Transcript Where are the electrons

Where are the electrons ?
• Rutherford found the nucleus to be in the center.
• He determined that the atom was mostly empty
space.
• So, how are the electrons arranged in that space?
• This was the shortcoming, of Rutherford's model
of the atom, the electron position could not be
explained.
What is the electromagnetic
spectrum?
• brainstorm
Electromagnetic radiation
• Is a form of energy that exhibits wavelike
behavior as it travels through space.
• Electromagnetic spectrum – all the forms of
electromagnetic radiation.
Electromagnetic spectrum
Electromagnetic spectrum
• Each line on the spectrum represents a
certain wave frequency of radiation.
• Each wave frequency is associated with a
certain amount of energy.
Electromagnetic spectrum

The elecromagnetic spectrum includes
all forms of radiation, one of which is visible
light -- the radiation to which our eyes are
sensitive.

Gamma-rays, X-rays, ultraviolet,
infrared and radio waves are also forms of
radiation. We divide the spectrum up
according to the wavelength of the radiation.
Electromagnetic spectrum
Electromagnetic spectrum
A wave is described by
• Wavelength - the distance between
corresponding points on adjacent waves
• Frequency- the number of waves that pass a
given point in a specific time, usually one
second. Unit is the hertz, (Hz).
• http://www.colorado.edu/physics/2000/wav
es_particles/
Wave description
• Mathematically – c= speed = 3 x108 m/s
– Speed of light= frequency x wavelength
– Since c is constant
– Frequency is inversely proportional to
wavelength
– All three, speed, wavelength, frequency are
related
example
• What is the frequency of a wave that has a
wavelength of 200 nm?
• 3 x 108/ 200 nm =
1900’s
• Wave model of light was accepted
Early 1900’s
Photoelectric experiments
• Experiment - Planck
• Photoelectric effect – refers to the emission of
electrons from a metal when light shines on the
metal.
• Shined light on a metal varying the frequency of
the light. Below a certain frequency the electrons
were not emitted.
• When the frequency was high enough there was
enough energy to knock electrons loose from the
metal.
Quantum
• Planck suggested that an object emits
energy in little packets of energy called a
quantum.
• Quantum – minimum quantity of energy
that can be lost or gained by an electron.
• Energy that is given off is related to wave
frequency.
Go to
• http://www.colorado.edu/physics/2000/quan
tumzone/lines2.html
• http://www.colorado.edu/physics/2000/wav
es_particles/
Planck in other words.
• Proposed that there is a fundamental
restriction on the amount of energy that an
object emits or absorbs; each piece of
energy is a quantum.
• E= hv
• h = 6.6262 x 10-34 Js
• v= the frequency of the wave
Dual wave-particle nature of light
• Einstein expanded on Planck’s theory and
said that electromagnetic radiation has a
dual wave-particle nature.
• While light exhibits wave like properties, it
can also be thought of as a stream of
particles.
• Each particle carries a quantum of energy –
these particles are called photons.
Vocab
• Photon – particle of electromagnetic
radiation having zero mass and carrying a
quantum of energy.
Ephoton= hv
• Quantum – the minimum quantity of
energy that can be gained or lost by an
electron.
Photoelectric effect explained
• Electromagnetic radiation is only absorbed
in whole number of photons.
• In order for an electron to be bumped off, it
must be struck by a photon of a certain
minimal amount of energy.
• The minimal energy corresponds to a
minimal frequency
continued
• Different metals require energy at different
frequencies to exhibit a photoelectric effect.
• Each metal has a certain required minimal
level of energy required for the electrons to
be knocked loose.
Ground state
• The lowest possible energy level
• Close to the nucleus
• The lowest energy state of an atom
Excited state
• When an atom absorbs energy and the
electrons move to a higher energy level.
• An atom has a higher potential energy than
its ground state.
Like a ladder
• Energy levels are like
the rungs on a ladder,
you can not step
between the rungs.
Electrons must jump
from level to level,
they can not reside
between the levels
•
electrons
• When an electron gains enough energy it
jumps up an energy level ( rung on a
ladder), and becomes excited.
• It then immediately returns to its ground
state.
• The energy is released as a particular
wavelength that corresponds to a particular
color. A photon of radiation.
Line emission spectra
• The spectra given off by the electrons of a
certain element.
• Each spectra is unique to the element or
compound.
• Line spectra are used to identify elements
and/or compounds
Line emission spectra
• The fact that specific frequencies are
emitted indicates that the energy differences
between an atoms energy states are fixed
• Each line on the spectra represents a photon
of energy.
Bohr model of atom
• The electron circles the nucleus in an
orbital.
• When the electron gains enough energy ( a
certain photon) it jumps to a higher level
orbital.
• When it returns to ground state it emits the
energy (photon). The frequency of the
photon is seen in the spectra.
Bohr successfully calculated the
hydrogen line spectra
• Bohrs model of the hydrogen atom accounted
mathematically for the energy of each of the
transitions of the Lyman, Balmer and Paschen
spectral series.
• Bottom line – Bohrs model of the H atom, with
the line spectra, could be mathematically
supported.
• summary and beyond