Chapter 7

Electron Configuration

Quantum Numbers

Each electron in an atom can be described by 4 quantum numbers.

1. Energy Level =n (Levels 1-7)

2. Sublevel = l (6 sublevels:spdfgh)

3. Spin = ms (spin up & spin down)

4. Orientation = ml (orientation in electron cloud)

Pauli Exclusion Principle

States that no 2 electrons in an atom can have the same 4 quantum numbers.

Here’s an example of an energy level diagram for Na:

Electron Configuration

Configures the most stable arrangement of electrons in sublevels and orbitals.

On the periodic table: Groups 1 & 2 are the s orbital Groups 13-18 are the p orbitals Groups 3-12 are the d orbitals Inner transition metals are the f orbitals

Orbitals

The s orbital holds a maximum of 2 electrons

The p orbital holds a maximum of 6 electrons

The d orbital holds a maximum of 10 electrons

The f orbital holds a maximum of 14 electrons

Orbitals Continued

When placing electrons into each orbital you must make sure all orbitals have a single electron first before assigning the second electron.

Once this is done, you may go back and fill in the orbitals with the remaining electrons.

Example 1

Write the electron configuration of N.

1. Find the atomic number of nitrogen.

2. How many electrons does it have?

3. Begin with the lowest energy level and work your way up.

N = 1s2 2s2 2p1 2p1 2p1

Notice that all 3 p orbitals were filled up first instead of just the first one.

Each block or element represents 1 electron. Each time you move to the next element (from left to right) you are adding another electron to the configuration.

On Your Own

EX 1: Electron configuration of Li

EX 2: Electron configuration of Ne

EX 3: Electron configuration of Ti

Noble Gas Configuration

As you progress to higher atomic numbers, it becomes difficult to write out the electron configuration.

A shortcut to electron configuration of higher atomic elements is called noble gas configuration.

EX: Ti EX: Br

Noble Gas Configuration

It is called the noble gas configuration because you will use the electron configuration of the noble gases as an abbreviation.

The electron configuration of the noble gases is used because they fill up all of their outermost energy levels making them stable.

Example 1 Cl

1. Find the atomic number of Cl.

2. How many electrons does it have?

3. What is the nearest noble gas element? (Keep in mind it could have an atomic number less that Cl)

4. Put that noble gas in [brackets] then continue where it has left off.

5. Cl = [Ne] 3s2 3p2 3p2 3p1

On Your Own

Write the noble gas configuration of Zn

Write the noble gas configuration of Pb

Write the noble gas configuration of Au

Exceptions

There are exceptions to the rules of electron configuration.

3 elements violate the electron configuration rule:

1. Cr & Cu violate the rule because their s & d orbitals are so close together.

2. Pd violates the rule for stability reasons.

Orbital Size

As the number or the outermost energy level increase, the size and energy of the orbital increases.

As you move down the columns, the energy of the outermost sublevel increases. The higher the energy level, the farther the outermost electrons are from the nucleus

As the valence electrons gets farther from the nucleus, the s orbital it occupies gets larger.

Elements within the same group share similar valence electron structures but do not have the same number of energy levels and thus does not yield the same amount of energy.

Key Concepts

The position of an element in the periodic table reveals the number of valence electrons.

The outermost valence electrons determine the properties of an element.

Electrons are found only in levels of fixed energy in an atom.

Energy levels have sublevels. Each sublevel can hold a specific number of electrons.

Sublevels can be divided into s, p, d, and f orbitals. Sublevels hold 2, 6, 10, 18 electrons respectively.

The organization of the periodic table reflects the electron configuration of the elements.

The active metals occupy the s block of the periodic table while metals, metalloids, and non metals occupy the p block.

Within a period of the periodic table, the number of valence electrons for main group elements increases from 1 to 8.

The transition elements, groups 3-12, occupy the d block of the periodic table. These elements can have valence electrons in both s & d orbitals.

The lanthanides and actinides, called the inner transition elements, occupy the f block of the periodic table. Their valence electrons are in the s and f sublevels.

References

Phillips, John S., Victor S. Strozak, and Cheryl Wistrom. "Chapter 7: Completing the Model of the Atom." Chemistry: concepts and applications. New York: Glencoe/McGraw-Hill, 2005. 228 - 251.