1. Electrons in Atoms

2. In this chapter…
Scientists pursue an understanding of how electrons are arranged within atoms
Electron arrangement plays a role in chemical behavior

3. Early 1900’s- scientists observed that certain elements emit visible light when heated in a flame

4. Wave Nature of Light
Electromagnetic Radiation- form of energy that exhibits wavelike behavior as it travels through space

5. Vocabulary to know..
Wavelength- shortest distance between equivalent points on a wave
Symbol- λ
Unit- meters, centimeters, or nanometers (1 nm= 1x10-9m)
Frequency- the number of waves that pass a given point per second
Symbol- ν
Unit- Hertz (SI Unit)= (1/s)= (s-1) cycle per second

6. How are they related? C= λν c= 3.00x108 m/s
ALL electromagnetic waves,
including visible light,
travel at the speed of light
c= 3.00x108 m/s
Wavelength must be in meters!
C= λν

7. Electromagnetic Spectrum
Encompasses all forms of electromagnetic radiation
The only differences in the types of radiation being their wavelengths and frequencies

8. ROYGBIV

9. As the wavelength increases, the frequency decreases.
As the frequency increases, the energy increases.

11. Calculations 
1. Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 3.44 x109 Hz?
C= λν
C= 3.00x108 m/s
ν = 3.44 x109 Hz
λ = ???

12. 3.00x108 = λ(3.44x109 Hz)
λ= 8.72 x10-2 m

13. 2. Yellow light has a wavelength of 589 nm. What is the frequency. 5
2. Yellow light has a wavelength of 589 nm. What is the frequency? 5.09x1014 Hz

14. Particle Nature of Light (Honors)
Quantum Concept
Explained why colors of heated matter correspond to different frequencies and wavelengths
Max Plank- “matter can gain or lose only in small, specific amounts called quanta”
Quantum- the minimum amount of energy that can be gained or lost by an atom

15. Energy of a quantum is related to the frequency of the emitted radiation by the equation
Equantum= hv
E= energy
h = Plank’s Constant (6.63x10-34Js)
v= frequency
Joule (J)= SI unit for energy

16. Photon- a particle of EM radiation with no mass that carries a quantum of energy
Ephoton= hv

17. Example
Calculate the quantum of energy that an object can absorb from light with a wavelength of 477 nm.
4.17x10-19 J

18. Atomic Emission Spectra
Set of frequencies of the electromagnetic waves emitted by atoms of the element
Example- The light of neon sign is produced by passing electricity through a tube filled with neon gas. Neon atoms release energy by emitting light.

19. An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element

20. Bohr Model of the Atom
Proposed that the hydrogen atom has only certain allowable energy states
Ground State- lowest energy state of an atom
Excited State- higher energy state

21. An electron must absorb energy to move from a lower energy level to a higher level.
Electrons do not stay in the excited state. When the electrons return to lower energy levels, energy is emitted.

22. The Heisenburg uncertainty principle - states that it is impossible to know precisely both the velocity and position of a particle at the same time

23. The Bohr Model
Using the Bohr Model from your packet, what is the wavelength of energy that is emitted when an electron falls from n= 6 to n=3?
wavelength = 1094 nm

24. B) What is the frequency of this radiation. 2
B) What is the frequency of this radiation? 2.75x1014 Hz C) What is the energy of a photon of this radiation? (Honors) 1.82x10-19 J

25. Atomic Orbital- a 3D region around the nucleus describing the electron’s probable location

26. Atomic Orbitals Energy Levels (n)- the major energy levels of an atom
Ex: n = 1 energy level closest to the nucleus
Energy level → sublevel → orbital
Every orbital can hold up to 2 e-

27. Sublevels are represented by the letters s, p, d, f
lowest energy highest energy

28. First 4 Principal Energy Levels
Sublevel
Orbital
Number of Electrons 1 s 1 2 s 1 2 2 p 3 6
(8 total e-) s 1 2 p 3 6 3 d 5 10
(18 total e-) s 1 2 p 3 6 4 d 5 10 f 7 14
(32 total e-)
2n2 = maximum # of electrons in energy level

29. Electron Arrangement in Atoms
Electron Configurations- the arrangement of electrons in an atom

30. Aufbau Principle : Electrons enter orbitals from lowest to highest energy

31. Writing Electron Configurations
H (1e-)
1s1
energy level sublevel # e-

32. Writing Electron Configurations
He (2e-)
1s2
Li (3e-)
1s2 2s1
Be (4e-)
1s22s2

33. Writing Electron Configurations
B (5e-)
1s22s22p1
C (6e-)
1s22s22p2
Ne (10e-)
1s22s22p6

34. Writing Electron Configurations
Na (11e-)
1s22s22p63s1
Si (14e-)
1s22s22p63s23p2
Cl (17e-)
1s22s22p63s23p5

35. s p d f

36. Noble Gas Configuration
Used to shorten electron configurations
Sodium: #11- instead of 1s22s22p63s1 can be shortened to [Ne] 3s1

37. Examples Write the shorthand electron configuration of Mn. [Ar]4s23d5
[Xe]6s24f145d106p5

38. Big Bang – Sheldon
Video

39. Valence Electrons (V.E.)
Electrons in the atom’s outermost energy level
Determine the chemical properties of an element
V.E. are used in forming chemical bonds

40. Examples
Write the electron configuration and give the number of valence e-. Mg
2 valence e- Br
7 valence e- V

41. Exceptions 1. Cu not [Ar]4s23d9 but [Ar]4s13d10 2. Ag [Kr]5s14d10
3.  Au
[Xe]6s14f145d10

42. Exceptions
4. Cr [Ar]4s13d5 5. Mo [Kr]5s14d5

43. Ions Cations (+ ions) –remove e- Anions (- ions) - add e- O: 1s22s22p4
O2- is isoelectronic with ________. Ne

44. Examples Write the electron configuration for: P3-: 1s22s22p63s23p6
Al3+:
 1s22s22p6
Ba2+:
[Xe]

45. Examples Pb: [Xe]6s24f145d106p2 Pb2+: [Xe]6s24f145d10 Pb4+:
[Xe]4f145d10

46. Transition Metals
Fe:
 [Ar]4s23d6
Fe2+:
 [Ar]3d6
Fe3+:
[Ar]3d5

47. Transition Metals Mn: [Ar]4s23d5 Mn2+: [Ar]3d5 Mn4+: [Ar]3d3
What is the highest possible charge for Mn? +7

48. Excited state: e- jumps to higher energy level
Ex: 1s22s22p63p6
Ground state: normal e- configuration (lowest energy)
Ex: 1s22s22p63s23p1
Blue Book: pg 358 # 37-39

49. Orbital Diagrams Use arrows to represent electrons
Use lines to represent orbitals
Every orbital can hold up to 2 e-

50. s ____
p ____ ____ ____
d ____ ____ ____ ____ ____
Lines represent orbitals.

52. Orbital Diagram
Draw the orbital diagram for carbon

53. Hund’s Rule- atoms contain the maximum number of unpaired electrons

54. p s