1. Chapter 5: Arrangement of Electrons in Atoms

2. Niels Bohr (1885 – 1962) – Danish Physicist
Improved Rutherford’s work by saying electrons do not lose energy as they emit light, so they will stay in orbit
Stated there are definite levels and the electrons follow set paths without gaining or losing energy (Planetary Model 1913)

3. Niels Bohr (1885 – 1962) – Danish Physicist
Each level has a specific amount of energy associated with it
The electrons can only jump levels if they gain or lose that specific amount of energy

4. Energy Levels
In the ground state, electrons are at their lowest, most stable energy levels.
In the excited state, electrons require a specific amount of energy to move to a higher energy level.

5. Max Planck (1858 – 1947) – German Physicist
Proposed Planck’s Theory which says that energy is given off in specific energy amounts called quanta which is based on the particle nature of light.
Each quantum of energy corresponds to the different energy levels for electrons.
Proposed the equation: E=hf, E = energy
h = Planck’s constant (6.626 x 10 −34 J∙s)
f = frequency

6. Photoelectric Effect
When light shines on the surface of a metal, the metal surface emits electrons

7. Louis de Broglie (1892 – 1987) – French Physicist
Suggested that if waves can have a particle nature, particles can have a wave nature, known as the “wave-particle duality” principle
Wondered why the positive nucleus and negative electrons do not attract.
Proposed that electrons moved so fast (speed of light) that they had properties of waves instead of particles.

8. The Study of Waves
Wave: a progressive disturbance propagated from point to point in a medium or space without progress or advance by the points themselves

9. Mechanical Waves
Mechanical wave: a wave that requires an energy source and an elastic material medium to travel.

10. Electromagnetic Waves
a wave that does not require a material medium to travel
it propagates by electric and magnetic fields

11. Transverse Waves
Transverse wave: displacement of the medium is perpendicular to the direction of propagation of the wave.

12. Longitudinal Waves
Longitudinal waves: displacement of the medium is parallel to the direction of propagation of the wave (also called compressional waves)

13. Properties of Waves
Wavelength (λ): The distance between any part of the wave (peak) and the nearest part that is in phase with it (another peak). Standard unit is meters (m).
Frequency (f ): The number of peaks which pass a given point each second. Standard unit is cycles per second which is a hertz (Hz).
Amplitude (A): The maximum displacement of a vibrating particle from its equilibrium position. Standard unit is meters (m).
Velocity (v): the distance a wave travels in a given time. Standard unit is meters per second (m/s).
Energy (E): The energy of a single photon of radiation of a given frequency. Standard unit is joules (J).

14. Equations E=hf c=λf λ = wavelength (m) E = energy (J)
h = Planck’s constant (6.626x 10 −34 J∙s)
f = frequency (Hz)
c=λf
c = speed of light (m/s)
λ = wavelength (m)
f = frequency (Hz)

15. Transverse Wave

16. Werner Heisenberg (1901 – 1976) – German Physicist
“Heisenberg uncertainty principle”:
states that the position and momentum of an electron cannot simultaneously be measured and known.
The arrangement of electrons is discussed in terms of the probability of finding an electron in a certain location.

17. Erwin Schrödinger (1887 – 1961) – Austrian Physicist
Studied the wave nature of the electron and developed mathematical equations to describe their wave-like behavior.
The most probable location of the electrons can be found and the plot of this probability is called the charge cloud model.

18. the four quantum numbers
Principal Quantum Number (n)
Refers to the energy levels in the atom which is the distance from the nucleus and designated with a positive whole number (n=1,2,3,4,5,6,7)
Wavelength of emitted photon is determined by the “energy jump” between energy levels
Energy levels (or shells) means electrons are contained in an area where the probability of finding the electron is 90%
Angular Momentum Quantum Number (l )
Refers to the sublevel (within an energy level) which is one or more “partitions” each with a slightly different energy. (l = 0,1,2,3)
The types of sublevels:
l = 0 (s sublevel)
l = 1 (p sublevel)
l = 2 (d sublevel)
l = 3 (f sublevel)

19. the four quantum numbers
Magnetic Quantum Number (m)
Refers to the orientation in space of the electrons in a sublevel
Can have any whole number value from – l to + l which will tell how many orbitals are in a sublevel.
A maximum of 2 electrons per orbital.
Sublevel # of Orbitals Total # of electrons s p d f

20. the four quantum numbers
Spin Quantum Number (s) or – 1 2
Refers to the spin of the electron.
Pauli Exclusion Principle: if two electrons occupy the same orbital, they must have opposite spin.
Half-filled orbital: _____
Filled orbital: _____

21. Permissible Values of Quantum Numbers for Atomic Orbitals
n l m Orbital # of Subshells #of Orbitals Max # of Electrons s s
1 -1,0, p s
,0, p
2 -2,-1,0,1, d s
,0, p
,-1,0,1, d
3 -3,-2,-1,0,1,2,3 4f

22. Electron Orbitals

23. Distribution of Electrons for Different Elements (Electron Configuration)
Electrons will occupy the lowest energy levels and sublevels first.
Notation:
Principal Quantum Number, n (energy level)
Number of electrons 2p 2 y
Orientation of Orbital
Type of Orbital (sublevel)

24. Aufbau principle
Aufbau principle: an electron occupies the lowest-energy orbital that can receive it

25. Hund’s Rule
Hund’s Rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state

26. Give the long notation electron configuration for:Ca
Ca2+ Ag

27. Give the long notation electron configuration for:
O 1s22s22p4
O2– 1s22s22p6
Ca 1s22s22p63s23p64s2 
Ca2+ 1s22s22p63s23p6
Ag 1s22s22p63s23p64s23d104p65s24d9

28. Give the short notation (noble gas notation)Ca
Ca2+ Ag

29. Give the short notation (noble gas notation)
O [He]2s22p4
O2– [He]2s22p6
Ca [Ne]4s2 
Ca2+ [Ne]
Ag [Kr]5s24d9

30. Orbital Diagrams
Usually only done for the outer shell electrons, which always includes the s and p orbitals.

31. Give the orbital diagrams (arrow notation)2p
O2– ____ ____ ____ ____
  2s

32. Give the orbital diagrams (arrow notation)
Ca ____
  4s
Ca2+ ____ 4s
Ag ____ 5s

33. Shows the outer shell or valence electrons for elements.
Electron Dot Diagrams
Shows the outer shell or valence electrons for elements.

34. Give the electron dot diagramsCa
Ca2+ Ag