1. The Atom
(Chapter 9)
Just How Small is an Atom?

2. Student Learning Objectives
Recall the defining properties of an atom
Apply the dual nature of light to atoms
Analyze the quantum nature of the atom

3. What are the defining properties of regular atoms?
All atoms contain protons, neutrons, electrons. There some exceptions.
What makes each atom unique is the number of protons.
Regular, neutral atoms have an equal number of electrons, protons, and neutrons.
Losing or gaining electrons causes an atom to become an ion.
Losing or gaining neutrons causes an atom to become an isotope
Ions = electrons lost or gained
Isotopes = different number of neutrons
The 2,400-year search for the atom - Theresa Doud

4. Practice Where is most of the mass in an atom?
2) What contributes most to the size of an atom?
3) What is an ion? What is an isotope?
4) What prevents the atoms of one object from passing through the atoms of another object?
5) Imagine a world with elements classified according to numbers of electrons.
Nucleus
Electron orbits
Electric force – electrons repelling
Electrons can be lost or gained

5. What is meant by the dual nature of light?
Quantum theory predicts that objects emit individual packets of energy (photons).
The dual nature of light is that sometimes light behaves as a wave (wavelengths) and sometimes light behaves as a particle (photons).
Electromagnetic radiation exhibits wave-particle duality.
Half way down link page
Is light a particle or a wave?

6. de Broglie Wavelengths
All subatomic particles exhibit wave-particle duality.
A standing wave is set up by the motion of electrons.
de Broglie Wavelengths
lDB = h mv
Seeing the Smallest Thing in the Universe

7. How do atoms emit and absorb light?
Each atom has its own unique pattern of allowed orbits, based on the number of protons.
When electrons change orbits, light is emitted or absorbed.
Bohr Model-Early Quantum theorem

8. Concept of Quantized Energy four specific potential energy values
Continuous Energy
The uncertain location of electrons - George Zaidan and Charles Morton
Section 9.2

9. Each orbit has a specific energy associated with it.
The farther an electron is from the nucleus, the higher its energy level.
As long as an electron remains in a quantized orbit, the atom emits no radiation.
Quantum Jump: an abrupt transition of a system described by quantum mechanics from one discrete state to another
Particles and waves: The central mystery of quantum mechanics - Chad Orzel

10. Electrons quantum jump (instantaneous transition) from one allowed orbit to another, changing the energy state of the atom.
Ground State: lowest energy levels full
Excited State: energy levels added
Each type of atom has its own set of energy levels, and emits its own specific wavelengths of photons. (Atomic Spectrum)

11. E(ni) – E(nf) = E(photon)
Photon Energy
Each photon has a particular energy, based on the quantum jump.
E(ni) – E(nf) = E(photon)
Each color of light is associated with a frequency/wavelength as well as a particular energy.
h = 6.63 x Js
E = hc l
E = hn
What is the Heisenberg Uncertainty Principle? - Chad Orzel

12. Line Emission Spectrum for Hydrogen
When light from a gas-discharge tube is analyzed only spectral lines of certain frequencies are found
Section 9.3

13. Line Absorption Spectrum for Hydrogen
Results in dark lines (same as the bright lines of the line emission spectrum) of missing colors.
Section 9.3

14. Spectral Lines for Hydrogen
Transitions among discrete energy orbit levels give rise to discrete spectral lines within the UV, visible, and IR wavelengths
Section 9.3

15. Practice
What is the energy of a photon that results from the 4  2 transition in the hydrogen atom?
2) What is the energy of one yellow photon (l = 517 nm) from the Sun? How does this compare to the total energy output of the Sun?
Hydrogen Orbit eV
n = 1 8
n = 2 15
n = 3 30
n = 4 80
80 – 15 = 65 eV (1 eV = × joules)
E = 3.84 x Joules
Schrödinger's cat: A thought experiment in quantum mechanics - Chad Orzel