ch 7electronconfiguration 111105214915 phpapp02

7/30/2019 Ch 7electronconfiguration 111105214915 Phpapp02

1/21

Chapter 7

Electron Configuration


2/21

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)


3/21

Pauli Exclusion Principle

States that no 2 electrons in an atom can

have the same 4 quantum numbers.

Heres an example of an energy leveldiagram for Na:


4/21

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


5/21

Orbitals

The s orbital holds a maximum of 2

electrons

The p orbital holds a maximum of 6electrons

The d orbital holds a maximum of 10

electrons

The f orbital holds a maximum of 14

electrons


6/21

Orbitals Continued

When placing electrons into each orbital

you must make sure all orbitals have a

single electron first before assigning thesecond electron.

Once this is done, you may go back and

fill in the orbitals with the remaining

electrons.


7/21

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.


8/21

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.


9/21

On Your Own

EX 1: Electron configuration of Li

EX 2: Electron configuration of Ne

EX 3: Electron configuration of Ti


10/21

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


11/21

Noble Gas Configuration

It is called the noble gas configuration

because you will use the electron

configuration of the noble gases as anabbreviation.

The electron configuration of the noble

gases is used because they fill up all of

their outermost energy levels makingthem stable.


12/21

Example 1Cl

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


13/21

On Your Own

Write the noble gas configuration of Zn

Write the noble gas configuration of Pb

Write the noble gas configuration of Au


14/21

Exceptions

There are exceptions to the rules of

electron configuration.

3 elements violate the electronconfiguration rule:

1. Cr & Cu violate the rule because their s

& d orbitals are so close together.

2. Pd violates the rule for stability

reasons.


15/21

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 arefrom the nucleus


16/21

As the valence electrons gets farther from the

nucleus, the s orbital it occupies gets larger.

Elements within the same group share similarvalence electron structures but do not have the

same number of energy levels and thus does

not yield the same amount of energy.


17/21

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.


18/21

Energy levels have sublevels.

Each sublevel can hold a specificnumber 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.


19/21

The active metals occupy the s block of

the periodic table while metals,

metalloids, and non metals occupy the pblock.

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 valenceelectrons in both s & d orbitals.


20/21

The lanthanides and actinides, called the

inner transition elements, occupy the fblock of the periodic table. Their valence

electrons are in the s and f sublevels.


21/21

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.

ch 7electronconfiguration 111105214915 phpapp02

Documents