Download The Development of a New Atomic Model:

The Development of a New Atomic Model:
Properties of Light:
- Light behaves as a wave
- Light also has particle-like characteristics.
********************The Wave Description of Light************
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Electromagnetic radiation: a form of energy that exhibits wavelike behavior as it
travels through space.
o Ex: visible light, X-rays, ultraviolet, infrared light, microwaves, radio waves.
o Explain all forms of the electromagnetic spectrum and there applications
(Figure 4-1 Pg. 92)
o Light moves at a constant speed of 3.0 * 10 8 m/s through air.
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Light behaves like a wave
o Wavelength (λ)= the distance between corresponding points on adjacent wave.
 Units of meter, centimeter, or nanometers
o Frequency (v)= number of wave that pass a given point in a specific time (one
second).
 Units of wave/second or Hertz (Hz)
o Trough = bottom of the peak
o Amplitude =
Compare light waves to the waves of water
Frequency and wavelength are mathematically related
c = λv (product is a constant)
o c, is the speed of light
o λ is inversely proportional to v.
o Wavelength decreases as its frequency increases, and vice versa.
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WORK PROBLEMS ON THE BOARD!!!!!!
The Photoelectric Effect:
- In the early 1900’s, scientist conducted two experiments involving interactions of light
and matter that could not by explained by the wave theory of light.
o Photoelectric Effect = the emission of electrons from a metal when light shines
on the metal (Figure 4-3)
 Involved the frequency of light striking the metal.
 No electrons were emitted if the light’s frequency was below a certain
minimum.
 Light, a form of energy, can knock loose an electron from a metal.
 Wave theory of light predicted that light of any frequency could supply
enough energy to eject an electron.
 Scientists couldn’t explain why the light had to be of a minimum
frequency in order for the photoelectric effect to occur.
*****************The Particle Description of Light************
The Particle Description of Light:
- Max Planck- studied emission of light by hot objects
- Proposed that a hot object does not emit electromagnetic energy continuously.
- The object emits energy in small, specific amounts called quanta.
o Quantum = the minimum quantity of energy that can be lost or gained by an
atom.
There is a relationship between a quantum of energy and the frequency of radiation:
E = hv
E is energy of a quantum of radiation, in joules
V is the frequency of the radiation emitted
h is Planck’s constant (6.626 * 10 –34 Joules * seconds)
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1905, Albert Einstein expanded on Plack’s theory by introducing the radical idea that
electromagnetic radiation has a dual wave-particle nature.
o Light exhibits wavelike particles
o Light is a stream of particles, where each particle carries a quantum of energy
o Photons = a particle of electromagnetic radiation having zero mass and carrying a
quantum of energy. (energy depends on the frequency)
Ephoton = hv
WORK PROBLEMS ON BOARD!!!!!!
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Electromagnetic radiation is absorbed by matter only in whole numbers of photons.
In order for an electron to be ejected from a metal surface, the electron must be struck by
a single photon possessing at least the minimum energy required to knock the electron
loose.
The minimum of energy corresponds to the minimum of frequency.
o Photon’s frequency is below the minimum, then the electron remains bound to the
metal surface
o Different metals require different minimum frequencies to exhibit the
photoelectric effect.
Hydrogen-Atom Line-Emission Spectrum:
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When current is passed through a gas at low pressure, the potential energy of some of the
gas atoms increases.
o Ground state = lowest energy state of an atom
o Excited state = a state in which an atom has higher potential energy than it has in
its ground state.
When an excited electron returns to its ground state, if gives off the energy it gained in
the form of electromagnetic radiation. (Visible light  what we can see)
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o Explain the production of colored light in neon signs (Figure 4-40
 When a narrow beam of the emitted light was shined through a prism, it
was separated into a series of specific frequencies (specific wavelengths)
of visible light.
 Bands of light = hydrogen’s line-emission spectrum
 Expected to see a continuous range of frequencies = continuous spectrum
 Hydrogen atoms only gave off specific frequencies of light
(Photoelectric effect)
WHY?  Later discussion of Quantum Theory
o When an excited atom falls back from an excited state to its ground state, or a
lower-energy excited state, it emits a photon of radiation only at the appropriate v.
o Energy of this photon is equal to the difference in energy between the atom’s
initial state and its final state (figure 4-7)
o Hydrogen atoms emit only specific frequencies of light so the energy differences
between the atom’s energy states were fixed.
 The electron of a hydrogen atom exists only in very specific energy states.
Bohr Model of the Hydrogen Atom:
- How can we provide a model of the hydrogen atom that accounted for the various
wavelengths of its line-emission spectrum?
- 1913 = Niels Bohr proposed a model of the hydrogen atom that linked the atom’s
electron with photon emission.
- Electrons can orbit the nucleus only in allowed paths or orbits that has a definite
fixed energy.
o Low energy state = ground level v/s Higher energy state = higher energy levels
o Not only nuclear forces that keep electron from smashing into the nucleus, but it
is also the definite fixed energy of electrons in their designated energy levels
(orbits)
- Potential energy of electrons increases the further away the electron is from the
nucleus.
- Electrons can be in one orbit or another, not in between orbits.
So how does the Bohr model of the hydrogen atom explain the observed spectral lines?
-While in orbit, the electron can neither gain nor lose energy, but it can move to a higher energy
orbit by gaining an amount of energy that is equal to the difference in energy.
-Excited ground = photon of light is emitted that is equal to the difference in energy.
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Energy corresponds to a certain frequency (E = hv)
Based on the wavelength, Bohr could calculate the energies that an electron would have
in the allowed energy levels for the hydrogen atom. (Waves WS)
* Bohr’s model of the hydrogen atom did not explain the spectra of atoms with more that
one electron. Nor did Bohr’s theory explain the chemical behavior of atoms.