1. Development of a New Atomic Model
Chapter 4
Development of a New Atomic Model

2. Rutherford Atomic Model
Rutherford’s nuclear model was incomplete because he never explained how the electrons are arranged and why they are not pulled into the nucleus.
It lacked an explanation for the differences in chemical behavior among elements.

3. Electrons in Atoms
In the 1900s, scientists began to unravel the puzzle of chemical behavior by observing that some elements emitted visible light when heated. The emitted light is related to the arrangement of the electrons in the atom.

4. Electromagnetic Radiation
Electromagnetic radiation-a form of energy traveling through space in the form of waves.

5. Characteristics of Waves
Wavelength-length of a wave from peak to peak, unit is meter.
Frequency-number of waves that pass a point in one second,
waves/second = hertz (Hz)

6. Amplitude-wave’s height
Speed-All waves travel at the speed of light X 108 m/s in a vacuum
Illustrate wavelength and frequency on a wave.

7. Electromagnetic Radiation Spectrum
Memorize the relative positions of the following types of electromagnetic radiations:

8. Relationships
Each type of radiation has a different wavelength, a different frequency.
How is wavelength related to frequency?
How is the energy related to the frequency?
How is wavelength and energy related?

9. Formulas
Memorize the following formulas that relate the above relationships to each other
c = w * f c = 3.0 X108 m/s= speed of light
w = wavelength (m)
f = frequency (Hz)
E = h * f E = energy (J)
h = Planck’s constant = X10-34 Js

10. 1. Calculate the frequency of a wave with a wavelength of 0.01 cm.
Example
1. Calculate the frequency of a wave with a wavelength of 0.01 cm.

11. Example
2. Calculate the energy of a wave that is 88.1 megahertzes.

12. Photoelectric effect

13. Light as a Particle
Light can knock loose electrons from a metal. This is known as the photoelectric effect.
A quantum is the minimum quantity of energy that can be lost or gained by an atom.
It is Albert Einstein that proposed the idea that light has both wave and particle properties.
Photon = quantum of energy = packets of energy

14. Spectra
Observe visible light through a prism to see a continuous spectrum.
Observe the line-spectrum of hydrogen.
Illustrate the 4 colored lines in the hydrogen line spectrum.

15. Energy States
Observe the silent lecture that explains the line spectrum of hydrogen.
Ground state-lowest energy state
Excited state- higher state than ground state

16. Hydrogen Line Spectra
Neil Bohr’s (1913) explanation of the line spectra of hydrogen:
1. Electrons orbit around the nucleus.
2. When an electron is excited, it moves into a higher energy level (orbit).

17. Hydrogen Line Spectra
When the electron drops back to the ground state, it loses the energy in the form of light.
4. Bohr predicted lines in the UV and IR range that was discovered much later

18. Bohr’s Model of Atom
Illustrate Bohr’s model of the atom:

19. Colors in the Line Spectra
Each line in the line spectrum of hydrogen represents an electron jumping from a higher energy level down to a lower energy level by releasing a photon of light.

20. Problems in Bohr’s Model
Observe the line spectrum of elements with more than one electron.
Bohr could not explain the extra lines for atoms larger than hydrogen
Schrödinger was able to account for the other lines in the line spectrum of large atoms.

21. Schrodinger’s Quantum Mechanical Model of Atom
The model is an equation that describes the probability of finding the electron at a given place around the nucleus.

22. Schrodinger’s Equation
υ2ψ + υ2ψ + υ2ψ + 8¶2m(E-V) ψ = 0
υx υy υz2
υ-partial derivative m-mass of electron
Ψ-amplitude of light wave V- potential energy
¶-pi E-energy

23. Orbital- the 3-D region around the nucleus where there is high probability of finding the electron.

24. Schrodinger’s Equation
Describes the behavior of subatomic particles.
He related the amplitude of the electron wave to any point in space surrounding the nucleus. (has only been solved for hydrogen)

25. Quantum numbers - Represents different energy levels of the electron

26. 4 Variables of Model
1. Principal quantum number (n)-main energy level (distance from the nucleus)

27. 2. Angular momentum quantum number, l, shape of orbital (sublevels),
4 shapes: s p d f

28. S sublevel

29. P sublevel

30. d sublevel

31. F sublevel

33. 3. Magnetic quantum number, m, orientation(#) of orbital
s has one orientation only; p has 3; d has 5; f has 7 different orientations
4. Spin quantum number, s, spin of electrons-clockwise or counterclockwise

34. Electron Configurations
Electron Configuration-arrangements of electrons in atoms
Rules governing electron configuration:
1. Aufbau principle – electrons occupies lowest energy orbital it can get.
Order:
1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d
Sequence Aid:

35. Use the diagonal rule or Aufbau
series.

36. Electron Configuration
Pauli Exclusion Principle – no two electrons can have the same four quantum numbers. Therefore, two electrons in the same orbital must have opposite spins.
Example:
Hund’s Rule-orbitals of the same energy must have one electron each before the second is added in.
Orbital Notation: I I I I I I Electron Configuration: 2p6

37. Example 2
Write the orbital notation and the electron configuration notation of boron, neon, and nickel.
Prac p107

38. Questions
How many sublevels are in the second energy level? Fourth energy level?
2. How many electrons are in the fourth energy level?

39. Valence Electrons
Valence Electrons- number of electrons in the highest energy level
Highest occupied level – highest energy level electrons
Inner-shell electrons – electrons not in highest energy level

40. Electron Dot Structures
Electron Dot Structures (Lewis Structures) used as shorthand since valence electrons are the one used to make compounds.

41. Steps: 1. Determine the number of valence electrons for the element.
2. Place one valence electron on the four sides of the symbol, and then pair the valence electrons until all are used.
The symbol represents the inner electrons and the dots represent the valence electrons.

42. Valence Electrons Example – Draw the Lewis dot structures for each.
1. Boron
2. Oxygen
3. Fluorine
4. Neon

43. Noble Gas Notation
Noble-gas notation: [symbol of the noble gas in the row above] and finish the electron configuration
Example 4: Write B, Ne, and Ni in the noble gas notation