Transcript chemistry - Central Lyon

chemistry
Chapter 5
5.1
Models of the Atom
• The scale model shown is a
physical model. However, not
all models are physical. In fact,
several theoretical models of
the atom have been
developed over the last few
hundred years. You will learn
about the currently accepted
model of how electrons
behave in atoms.
5.1
The Development of Atomic Models
•The Development of Atomic Models
– What was inadequate about Rutherford’s atomic
model?
5.1
The Development of Atomic Models
•Rutherford’s atomic model could not explain the
chemical properties of elements.
–Rutherford’s atomic model could not explain why objects
change color when heated.
5.1
The Development of Atomic Models
• The timeline shoes the development of atomic models
from 1803 to 1911.
5.1
The Development of Atomic Models
• The timeline shows the development of atomic models
from 1913 to 1932.
5.1
The Bohr Model
•The Bohr Model
– What was the new proposal in the Bohr model of
the atom?
5.1
The Bohr Model
–Bohr proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
5.1
The Bohr Model
• Each possible electron orbit in Bohr’s model has a fixed
energy.
– The fixed energies an electron can have are called energy
levels.
– A quantum of energy is the amount of energy required to
move an electron from one energy level to another energy
level.
5.1
The Bohr Model
• Like the rungs of the
strange ladder, the energy
levels in an atom are not
equally spaced.
• The higher the energy
level occupied by an
electron, the less energy it
takes to move from that
energy level to the next
higher energy level.
5.1
The Quantum Mechanical Model
•The Quantum Mechanical Model
– What does the quantum mechanical model
determine about the electrons in an atom?
5.1
The Quantum Mechanical Model
–The quantum mechanical model determines the
allowed energies an electron can have and how
likely it is to find the electron in various locations
around the nucleus.
5.1
The Quantum Mechanical Model
• Austrian physicist Erwin Schrödinger (1887–1961) used
new theoretical calculations and results to devise and
solve a mathematical equation describing the behavior
of the electron in a hydrogen atom.
• The modern description of the electrons in atoms, the
quantum mechanical model, comes from the
mathematical solutions to the Schrödinger equation.
5.1
The Quantum Mechanical Model
• The propeller blade has the same probability of being
anywhere in the blurry region, but you cannot tell its
location at any instant. The electron cloud of an atom
can be compared to a spinning airplane propeller.
5.1
The Quantum Mechanical Model
• In the quantum mechanical model, the probability of
finding an electron within a certain volume of space
surrounding the nucleus can be represented as a fuzzy
cloud. The cloud is more dense where the probability
of finding the electron is high.
5.1
Atomic Orbitals
•Atomic Orbitals
– How do sublevels of principal energy levels differ?
5.1
Atomic Orbitals
• An atomic orbital is often thought of as a region of
space in which there is a high probability of finding an
electron.
–Each energy sublevel corresponds to an orbital of a different
shape, which describes where the electron is likely to be
found.
5.1
Atomic Orbitals
• Different atomic orbitals are denoted by letters. The s
orbitals are spherical, and p orbitals are dumbbellshaped.
5.1
Atomic Orbitals
• Four of the five d orbitals have the same shape but
different orientations in space.
5.1
Atomic Orbitals
• The numbers and kinds of atomic orbitals depend on
the energy sublevel.
5.1
Atomic Orbitals
• The number of electrons allowed in each of the first four
energy levels are shown here.
Atomic Orbitals
– Animation 5
– Observe the characteristics of atomic orbitals.
5.2
Electron Arrangement in Atoms
• If this rock were to tumble
over, it would end up at a
lower height. It would have
less energy than before, but
its position would be more
stable. You will learn that
energy and stability play an
important role in determining
how electrons are configured
in an atom.
5.2
Electron Configurations
•Electron Configurations
– What are the three rules for writing the electron
configurations of elements?
5.2
Electron Configurations
• The ways in which electrons are arranged in various
orbitals around the nuclei of atoms are called electron
configurations.
–Three rules—the aufbau principle, the Pauli exclusion
principle, and Hund’s rule—tell you how to find the electron
configurations of atoms.
5.2
Electron Configurations
–Aufbau Principle
• According to the aufbau principle, electrons occupy the
orbitals of lowest energy first. In the aufbau diagram
below, each box represents an atomic orbital.
5.2
Electron Configurations
–Pauli Exclusion Principle
• According to the Pauli exclusion principle, an atomic
orbital may describe at most two electrons. To occupy
the same orbital, two electrons must have opposite
spins; that is, the electron spins must be paired.
5.2
Electron Configurations
–Hund’s Rule
• Hund’s rule states that electrons occupy orbitals of the
same energy in a way that makes the number of
electrons with the same spin direction as large as
possible.
5.2
Electron Configurations
• Orbital Filling Diagram
Electron Configurations
– Simulation 2
– Fill atomic orbitals to build the ground state of
several atoms.
for Conceptual Problem 1.1
Problem Solving 5.9 Solve Problem 9
with the help of an interactive guided
tutorial.
5.2
Exceptional Electron Configurations
•Exceptional Electron Configurations
– Why do actual electron configurations for some
elements differ from those assigned using the
aufbau principle?
5.2
Exceptional Electron Configurations
–Some actual electron configurations differ from
those assigned using the aufbau principle because
half-filled sublevels are not as stable as filled
sublevels, but they are more stable than other
configurations.
5.2
Exceptional Electron Configurations
• Exceptions to the aufbau
principle are due to subtle
electron-electron interactions
in orbitals with very similar
energies.
• Copper has an electron
configuration that is an
exception to the aufbau
principle.
5.3
Physics and the Quantum Mechanical Model
• Neon advertising signs are
formed from glass tubes
bent in various shapes. An
electric current passing
through the gas in each glass
tube makes the gas glow
with its own characteristic
color. You will learn why each
gas glows with a specific
color of light.
5.3
Light
•Light
– How are the wavelength and frequency of light
related?
5.3
Light
– The amplitude of a wave is the wave’s height from zero to the
crest.
– The wavelength, represented by  (the Greek letter lambda),
is the distance between the crests.
5.3
Light
– The frequency, represented by  (the Greek letter nu), is the
number of wave cycles to pass a given point per unit of time.
– The SI unit of cycles per second is called a hertz (Hz).
5.3
Light
–The wavelength and frequency of light are inversely
proportional to each other.
5.3
Light
• The product of the frequency and wavelength always
equals a constant (c), the speed of light.
5.3
Light
• According to the wave model, light consists of
electromagnetic waves.
– Electromagnetic radiation includes radio waves, microwaves,
infrared waves, visible light, ultraviolet waves, X-rays, and
gamma rays.
– All electromagnetic waves travel in a vacuum at a speed of
2.998  108 m/s.
5.3
Light
• Sunlight consists of light with a continuous range of
wavelengths and frequencies.
– When sunlight passes through a prism, the different
frequencies separate into a spectrum of colors.
– In the visible spectrum, red light has the longest wavelength
and the lowest frequency.
5.3
Light
• The electromagnetic spectrum consists of radiation
over a broad band of wavelengths.
Light
– Simulation 3
– Explore the properties of electromagnetic
radiation.
5.1
5.1
5.1
for Sample Problem 5.1
Problem-Solving 5.15 Solve
Problem 15 with the help of an
interactive guided tutorial.
5.3
Atomic Spectra
•Atomic Spectra
– What causes atomic emission spectra?
5.3
Atomic Spectra
–When atoms absorb energy, electrons move into
higher energy levels. These electrons then lose
energy by emitting light when they return to lower
energy levels.
5.3
Atomic Spectra
• A prism separates light into the colors it contains. When
white light passes through a prism, it produces a rainbow of
colors.
5.3
Atomic Spectra
• When light from a helium lamp passes through a prism,
discrete lines are produced.
5.3
Atomic Spectra
• The frequencies of light emitted by an element
separate into discrete lines to give the atomic
emission spectrum of the element.
Mercury
Nitrogen
5.3
An Explanation of Atomic Spectra
•An Explanation of Atomic Spectra
– How are the frequencies of light an atom emits
related to changes of electron energies?
5.3
An Explanation of Atomic Spectra
• In the Bohr model, the lone electron in the
hydrogen atom can have only certain specific
energies.
– When the electron has its lowest possible
energy, the atom is in its ground state.
– Excitation of the electron by absorbing energy
raises the atom from the ground state to an
excited state.
– A quantum of energy in the form of light is
emitted when the electron drops back to a
lower energy level.
5.3
An Explanation of Atomic Spectra
–The light emitted by an electron moving from a
higher to a lower energy level has a frequency
directly proportional to the energy change of the
electron.
5.3
An Explanation of Atomic Spectra
• The three groups of lines in the hydrogen spectrum
correspond to the transition of electrons from higher
energy levels to lower energy levels.
An Explanation of Atomic Spectra
– Animation 6
– Learn about atomic emission spectra and how
neon lights work.
5.3
Quantum Mechanics
•Quantum Mechanics
– How does quantum mechanics differ from
classical mechanics?
5.3
Quantum Mechanics
• In 1905, Albert Einstein successfully explained
experimental data by proposing that light could be
described as quanta of energy.
– The quanta behave as if they were particles.
– Light quanta are called photons.
• In 1924, De Broglie developed an equation that predicts
that all moving objects have wavelike behavior.
5.3
Quantum Mechanics
• Today, the wavelike properties of beams of electrons
are useful in magnifying objects. The electrons in an
electron microscope have much smaller wavelengths
than visible light. This allows a much clearer enlarged
image of a very small object, such as this mite.
Quantum Mechanics
– Simulation 4
– Simulate the photoelectric effect. Observe the
results as a function of radiation frequency and
intensity.
5.3
Quantum Mechanics
–Classical mechanics adequately describes the
motions of bodies much larger than atoms, while
quantum mechanics describes the motions of
subatomic particles and atoms as waves.
5.3
Quantum Mechanics
• The Heisenberg uncertainty principle states that it is
impossible to know exactly both the velocity and the
position of a particle at the same time.
– This limitation is critical in dealing with small particles such as
electrons.
– This limitation does not matter for ordinary-sized object such
as cars or airplanes.
5.3
Quantum Mechanics
» The Heisenberg Uncertainty Principle