point on a wave.
 A trough is the lowest point on a wave.
Crest
Trough
Wavelength
 Wavelength is simply the length of a
wave. It is the distance between two
crests or two troughs.
 Wavelength is measured in m, mm, or nm.
Wavelength
Crest
Trough
Amplitude
 Amplitude is simply the height of a wave. It
is the distance between the crest and
trough of a wave.
 Amplitude is measured in units of distance.
Amplitude
Frequency
 Frequency is the
number of waves
passing a given
point in a given
time.
 Frequency
describes the
energyof a wave.
Frequency
 Frequency describes
the energy of a
wave: the higher
the frequency, the
greater the
energyof that wave.
Frequency
 Frequency is
measured in hertz
or cycles per
secondor vibrations
per second or 1/sec
or sec-1 - they all
mean the same
thing.
Frequency
 As the wavelength
increases, frequency
decreases. This is
called an inverse
relationship.
Wave Properties
 Wavelength and amplitude give waves
their distinctive properties. For example,
the loudness of a sound wave is its
amplitude, the color of visible light is its
wavelength.
Types of waves
 Electromagnetic waves do not require a
medium or matter in order to travel.
Light is an example.
Light
 Light is an electromagnetic wave.
 Visible light is a small part of the
electromagnetic spectrum that humans are
able to see.
Light
 The electromagnetic spectrum consists of
different kinds of light of different
wavelengths.
EM Spectrum
EM Spectrum
EM Spectrum
EM Spectrum
EM Spectrum
EM Spectrum
EM Spectrum
Light Interactions
 White light is light consisting of all
colorsof visible light. These colors are
visible in a rainbow or through a prism.
Velocity
 The velocity of a wave is a product of its
frequency and wavelength.
v= fl
v = velocity
f = frequency
l = wavelength
Velocity
 The velocity of light through a vacuum(c) is
about 3.0 x 108 m/sec. It is slightly slower
through matter.
Photoelectric Effect
 Photoelectric effect refers to the emission
of electrons from a metal when light shines
on the metal.
Photoelectric Effect
 It was found that light of a certain
frequency would cause electrons to be
emitted by a particular metal. Light below
that frequency had no effect.
Emission Spectra
 If an object becomes hot enough it will
begin to emit light.
Emission Spectra
 Max Planck
suggested that hot
objects emit light in
specific amounts
called quanta (sing.
quantum).
Emission Spectra
 Planck showed the relationship between a
quantum of energy and the frequency of
the radiation.
Equantum= hf
Equantum= energy of a quantum in joules
h = Planck’s constant
f = frequency
Wave-Particle Duality
 Einstein later said
that light had a dual
nature – it behaved
as both a particle
and a wave.
Wave-Particle Duality
 Each particle of light, Einstein said,
carries a particular quantum of energy.
Wave-Particle Duality
 Einstein called the “particles” of light
photons which had zero mass and carried a
quantum of energy. The energy is
described as:
Ephoton= hf
Photoelectric Effect
 Einstein explained photoelectric effect by
saying in order for an electron to be ejected
from a metal, the photon striking it must
have enough energy to eject it.
Attraction
 Different metals have stronger attraction for
their electrons than other. Therefore, some
must absorb more energy than others to
emit electrons.
Ground State
 The lowest energy state of an atom is
called its ground state.
Excited State
 When a current is passed through a gas at
low pressure, the atoms become “excited.”
Excited State
 Atoms in an excited state have a higher
potential energy than their ground state.
Excited State
 An “excited” atom will return to its ground
state by releasing energy in the form of
electromagnetic radiation.
Emission Spectra
 Elements will emit radiation of certain
frequencies. This reflects the energy
states of its electrons and is called a
bright-line or emission spectrum.
Emission Spectra
 The emission spectrum of an element is
like its “fingerprint”.
Sodium
Helium
Mercury
Energy Levels
 Studying the
emission spectrum of
hydrogen lead Niels
Bohr to the idea of
energy levels.
Energy Levels
 The spectrum Bohr and others observed
was the result of excited electrons
releasing photons as they returned to
their ground states.
Energy Levels
 The difference in the energy of photons was
reflected in the different frequencies of
light they observed.
Section 2 Objectives
 Be able to define: diffraction, interference,
Heisenberg Uncertainty Principle, Quantum
Theory, quantum numbers, principal
quantum number, angular momentum
quantum number, magnetic quantum
number, spin quantum number.
 Be able to distinguish between the Bohr
model and the quantum model of the atom.
Section 2 Objectives
 Be able to explain how the Heisenberg
Uncertainty Principle and the Schroedinger
Wave Equation led to the idea of atomic
orbitals.
 Be able to list the four quantum numbers
that describe each electron in an atom.
Section 2 Objectives
 Be able to relate the number of sublevels
corresponding to each of an atom’s main
energy levels, the number of orbitals per
sublevel, and the number of orbitals per
main energy level.
Electrons as Waves
 French scientist
Louis De Broglie
demonstrated that
electrons had a dual
nature also.
Electrons as Waves
 De Broglie showed that electrons behaved
as waves confined to the atom. The energy
of those electrons could be found like that
of waves:
E = hf
Electrons as Waves
 Electron beams were
shown to exhibit the
wave properties of
diffraction and
interference.
Heisenberg Uncertainty
 Werner Heisenberg
tried to find the
location and velocity
of electrons in the
atom.
Heisenberg Uncertainty
 Heisenberg found that it is
impossible to
simultaneously determine
the position and velocity of
an electron or any other
particle (The Heisenberg
Uncertainty Principle).
Schrödinger Wave Equation
 Erwin Schrödinger
said that electrons
had a dual
nature(like light) and
treated them as
waves.
Quantum Theory
 Schrödinger’s wave equation and
Heisenberg’s Uncertainty Principle laid
the foundation of modern quantum
theory.
Quantum Theory
 Quantum theory describes mathematically
the wave properties of electrons and
other very small particles.
Quantum Theory
 According to the Bohr model we should be
able to predict the location and velocity
of an electron at any time.
Quantum Theory
 Quantum theory disagrees
with the Bohr model and
says that electrons can be
found in regions of high
probability but cannot
be pinpointed.
Orbitals
 Quantum theory describes
electrons as inhabiting a
three-dimensional region
around the nucleus that
indicates their probable
locations. These regions are
called orbitals.
Orbitals
 Scientists use quantum numbers to
describe orbitals. These numbers describe
the properties of the orbitals and the
electrons which occupy them.
s and p orbitals
d orbitals
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