1. SCH12U February 7 2011 Mr. Dvorsky
Wave Mechanical Model
SCH12U
February
Mr. Dvorsky

2. Recall.... -When a substance is exposed to a certain intensity of light, the atoms absorb some of the energy. Such atoms are said to be excited. -When atoms are in an excited state, unique changes occur that can be used to identify the atom. -The method of studying substances exposed to exciting energy is called spectroscopy. A pattern of radiant energy studied in spectroscopy is called a spectrum.

3. Excited atoms soon lose the energy they’ve gained.
The energy emitted by gaseous atoms occurs at specific points in a spectrum.
-The lines of an
emission spectrum are
characteristic of the
element being
excited.

4. -In addition to emission spectrum: if these same atoms are exposed to light of all wavelengths, they will absorb energy. If you pass this light through the gaseous form of the element, some of the incident light will be missing.

5. Emission and absorption spectra are
complementary. The colours of light that the
atoms of an element will absorb are the same
ones it will emit when it is excited.
So: the same mechanism that operates to emit
light from an atom operates in reverse to
Absorb light.

6. -electron position in the atom is used to store energy
-electron position in the atom is used to store energy. -there are only so many positions it can occupy. -so if a hydrogen atom is hit by light, and the energy of that light is not the exact amount necessary to shift its position to another legal position, nothing happens. If it is, the atom accepts that energy and the electron moves up (absorption). -later when the electron drops back into its original position, the same amount of light is given off that it took in (emits at same wavelength).

7. How did Bohr put his model of the atom together
How did Bohr put his model of the atom together? Important Idea Max Planck – energy instead of being given off continuously is given off in little packets called quanta. –We refer to quanta of light as photons. -also, the amount of energy given off is directly related to the frequency of the light emitted. E = hf, where h is a constant known as Planck’s constant, x joules per hertz.

8. Bohr now had the idea that energy could be given off in quanta instead of continuously. -absorption of light by hydrogen at definite wavelengths correspond to the definite changes in energy of the electron that we discussed earlier. Also, electrons could only occupy certain orbits whose differences in energy = the energy absorbed when the atom was excited.

9. Summary of the relationship between electromagnetic energy and an electron: An electromagnetic wave of a certain frequency has only one possible wavelength given by: wavelength = c/f. It has only one possible amount of energy given by: E = hf -since h and c are constants, it frequency, wavelength, or energy is known, we can calculate the other two.

10. Weakness of Bohrs Model -works for only 1 electrons systems such as hydrogen. -the frequencies predicted by Bohr in the hydrogen spectrum are essentially correct but not exactly correct. –better instrumentation has shown the bands predicted by Bohr are not single lines but several lines closely spaced. -scientists now had to account for this “fine structure” in the hydrogen spectrum.

11. De Broglie knew of Planck’s ideas concerning light being made of discrete packets called quanta-this gives waves the properties of particles.
De Broglie thought that if Planck were correct, it might be possible for particles to also have some properties of waves.
His idea of “matter-waves” were later proven experimentally.
The wave-particle duality of nature

12. Schrodinger -treated electron as a wave and developed a mathematical equation to describe its wave-like behaviour. -this equation can be used to describe the location in space (x,y,z) where an electron can be found. -these regions are called orbitals

13. Heisenberg -impossible to determine the exact location of a single subatomic particle (Heisenberg’s uncertainty principle). -therefore Schrodinger’s orbitals are actually probability distributions of where an electron MAY be found.

14. Orbitals
When a planet moves around the Sun, you can plot a definite path for which it is called an orbit.
Similarly, a simple view of the atom has electrons orbiting around the nucleus.
In fact, electrons inhabit regions of space known as orbitals

15. To plot a path for something you must know where it is and where it will be an instant later. This cannot be done for electrons
Instead we use probability distributions

16. Hydrogen’s electron, the 1s orbital.
Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soon afterwards, you do the same thing, and find that it is in a new position. You have no idea how it got from the first place to the second.You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the electron is likely to be found.

17. For the electron of hydrogen, it can be found anywhere in a spherical region around the nucleus. Most of the time it will be found within that region. –this region is called an orbital, it is where electrons live. –cannot say what the electron is doing in that space, just that it will have a particular energy level.

18. Each orbital has a name.
The orbital occupied by the hydrogen electron is called 1s.

20. Not all electrons inhabit s orbitals. In fact, most don’t.
At the second energy level, not only is there a 2s orbital, but also 3 2p orbitals
p orbitals resemble two balloons held together at the nucleus
Why 3? At any level it is possible to have 3 identical p orbitals, just that they are pointing in different directions.

21. -there are also d and f orbitals which become available for electrons to inhabit at higher energy levels
-at the third level there are 5 d orbitals, as well as the 3s orbital, and the three 3p orbitals (for a total of 9 orbitals at that level)
-at the fourth level there is the 4s, plus the three 4p orbitals, plus the five 4d orbitals, and seven 4f orbitals

22. The inverted pyramid analogy
You can think of an atom
as a strangely arranged
house with the nucleus living o
on the ground floor. There
are various rooms (orbitals) on the higher floors
(energy levels) occupied by the electrons.
-On the first floor there is only 1 room (1s), on the
second floor there are 4 rooms (2s, 2px, 2py, 2pz),
on the third floor there are 9 rooms (1s, 3 3p
orbitals, 5 3d orbitals). The rooms aren’t big, each
can hold only two electrons.

24. The order of filling orbitals
Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones (Aufbau principle)
Due to electron repulsion, all orbitals of equal energy acquire one electron before any orbital accepts two electrons (Hund’s rule)

25. Diagram summarizing up to the 4p level
-notice that s orbitals have lower energy than p at the same level so s orbital fills with electrons before the corresponding p orbitals
Notice 3d orbitals. They are slightly higher level than the 4s and so it is the 4s which fill first followed by 3d and then 4p.
-overlap causes some confusion!

26. Orbital filling and the periodic table
The first period:

27. The second period:

28. The second period (continued):
Make life easier. Use shortcuts!
1. You can just lump together all the various p orbitals

31. Once we get to level 4, we must recall this: -4s is at a lower energy than 3d and will therefore fill before the orbitals of 3d