Light and Planck's Constant

Notes for Page 29.

Review: (all Review Points taken from Page 17)
1) Energy of electrons increases as you move away from the nucleus.
2) You cannot pinpoint the exact location of an electron - only the general area (Heisenburg Uncertainty Principle - loosely)
3) Bohr put electrons in orbitals.

Waves:

Frequency is inversly proportional to wavelength

- As frequency increases, wavelength decreases (gets shorter)
- As frequency decreases, wavelength increases (gets longer)

The formula for the relationship between frequency and wavelength is:

Why do I care about waves?

 1) Light travels in waves

 2) Electrons are made of light

Light

- Travels in waves

- Very fast

- Electromagnetic Spectrum



How does this apply to Chemistry?

- Electrons are photons of light

- Atoms have an emission spectrum!

   * called the Atomic Emission Spectrum

   * Unique to each element

      (this is how they determine what stars are made of)

   * Seen when an excited atom passes through a gas

How does it work?

1) Atom absorbs a certain amount of energy

2) Atom moves to a higher energy level (now called "excited")

3) Moves back down to its normal energy level (Ground State)

    - emits a packet of energy (quanta or photon) as it does this

4) We can see this emission of energy

    - comes out as color!

Photoelectric Effect

- Metals heat up and change color

- Max Planck wanted to explain this change of color

  * determined that energy changes in set units (quanta)

- Planck's Constant

  * Pertains to amount of energy released when an excited electron goes back to its Ground State

  * States: The amount of radiant energy (E) absorbed or emitted is proportional to the frequency of the

     radiation absorbed.

Planck's Constant = 6.6262 x 10 ^-34 Js

Websites:
Click Here for an explanation of quanta, or photon.
Click Here for an explanation of Planck's Constant.
Click Here for the history and explanation of Planck's Constant.

Bohr Diagram

Things you need to know before you can draw Bohr Structures:

Atoms contain protons, neutrons, and electrons.

Protons

• Positively charged
• Located in the Nucleus
• Contribute to Mass
• Contribute to Charge
• Tell you which element you have
• # CANNOT change
• Are = to the Atomic Number

Neutrons

• Neutrally charged
• Located in the Nucleus
• Contribute to mass
• # CAN change (isotopes)

Electrons

• Negatively charged
• Located in the orbitals
• No mass
• Contribute to charge
• # CAN change

***Assume, for now, that all atoms are neutral. This means that #protons = #electrons***

Bohr Structure

• Shows ALL electrons
• Maps electrons in their orbitals

Orbitals can hold electrons. (n = orbital #)

• Orbital #1 can hold 2 electrons.
• Orbital #2 can hold 8 electrons.
• Orbital #3 can hold 18 electrons.
• Orbital #4 can hold 32 electrons.

Steps:

1. Ask yourself "How many protons are there in this neutral atom?"
2. Based on Mass #, determine number of nuetrons.
3. Since it's a neutral atom: # p = # e
4. Draw orbitals (# orbitals = Period #)
5. Place electrons in orbitals

How to place electrons:

1. Orbital #1 - electrons pair together at the top of the circle
2. All other Orbitals - electrons are place one at a time starting at the top, moving clockwise
3. Once you have four electrons in an orbital the electrons start pairing up.
4. Once you reach the max number that an orbital can hold you move to the next orbital.

Average Atomic Mass

Every element has at least 2 naturally occurring isotopes.

This begs the question, "If elements can have different masses, why is there only one mass on the Periodic Table?"

The mass on the Periodic Table is actually the Average Atomic Mass of all of that element's isotopes.

Quick notes on Average Atomic Mass:

• Units: amu (Atomic Mass Unit)
• The mass on the PT
• Comprised of the average of all isotopic masses

This is the formula for Average Atomic Mass:



This is what the formula means:

1. Isotopes don't exist in equal amounts
2. Take each isotope's mass
3. Multiply it by its % abundance (in decimal form)
4. Do this for all isotopes of the element
5. Add them together

Ex.

Carbon has 2 main isotopes: Carbon-12 and Carbon-13.

C-12 makes up about 98% of all of the Carbon in the world.
C-13 makes up about 1.1% of all of the Carbon in the world.

To find the Average Atomic Mass of Carbon:

1) Change % to decimal

98% = .98
1.1% = .011

2) Multiply decimal by corresponding mass

.98 x 12 = 11.76 amu
.011 x 13 = .143 amu

3) Add the products together

11.76 amu + .143 amu = 11.903 amu

4) The Average Atomic Mass of Carbon is 11.903 amu

Websites
Here is a video explaining how to calculate Average Atomic Mass.
This website gives a brief explanation and a few examples.
Worksheets
This is a great worksheet on Average Atomic Mass.

Isotopes and Ions

Notes for Page 23.

Atoms are identified by their # of protons.

• The # of protons DOES NOT change
• The # of neutrons CAN change
• The # of electrons CAN change

When the # of neutrons and electrons change the atomic structure/description also changes.

Isotopes

• Atoms of the SAME element with DIFFERENT atomic masses
• Occur when a change in neutrons occurs
• All elements have at least two naturally occurring elements

Two ways to write isotopes

1)

Ex.

2)

Ex.

Ions

• Atoms with a charge
• Occur when there is a change in # of electrons

Two types:

1) Cation - positively charged ion

2) Anion - negatively charged ion

Charge is written in the top right corner of the chemical symbol

Quizzes

Quiz #1 is a great quiz!

Quiz #2 is another great quiz!

Quiz #3 is a very basic matching quiz.

Quiz #4 is a really great quiz.

Worksheets

This is a great worksheet on the Subatomic Chart that includes Isotopes and Ions.

This is a great worksheet on atomic structure with isotopes and ions.

PT Review and Intro to Atomic Structure

Notes for Page 21.

In neutral (n) atoms
       Proton # is = to Electron #.

In all atoms
     Mass # is = p + n

Atoms
- Smallest unit of an element that retains elemental properites
   (Dalton's Postulates)
- Are made of Subatomic particles (sub = smaller than, atomic = atom: smaller than an atom)
    1) Protons
         * found in the nucleus
         * charge of +1
         * mass = 1 amu
    2) Neutrons
         * found in the nucleus
         * no charge (neutral)
         * mass = 1 amu
    3) Electrons
         * found in the electron cloud
         * charge of -1
         * negligible (no) mass

History of the Periodic Table

Notes for Page 19.

Click here to download the PowerPoint presentation from class.

If you are unable to download the PowerPoint, please email me at EORoark713@gmail.com or at Elizabeth.Roark@aliefisd.net.

*Please allow 24 hours for a response.

History of the Atomic Theory

Notes for Page 17.

Click here to download the PowerPoint presentation from class.

If you are unable to download, please email me at EORoark713@gmail.com or at Elizabeth.Roark@aliefisd.net.

*Please allow 24 hours for a response.