Download AP Chapter 5

AP Chapter 5
Structure of the Atom
Review Quiz Chapter 5
• Net Ionic Equations
John Dalton (1766 – 1844)
• Proposed the first
scientifically supported
atomic theory.
Dalton’s Model of the Atom
• Dalton's model was that
the atoms were tiny,
indivisible, indestructible
particles and that each
one had a certain mass,
size, and chemical
behavior that was
determined by what kind
of element they were.
• What made us change
Dalton’s Atomic Model?
J. J. Thomson (1856 – 1940)
• Credited for the
discovery of the
electron.
• Invented the mass
spectrometer which
led to his discovery
of isotopes.
Scientific Inquiry
• Scientists use experimental results to test
scientific models, such as the model of the
atom.
• When experimental results are not
consistent with the predictions of a
scientific model, the model must be
revised or replaced.
Lord Ernest Rutherford (1871 – 1937)
• Discovered the
nucleus of the atom.
Rutherford’s Gold Foil Experiment
Rutherford’s Gold Foil Experiment
Rutherford’s Atomic Model
The “Planetary Model” of the Atom
• The nucleus is very small, dense, and positively
charged.
• Electrons surround the nucleus.
• Most of the atom is empty space
Subatomic Particles
PARTICLE
SYMBOL CHARGE
MASS
(amu)
LOCATION
electron
e-
-1
0
orbit nucleus
proton
p+
+1
1
inside nucleus
neutron
n0
0
1
inside nucleus
Niels Bohr
• In 1913 Bohr published a
theory about the structure of
the atom based on an earlier
theory of Rutherford's.
• Bohr expanded upon this
theory by proposing that
electrons travel only in certain
successively larger orbits.
Bohr Model of the Atom
The Bohr Atom
• Electrons orbit the
nucleus in orbits that
represent specific
quantities of energy.
• The energies of the
electrons in the atom are
quantized.
• Only certain electron
orbits (energy levels) are
allowed.
Electromagnetic Waves
• Properties of waves include speed, frequency,
wavelength and energy
• All electromagnetic waves including light travel at a
speed of 3 x 108 m/s.
• However the frequency, wavelength and energy of
the waves vary.
Wavelength ()
Measured in
units of length:
m, nm, Aº
Frequency ()
Measured in cycles/second = hertz (Hz)
Visible Light
Electromagnetic Radiation
• For all waves
• = c
c = the speed of light = 3.00 x 10
8
m/s
A photon of red light has a wavelength of
665 nm. What is the frequency of this light?
•
•
•
•
•
665 nm = 665 x 10-9 m
c=•
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = 4.511278 x 1014/s or Hz
 = 4.51 x 1014/s or Hz
An x-ray has a frequency of 7.25 x 1020 Hz.
What is the wavelength?
•
•
•
•
•
7.25 x 1020 Hz = 7.25 x 1020/s
c=•
 = c ÷  = 3.00 x 108 m/s ÷ 7.25 x 1020/s
 = 4.137931 x 10-13 m
 = 4.14 x 10-13 m
Energy of Electromagnetic
Radiation
• For all waves:
h = 6.63 x 10
E = h•
-34
J • s or J/Hz
A photon of red light has a wavelength of
665 nm. What is the energy of this light?
A photon of red light has a wavelength of
665 nm. What is the energy of this light?
•
•
•
•
•
665 nm = 665 x 10-9 m
c=•
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = 4.511278 x 1014/s or Hz
 = 4.51 x 1014/s or Hz
A photon of red light has a wavelength of
665 nm. What is the energy of this light?
•
•
•
•
•
•
•
665 nm = 665 x 10-9 m
c=•
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = 4.511278 x 1014/s or Hz
 = 4.51 x 1014/s or Hz
E = h •  = (6.63 x 10-34 J • s)(4.51 x 1014/s)
E = 2.99013 x 10-19 J
A photon of red light has a wavelength of
665 nm. What is the energy of this light?
•
•
•
•
•
•
•
665 nm = 665 x 10-9 m
c=•
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = 4.511278 x 1014/s or Hz
 = 4.51 x 1014/s or Hz
E = h •  = (6.63 x 10-34 J • s)(4.51 x 1014/s)
E = 2.99013 x 10-19 J = 2.99 x 10-19 J
An x-ray has a frequency of 7.25 x 1020 Hz.
What is it’s energy?
An x-ray has a frequency of 7.25 x 1020 Hz.
What is it’s energy?
• E = h •  = (6.63 x 10-34 J/Hz)(7.25 x 1020Hz)
• E = 4.80675 x 10-13 J
• E = 4.81 x 10-13 J
wavelength, frequency and energy
Red Light
X-ray
•  = 665 x 10-9 m
•  = 4.51 x 1014 Hz
• E = 2.99 x 10-19 J
•  = 4.14 x 10-13 m
•  = 7.25 x 1020 Hz
• E = 4.81 x 10-13 J
Wavelength, frequency and energy
• Wavelength and frequency have an indirect
relationship.
• Energy and frequency have a direct
relationship.
• Electromagnetic radiation of short wavelength
will have high frequency and high energy.
• Electromagnetic radiation of long wavelength
will have low frequency and low energy.
Niels Bohr
• Bohr also described the way
atoms emit radiation by
suggesting that when an
electron jumps from an outer
orbit to an inner one, that it
emits light.
Bohr Model
Slide 9
• Electrons move around the
nucleus in orbits of definite
energies.
• The energy of the orbit is
related to its distance from
the nucleus. The lowest
energy is found in the orbit
closest to the nucleus.
• Radiation is absorbed or
emitted when an electron
moves from one orbit to
Fig. 10-6, p. 269
another.
Bohr Model of the Atom
• The Bohr atom
The Bohr Atom
• Electrons orbit the
nucleus in orbits
that represent
specific quantities
of energy.
• Electrons closer to
the nucleus have
less energy.
Ground State
• The lowest energy
state of an atom.
Excited State
• Any energy state of
an atom that is of
higher in energy
than the ground
state.
Energy Absorbed
Absorption (Dark – Line) Spectra
Energy Emitted
Electron jumps
to a lower orbit
Emission (Bright – Line) Spectra
Emission Spectra
The lines present in an emission
spectrum are the lines missing in
an absorption spectrum.
The Heisenberg Uncertainty Principle
The Uncertainty principle
• Heisenberg determined that it was
impossible to experimentally determine both
the position and the speed of the electron at
the same time.
• This became known as the Heisenberg
Uncertainty Principle.
• It simply means that the electron is so small
and moving so fast, that the simple act of
trying to measure its speed or position would
change either quantity.
Problems with the Bohr Model
• It violates the
Heisenberg
Uncertainty Principle
because it considers
electrons to have
known orbits.
• It makes poor
predictions regarding
the spectra of atoms
larger than hydrogen.
Schrodinger
• The Austrian scientist,
Erwin Schrödinger,
pursued the of the
electron having wave
properties and it
seemed to him that
the electron might be
like a standing wave
around the nucleus.
Schrodinger’s Model
• This idea agreed very well with Bohr's idea
of quantized energy levels: only certain
energies and therefore, wavelengths would
be allowed in the atom.
• This explained why only certain colors
(wavelengths) were seen in the spectrum of
the hydrogen atom.
Schrodinger’s Model
• Schrodinger set out to make a mathematical
model that assumed the electron was a
standing wave around the nucleus.
• His solutions to that model agreed not only
with the experimental evidence for hydrogen
(as Bohr’s did too), but gave excellent
results for all atoms when compared to their
actual spectrum.
Quantum Mechanics: The
Structure of Atoms
Schrodinger’s Model
• This is our modern model of the atom and is
known as the Quantum Mechanical Model.
• Calculations involving the QM model of the
atom can be approximated using computers.
• The solutions to these calculations is the
basis for modern software that calculates
the structure and reactivity of molecules.
The Quantum Mechanical Model
The Quantum Mechanical Model
• The quantum
mechanical model is
based on quantum
theory, which says
matter also has
properties associated
with waves.
The Quantum Mechanical Model
• According to quantum
theory, it’s impossible
to know the exact
position and
momentum of an
electron at the same
time. This is known as
the Uncertainty
Principle.
The Quantum Mechanical Model
• The quantum
mechanical model of
the atom uses
complex shapes
of orbitals (sometimes
called electron
clouds), volumes of
space where an
electron is likely to be
found.
The Quantum Mechanical Model
• Therefore the
quantum mechanical
model is based on
probability rather than
certainty.
Absorption and Emission of Energy
by Molecules
• A photon is a particle representing a
specific amount of electromagnetic
energy.
• When a photon is absorbed by a molecule,
the energy of the molecule is increased.
• When a photon is released by a molecule,
the energy of the molecule is decreased.
Absorption and Emission of Energy
by Molecules
• Different types of molecular motion lead to
absorption or emission of photons in
different spectral regions.
Absorption and Emission of Energy
by Molecules
• Ultraviolet and visible radiation is
associated with electron transitions
between energy levels and so can be used
to study electronic structure.
• This is the same as atomic spectroscopy.
• Molecules however have bonds which
allow for different types of absorption and
emission.
Absorption and Emission of Energy
by Molecules
• Infrared radiation is associated with molecular
vibrations and so can be used to detect the type
of bonds present within the molecule.
Absorption and Emission of Energy
by Molecules
• Infrared radiation is associated with molecular
vibrations and so can be used to detect the type
of bonds present within the molecule.
Absorption and Emission of Energy
by Molecules
• The microwave
region of the
spectrum is used
for rotational
spectroscopy.
World of Chemistry:
Molecular Fingerprints
• ..\..\..\Videos\World of Chemistry\Molecular
Fingerprints.mpg
Energy Levels – Sublevels - Orbitals
• Electrons in an atom are within atomic
orbitals which are within sublevels which
are within energy levels.
Electron Configuration
7
4d
7 electrons are in the d sublevel in the 4th energy level
SUBLEVEL
s
NUMBER OF
ORBITALS
1
MAX. # OF
ELECTRONS
2
p
3
6
d
5
10
f
7
14
Arrow Diagram
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
Write the
electron
configuration
for lead
(Z = 82).
The periodic
table and
electron
configuration.
Periodic Table and Electron Configuration
s
p
1
2
d (period-1)
3
4
5
6
7
6
7
6
7
f (period-2)
© 1998 by Harcourt Brace & Company
C. Periodic Patterns
• Example - Germanium
1
2
3
4
5
6
7
[Ar]
2
10
4s 3d
2
4p
Write the abbreviated electron configuration
for lead (Z = 82) using the periodic table.
O
6
7
© 1998 by Harcourt Brace & Company
A. General Rules
• Pauli Exclusion Principle
– Each orbital can hold TWO electrons with
opposite spins.
A. General Rules
• Hund’s Rule
– Within a sublevel, place one e- per orbital
before pairing them.
– All electrons in singly filled orbitals have the
same direction of spin.
Incorrect
Correct
Orbital Filling Diagrams
B. Notation
• Orbital Diagram
O
8e-
1s
2s
• Electron Configuration
2
2
4
1s 2s 2p
2p
Magnetism
• Magnetism can be a complicated concept.
• We will only deal with magnetism as it is
predicted by electron spin which is an
oversimplified way of dealing with this
subject.
Electron Spin & Magnetism
• Opposite spins
produce opposite
magnetic fields.
Electron Spin & Magnetism
• Most materials with one or
more unpaired electrons are
at least slightly magnetic.
These substances are said
to be paramagnetic.
• The overall magnetic field
strength of atoms with all
paired electrons is zero.
These substances are said
to be diamagnetic.
Molecular Orbital Diagrams
• One of the reasons that magnetism can be
complicated is that often times the bonding
within atoms creates situations where
multiple atoms bond together creating
molecular orbitals which greatly
complicates the ability to predict the
pairing of electrons.
Molecular Orbitals
• A molecular orbital is an orbital within a
molecule whereas an atomic orbital is an
orbital within an individual atom.
• Both types of orbitals can hold no more
than a pair of electrons.
• We can draw a molecular orbital diagram
to help explain the magnetic character of
some simple diatomic molecules.
Molecular Orbitals Diagrams
See the blank
MO diagram in
your notebook.
Molecular Orbitals Diagram for H2
Molecular Orbitals Diagram for C2
MO Diagrams Can Be Complicated
.
1. Write the molecular
orbital diagrams for
O2 and F2.
2. Classify each
substance as
paramagnetic or
diamagnetic.
3. Which substance
will be attracted to a
magnetic field?
Explain.
Paramagnetism of O2
..\..\..\Videos\Paramagnetism of Liquid Oxygen.wmv
Photoelectric Effect
• The photoelectric effect is the emission
of electrons from substances that are
exposed to light.
• Ionization energy is the amount of energy
it takes to remove an electron from an
atom.
Photoelectric Effect
• More energy is required to remove
successive electrons from atoms.
• This is due to Coulomb’s Law.
Coulomb’s Law
• Coulomb’s Law quantifies the general rule
of electrostatics that opposites charges
attract and like charges repel.
• The electrostatic force between two
charged bodies is proportional to the
product of the amount of charge on the
bodies divided by the square of the
distance between them.
Coulomb’s Law
F= force of attraction or repulsion
k = constant
q1 and q2 = the charges of the two bodies
r = radius between the charges
Coulomb’s Law
• If the bodies are oppositely charged, one
positive and one negative, they are
attracted toward one another; if the bodies
are similarly charged, both positive or both
negative, the force between them is
repulsive.
Coulomb’s Law
Coulomb's law helps describe the forces
that bind electrons to an atomic nucleus.
• Based on Coulomb’s
Law, the force between
two charged particles is
proportional to the
magnitude of each of
the two charges and
inversely proportional
to the square of the
distance (radius)
between them.
Photoelectric Effect
• Light consists of photons of a certain
energy (E = h • ).
• In the photoelectric effect, light ejects
electrons from a material. This requires
the photon to have sufficient energy to
eject the electron.
• This is known as Photoelectron
Spectroscopy (PES)
Photoelectron Spectroscopy
(PES)
• Photoelectron Spectroscopy (PES) determines
the energy needed to eject electrons from a
substance.
Photoelectron Spectroscopy
(PES)
• Photoelectron Spectroscopy (PES) supports the
shell, subshell model of the atom.
• PES data for multielectron atoms show that
certain electrons have ionization energies that
are relatively close to one another. We therefore
group these electrons into shells and subshells
as a result.
Photoelectron Spectroscopy
(PES)
• Photoelectron Spectroscopy (PES) data for
atoms also helps us to confirm the number of
electrons within particular shells and subshells.
• The intensity of the photoelectron signal at a
given energy is a measure of the number of
electrons in that shell or subshell.
All the previous PES slides all
come from this website
• http://www.chem.arizona.edu/chemt/Flash/
photoelectron.html
Spectrophotometer Lab:
Preparing and Diluting Solutions
Spectrophotometer Lab: Preparing and Diluting Solutions