TAKING A CLOSER LOOK AT ELECTRON ORBITALS

Mind Catalyst E- Analogy Share-Out

• Assuming that each pea represents an electron, the bull’s eye represents the nucleus, and the area around the bulls’ eye represents the area around the nucleus, what happens to the probability of finding an electron as the distance from the nucleus increases?

• Where is there zero probability of finding an electron?

• What is the overall shape that the spots collectively made on the target sheets?

• What would this overall shape (in 3-D) correspond to in an atom?

• Did this shape look familiar to you as you completed the webquest for homework?

So, What Did We Cover Last Class? • Louis de Broglie proposed that small particles of

matter (like electrons) exhibit observable wave-like properties

• Furthermore, Werner Heisenberg formulated the Uncertainty Principle that states it is impossible for us to know an electron’s exact position (where it is) and momentum (where it is going)

• As a result, we cannot identify specific orbits that electrons travel in as referenced in Bohr’s atomic model

• Instead, Erwin Schrodinger developed a complex math equation based on the Heisenberg Uncertainty Principle and Louis de Broglie’s work that described the energy and position of electrons in an atom

So, What Did We Cover Last Class? • The solutions to the equation tell us:

• The energy level (n) where the electron “resides”

• n = 1, n = 2, n = 3, etc...

• The regions of space within an atom where an electron is most likely to be found – ORBITALS!

• AND the shape of these orbitals

• In the webquest completed for homework, you explored these four orbitals and their energies!

The s-Orbital

• Spherical shape • Seen in all energy levels • Lowest in energy

• Every energy level has one s orbital • Each s-orbital can hold up to 2 electrons

The p-Orbital

p (z)

p (y)

p (x)

x-axis

y-axis

z-axis

• There are three p orbitals: px, py and pz • All three are dumbbell-shaped • Seen in 2nd energy level and above • Can hold up to 2 electrons PER SUBORBITAL (6 electrons total)

The d-Orbital

• Five clover-shaped orbitals • Can hold up to 2 electrons per suborbital (10 electrons total) • Seen in 3rd energy level and above

The f-Orbital

• Seven equal energy orbitals • Shape is not well-defined • Each suborbital can hold up to 2 electrons (14 electrons total) • Seen in 4th energy level and above

s p d

Arrangement of Atomic Orbitals • The orbitals of an atom are LAYERED!

In Summary… • Each energy level contains orbitals that layer

• n = 1 made of 1s

• n = 2 made of 2s and 2p

• n = 3 made of 3s, 3p, and 3d

• n = 4 made of 4s, 4p, 4d, and 4f

• n = 5 and above – made of s, p, d, and f-orbitals as well

• AND each orbitals are made of sub-orbitals • s-type orbitals are made of 1 sub-orbital

• Holds 2 electrons

• p-type orbitals are made of 3 sub-orbitals • Holds 6 electrons TOTAL (2 electrons per sub-orbital)

• d-type orbitals are made of 5 sub-orbitals • Holds 10 electrons TOTAL (2 electrons per sub-orbital)

• f-type orbitals are made of 7 sub-orbitals • Holds 14 electrons TOTAL (2 electrons per sub-orbital)

HOW DO THE ORBITALS FILL UP WITH ELECTRONS?

An Introduction to Electron Configurations

The Big Questions

• Now that we know how electrons fit into the atomic structure puzzle, we can now answer the questions:

• How are they arranged into orbitals and sub-orbitals?

• How can we communicate the arrangement of atoms in orbitals?

• What are valence electrons and why are they more important than other electrons?

• Complete the “Electron Configurations” activity

Assigning an Electron’s Address Explore

ELECTRON CONFIGURATIONS

• Before we begin… • How do you determine the number of electrons in

an atom again? • In a NEUTRAL atom, # of protons from periodic

table = # of electrons

• We show the way electrons are arranged in atoms by writing electron configurations • The electron configuration of an atom is the

complete description of the orbitals occupied by all of its electrons

• There are rules to follow!

Predicting Electron Locations

• Electrons are added one at a time to the lowest energy orbitals, or subshells, available until all the electrons of the atom have been accounted for

• “aufbau” is German for ‘build up or construct’

Rule #1 – Aufbau Principle

1s

3s

2s 2p

3d 3p

4s 4p 4d 4f

5s 5p 5d 5f

Energies of Electron Orbitals

• An orbital can hold only two electrons

• In other words, no two electrons can ever be in the same place at the same time

• The electrons MUST HAVE OPPOSITE SPINS

• Electrons are associated with “spin”, either one way or the other – like a top

• These spins are called “spin up” and “spin down”

• So, maximum number of electrons held in each orbital is as follows: • 2 for s

• 6 for p (2 x 3 p-orbitals)

• 10 for d (2 x 5 d-orbitals)

• 14 for f (2 x 7 f-orbitals)

Rule #2 – Pauli’s Exclusion Principle

• “Electrons must fill a subshell such that each orbital has a spin up electron before they are paired with spin down electrons”

• More energetically favorable

• A bus analogy:

• If you enter a bus and don’t know anyone on it, you will pick a seat that is completely empty rather than one that already has a person in it

• i.e. Electrons are unfriendly!

Rule #3 – Hund’s Rule

• Electrons fill in order from lowest to highest energy • Ground floor first then up

• The Pauli exclusion principle holds • An orbital can hold only two electrons

• Two electrons in the same orbital must have opposite signs (spins)

• You must know how many electrons can be held by each orbital

• Hund’s rule applies • The lowest energy configuration for an atom is the one having the

maximum number of unpaired electrons for a set of orbitals

• By convention, all unpaired electrons are represented as having parallel spins with the spin “up”

In Summary…

Applying the Electron Rules • Lucky for you, you don’t have to memorize the order in which electrons fill

the orbitals (subshells)

• You can just use the periodic table!

• The PT follows the Aufbau principle – notice (n-1) d-orbitals are filled after ns and before np orbitals AND (n-2) f-orbitals are filled after ns!

• Each element square represents ONE electron in that particular orbital

• Start with H and move through the table in order until the desired element is reached!

Illustrating Electron Configurations • There are several ways to represent

electron configurations: • Full electron configurations

• Condensed (noble gas) electron configurations

• Orbital diagrams

Full Electron Configurations

• In most cases, it is sufficient to write a list of all of the occupied subshells and indicate the number of electrons in each subshell with a superscript.

H 1s1

C 1s2 2s2 2p2

Ar 1s2 2s2 2p6 3s2 3p6

Practice!

• #7 and #10 on page ___________

• Noble gases have full valence shells, so you can condense electron configurations to “eliminate” the electrons already accounted for by the closest noble gas

It cuts down on a lot of writing, and that’s a good thing!

• To use noble gas notation:

• Write the symbol for the preceding noble gas [in brackets] to represent all of the electrons in its electron configuration

• Add the rest of the electrons at the end • Example -the full configuration for As-(Arsenic) is:

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3

• Notice, the part in red is the same as Argon’s configuration:

1s2 2s2 2p6 3s2 3p6

• The noble gas configuration will start with the gas in the row before it

[Ar] 4s2 3d10 4p3

There’s a Quicker Way…Introducing Condensed (Noble Gas) Notation

Practice!

• #13 and #17 on page ___________

Orbital Diagrams

• An orbital box diagram goes one step further by also illustrating the spins of the electrons

• This notation uses boxes to represent orbitals

• One arrow (↑) represents 1 e-

• 2 arrows (↑↓) represent 2 e-

• Same rules as full electron configurations apply!

Phosphorus Full Electron Configuration

1s2 2s2 2p2 2p2 2p2 3s2 3p1 3p1 3p1

Phosphorus Orbital Diagram

• The orbital box diagram indicates that the three electrons in the 3p subshell all have parallel (unpaired) spins

1s 2s 2p 3s 3p

Example

Order of Orbital Filling

Electron Configuration for Platinum (Element 78)

1s2 2s2 2p6 3s2 3p6 4s2 4p6 3d10 5s2 4d10 5p6 6s2 4f14 5d8 1s

2s

3s 4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

6d

4f

5f

7p

• When considering the principal electron shells (n = 1,2,3,…), there are two types of electrons:

• Core Electrons: electrons in the filled “inner” shell(s) of an atom

• Valence Electrons: electrons in the unfilled “outer” shell of an atom

• All elements in the same group have similar chemical properties because they have the same number of valence electrons in their outer shell!

What About Valence Electrons?

What Orbitals Correspond to Valence Electrons Location? • For elements in the first three periods: • The core electrons are those in the preceding noble gas

configuration

• The additional electrons in the outer shell are the valence electrons

• Example:

Full Configuration - B 1s2 2s2 2p1

Noble Gas Config - B [He] 2s2 2p1

Core: 1s2 Valence: 2s2 2p1

(Shell with n = 1) (Shell with n = 2)

Location of Valence Electrons • For elements in the fourth period and below in groups

3A – 7A, the filled d subshells are also part of the core, even though they are not included in the noble gas configuration

Full Config - Se 1s2 2s2 2p6 3s2 3p6 4s2 3d104p4

Noble Gas Config - Se [Ar] 4s2 3d104p4

Core: Valence:

1s2 2s2 2p6 3s2 3p63d10 4s2 4p4

• ONLY ELECTRONS IN HIGHEST NUMBERED s- AND p- ORBITALS ARE VALENCE ELECTRONS!