1. Electrons In Atoms

2. Historical View of Light
Early 1900s, light thought to behave as a wave
Later, light was discovered to have particle-like characteristics
Today, light is thought to have both wave and particle properties

3. Wave Description of Light
Electromagnetic Radiation – form of energy that exhibits wavelike behaviors as it travels through space
Travels at the speed of light in air
3.0 E 8 m/s
Ex: visible light, microwaves, x-rays, radio, etc.
Measures wavelength and frequency

4. Wavelength λ λ λ Symbol – λ (lambda)
Definition – distance between equivalent points on a adjacent waves λ λ λ

5. Frequency = 1/s = s-1 Symbol –  (sometimes f)
Definition – number of waves that pass a given point in a one second
Unit: Hertz (Hz)
Ex: 300 Hz = 300 1/s = 300 s-1
= 1/s
= s-1

6. Electromagnetic Waves
All EM waves travels at the speed of light (c)
c = speed of light (3 x 108 m/s)
λ = wavelength
v = frequency (sometimes f)
c = λ
= 3 x 108 m/s

7. Wavelength and Frequency Relationship
Inversely related: as one increases, other decreases
short λ
high frequency
long λ
low frequency

8. Electromagnetic Spectrum
Definition – shows all forms of electromagnetic radiation arranged according to wavelength/frequency

9. Ex problem #1 = λ = = c 4.76 x 10-8 m  c = λ 3 x108 m/s
What is the wavelength of a radiation with a frequency of 6.30 x 1015 Hz?
c = λ c
3 x108 m/s =
4.76 x 10-8 m
λ = = 
6.30 x 1015 s-1

10. Ex problem #2 = = c  = λ 3.108 x 1015 Hz c = λ 3 x108 m/s
What is the frequency of a radiation, which has a wavelength of x 10-8 m?
c = λ c
3 x108 m/s =
 = =
9.653 x 10-8 m λ
3.108 x 1015 Hz

11. Particle Nature of Light
Photoelectric Effect – emission of electrons from metal’s surface when light of specific frequency shines on surface
Radiation emitted from the object is in small, specific amounts called quanta
Quantum – minimum amount of energy that can be gained or lost by an atom
light e-
METAL

12. Ephoton= h E= Energy (measured in Joules)
Photon – particle of light(electron magnetic radiation) having zero mass and carrying a quantum of energy
Photon Energy
Energy of photon is directly proportional to the frequency of radiation
Ephoton= h
E= Energy (measured in Joules)
h = Planck’s constant = x J·s
 = frequency (s-1)

13. Ex problem E = h = (6.626 x 10-34 J·s) (6.582 x 1014 s-1)
#2) How much energy does a photon of light have if its frequency is x 1014 Hz?
E = h
= (6.626 x J·s)
(6.582 x 1014 s-1)
= x J

14. Hydrogen Line-emission Spectrum
When atoms in the gaseous state are heated, their energy increases.
Ground state- state of lowest energy
Excited state- higher potential energy than ground state

15. Hydrogen Line-emission Spectrum
When light shines through a prism, it is separated into a series of specific frequencies and wavelengths of visible light.
The bands of light are part of hydrogen’s line emission spectrum

16. Modern View of Light Wave Theory- waves are forms of energy
Particle Theory- particles are pieces of matter
Modern view- combines both
Einstein’s Theory of Relativity combines both matter and energy into one formula containing the speed of light
E= mc^2

17. Electromagnetic Spectrum

18. Bohr Model of an Atom Ground State – lowest energy state of an atom
Excited State – state when atom gains energy
**pay attention to the electrons**
Bohr Model – shows electron orbit and energy level of an electron

19. Bohr Model E1 E2 E3 E3 > E2 > E1 E1 = lowest energy level
The electron of the hydrogen atom can circle the nucleus in paths called orbits
In an orbit- electron has a definite fixed energy
Lowest energy state- closest to nucleus
Total energy of electron increases as it moves farther from nucleus E1 E2 E3
E3 > E2 > E1
E1 = lowest energy level

20. Ground State to Excited State4 6
in ground state, no energy radiated in excited state, electrons jump to higher energy level electron go from high E level to low E level photon emitted 1) 5 4 2) 3 3
Energy of atom 2 2 3) 1 4) 1

21. Quantum Theory
Describes mathematically the wave properties of electrons and is based on:
Heisenberg Uncertainty Principle: impossible to determine position and velocity of a particle at the same time
Schrödinger Wave Equation: equation that is used to describe electrons as waves

22. Quantum Numbers
Definition- numbers that specify the properties of atomic orbitals and their electrons
1st quantum number:
Main energy level or distance from nucleus
2nd quantum number:
Orbital shape
3rd quantum number:
Orbital orientation

23. Principal Quantum Number (n)
Definition – indicates the energy level (shells) surrounding the nucleus
- use periodic table to tell (look at rows)

24. Principal Quantum Number (n)
n = 1,2,3,…..(values of n are positive)
n=1 (closest to nucleus)
As n increases, the distance of the energy levels from the nucleus increases
2n2 = number of electrons in each level

25. Principal Quantum Number

26. Angular Momentum Quantum Number (l)
Definition – indicates shape of orbital that tells the path of the electrons
In orbitals- different shapes occupy different regions called sublevels or subshells
4 sublevels: s, p, d, f
S= lowest energy

27. s orbital
Shape: electrons travel in a sphere

28. s orbital 3s 1s 2s
The greater the energy level, the bigger the orbital

29. p orbital
Shape: dumbbell or figure 8 shaped

30. D orbital
Shape: double dumbbell

31. Types of Orbitals

32. Magnetic Quantum number (m)
Orientation of an orbital about the nucleus
S= 1
P= 3
D= 5
F=7

33. Spin Quantum Number (s)
+ ½ = clockwise turn
-1/2= counterclockwise turn

34. First level
S orbital= 2 electrons

35. Second level S orbital= P orbital= 1 orbital, 2 electrons
3 orbitals, 6 electrons

36. Third level S orbital= P orbital= D orbital= 1 orbital, 2 electrons
3 orbitals, 6 electrons
D orbital=
5 orbitals= 10 electrons

37. Fourth level S orbital= P orbital= D orbital= F orbital=
1 orbital, 2 electrons
P orbital=
3 orbitals, 6 electrons
D orbital=
5 orbitals= 10 electrons
F orbital=
7 orbitals, 14 electrons

38. Electron Configuration
Definition – arrangement of electrons in an atom
Atoms of different elements have different numbers of electrons
Electrons will assume lowest energy
Where are certain electrons located?

39. Rules Governing Electron Configurations
1) Pauli Exclusion Principle – no 2 electrons in the same atom can have the same set of 4 quantum numbers

40. Rules Governing Electron Configurations
2) Hund’s Rule – if orbitals have equal energy, one e- will go in each orbital before doubling up
all electrons in singly occupied orbitals must have same spin 1 2 3 5 6 4

41. Rules Governing Electron Configurations
3) Aufbau Principle – electrons occupy lowest energy orbital available
- fill up level 1 first, then level 2, etc.

42. Orbital filling table

43. Electron Configurations
Orbital notation:
An unoccupied orbital is represented by a line. An orbital with one electron is represented by a
An orbital containing 2 electrons is represented as
Sodium

44. Orbital Notation 7 n = 2 1s 2s 2p Nitrogen How many electrons?
What energy level is nitrogen on? 7
n = 2 1s 2s 2p

45. Orbital Notation 14 n = 3 1s 2s 2p 3s 3p Silicon How many electrons?
What energy level is silicon on? 14
n = 3 1s 2s 2p 3s 3p

46. Orbital Diagram 29 n = 4 1s 2s 2p 3s 3p 4s 3d Copper
How many electrons?
What energy level is copper on? 29
n = 4 1s 2s 2p 3s 3p 4s 3d

47. Electron Configuration Notation
Definition: the number of electrons in a sublevel is represented by adding superscripts to the sublevel designation
sodium

48. Electron Configuration Notation2 2 4
Oxygen (8 e-)
Sulfur (16 e-)
Vanadium (23 e-)
Zirconium (40 e-) 1s 2s 2p 2 2 6 2 4 1s 2s 2p 3s 3p 2 2 6 2 6 2 3 1s 2s 2p 3s 3p 4s 3d 2 2 6 2 6 2 10 6 2 2 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d

49. Noble Gas Notation Shortcut for electronic notation with a noble gas
Noble gases are used because they have a filled outer shell of electrons (octet)

50. Noble Gas Notation [He] 2s 2p [Ne] 3s 3p [Ar] 4s 3d [Kr] 5s 4d 2 4 2 4
Rule: start from previous noble gas, then write the configuration
Oxygen
Sulfur
Vanadium
Zirconium 2 4
[He] 2s 2p 2 4
[Ne] 3s 3p 2 3
[Ar] 4s 3d 2 2
[Kr] 5s 4d

51. Valence Electrons 1s 2s 2p 5 valence e- 2 valence e- 1s 2s 2p 3s 1s 2s
definition – electrons in outer most energy level
- located in highest s & p orbitals (max 8) N:
Mg:
Se: 2 2 3 1s 2s 2p
5 valence e- 2 2 6 2
2 valence e- 1s 2s 2p 3s 2 2 6 2 6 2 10 4 1s 2s 2p 3s 3p 4s 3d 4p
6 valence e-

52. Electron Dot Structure
Definition – shows number of valence e- by diagram, which are in the highest principal quantum number
Nitrogen (5 v.e.)
Magnesium (2 v.e.)
Selenium (6 v.e.) N Mg Se

53. Irregular confirmations of Cr and Cu
Chromium “promotes” a 4s electron to half
fill its 3d sublevel
Copper “promotes” a 4s electron to FILL
its 3d sublevel