1. Development of the Atomic Model

2. Remember Rutherford?
Many scientists in the early 20th century found Rutherford’s model to be incomplete.
It did not explain how the electrons occupied the space around the nucleus.
Nor did it explain why the negatively charged electrons were not pulled into the atom’s positively charged nucleus.
Nor did it account for the differences in chemical behavior among the various elements.

3. Light: is it a wave? Or is it a particle?
Before 1900, light was thought to be a wave.
But then it was found to have some particle-like character.
Still, many properties can be described in terms of waves, and thus an understanding of the wave nature of light is needed.

4. Wave Description of Light
Electromagnetic Radiation- a form of energy that travels thru space like a wave
Examples: visible light, microwaves that warm and cook your food, X rays that doctors and dentists use to examine bones and teeth, and waves that carry radio and television programs to your home.
Moves at a constant speed (3.0 x 108m/s)

5. Picture of the Electromagnetic Spectrum

6. Waves are measured by:
Wavelength (l)- distance between corresponding points of adjacent waves (range from m to nm)
Frequency (n)- number of waves that passes a given point in a specific amount of time
Units: 1 wave/1 second= Hz (hertz)

7. A longer wavelength has a lower frequency and lower energy.
A shorter wavelength has a higher frequency and higher energy.
UV light is so dangerous because it is high in energy.

8. A Typical Wave
Include
Wavelength
Trough
Crest
Amplitude
Amplitude

9. Wavelength, frequency and the speed of light are all related…
c=ln
l and n are inversely proportional
c is the speed of light (3.0 x 108 m/sec)

10. Limitations of the Wave Model
Does not describe important aspects of light’s interactions with matter.
Cannot explain why heated objects emit only certain frequencies of light at a given temperature.
Does not explain why some metals emit electrons when colored light of a specific frequency shines on them.

11. The Photoelectric Effect
Photoelectric Effect- emission of electron from a metal when light shines on the metal.
This stream of electrons creates an electric current.
For a given metal, no electrons will be emitted if the light’s frequency is below a certain minimum—no matter how long the light was shined on it.

12. Max Planck
Suggested that hot objects do not emit electromagnetic energy (light) continuously (as a wave would) but in small, specific amounts he called quanta.
Quanta: minimum amount of energy that can be gained/lost by an atom

13. Energy= Planck’s constant  frequency of radiation
Planck summed up the relationship between quantum of energy (E) and frequency (n) of radiation in an equation…
E= hn
Energy= Planck’s constant  frequency of radiation
(Joules) (6.626x10-34Js) Hz

14. Practice Problems
What is the energy of a quantum of light with a frequency of 4.31 x Hz?
A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

15. Albert Einstein
Einstein expanded on Planck’s theory by introducing wave-particle duality of light.
That is, while light has many wave like characteristics, it can also be thought of as a stream of particles or bundles of energy. (each of which carries a quantum of energy).
Einstein called these particles photons.
Photon: a particle of electromagnetic radiation having no mass and carrying a quantum of energy.

16. The energy of a particular photon depends on the frequency of the radiation.
Einstein explained the photoelectric effect by proposing that electromagnetic radiation is absorbed by matter only in whole numbers of photons.
In order for an electron to be emitted from a metal surface, the electron must be struck by a single photon possessing at least the minimum energy needed to knock the electron loose.

17. According to E=hn, minimum energy corresponds to frequency; so, if the photon’s frequency is below minimum, the electron won’t be released.
Electrons of different metals require different minimum frequencies because they are bound more or less tightly.

18. So what does this have to do with atoms?

19. Have you ever wondered how light is produced in the glowing tubes of neon signs?
The light of the neon sign is produced by passing electricity through a tube filled with neon gas.
Neon atoms in the tube become excited.
These excited and unstable atoms then release energy by emitting light.
                              

20. Line-Emission Spectrum
Ground state- lowest energy state of an atom.
Excited state- a state in which an atom has a higher potential energy than in its ground state.
When an excited atom returns to its ground state, the energy is given off in the form of electromagnetic radiation.
This is the light we see.

21. When that light is passed thru a prism, it separates into a series of specific frequencies of visible light called the line-emission spectrum.
Hydrogen
Helium
Carbon

22. Scientist had expected to see a continuous range of frequency…so why only specific intervals???
This lead to a new theory of the atom…

23. Quantum Theory
In order for an atom to fall from its excited state back to its ground state, it must emit the energy as a photon.
That energy is equal to the difference in energy between the atom’s initial state and its final state.

24. Bohr Model
Linked the electron in his hydrogen atom with its photon emission.
The electron can only circle the nucleus in a certain path (called an orbit) which has a fixed amount of energy.
When the electron is in that orbit, the atom has a definite, fixed energy.

25. The electron stays in lowest energy orbit (close to the nucleus) unless it gains enough energy to move to the next orbit; then it has to release the energy as a photon to return to lower energy orbit.
This lowest energy orbit is separated from the nucleus by a large empty space where the electron cannot exist.

26. This is comparable to being on a ladder, you must be on one of the rungs, not in between because you can’t stand in midair.

27. In the same way, an electron can be in one orbit or another, but not in between.
But how does this explain the observed spectral lines?????

28. While in orbit, the electron can neither gain nor lose energy
While in orbit, the electron can neither gain nor lose energy. It can however, move to a higher energy orbit by gaining an amount of energy equal to the difference in energy between the higher energy orbit and the initial lower energy orbit.

29. When an atom is in an excited state, its electron is in a higher energy orbit.
When the atom falls back from the excited state, the electron drops to a lower energy orbit.
A photon is emitted equal in energy to the energy difference between the levels.
The energy of each emitted photon corresponds to a particular frequency of emitted radiation and therefore a particular color.

30. Limits of Bohr’s model:
Didn’t explain elements with more than one electron nor the chemical behavior of atoms.
To most scientists of the time, the Bohr model contradicted common sense. They didn’t see why electrons couldn’t exist in a limitless number of orbitals with slightly different energies.

31. Why were electrons allowed only in certain orbitals with definite energy?
In order to understand the answer, scientists had to change their view of the electron.

32. Electron Configuration
the arrangement of electrons in an atom

33. The quantum model of the atom improves on the Bohr model because it describes the arrangements of electrons in atoms other than hydrogen.
Because atoms of different elements have different numbers of electrons, a distinct electron configuration exists for the atoms of each element.

34. But why do we care???
We care because the behavior of electrons is what Chemistry is all about.
All the interactions of elements and characteristic properties of elements deal with the electrons.

35. Electrons line up around the nucleus according to 3 basic rules:
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule

36. Aufbau principal
Electrons occupy the lowest-energy orbital that can receive it.
Orbital with lowest energy is 1s.
Sometimes sublevels of one energy level might be higher than sublevels of the next energy level.
(Beginning with the 3rd energy level, energies of the sublevels in different main energy levels begin to overlap.)
Represent orbitals with ____

37. Pauli Exclusion Principal
No two electrons in the same atom can have the same set of 4 quantum numbers.
What that means is: in order for two negative electrons to occupy the same orbital, they must have opposite spins.
Written like: _____

38. Hund’s Rule
Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin.
___ ___ ___

39. Electron configuration notation
Eliminates the lines and arrows of orbital notation. Instead, the number of electrons in a sublevel is shown as a superscript of the sublevel letter
Hydrogen’s orbital notation:_____
1 s
Hydrogen’s configuration: 1s1

40. Let’s put it all together…
Take a look at the periodic table.
Give the electron configuration and orbital (arrow) diagram for the following elements.
Na, N, Y, Br, Xe, Ca, Mn

41. Nobel Gas Notation Shortcut for electron configuration.
Each noble gas ends with the period number and 6 electrons.
Use { } and the closest noble gas.
Examples: Ga, Zr, Cl
Nobel gas configuration shows an outer main energy level fully occupied by 8 electrons.