1. Let’s Practice
1.Calculate the speed of a wave whose wavelength is 0.25m and whose frequency is 760 Hz.
2.Calculate the frequency of a blue light wave whose wavelength is 5.22 x 10-7m.
3. Calculate the energy of the blue light in #2.

2. 1. Calculate the speed of a wave whose wavelength is 0
1.Calculate the speed of a wave whose wavelength is 0.25m and whose frequency is 760 Hz.
C = λν = (0.25m)(760Hz) = 19 m/s
2.Calculate the frequency of a blue light wave whose wavelength is 5.22 x 10-7m.
Ν = c/λ = (3.00 x 10-8m/s)/ 5.22 x 10-7m
= x 1014Hz
3. Calculate the energy of the blue light in #2.
E = hν = (6.626 x 10-34Js)(5.75 x 1014Hz)
= x 1019J

3. GLOW IN THE DARK STARS
laser light nm wavelength stars glow? _______
flashlight nm wavelength stars glow? _______
UV light nm wavelength stars glow? _______
glowing star light nm wavelength
On a separate piece of paper, do the following:
A. Convert wavelength in nm to meters.
B. Using Planck’s energy formula and the speed of light formula, determine the energy for each.
C. Answer the questions:
1. Would infrared light cause the stars to phosphoresce?
2. Would microwave light cause the star to phosphoresce?
3. Give another example of minimum energy photons.

4. ATOMIC EMISSION SPECTRA AND QUANTUM THEORY

5. Atomic Emission Spectra
We talked before about flame tests and why the elements produce a color when heated. These elements produce a continuous emission of light.
The electrons in the atoms absorb energy and become excited. Excited and unstable electrons then drop back to their stable level, releasing energy by emitting light.

6. Atomic Emission Spectra
A neon light works the same way.
We can turn a continuous emission spectra into a discontinuous one by refracting the light.
An Atomic Emission Spectrum is a set of frequencies of electromagnetic waves emitted by atoms of the element. They are usually distinct color lines. Each element’s atomic emission spectrum is unique, can be used to identify the element, and can be used to determine if the element is part of an unknown compound.

7. Atomic Emission Spectra - Hydrogen
Page 144 shows the emission spectrum of hydrogen’s one electron. Notice that it is discontinuous – it is made up of only certain frequencies of light – Figure (b).

9. Spectra of Four Elements
Hydrogen: which emits 4 colors of light that's in the visible light range. Note that other frequencies, such as UV light might be emitted, but we can't see them.
Helium: It has 2 electrons and we see 7 colors.
Mercury: spectra shows 8 colors. Mercury also produces a lot of UV light which in fluorescent bulbs is normally converted to visible light by the use of certain minerals that capture UV light and emit visible light.
Uranium: Uranium emits many frequencies of colors. It appears that the elements that have more electrons emit more colors. So there seems to be a connection.

10. Atomic Emission Spectra

11. Atomic Emission Spectra
Acts as a fingerprint for elements.
Is the same everywhere in the universe for a particular element.
Allows us to determine the elements in a star far away.

12. QUANTUM THEORY
1913: Neils Bohr comes up with the quantum model of the hydrogen atom. He also correctly predicts the frequencies of the spectral lines in the hydrogen atomic emission spectrum.
He theorized that the quantized energy that Max Planck suggested, Einstein proposed, and Rydberg calculated could be the reason that certain frequencies of light were seen coming from hydrogen. He thought hydrogen’s electron could only occupy certain energy levels within the atom. Light that equaled the difference in the levels could cause the electron to jump to the higher level - this is called the EXCITED STATE. Go to the website ( to see this. Go to the topic “Bohr’s Atom” about halfway down.
When the electron fell back into the original orbit that same frequency of light would be emitted. This would appear as a single line on the atomic emission spectra.

13. BOHR’S THEORY
Electrons were now viewed to be in orbit around the nucleus. The electrons could only orbit at certain distances which represent distinct energy (quantum) levels. These energy levels were labeled n=1, n=2, and so forth. They were called the principle quantum numbers.
The smaller the orbit, the lower the atom’s energy state or energy level. Each orbit is assigned a quantum number, n.
Bohr said only certain atomic energies are possible for each atom’s electron and so only certain frequencies of electromagnetic radiation can be emitted.
See page 147, figure 5-11.

14. Hydrogen Atom Emissions
Balmer Series – visible light
Lyman Series – ultraviolet light
Paschen Series – Infrared light
Bracket Series – Infrared light

15. QUANTUM THEORY HISTORY
The spectra of elements showed that light waves also behaved like particles. Who was brave enough to ask, "If waves could behave like particles, can particles behave like waves?"

16. Quantum History
Louis de Broglie stated that Bohr’s quantized electron orbits had characteristics similar to waves. Particles could have wavelike behaviors. The energy had wavelike characteristics. (You do not need to know the equation.)

17. How can something be a particle and a wave?
Go to the website: ( to see this. Go to the topic “Electron: Wave and Particle” about halfway down.
This wave motion would be n=1. The electron starts out as a particle orbiting but also defining a wave-like oscillation. Next, the oscillation is shown without the electron particle. Finally the wave is changed into an electron cloud representing the probability of the electron’s position. In textbooks you might see the electron represented by a particle in orbit or as a cloud. That's because of electron's dual nature of being both wave and particle.

19. Quantum History
Werner Heisenberg came along and concluded that it is impossible to make any measurement on an object without disturbing the object. It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
In other words, the light used to measure the particle changes it.
Now watch the Captain Quantum video. You can find it on my website on the assignments page.

20. Quantum History
Erwin Schödinger, in 1926, developed an atomic model in which the electrons are treated as waves. This model is called the quantum mechanical or wave mechanical model.

21. Quantum Mechanical Model
This model does not describe the path of the electron around the nucleus.
The three dimensional region around the nucleus called an atomic orbital describes the electron’s probable location.
It is a fuzzy cloud in which the density of the cloud at a given point is proportional to the probability of finding the electron at that point.
The electron cloud has no definite boundary – it is arbitrarily drawn at 90%. See page 152, figure 5-15.

22. Quantum Mechanical Model

23. Quantum Mechanical Model
The atom has a dense, centrally located, positively charged nucleus full of protons and neutrons surrounded by mostly empty space where the electrons are.
The energy of electrons is quantized.
Electrons exhibit both wave and particle behaviors.

24. Quantum Mechanical Model
4. The absolute location of an electron is impossible to determine – its location and velocity cannot be determined at the same time.
5. The electrons travel in orbitals that have characteristic sizes, shapes, and energies, but do not describe how the electrons move.

25. DO NOW Pick up two handouts
Turn in corrected paperwork – if you . Make sure you have the old paperwork, the corrected pages, and the yellow sheet all paper clipped together.
Get out problem set homework and Atomic Emission Spectra notes.

26. DO NOW Pick up two handouts Turn in “Glow in the Dark Stars” homework.
Get out problem set homework and Atomic Emission Spectra notes.

28. Quantum Mechanical Model
So how do electrons position themselves outside the nucleus?
As the progression of elements were built by adding one proton and one electron at a time, the position of the protons was always in the center of the atom in the nucleus.
However, electrons repelled each other, so as each electron got added for each new element, they would find a position and shape that maximized their distance from each other. Amazingly, the way they positioned themselves followed a fairly basic pattern.

29. Quantum Mechanical Model
Quantum Numbers are used to describe the most probable location of an electron.
The electrons fill the orbitals in a pattern.
No matter what element, the pattern is the same.

30. Quantum Numbers There are four Quantum numbers.
They each tell something different about the electron’s location

32. Principle quantum number, n
Is the energy level number.
Gives information about the relative size and energy of the atomic orbitals.
It can have values of 1, 2, 3, 4….
The greatest number of electrons possible in any one level is 2n2.
Example: The maximum number of electrons that can occupy the first level is 2(1)2 = 2; the fourth level is 2(4)2 = 32.

33. Angular Momentum Quantum Number, l
Is the energy sublevel number.
It gives information as to the shape of the orbitals.
The first four levels are s, p, d, and f.
Example: The first energy level has only an s sublevel. The second energy level has an s and p sublevel. This third energy level has s, p, and d sublevel.

34. Angular Momentum Quantum Number, l

36. Angular Momentum Quantum Number, l
The quality of the spectroscopic lines were labeled sharp, principle, diffuse, and fundamental.
It was believed that the different orbitals were responsible for the quality of lines; for example, orbitals that created “sharp” lines were given the name “s”. “p” orbitals made the principle lines, etc.
This turned out not to be true, but the names stuck.

37. Magnetic Quantum Number, m
Gives information about the orientation in space of an orbital.
The s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has 5 orbitals, and the f sublevel has 7 orbitals.
This number determines which p, d, or f orbital the electrons are in.

42. Electron Spin Quantum number, s
Indicates the direction of the electron spin.
Spin is either clockwise or counterclockwise.
Is designated with a +1/2 or a –1/2 .

43. Quantum Numbers
For example, what are the four quantum numbers for the last electron in an oxygen atom? The last electron is located in the 2p orbital with a down spin. n = 2 l = 1 m = -1 s = -1/2

45. What’s Next?
Thankfully, we will not learn how to write the quantum numbers for each electron in an atom.
We will learn electron configuration and orbital diagrams to depict the probable location of electrons.
Stay tuned…..