1. Chapter 13 Electrons in Atoms

2. Section 13.1 Models of the Atom
OBJECTIVES:
Summarize the development of atomic theory.

3. Section 13.1 Models of the Atom
OBJECTIVES:
Explain the significance of quantized energies of electrons as they relate to the quantum mechanical model of the atom.

4. J. J. Thomson’s Model Discovered electrons
Negative electron floating around
“Plum-Pudding” model

5. Ernest Rutherford’s Model
Discovered dense positive piece at the center of the atom- nucleus
“Nuclear model”

6. Niels Bohr’s Model Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.
“Planetary model”

7. Bohr’s planetary model
electron cannot exist between energy levels, just like you can’t stand between rungs on ladder
Quantum of energy required to move to the next highest level

8. The Quantum Mechanical Model
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Erwin Schrodinger
Mathematical solution

9. The Quantum Mechanical Model
Has energy levels for electrons.
Orbits are not circular.
It can only tell us the probability of finding an electron a certain distance from the nucleus.

10. The Quantum Mechanical Model
The atom is found inside a blurry “electron cloud”
An area where there is a chance of finding an electron.
Think of fan blades

11. Atomic Orbitals
Principal Quantum Number (n) = the energy level of the electron.
Within each energy level, the complex math of Schrodinger’s equation describes several shapes.
Atomic orbitals - regions where there is a high probability of finding an electron.
Sublevels

12. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6d 5 10 3 7 14 4 f

13. By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p
2s22p6
8 total electrons

14. By Energy Level Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, ahd 14 in f
4s24p64d104f14
32 total electrons

15. Section 13.2 Electron Arrangement in Atoms
OBJECTIVES:
Apply the aufbau principle, the Pauli exclusion principle, and Hund’s rule in writing the electron configurations of elements.

16. Section 13.2 Electron Arrangement in Atoms
OBJECTIVES:
Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.

17. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
Aufbau diagram - page 367

18. Electron Configurations
Aufbau principle- electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

19. Electron Configuration
Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.
Phosphorus (15 electrons)

20. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The first two electrons go into the 1s orbital
Notice the opposite spins
only 13 more to go...

21. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s
The next electrons go into the 2s orbital
only 11 more...

22. Increasing energy The next electrons go into the 2p orbital3d 4d 5d 7p 6d 4f 5f
The next electrons go into the 2p orbital
only 5 more...

23. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s
The next electrons go into the 3s orbital
only 3 more...

24. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into separate shapes
3 unpaired electrons
= 1s22s22p63s23p3

25. The easy way to remember
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2
2 electrons

26. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2
4 electrons

27. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2
12 electrons

28. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2
20 electrons

29. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
38 electrons

30. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
56 electrons

31. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2
88 electrons

32. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
108 electrons

33. Exceptional Electron Configurations

34. Orbitals fill in order Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order

35. Write these electron configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 expected
But this is wrong!!

36. Chromium is actually: 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.

37. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions: d4, d9

38. Section 13.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Calculate the wavelength, frequency, or energy of light, given two of these values.

39. Section 13.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Explain the origin of the atomic emission spectrum of an element.

40. Light
The study of light led to the development of the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many kinds of waves
All move at 3.00 x 108 m/s = c

41. Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough

42. Parts of Wave - p.372 Origin - the base line of the energy.
Crest - high point on a wave
Trough - Low point on a wave
Amplitude - distance from origin to crest
Wavelength - distance from crest to crest
Wavelength is abbreviated by the Greek letter lambda = l

43. Frequency
The number of waves that pass a given point per second.
Units: cycles/sec or hertz (hz or sec-1)
Abbreviated by Greek letter nu = n
c = ln

44. Frequency and wavelength
Are inversely related
As one goes up the other goes down.
Different frequencies of light are different colors of light.
There is a wide variety of frequencies
The whole range is called a spectrum, Fig , page 373

45. Low energy
High energy
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low Frequency
High Frequency
Long Wavelength
Short Wavelength
Visible Light

46. Atomic Spectrum Each element gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of.

47. These are called discontinuous spectra, or line spectra
unique to each element.
These are emission spectra
The light is emitted given off
Sample 13-2 p.375

48. Light is a Particle Energy is quantized. Light is energy
Light must be quantized
These smallest pieces of light are called photons.
Photoelectric effect?
Energy & frequency: directly related.

49. Energy and frequency E = h x  E is the energy of the photon
 is the frequency
h is Planck’s constant
h = x Joules x sec.
joule is the metric unit of Energy

50. The Math in Chapter 13
2 equations so far:
c = 
E = h
Know these!

51. Examples
What is the wavelength of blue light with a frequency of 8.3 x 1015 hz?
What is the frequency of red light with a wavelength of 4.2 x 10-5 m?
What is the energy of a photon of each of the above?

52. Explanation of atomic spectra
When we write electron configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level.

53. Changing the energy
Let’s look at a hydrogen atom

54. Changing the energy
Heat or electricity or light can move the electron up energy levels (“excited”)

55. Changing the energy
As the electron falls back to ground state, it gives the energy back as light

56. Changing the energy May fall down in steps
Each with a different energy

57. { { {

58. Ultraviolet
Visible
Infrared
Further they fall, more energy, higher frequency.
This is simplified
the orbitals also have different energies inside energy levels
All the electrons can move around.

59. What is light? Light is a particle - it comes in chunks.
Light is a wave- we can measure its wavelength and it behaves as a wave
If we combine E=mc2 , c=, E = 1/2 mv2 and E = h
We can get:  = h/mv
called de Broglie’s equation
Calculates the wavelength of a particle.

60. Sample problem
What is the approximate mass of a particle having a wavelength of 10-7 meters, and a speed of 1 m/s?
Use  = h/mv
= 6.6 x 10-27
(Note: 1 J = N x m; 1 N = 1 kg x m/s2

61. Matter is a Wave Does not apply to large objects
Things bigger than an atom
A baseball has a wavelength of about m when moving 30 m/s
An electron at the same speed has a wavelength of 10-3 cm
Big enough to measure.