Transcript Chapter 13 Notes Electron Models Evolution of Electron Models • The first model of the electron was given by J.J.

Chapter 13 Notes
Electron Models
Evolution of Electron Models
• The first model of the electron was given
by J.J. Thompson—the electron’s
discoverer. His was the “plum pudding”
model.
The Rutherford Model
• With Rutherford’s discovery of the
nucleus of an atom, the atomic model
changed.
The Bohr Model
• Niels Bohr introduced his model, which
answered why electrons do not fall into the
nucleus.
• He introduced the concept of energy levels,
where the electrons orbited similar to the
way the planets orbit the sun.
Bohr Model and Energy Levels
• In the Bohr model, electrons are in energy
levels, or regions where they most probably
are orbiting around the nucleus.
• The analogy is that energy levels are like
the rungs of a ladder—you cannot be
between rungs, just like an electron
cannot be between energy levels.
• A quantum of energy is the amount of
energy it takes to move from one energy
level to the next.
Bohr Model and Energy Levels
• The Bohr model worked well for explaining
the behavior of electrons in hydrogen, but
for all other elements, the equations he
used to predict the electron location did
not work.
Quantum Mechanical Model
• In 1926, Erwin Schrodinger used the new
quantum theory to write and solve
mathematical equations to describe electron
location.
The Quantum Mechanical Model, cont.
• Today’s model comes from the solutions to
Schrodinger’s equations.
• Previous models were based on physical
models of the motion of large objects.
• This model does not predict the path of
electrons, but estimates the probability of
finding an electron in a certain position.
• There is no physical analogy for this model!
Where are the electrons?
• In an atom, principal energy levels (n) can
hold electrons. These principal energy levels
are assigned values in order of increasing
energy (n=1,2,3,4...).
• Within each principal energy level, electrons
occupy energy sublevels. There are as many
sublevels as the number of the energy level
(i.e., level 1 has 1 sublevel, level 2 has 2
sublevels, etc.)
Where are the electrons?
• There are four types of sublevels we will talk
about—s,p,d and f. Inside the sublevel are
atomic orbitals that hold the electrons. Every
atomic orbital can hold two electrons.
• S has one orbital, P has three, D has five and
F has seven. How many electrons can each
one hold?
Orbital Shapes
s orbital = s sublevel
+
px orbital
+
py orbital
= p sublevel
pz orbital
• http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html
Where are the electrons?
• So how many electrons can each energy level
hold?
– Level 1 has an s sublevel=2 e– Level 2 has an s and a p sublevel=8e– Level 3 has an s, p and d sublevel=18e– Level 4 has an s, p, d and f sublevel=32e-
Electron Configuration
Electron Configuration
• In the atom, electrons and the
nucleus interact to make the
most stable arrangement
possible.
• The ways that electrons are
arranged around the nucleus
of an atom is called the
electron configuration.
Aufbau Principal
• Electrons enter orbitals of the lowest energy
first.
Pauli’s Exclusion Principal
• An atomic orbital may describe at most two
electrons.
Hund’s Rule
• When electrons occupy orbitals of equal
energy, one electron enters each orbital until
all orbitals contain one electron with parallel
spins.
1sHe
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
EMR and Quantum Theory
What does a wave look like?
• With your partner, label all the parts
of a wave you can remember.
A Quick Look at Waves—Parts of Waves
A Quick Look at Waves
• The number of waves to pass a point
in a given time is called frequency (n)
and is measured in 1/s or Hertz (Hz).
Electromagnetic Radiation (EMR)
• According to the wave model, visible
light consists of electromagnetic
waves and is just a small fraction of
waves classified as electromagnetic
radiation.
• Other EMR includes radio waves,
microwaves, infrared, ultraviolet,
X-rays, gamma rays and cosmic rays.
Electromagnetic Radiation (EMR)
• All of these waves travel at the same
speed, 3.0x108m/s!
• The waves differ in their frequencies
and wavelengths, and obey the
equation: c = l n
• This is an inverse relationship—as
the frequency increases, the
wavelength decreases.
Practice Problems
• What is the wavelength of an
electromagnetic wave with a frequency of
4.45x1015Hz?
• What is the frequency of a light wave with
a wavelength of 497nm?
• What is the wavelength in nanometers of
an electromagnetic wave with a frequency
of 2.97x1014Hz?
Atomic Emission Spectra
• When sunlight is broken down into
the waves it is made of, it creates a
continuous spectrum
• Scientists used a hydrogen lamp to
produce light, they expected a
continuous spectrum but it wasn’t!
They had an atomic emission
spectrum.
• When we previously found the
electron configuration for elements, it
was for electrons at ground state (the
lowest energy possible).
• As energy is added to atoms, they
absorb the energy by electrons going
from ground state to an excited state,
where electrons are no longer in the
lowest energy orbitals.
• Electrons can then only go back to
ground state by releasing the energy,
usually in the form of light in discreet
packets called photons.
• These packets defied classical
physics, that said electrons would go
back to ground state continuously.
Max Planck
• To understand why this points towards
the concept of energy levels, we need
to know about Max Planck’s discovery:
•E = h n
• Planck’s constant (h=6.6262x10-34Js)
Practice Problems
• How much energy is associated with a
wave with a frequency of 4.4x1014Hz?
• An electromagnetic wave is found to have
1.18x10-19J of energy. What is its
frequency?
• How much energy is associated with a
wave of red light with a wavelength of
697nm?
Putting It Together
• So, if only specific frequencies of
light are emitted when electrons fall
back to ground state from being
excited, then there are only certain
energies that electrons can have.
This explains atomic emission
spectra!
Even Stranger…
• Louis de Broglie predicts yet another
property of electrons—that they have
both a wave nature and a particle
nature.
• Any moving particle can be described
to have a wave nature described by
de Broglies equation:
•
l = h / mv
• Even stranger still is the Heisenberg
Uncertainty Principle.
• It states that you cannot know both a
particles exact position and exact
velocity (the more you know about
one the less you know about the
other).