1. Electronic Structure of Atoms
Ch 6

2. The wavelength of a wave.

3. 3 Characteristics of Electromagnetic Radiation
Wavelength (l - lamda) distance between 2 consecutive peaks measured in meters
Frequency (n - nu) the number of waves per second that pass a given point in space measured in 1/s or hertz
Speed (c) all electromagnetic radiation travels at the same speed.
2.998 x m/s = approximately 3 x 108 m/s
c = n l

4. See page 215

5. FM Radio broadcasts at 88 to 108 megahertz
FM Radio broadcasts at 88 to 108 megahertz. An FM radio station broadcasts at a frequency of 91.5 megahertz. What is it’s wavelength? AM Radio broadcasts at 500 – 1600 kilohertz. KDLM broadcasts at 1340 Kh or 1,340,000 hertz. What is the wavelength?

6. 3 Observations
Blackbody radiation is the emission of light from a hot object.
Photoelectric effect is the emission of e- from a metal surface on which light is shining.
Spectra emission is light from electronically excited gaseous atoms.

7. Quantum Theory - Max Planck found that energy could only be gained or lost in whole number quantities. DE = h n DE = change in energy (joules) h = Planck’s constant x joule second or x Kg m2/s n = frequency in 1/s or hertz It had been thought that Energy was continuous. We now know that it is quantized, in discrete units called quantum. Albert Einstein purposed the idea all electromagnetic radiation is quantized in streams of particles called photons.

8. For electromagnetic radiation of wavelength 242
For electromagnetic radiation of wavelength nm, which is the longest wavelength that will bring about the photo dissociation of O2 , the photoelectric effect. What is the energy of one photon? Of 1 mole of photons?

9. The dual nature of light
Wave evidence
Reflection as in a mirror
Refraction: the bending of light as it passes into a new medium with a different index of diffraction.
Diffraction: the spreading out of light into the different wavelengths of light (different colors of the rainbow)
Particle evidence
light exerts pressure
Photoelectric is the emission of e- from a metal surface on which light is shining
Black body radiation is the emission of light from a hot object.
Spectra emissions is light from electronically excited gaseous atoms.

10. Bohr’s model – The Quantum model of the Atom The e- in the hydrogen atom move around the nucleus only in certain allowed circular orbits.
e- move in a circular orbit
2. e- has only a fixed set of orbitals. As long as an e- remains in a fixed orbital, it’s energy is constant.
3. e- can pass only from 1 allowed orbital to another. Discrete quantities of energy.

11. When an e- gains energy, it will jump to a higher energy level (excited). They will always come down to the ground state (lowest possible energy state or level available) The energy difference between the two energy levels will be given off as light.
Colors and wavelengths of photons in the visible region. Hydrogen bright line spectra.

13. An excited H atom returns to
a lower energy level.

14. DE = Eenergy of final state – E energy of initial state. EN = -2
DE = Eenergy of final state – E energy of initial state EN = x J EN = energy of specific energy level N = energy level N2
Determine the wavelength of the light emitted from a hydrogen atom when an e- drops from the 5th energy level to the 2nd energy level.

15. The color of the photon emitted depends on
the energy change that produces it.

16. De Broglie’s equation is used to determine the mass of a particle traveling as a wave. l = h / mv l = wavelength in m h = Planck’s constant x Kg m2/s m = mass in Kg v = velocity in m/s

17. What is the wavelength associated with an e- traveling 1/10 the speed of light? me- = x Kg

18. Heisenberg’s Uncertainty Principle
X – D ( m n ) > h/ 4p
X = position uncertainty
m n = momentum uncertainty
There is a fundamental uncertainty as to the position and momentum of an e- at any given time. The more accurate we know the position, the less we know about the momentum.

19. Bohr’s model works for the hydrogen atom, but it became apparent quickly that it would not work for more complex atoms. Heisenberg, de Broglie, and Schrodinger used wave mechanics or quantum mechanics to approach the structure of the atom. They showed that e- act as a particle and a wave. They thought of the e- as a standing wave, as in musical instruments. Schrodinger worked out the math to explain the e- wave function y (psi) = orbital

20. We know the probability of e- in a region y2
The orbital that describes the e- density of the hydrogen electron in its lowest possible energy state.

21. Quantum Numbers : define position of the e- in the atom
Principle quantum number (n) = energy level
integral value = 1, 2, 3, …
Angular Momentum quantum number (l) = sublevel or the shape of the orbitals
integral value of 0 to n-1
l = 0  s sublevel
l = 1  p sublevel
l = 2  d sublevel
l = 3  f sublevel

22. Magnetic quantum number (ml ) = orbital. integral value = - l to + l
Magnetic quantum number (ml ) = orbital integral value = - l to + l l = 0  s sublevel  ml = l = 1  p sublevel  ml = -1, 0, l = 2  d sublevel  ml = -2, -1, 0, +1, l = 3  f sublevel  ml = Electron Spin quantum number (ms) = the 2 e- in an orbital will spin the opposite direction. Value of +1/2 or -1/2

23. Charts show the radii probability
Charts show the radii probability. This shows the distance that e- are most likely found from the nucleus.

24. The hydrogen 1s orbital.
The 1s orbital is a sphere that encloses 90% of the total e- probability.

26. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

27. The three 2p orbitals.

28. Diagram of principal energy levels 1 and 2.

31. Rules for assigning the principle quantum numbers.
1. e- occupy the orbitals in a way to minimize the energy of the atom.
2. No 2 e- have all 4 quantum numbers the same
(Pauli Exclusion Principle)
3. e- will occupy the lowest energy configuration available. When orbitals of identical energy are available, e- occupy these orbitals singly (Hund’s rule). Atoms tend to have as many unpaired e- as possible.
Aufbau principle: As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen – like orbitals.

33. Unusual configurations
Expect to see What actually occurs
Cr: [Ar] 4s2 3d4 Cr: [Ar] 4s1 3d5
Cu: [Ar] 4s2 3d9 Cu: [Ar] 4s1 3d10

34. 1. Helium 4s
3px 3py 3pz 3s
2px 2py 2pz
Electron configuration ___________________________________ 2s
1s Core notation ____________________________________

35. 2. Nitrogen px 4py 4pz
3d 3d 3d 3d 3d 4s
3px 3py 3pz 3s
2px 2py 2pz
Electron configuration _______________________________________ 2s
1s Core notation _____________________________________________

36. Orbitals being filled for elements in various parts of the periodic table.

37. Filling order for placement of e- in the orbitals.
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p
Filling order for placement of e- in the orbitals.