1. Unit 5 Electrons in Atoms
Chemistry I
Mr. Patel
SWHS

2. Topic Outline Continue Learning Major Ions Atomic Models (5.1)
Electron Configurations (5.2)
Light and Quantum Mechanics (5.3)
Lewis Dot Structures (7.1)

3. Atomic Models Democritus’s Model Thomson’s Plum Pudding Model
Rutherford’s Model
Electrons travel in orbit around nucleus
Could NOT explain chemical properties of elements
Need a model for electrons

4. Bohr Model
Niels Bohr – electrons is found in a specific orbit around nucleus
Each orbit has a specific energy = energy level
The further away from the nucleus, the higher the energy

5. Bohr Model

6. Bohr Model An electron can move between levels
Can not be between levels
Think of a ladder
An electron must gain or lose energy to change levels
A quantum of energy – energy to move to another level

7. Bohr Model More energy between levels when closer to nucleus
Less energy between levels when farther
Energy levels get closer together

8. Bohr Model Ground state – lowest energy state for an electron
Excited state – any higher energy state

10. Electron Excitation

11. Bohr Model Each ring on a Bohr Model is labeled as “n”
n must be a whole number
n=1, n=2, n=3, etc. (period number)
Each ring (n) can hold a specific number of electrons
n=1 2 electrons
n=2 8 electrons
n=3 18 electrons
n=4 32 electrons

13. Drawing Bohr (Rutherford) Diagrams

14. Ex: Draw the Bohr Diagram for Hydrogen.

15. Ex: Draw the Bohr Diagram for Neon.

16. Ex: Draw the PEL Diagram for Calcium.

17. Ex: Draw the PEL Diagram for Argon.

18. Bohr Model Correct: Electrons have energy levels and can move
Incorrect: Electrons move in orbits
Matter has a Wave-Particle Duality

19. Dual Nature of Electrons
Electrons as Particles
Electrons as Waves
Photoelectric Effect
Young's Double Slit Experiment

20. Modern Theory
Rutherford and Bohr based models on behavior of large objects
Small objects behave differently – quantum mechanics
Schrödinger Equation solutions  quantum mechanical model of the atom

21. Schrödinger Equation

22. The Cat – A Thought Experiment
Schrodinger Cat 1
Schrodinger Cat 2

23. Quantum Mechanical Model
Determines the allowed energies of the electrons
The probability of where an electron is – electrons housed in electron clouds

24. Atomic Orbitals
Region in space where there is a high probability of finding an electron
Principal quantum number (n) – energy level
think of the ring labels of the Bohr model
Each energy level can be made up of sublevels – orbitals of similar energy but different shapes

25. s orbital
Shape: sphere

26. p orbital
Shape: Dumbbell

27. d orbital
Shape: clover (mostly)

28. f orbital
Shape: multiple clover

29. Atomic Orbitals

30. Electron Configurations
Electrons found in orbitals
Electron configuration – ways in which various electrons are arranged in orbitals
4 orbitals: s (2 electrons), p (6 electrons) d (10 electrons), f (14 electrons)

31. Three Rules to find Elec. Config
Aufbau Principle
Electrons occupy orbitals of lower energy first
For same n, low to high energy: s, p, d, f

32. Three Rules to find Elec. Config
Pauli Exclusion Principle
Each atomic orbital can have at most 2 electron
Each electron in an orbital must have opposite spins
2 spins: spin up or spin down
How we draw: 1 electron in s orbital: ____
2 electrons in s orbital: ____
We use arrow with “half head”

33. Three Rules to find Elec. Config
Hund’s Rule
Electrons occupy orbitals to maximize spin
For same n, place electrons spin up first then pair them with spin down
1 electron in p orbital ____ ____ ____
2 electrons in p orbital ____ ____ ____
3 electrons in p orbital ____ ____ ____
4 electrons in p orbital ____ ____ ____
5 electrons in p orbital ____ ____ ____
6 electrons in p orbital ____ ____ ____

34. Orbital Blocks on PT s-block: Groups 1A and 2A (exception: He)
p-block: Groups 3A-8A (exception: He)
d-block: transition metals
f-block: inner transition metals
Remember, the period number is n = principal energy level

35. Orbital Blocks on PT

36. How to write electron configuration
Ex: What is the electron configuration for O?
O = oxygen, atomic number 8 = 8 electrons
Draw spaces:
____ ____ ____ ____ ____ s s p
Fill spaces according to rules:
Write: 1s22s22p4

37. How to write electron configuration
Ex: What is the electron configuration for C?
C = carbon, atomic number 6 = 6 electrons
Draw spaces:
____ ____ ____ ____ ____ s s p
Fill spaces according to rules:
Write: 1s22s22p2

38. 3 ways to write electron configurations
Using boxes and arrows
____ ____ ____ ____ ____ s s p
Long EC: Cl: 1s22s22p63s23p5
Short EC: Cl: [Ne] 3s23p5
Put last noble gas in brackets and write electrons from there

39. Writing EC This is much easier than it looks.
Simply, start at hydrogen and walk to the desired element counting all the elements you pass

40. Ex. Write EC (all three ways) for Boron.

41. Ex. Write EC (all three ways) for Mg.

42. Ex. Write EC (all three ways) for V.

43. Ex. Write EC (long and short) for Fr.

44. A Look Back… So far we have covered (and mastered):
Evolution of the Atomic Model
Democritus, Thomson, Rutherford, Bohr, QM
Bohr Model and Bohr Diagram
Quantum Mechanical Model and Orbitals
Rules of Electron Configuration
Writing Electron Configurations

45. Electrons Chemical reactions are the breaking and forming of bonds
There are two types of bonds: covalent and ionic (and metallic) = next unit
Bonding involves the movement of electrons

46. Valence Electrons
Valence Electron: electrons in the highest occupied level
These are the electrons that participate in bonding!!!

47. Valence Electrons
You do not have to draw a Bohr Model every time you need to determine the VE’s
The valence electrons (valency) for an atom is the same as the group number
Note: In general, transition metals have two valence electrons.

48. Determine the valence electrons for:Ca Be O Si H Ne Ar 2 6 4 1 10

49. N B Lewis Dot Structures Show bonding electrons
These structures show only valence electrons.
How to draw:
Write Symbol for element
Determine group number
Place that many (group number) dots around symbol N B

50. Lewis Dot Structures

51. Lewis Structures (Future)

52. Draw the Lewis Dot Structure for: Cs Al Ge Br

53. Practice Time!!!

54. The Octet Rule Remember that Noble Gases were very stable
They all have 8 valence electrons (2 for He)
FULL outer shell of electrons
Every element will try to become like a noble gas
The Octet Rule – atoms will try to have a full outer shell (= 8 electrons) when bonding

55. Cations Metals tend to lose electrons to have a full outer shell
Cation – positively charge ion
Results from metals losing electrons
Naming: element name + ion
Ex: Na = Sodium but Na1+ = sodium ion

56. Anions
Nonmetals/Metalloids tend to gain electrons to have a full outer shell
Anion– negatively charged ion
Results from nonmetals gaining electrons
Naming: element name with –ide ending + ion
Ex: Br = Bromine but Br1- = bromide ion

57. Cations/Anions
To determine the charge of an element’s ion, look at the group/column that it is in
Group 1: 1+
Group 2: 2+
Group 3: 3+
Group 4: 0
Group 5: 3-
Group 6: 2-
Group 7: 1-

58. Lewis Dot Structures for Ions
Draw the normal Lewis Dot structure for the neutral element
Add electrons if gained or remove electrons if lost
Place the appropriate charge 3- N N
Nitride ion:
Nitrogen:

59. Draw the Lewis Dot Structure for Phosphorus.
Will this element for a cation or anion?
What charge will it have?
What will be the name of the ion?
What noble gas is it similar to?
Draw the Lewis Dot Structure for the ion.

60. Draw the Lewis Dot Structure for Barium.
Will this element for a cation or anion?
What charge will it have?
What will be the name of the ion?
What noble gas is it similar to?
Draw the Lewis Dot Structure for the ion.

61. Light – A Wave Newton tried to prove light to be a particle
However, experimental data showed that light was actually behaving as a wave
The study of light led to the quantum mechanical model of the atom

62. Properties of Waves Wavelength (λ) Frequency (ν) Amplitude “lambda”
Distance between crests
Frequency (ν)
“nu”
Cycles per second
Hertz (Hz)
Amplitude
Height from zero to the crest
Crest
Trough

63. Properties of Waves The Wave Equation: c = λ∙ν
c = speed of light = 3 x 108 m/s
λ = wavelength – must be in meters
ν = frequency – in Hertz
Energy of Light: E = h∙ν
E = energy – in Joules
h = Planck’s constant = x J∙s

64. Electromagnetic Spectrum
When light passes through a prism, it is separated into different frequencies
You need to know:
Name and order of each regions
Order based on wavelength
Order based on frequency
Order based on energy
Details of the Visible Region

65. Electromagnetic Spectrum EM Spectrum Song
Long Wavelength Low frequency Low Energy
Short Wavelength High frequency High Energy

66. Review: Properties of Waves
The Wave Equation: c = λ∙ν
c = speed of light = 3 x 108 m/s
λ = wavelength – must be in meters
ν = frequency – in Hertz
Energy of Light: E = h∙ν
E = energy – in Joules
h = Planck’s constant = x J∙s

67. Equations: c = λ∙ν E = h∙ν
c = speed of light = 3 x 108 m/s
h = Planck’s constant = x J∙s
Ex: What is the frequency of a wave with a wavelength of 3.68 x 10-9 m? The energy?

68. Equations: c = λ∙ν E = h∙ν
c = speed of light = 3 x 108 m/s
h = Planck’s constant = x J∙s
Ex: What is the frequency of a wave with a wavelength of 700 nm? The Energy?

69. Practice!!!

70. Atomic Spectra
When atoms absorb energy, they move into higher energy levels
These electrons then return back to a lower level and release energy as light
Each atom releases light in a special way

71. Atomic Spectra
Atomic Emission Spectrum – the frequencies of light released by an element split into separate discrete lines (unlike light)

72. Hydrogen Spectrum
Balmer Series
Visible Region
Lyman Series
UV Region

73. Light – A Particle Light deserves a quantum mechanical treatment
Light also behaves as a particle and a wave (Particle-Wave Duality)
Light particles called photons – packets or quanta of light (E = h∙ν)

74. deBroglie Relation
Louis deBroglie determined that all matter that is moving can be considered as waves
Large object have such a small wavelength that it can not be observed
His math showed that as mass decreases, the wave function becomes more important (λ=h/mv)

75. Heisenberg Uncertainty principle
Heisenberg Uncertainty Principle – it is impossible to know the exact velocity and position for a particle at the same time
Heisenberg Uncertainty principle