an investigation into electrons and their location and behavior within the atom learning targets: ...

Energy Levels and Orbitals

An investigation into electrons and their location and behavior within the atom

Learning Targets: Describe the process of excitation and emission

of energy by an electron. Understand what each quantum number

represents and how they are determined – energy level, subshell, orbital, and spin.

Emission Spectroscopy

The spectra that were shown through emission spectroscopy led Niels Bohr to question the structure of the atom.


With white light, all of the colors of the visible spectrum are shown.


Since that was NOT what the spectra of elements looked like, Bohr began to look at why only certain wavelengths of color appeared.


E = hc λ

Energy h = 6.63 x 10-34 Js

wavelengthc = speed of light

This equation shows that larger wavelengths indicate lower amounts of energy and smaller wavelengths indicate higher amounts of energy... an inverse relationship.

Bohr realized that the specific wavelengths revealed specific amounts of energy.


Specific amounts of energy!!

That inferred that energy within the atom existed at specific amounts. Bohr called these orbits, or energy levels.

An electron cannot be in-between energy levels, it can only be within an energy level.

Therefore, energy is quantized.

The Bohr Model


Bohr realized that the spectra were being created as electrons moved between these energy levels:

If an electron absorbs energy, it may jump to a higher energy level.

When an electron is at a higher energy level we say that the electron is in its “excited” state.

When the electron releases energy in the form of radiation, we say that the electron has returned to its “ground” state.

The type of radiation that is emitted depends on the amount of energy released.


Nucleus

1st Energy Level

3rd Energy Level2nd Energy Level

4th Energy Level

Energy Coming In!

The Bohr Model

When energy enters the atom, an electron

(shown in red) can absorb the energy becoming

excited, AND jumping to

higher energy levels.


Nucleus

1st Energy Level


4th Energy Level

Energy emitted (infrared)

Energy emitted (red light)

Energy emitted (ultraviolet light)

The Bohr ModelWhen the electron

releases the energy, the

electron returns to

lower energy levels. Other

forms of electromagnet

ic radiation, besides visible light, can be emitted.


Nucleus

1st Energy Level


4th Energy Level

Energy emitted (blue/green light)

Energy emitted (ultraviolet light)

The Bohr ModelWhen the electron

returns to its ground state,

it has the option of

jumping down multiple

energy levels, rather than

one at a time.


Nucleus

1st Energy Level


4th Energy Level

Energy emitted (blue/green light)

The Bohr ModelSince a

sample of gas has many

atoms, there are many

electrons. This is why Bohr saw multiple

colors.But there were

other electromagnetic waves, too.

Energy emitted (red light)


This is the full electromagnetic spectrum.


Bohr saw Visible Light: wavelength is in the

range of 400 to 700 nanometers (4 x 10-7 meters)

ROY G. BIV White light is made of all

the colors of light

Electromagnetic Waves


Gamma rays: cosmic radiation, very high energy

Ultraviolet rays (UV): solar radiation, high energy

Infrared rays (IR): thermal radiation, remote controls, low energy

Microwave rays: microwave oven, very low energy

Electromagnetic Waves


2 --> 1Ultraviolet3 --> 1Ultraviolet4 --> 1Ultraviolet3 --> 2 Visible Red4 --> 2 Visible Blue/Green5 --> 2 Visible Blue4 --> 3 Infrared

Energy LevelChange

Spectra Emission

Electrons release certain types of electromagnetic radiation as they fall to specific lower energy levels.

Quantum Mechanical Model

In addition to knowing that there were energy levels in the atom, three scientists began to notice other things... Heisenberg – impossible to know the exact

position and exact speed of an electron at the same time

De Broglie – electrons have wave-like properties, as in they move in wave patterns

Schroedinger – developed probability of finding each electron in a given location


Heisenberg Bohr suggested that the

electrons move in perfect circles around the nucleus.

Heisenberg showed that, instead, the electron moves in a three dimensional cloud of probability that is smeared out over the orbit – Heisenberg uncertainty principle


DeBroglie Bohr suggested that

the electrons move in perfect circles around the nucleus.

DeBroglie showed that there were other shapes because the electrons moved like waves – wave-particle duality.


DeBroglie Watch this YouTube video.

http://www.youtube.com/watch?v=-YYBCNQnYNM&feature=related


Schrodinger Schroedinger realized

how to put the theories of Bohr, Heisenberg, and DeBroglie together by creating a mathematical equation to find the most likely location for each electron within an atom – wave equation.

Watch this YouTube video.

http://www.youtube.com/watch?v=7GTCus7KTb0&feature=channel&list=UL


Every electron within an atom has “coordinates”. Schrodinger gave these coordinates numerical values, known as quantum numbers. Each quantum number describes part of the coordinates that determine the energy and probable location of any electron for any atom.

Energy levels begin at the number 1.

Each level is higher in energy than the next.

The higher in energy, the farther away from the nucleus.


First Quantum Number Energy Level


Second Quantum Number Subshell Atoms are three

dimensional. Within the energy

levels exist different shapes, or subshells.

The shapes are determined by how much energy is required to create them.

• f – think fireburst

• s – think sphere

• p – think peanut

• d – think daisy


Second Quantum Number Subshell There are four main shapes: s, p, d, and f.

Second Quantum Number Subshell


Second Quantum Number Subshell Since the subshells

are determined by how much energy is required to create them, lower energy levels have fewer subshells. (The lower the energy level, the lower the energy.)

The 1st energy level can only contain the s subshell.

A simple sphere does not take a lot of energy to create.


Second Quantum Number Subshell The higher the

energy level, the more subshells can be held. The 2nd energy level

can contain the s and the p subshell.

As Bohr suggested, these subshells are further away from the nucleus.


Second Quantum Number Subshell The 3rd energy

level must contain three subshells - the s, p, and d.

In effect, the numeric value that represents the energy level also represents the number of subshells within that energy level.

Quantum Mechanical Model To recap:

Energy level 1 = 1 subshell (s) Energy level 2 = 2 subshells (s and p) Energy level 3 = 3 subshells (s, p, and d) Energy level 4 = 4 subshells (s, p, d, and f) etc.

Why are more subshells present? Each energy level is larger than the previous.

As a result, there are more possible locations for where an electron could reside.

Nucleus

1s subshell

2s subshell

2p subshell

3s subshell3p subshell


3d subshell

4s subshell



Watch this YouTube video.

http://www.youtube.com/watch?v=VfBcfYR1VQo&feature=related

Quantum Mechanical ModelThird Quantum Number Orbitals

Did you notice that there were different positions of some of the subshells?

The different positions, or orientations, are called orbitals, not orbits.

The orbitals are determined by which subshell they are in and in which positions they are.

The s orbital does NOT have a different position.

The p orbital has THREE different orientations – x,

y, and z.

Each orbital has a specific number of locations on the x, y, z axes.

- s has 1 orbital orientation (just s)

- d has 5 orbital orientations (dxy

, dxz

, dyz

, dz

2, dx

2-y

2)

- p has 3 orbital orientations (px, p

y, p

z)

- f has 7 orbital orientations (too complex to list)


Third Quantum Number Orbitals


If the next subshell is called “g”, how many orbital orientations should it have?

●_________After “g” the next subshell would be “h”. How many orbital orientations should it have?

●_________


9

11

There is 1 s orbital

There are 3 p orbitals

There are 5 d orbitals

There are 7 f orbitals


Each electron can be spin up (+1/2) or spin down (-1/2)

No two electrons in the same orbital orientation can have the same spin.

With only one spin up and one spin down, the maximum number of electrons that can fit into any given orbital orientation is two.

This is called the Pauli Exclusion Principle.


Fourth Quantum Number Electron Spin


Let’s put it all together!

Energy LevelPossible

SubshellsAtomic Orbitals

Number of Electrons in

Each Subshell

Maximum Possible

Electrons in Energy Level

1 s 1 2 2

2sp

13

26

8

3spd

135

26

1018

4

spdf

1357

26

1014

32


an investigation into electrons and their location and behavior within the atom learning targets: ...

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