Transcript Electron Notes

Chapter 5
Electrons in Atoms
Wave Nature of Light
• Electromagnetic radiation which is a
form of energy that exhibits wavelike
behavior as it travels through space.
• Examples: light, radio waves, x-rays, etc
Parts of a Wave
wavelength
crest
amplitude
origin
amplitude
wavelength
trough
Wavelength
• Waves have a repetitive nature.
• Wavelength- ( lambda)
– shortest distance between corresponding points
on adjacent waves.
– Measured in units like meters, centimeters, or
nanometers depending on the size.
– 1 x 10-9 meters = 1 nanometer
Frequency
• # of waves that pass a given point per
second.
• Units are waves/sec, cycles/sec or
Hertz (Hz)
• Abbreviated n the Greek letter nu or
by an f
c = f
Frequency and wavelength
• Are inversely
related
• As one goes up
the other goes
down.
High frequency, Short Wavelength
Low frequency, Long Wavelength
Wave Formula
• All electromagnetic waves, including
visible light, travel at the speed of 3.00 x
10 8 m/s in a vacuum.
• Speed of light = c = 3.00 x 108 m/s
c=f
Speed of light = (wavelength) x (frequency)
Example Problem
• What is the wavelength of a microwave having
a frequency of 3.44 x 109 Hz?
Formula: c=f
 =?
f = 3.44 x 109 Hz
c = 3.00 x 108 m/s
3.00 x 108 m/s =  (3.44 x 109 s-1)
3.00E8 / 3.44E9 = 8.72 x 10-2 m
Practice
• What is the frequency of green light, which
has a wavelength of 5.90 x 10-7m?
• A popular radio station broadcast with a
frequency of 94.7MHz, what is the wavelength
of the broadcast? ( frequency needs to be is
Hz)
• Different frequencies produce different types
of waves.
• The entire range of frequencies is called the
electromagnetic spectrum
• We are only able to see with our eyes a small
portion of the spectrum = visible light
• ROY G BIV
• Different colors mean different
frequencies/wavelengths
Energy & The Spectrum
• The energy of a wave increases with
increasing frequency
• High Frequency = High Energy
• Low Frequency = Low Energy
• Blue light has more energy than Red light
Low
energy
Radio Micro Infrared
Ultrawaves waves .
violet
Low
Frequency
Long
Wavelength
Visible Light
High
energy
XGamma
Rays Rays
High
Frequency
Short
Wavelength
Quanta
• Max Planck suggested the idea of quanta
or packets of energy.
• Quanta is the minimum amount of energy
that can be lost or gained by an atom.
• Energy is quantized = it comes in packets
(like stairs or pennies only whole numbers)
Planck’s Constant
• h = 6.626 x 10-34 J.s (Joule seconds)
Energy = (Planck’s constant)(frequency)
E=hf
Example: What is the energy in Joules of a photon from the violet
portion of the rainbow if it has a frequency of 7.23 x 1014 Hz?
E=?
h = 6.626 x 10-34 Js
f = 7.23 x 1014 Hz (or s-1)
E = (6.626 x 10-34 Js)(7.23 x 1014 s-1)
E = 4.79 x 10-19 J
Photoelectric Effect
• In the 1900s, scientist studied interactions of light
and matter.
• One experiment involved the photoelectric
effect, which refers to the emission of electrons
from a metal when light shines on the metal.
• This involved the frequency of the light. It was
found that light was a form of energy that could
knock an electron loose from a metal.
Photon
• Light waves can also be thought of as streams
of particle.
• Einstein called these particles photons (He
won a Nobel Prize for this)
• A photon is a particle of electromagnetic
radiation having zero mass and carrying a
quantum energy.
Bohr’s Model
• Why don’t electrons fall into
nucleus?
• Bohr suggested that they move like
planets around sun.
• Certain amounts of energy separate
one level from another.
• Nucleus is
found inside a
blurry
“electron
cloud”
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s Model
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
• Further away
from nucleus
means more
energy.
• There is no “in
between”
energy
• Energy Levels
Bohr Model of the Atom
• Ground state- the lowest energy state of an
atom.
• Excited state – state in which an atom has a
higher potential energy than its ground state.
• Energy is quantized. It comes in chunks.
• quanta - amount of energy needed to move from
one energy level to another.
• Since energy of an atom is never “in between”
there must be a quantum leap in energy.
Bohr Energy Levels
•
•
•
•
K = 2 electrons – 1st
L = 8 electrons – 2nd
M = 18 electrons – 3rd
N = 32 electrons – 4th
Heisenberg Uncertainty Principle
• This is the theory that states that it is
impossible to determine simultaneously
both the position and velocity of an
electron or any other particle.
Quantum Theory
• Schrodinger derived an equation that
described energy & position of
electrons in atom
• Schrodinger along with other
scientists laid the foundation for the
modern quantum theory, which
describes mathematically the wave
properties of electrons and other very
small particles.