1. Atomic Structure and Periodicicity
Chapter 7
Atomic Structure and Periodicicity

2. Terms Atomic Number Groups Periods Quantum Numbers Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Atomic Radius
Ionization Energy
Electronegativity

3. Atomic Number, Z
The atomic number of an element is the number of protons in the nucleus of an atom of that element.
Atoms of an element that are chemically alike but differ in mass are called isotopes of the element.
Atoms of the same element (Z) but different mass number (A).

4. Number of neutrons = mass number – atomic number
The mass of a neutron is almost the same as the mass of a proton.
The sum of the protons and neutrons in the nucleus is the mass number of that particular atom.
Number of neutrons = mass number – atomic number

5. ATOMIC STRUCTURE Particle Charge Mass proton + charge 1 neutron
No charge 1
electron
- charge
Too small

6. n describes the electron energy shell.
The Principal
Quantum Number - n
n describes the electron
energy shell.

7. The farther an electron is from the nucleus, the greater the numeric
The Principal
Quantum Number - n
The farther an electron
is from the nucleus,
the greater the numeric
value for n.

8. For the known elements, n = 1 through 7
The Principal
Quantum Number - n
For the known elements,
n = 1 through 7

9. n = 1 through 7

10. The maximum number of electrons in n is 2n2 The Principal
Quantum Number - n
The maximum number
of electrons in n is 2n2

11. Energy level 1 can have a maximum of ____ electrons.
The maximum number
of electrons in n is 2n2
Energy level 1 can
have a maximum of
____ electrons.

12. 2 The maximum number of electrons in n is 2n2 Energy level 1 can
have a maximum of
____ electrons. 2

13. The maximum number
of electrons in n is 2n2
Energy level 2 can
have a maximum of
____ electrons.

14. 8 The maximum number of electrons in n is 2n2 Energy level 2 can
have a maximum of
____ electrons. 8

15. The maximum number of electrons in n is 2n2 Energy level 3 can
have a maximum of
____ electrons.

16. 18 The maximum number of electrons in n is 2n2 Energy level 3 can
have a maximum of
____ electrons. 18

17. The maximum number
of electrons in n is 2n2
Energy level 4 can
have a maximum of
____ electrons.

18. 32 The maximum number of electrons in n is 2n2 Energy level 4 can
have a maximum of
____ electrons. 32

19. The Angular
Quantum Number - l
l describes the electron
energy subshell or
shape.

20. The Angular
Quantum Number - l
Each electron energy
level n, has n-1 sublevels.

21. Energy level 1 has ____ sublevels. Each electron energy
level n, has n sublevels.
Energy level 1 has
____ sublevels.

22. Energy level 1 has ____ sublevels. Each electron energy
level n, has n sublevels.
Energy level 1 has
____ sublevels.

23. Energy level 2 has ____ sublevels. Each electron energy
level n, has n sublevels.
Energy level 2 has
____ sublevels.

24. Energy level 2 has ____ sublevels. 1 Each electron energy
level n, has n sublevels.
Energy level 2 has
____ sublevels. 1

25. Energy level 3 has ____ sublevels. Each electron energy
level n, has n sublevels.
Energy level 3 has
____ sublevels.

26. Energy level 3 has ____ sublevels. 2 Each electron energy
level n, has n sublevels.
Energy level 3 has
____ sublevels. 2

27. Energy level 4 has ____ sublevels. Each electron energy
level n, has n sublevels.
Energy level 4 has
____ sublevels.

28. Energy level 4 has ____ sublevels. 3 Each electron energy
level n, has n sublevels.
Energy level 4 has
____ sublevels. 3

29. The Angular
Quantum Number - l
Sublevels in the known
elements are represented
by the letters s, p, d, and f.

30. s p d f

31. The Magnetic
Quantum Number - m
m describes the orbital
within the sublevel or
orientation.

32. The Magnetic
Quantum Number - m
s sublevel = 1 orbital
p sublevel = 3 orbitals
d sublevel = 5 orbitals
f sublevel = 7 orbitals

33. The Magnetic
Quantum Number - m
Each orbital can hold
a maximum of 2 electrons.

34. Pairs of squares on the
periodic table are orbitals.

35. The s-block has 1 orbital.

36. The p-block has 3 orbitals.

37. The d-block has 5 orbitals.

38. The f-block has 7 orbitals.

39. The Spin
Quantum Number - s
s describes the spin
of the electron.

40. The Spin
Quantum Number - s
Electrons in the same
orbital must have
opposite spins.

41. The Spin
Quantum Number - s
Possible spins are
clockwise, + 1/2, and
counterclockwise, - 1/2.

42. Shape of 1s Orbital

43. Shape of 2p Orbital

44. Shape of 3d Orbitals

45. Quantum Numbers and Atomic Orbitals
An atomic orbital is specified by three quantum numbers.
n the principal quantum number - a positive integer
l the angular momentum quantum number - an integer from 0 to n-1
ml the magnetic quantum number - an integer from -l to +l
The three quantum numbers are actually giving the energy of the electron in the orbital and a fourth q.n. is needed to describe a property of electrons called spin. The spin can be … The spin q.n., ms can be + ½ or - ½ .

46. Written Notes for sections 7.1-7.5 Use Handout
Stop
Written Notes for sections
Use Handout

47. Electron Configuration
Three rules apply
The aufbau principle/rule
The Pauli exclusion principle/rule
Hund’s rule
These all define how an electron is arranged.

48. Aufbau Principle
States: each electron occupies the lowest energy orbital available.
So, the first step in determining the ground-state is learning the sequence of atomic orbital's from lowest to highest energy levels.
The aufbau Diagram

49. The aufbau Diagram

50. The aufbau Diagram: Features
In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, and f.

51. Noble Gas Notation
If the element is past the first noble gas of the Periodic Table you can use noble gas notation to write the electron configuration.
Sodium is past both He and Ne therefore we can write its electron configuration like…
[Ne] 3s1
This is saying you go to Neon then to the 1 element on the 3rd row.

52. Electron Configurations…
…are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how:
Aufbau principle - electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies – follow the diagram!
Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

53. Electron Configurations
Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to.
Let’s write the electron configuration diagram for Phosphorus
We need to account for all 15 electrons in phosphorus

54. Orbitals fill in an order
Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation.

55. Atomic Radius Move down a group: Increases
Move Left to Right: Decreases

56. Atomic Radii for Selected Atoms

57. Group trends: Atomic RadiusH Li Na Mg Al Si P S Cl Ar K Rb 57

58. Ionization Energy To form an ion you need ENERGY!
The energy needed to overcome the attraction between the positive energy in nucleus and the charge of the electron. This is the IONIZATION ENERGY
I.E. can be thought of as an indication of how strongly an atom’s nucleus holds onto the valance electrons.

59. Ionization Energy
A high I.E. means the atom has a strong hold on its valance electrons.
A low I.E. means the atom loses its valance electrons easily.
It is possible for atoms to lose of gain more than one electron. The amount of energy required to do this is called the second ionization energy. So the third is…

60. Ionization Energy First ionization energy:
increases from left to right across a period;
decreases going down a group.

61. Trends in Ionization Energies for the Representative Elements

62. Electronegativity
Electronegativity indicates the relative ability of its atom to attract electrons.
As you move from left to right in a period the electronegativity increases.
As you move down in a group the electronegativity decreases.
The lower left side of the periodic table has the lowest and the upper right has the highest!!

63. Electron Affinity
The energy change associated with the addition of an electron to a gaseous atom.
X(g) + e  X(g)
These values tend to be exothermic (energy released).
Adding an electron to an atom causes it to give off energy. So the value for electron affinity will carry a negative sign.

64. Trends in Electronegativity.

65. Summary
Quantum Numbers: Principal QN, Angular momentum QN, Magnetic QN, an Electron Spin QN.
Orbital Shapes and Energies: s, p, and d orbitals.
Pauli Exclusion Principle:
Aufbau Principle:
Hund’s Rule:
Periodic Table:
Periodic Trends:
Atomic Radius
Ionization Energy:
Electron Affinity:

66. STOP
End of Chapter 7