1. Orbital Diagrams
Total
Element Electrons
H 1
He 2
Li 3
Be 4 1s 2s

2. Orbital Diagrams
Total
Element Electrons
B 5
C 6
N 7
O 8
F 9 1s 2s 2p

3. Orbital Diagrams For Ne (10 e-)2p 3s
Filling of the 2p subshell is complete at neon.
The outermost shell (n = 2) contains an octet (8) of electrons.

4. Orbital Diagrams Every noble gas has a complete outer shell. He:
2 electrons in the outer shell
All other noble gases
an octet of electrons in the outer shell
This configuration is exceptionally stable.
Responsible for the unreactive nature of the noble gases.
Elements that ionize easily do so in a way that gives them the same octet of electrons

5. Orbital Diagrams Neon core For Sodium (Na) 11 electrons
1 more electron than the noble gas neon 1s 2s 2p 3s
Neon core

6. Orbital Diagrams
Electrons that are in shells that are not occupied by the nearest noble gas element are called valence electrons.
For Na, the 3s electrons are valence electrons
Valence electrons:
Used to form chemical bonds
The ones lost to form cations

7. Orbital Diagrams Example: Draw the orbital diagram for potassium.
Know: Z = atomic number =

8. Orbital Diagrams
Example: Draw the orbital diagram for Ti.

9. Orbital Diagrams A useful periodic trend:
For atoms in the 1st period, the electrons are being added to the n=1 shell.
For atoms in the 2nd period, the last electrons are being added to the n=2 shell.
Etc.

10. Orbital Diagrams S block p block d block f block
Another useful periodic trend:
p block
S block
d block
f block

11. Electron Configuration
Drawing orbital diagrams gives information not only about the orbitals that are/have been filled but also about the number of unpaired electrons.
Orbital diagrams can be cumbersome!!

12. Electron Configuration
A short-hand notation is commonly used in place of orbital diagrams to describe the electron configuration of an atom.
Electron configuration:
a particular arrangement of electrons in the orbitals of an atom

13. Electron Configuration
The electron configuration tells the number of electrons found in each subshell.
If there are three electrons in a 2p subshell, we would write:
2p3
where the superscript (3) indicates the number of electrons in that subshell

14. Electron Configuration
The orbital diagram for an O atom: 1s 2s 2p 3s
The electron configuration for an O atom:
1s22s22p4

15. Electron Configuration
To determine the electron configuration of an atom (or ion) without first writing the orbital diagram:
determine the number of electrons present
add electrons to each subshell in the correct order starting with the lowest energy subshell until all electrons have been added
use the “diagonal” diagram to help determine relative energy (i.e. filling order)

16. Electron Configuration
Example: Write the electron configuration of a Mn atom (Z = 25).

17. Electron Configuration
Example: Write the electron configuration of an O2- ion (Z = 8).

18. Electron Configuration
Example: Write the electron configuration of a krypton atom (Z = 36).
1s22s22p63s23p64s23d104p6
This is the Kr “core”  [Kr]
The noble gas “core” can be used to write the electron configuration of an element using core notation:
noble gas “core”
valence electrons

19. Electron Configuration
To write the electron configuration using core notation:
find the noble gas that comes before the atom
determine how many additional electrons must be added beyond what the noble gas has
Atomic number of atom minus atomic number of noble gas

20. Electron Configuration
To write the electron configuration using core notation (cont):
determine the period number of the element
this determines the value of n of the s subshell to start with when adding extra electrons
add electrons starting in the “n” s subshell

21. Electron Configuration
Example: Write the core electron configuration of Sr.

22. Electron Configuration
Example: Write the core electron configuration of Br.

23. Electron Configuration - Anomalies
Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

24. Electron Configuration - Anomalies
For instance, the electron configuration for chromium is
[Ar] 4s1 3d5
rather than the expected
[Ar] 4s2 3d4.

25. Isoelectronic Series
When atoms ionize, they form ions with the same number of electrons as the nearest (in atomic number) noble gas.
Na = 1s22s22p63s1 = [Ne]3s1
Na+ = 1s22s22p = [Ne]
Cl = 1s22s22p63s23p5 = [Ne]3s23p5
Cl- = 1s22s22p63s23p6 = [Ar]

26. Isoelectronic Series N3- (10 e-): 1s22s22p6 = [Ne] N (7 e-): 1s22s22p3
O (8 e-): 1s22s22p4
F (9 e-): 1s22s22p5
N3- (10 e-): 1s22s22p6 = [Ne]
O2- (10 e-): 1s22s22p6 = [Ne]
F- (10 e-): 1s22s22p6 = [Ne]

27. Isoelectronic Series Na+ (10 e-): 1s22s22p6 = [Ne]
Na (11 e-): 1s22s22p63s1
Mg (12 e-): 1s22s22p63s2
Al (13 e-): 1s22s22p63s23p1
Na+ (10 e-): 1s22s22p6 = [Ne]
Mg2+ (10 e-): 1s22s22p6 = [Ne]
Al3+ (10 e-): 1s22s22p6 = [Ne]

28. Ions of the highlighted elements are isoelectronic with Ne.Be B C N O F Ne Na Mg Al Si P S Cl Ar 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

29. Isoelectronic Series
Isoelectronic: having the same number of electrons
N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+ form an isoelectronic series.
A group of atoms or ions that all contain the same number of electrons

30. Isoelectronic Series Examples of isoelectronic series:
N3-, O2-, F-, Ne, Na+, Mg2+, Al3+
Se2-, Br-, Kr, Rb+, Sr2+, Y3+
Cr, Fe2+, and Co3+

31. Periodic Properties of Elements
Chemical and physical properties of the elements vary with their position in the periodic table.
Atomic size
Size of Atom vs. Ion
Size of Ions in Isoelectronic series
Ionization energy
Electron affinity
Metallic character

32. Periodic Properties--Atomic Size
The relative size (radius) of an atom of an element can be predicted by its position in the periodic table.
Trends
Within a group (column), the atomic radius tends to increase from top to bottom
Within a period (row), the atomic radius tends to decrease as we move from left to right

33. Periodic Properties--Atomic Size
Periodic Table
Increasing size
Increasing size
Lower “lefter” larger

34. Periodic Properties--Atomic Size
Example: Which element would have the larger atomic radius, Ar or Br?

35. Periodic Properties – Atom vs. Ion Size
Trends to know:
Cations (+) are smaller than their parent atoms.
Electrons are removed from the outer shell.
Anions (-) are larger than their parent atoms.
Electron-electron repulsion causes the electrons to spread out more in space.

36. Periodic Properties – Ion Size
Trends to know:
For ions in the same group (same charge), size increases from top to bottom.
Same trend as for the size of parent atoms
I- is larger than F-

37. Periodic Properties – Ion Size
Trends to know:
For an isoelectronic series of ions, the size decreases with increasing atomic number.
Na+ is smaller than O2-

38. Periodic Properties Ionization Energy
The ease with which an electron can be removed from an atom to form an ion is an important indicator of its chemical behavior.
Ionization energy: the minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion.
Formation of cation (+) or more positively charged cation

39. Periodic Properties Ionization Energy
Ionization of Gaseous Sodium:
Na (g)  Na+ (g) + e-
As the ionization energy increases, it becomes harder to remove an electron.

40. Periodic Properties Ionization Energy
Within each row, the ionization energy increases from left to right
Its easiest to remove an electron from an alkali metal and hardest to remove one from a noble gas.
Within each group, the ionization energy generally decreases from top to bottom
It’s easier to ionize K than Li.

41. Periodic Properties Ionization Energy
Example: Which element has the higher ionization energy, Br or Ca? Which one will lose an electron easier?

42. Periodic Properties Electron Affinity
The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity.
Cl (g) + e-  Cl- (g)
The electron affinity becomes increasingly negative as the attraction between an atom and an electron increases
more negative electron affinity = more likely to gain an electron and form an anion

43. Periodic Properties Electron Affinity
Trends:
Halogens have the most negative electron affinities.
Electron affinities become increasing negative moving from the left toward the halogens.
Electron affinities do not change significantly within a group.

44. Periodic Properties Electron Affinity
Trends:
Noble gases will not accept another electron.
To do so would require adding an electron to a new electron shell (significantly higher in energy)

45. Periodic Properties Metallic Character
Metals:
shiny luster
malleable and ductile
good conductors of heat and electricity
form cations
Metallic character
increases from top to bottom
Increases from right to left