CLASS CONCEPTS

1. Introduction to Chemistry

2. The Periodic Table

3. Quantum Numbers

4. Electron Configuration

5. Chemical Families

6. Oxidation Numbers

7. Chemical Formulas

8. Chemical Names

9. Formula Mass

10. Percentage Composition

11. Reaction Types

12. Balancing Equations

13. The Mole Concept

14. Solution Concentration

15. Stoichiometry

16. Kinetic Theory

17. The Gas Laws

18. Enthalpy & Heat

19. Reaction Rates

20. Acids & Bases

21. pH Scale

22. Salts

23. Net Ionic Equations

24. Redox Reactions

25. Organic Chemistry

26. Nuclear Chemistry

4. Electron Configuration and Orbital Diagrams

According to the Copenhagen interpretationWWW of quantum mechanics, the position of a particular electron is not well defined until an act of measurement causes it to be detected. Electrons are able to move from one energy level to another only if there is a vacancy in that level.

There are two ways of representing the electron distribution among the various orbitals of an atom:

1. Electron configuration

An electron configuration consists of the symbol for the occupied subshell with a superscript indicating the number of electrons in the subshell.

The electron configuration for sodium (atomic number 11) is

1s22s22p63s1
  • The large numbers represent the energy level.
     
  • The letters represent the sublevel.
     
  • The superscript numbers indicate the number of electrons in the sublevel.

NOTE: Electron configuration can be read directly from the periodic table. On the table outline at the right, helium has been moved into the s-block next to hydrogen, (because its 2 electrons are in the 1s orbital). Each square on the table represents an electron. The shape of the table identifies the subshells (the s, p, d, and f blocks) and each horizontal row represents a given value for n.

The square representing sodium has been colored in on the table above. Click on the table to see how counting the squares will give sodium's electron configuration.

There are two ways to check an electron configuration:

  • The last notation in the electron configuration represents the location of the element on the periodic table.

    Example: the 3s1 in the electron configuration for sodium above indicates its location as the first square in the s sublevel on the third row of the periodic table.

  • The total of the superscripts in an electron configuration equals the atomic number of the element.

    Example: the total of the superscripts in the electron configuration for sodium above is 11, its atomic number.

 
Electron configuration

 
2. Orbital diagram

An orbital diagram consists of a box representing each orbital and a half arrow representing each electron.

The orbital diagram below is for sodium (atomic number 11)

A half arrow pointing up, , represents an electron with ms = + 1/2

A half arrow pointing down,, represents an electron with ms = −1/2

The spin of an electron is often referred to as either "spin up" or "spin down".

NOTE: Orbital diagrams can be read directly from the periodic table as well as electron configurations.

Electrons are said to be paired when they are in the same orbital. An unpaired electron is one not accompanied by a partner of opposite spin.

Hund's rule states that for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.

Simply stated - electrons will occupy orbitals singly, if possible, and single electrons in the same subshell will all have the same spin.
The orbital diagram below represents nitrogen, with unpaired 2p electrons

Electrons arranged in this way are said to have parallel spins.

Hund's rule is based in part on the fact that electrons repel one another. By occupying different orbitals, the electrons remain as far as possible from one another, minimizing electron-electron repulsions.

 
Condensed Configurations

For large atoms, showing all the electrons with an electron configuration or orbital diagram can become quite complex. Since it is the outermost electrons that are largely responsible for chemical behavior, we can condense the electron configuration and orbital diagram to focus on those electrons.

Outer-shell electrons, those involved in chemical bonding, are called valence electrons. Those electrons below the outer shell, inner-shell electrons, are usually referred to as core electrons.
The electron configuration and orbital diagram can be condensed by beginning with the nearest (before the atom) noble gas symbol in brackets to represent the core electrons, then showing the valence electrons as usual.

Sodium's complete electron configuration is

1s22s22p63s1

The same electron configuration in condensed form becomes

[Ne]3s1

The complete orbital diagram for sodium is

The same orbital diagram in condensed form becomes

 

Orbital Overlaps

The s-block and p-block elements on the periodic table together are the main-group elements. For most atoms, it is the s and p electrons that act as valence electrons.

The transition metals begin with scandium (atomic number 21). This is the first element with d orbitals. While it would appear these are 4d from their location on the periodic table, they are actually 3d. Remember from the energy diagram for many-electron atom orbitals earlier, the 4s subshell actually has a lower energy than the 3d.

NOTE: The "d-block" electrons are actually n − 1 from their location on the periodic table.

The lanthanide series begins with lanthanum (atomic number 57). This is the first element with f orbitals. As with elements in the d-block, there is also an overlap in orbital energies in the f-block elements. Notice that the f-block elements are a part of the sixth row on the periodic table. However, because of the energy overlaps, the first f orbitals are 4f.

The actinide series begins with actinium (atomic number 89). While this series is part of the seventh row on the periodic table, the orbitals are 5f.

NOTE: The "f-block" electrons are actually n − 2 from their location on the periodic table.
 
Electron Configurations of Ions

When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest principal quantum number, n.

Removing one electron from a lithium atom produces:

Li (1s22s1) ⇒ Li+ (1s2)

Removing two electrons from an iron atom produces:

Fe ([Ar]3d64s2) ⇒ Fe+2 ([Ar]3d6)

Removing a third electron from iron produces:

Fe+2 ([Ar]3d6) ⇒ Fe+3 ([Ar]3d5)

When electrons are added to an atom to form an anion, they are added to the empty or partially filled orbital having the lowest value of n.

Adding one electron to a fluorine atom produces:

F (1s22s22p5) ⇒ F (1s22s22p6)
 

Anomalous Electron Configurations

Certain atoms appear to violate the orbital-filling rules:

  • Chromium's electron configuration is [Ar]3d54s1 rather than the expected [Ar]3d44s2
     
  • Copper's electron configuration is [Ar]3d104s1 rather than the expected [Ar]3d94s2
     
  • There are other examples in the heavier d-block metals and the f-block metals.
This anomalous behaviorWWW occurs when there are enough electrons to produce a half-filled set of degenerate orbitals (as in chromium) or to completely fill a d or f subshell (as in copper).

Although an interesting part of the atomic theory, these anomalies are not of great chemical significance.

 
Glass Tubing

Electron Distribution