1. Arrangement of Electrons in Atoms
Chapter 4

2. The New Atomic Model
Investigations  relationship between light and atom’s electrons
How are electrons arranged? Why don’t they fall into the nucleus?

3. Light a wave or particle?
Wave Description:
Electromagnetic Radiation: energy that acts like a wave in space
All forms create Electromagnetic Spectrum

4. Electromagnetic Spectrum

5. Electromagnetic Spectrum
All forms move at speed of light, c, 3.00x108 m/s
Forms identified by:
wavelength, , the distance b/ corresponding points on adjacent waves. Units: nm, cm, or m
Frequency, , # of waves that pass a given point in a specific time, 1 sec. Unit: 1/s = Hertz, Hz

6. Wavelength and Frequency

7. Wavelength and Frequency
Inverse proportion equation!!
Frequency, 1/s
speed of light, m/s
wavelength, m

8. Calculation
Calculate the wavelength of a radio wave with a frequency of x 106s-1
Determine the frequency of light whose wavelength is nm.

9. Particle Nature of Light
Photoelectric Effect: emission of electrons from a metal when light shines on the metal

10. Photoelectric Effect
Light had to be certain frequency to knock e- loose
Wave theory  any frequency should work (just might take a while)
Light must also be a particle!
Max Planck(1900) explanation: objects emit energy in small packets called quanta
Video - 16

11. Max Planck
Quantum of energy is the smallest amount of energy that can be lost or gained by an atom
E = h
Frequency, s-1
Energy of quantum, in joules, J
Planck’s constant, 6.626x10-34 Js

12. Energy Calculation
What is the energy of green light, with a wavelength of 500. nm?

13. Albert Einstein Light is both wave and particle!
Particle of light = photon, having zero mass and a quantum of energy
Photons hit metal and knock e- out, but photon has to have enough energy

14. H-atom Emission Spectrum
Pass a current through gas at low pressure it excites the atoms
Ground state: lowest energy state of an atom
Excited state: atom has higher potential energy than it has in ground state

15. H – Atom Spectrum
When atom jumps from excited state to ground state it gives off energy  LIGHT! E2
Ephoton = E2 – E1 = hv E1

16. Bohr Model of H-atom

17. H-atom Line Emission Spectrum

18. Element Emission Spectras
Helium – 23 lines
Neon – 75 lines
Argon lines
Xenon – 139 lines
Mercury – 40 lines

19. H-atom Line Emission Spectrum
More lines in UV (Lyman series) and IR(Paschen series)
Why did H-atom only emit certain colors of light?
Explanation led to new atomic theory  Quantum Theory

20. Bohr Model of H-atom 1913 – Niels Bohr
e- circles nucleus in certain paths, orbits or atomic energy levels
e- is higher in energy the farther away from nucleus
e- cannot be between orbits
Video - 23

21. Bohr Model of H-atom

22. Bohr Model of H-atom
From wavelengths of emission spectrum Bohr calculated energy levels of H-atom
Model worked ONLY for H-atom

23. Quantum Model of Atom Can e- behave as a wave?
Yes!
To find e- use a photon, but photon will knock the e- off course
Heisenberg Uncertainty Principle: impossible to determine position and velocity of a particle at the same time.

24. Schrödinger Wave Equation
1926 – developed equation and only e- waves of certain frequencies were solutions
Quantization of e-  probability of finding e- in atom
No neat orbits  probability clouds or orbitals

25. Electron Configurations

26. Atomic Orbitals
Def: 3-D region around nucleus that indicates the probable location of an electron
Energy levels or shells:
Numbered 1-7
Smaller number = closer to nucleus, lower energy

27. Sublevels Each shell has sublevels s p d f 1 – s orbital
3 – p orbitals d
5 – d orbitals f
7 – f orbitals

28. Shells and Sublevels Shells and sublevels together: 1s 2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f, etc.
s is the lowest energy and f is the highest

29. Orbitals Each orbital in a sublevel can hold a maximum of 2 e-
1 – s 2 e- max.
3 – p orbitals 6 e- max.
5 – d orbitals 10 e- max.
7 – f orbitals 14 e- max.

30. Electron Configurations
Arrangement of e- in atom
Orbital Notation:
H has 1e-
Rules:
Aufbau Principle: electron occupies lowest energy level that can receive it

31. Electron Configurations
2. Pauli Exclusion Principle: no two e- in an sublevel orbital can have the same spin
3. Hund’s Rule: orbitals of equal energy are occupied by one e- before pairing up e-. All single occupied orbitals must have same spin.
He – 2e-

32. Energy of sublevels

33. Electron Configurations

34. Electron Configuration NotationB Ni Hg

35. Noble Gas Notation
Use noble gas from previous row Al Pb

36. Special Cases
d sublevel more stable with half-filled or completely filled sublevel Cr Cu