Transcript 5.1

Electronic Configurations of 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
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
dumbbell-shaped.
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.
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.
5.2 Section Quiz.
• 5.2.
5.2 Section Quiz.
– 1.
Identify the element that corresponds
to the following electron configuration:
1s22s22p5.
•
•
•
•
F
Cl
Ne
O
5.2 Section Quiz.
– 2.
Write the electron configuration for the
atom N.
•
•
•
•
1s22s22p5
1s22s22p3
1s22s1p2
1s22s22p1
5.2 Section Quiz.
– 3.
The electron configurations for some
elements differ from those predicted by the
aufbau principle because the
• the lowest energy level is completely filled.
• none of the energy levels are completely filled.
• half-filled sublevels are less stable than filled
energy levels.
• half-filled sublevels are more stable than some
`
other arrangements.
chemistry
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
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.
The Planck Equation
• The energy of electromagnetic radiation is
directly related to the frequency
•
E = h
• h is the Planck constant = 6.626 x 10-34 J.s
  is the frequency in Hz
END OF SHOW