# Why did you pick the shape that you did?

## Presentation on theme: "Why did you pick the shape that you did?"— Presentation transcript:

Design a firework and draw/color it in the space for your “Do Now” for today.
Why did you pick the shape that you did? Why did you choose the colors that you did? No lab today =( Gas lines not working – we will do it on Thursday. Sorry guys. Do Now – 11/10/09

HW for this week Due: Nov 12 (Thu)
Create a poster of the entire Electromagnetic Spectrum: (10 pts) * Visible light must be in color (Think Roy G. Biv (Red, Orange, Yellow, Green Blue, Inidgo, Violet) * Label all EM Radiation clearly * Name, prd, date * Extra 2 points for artictic creativity * Look online and in your book for the necessary information.

New-ish Cornell Note format
Copy the light blue headings of the slide to the left side. Copy the red writing onto the right side. ONLY copy information that is missing from your notes from last week.

Defined as energy that exhibits wavelike behavior. Waves are characterized by: Wavelength (λ): The distance between successive crest or troughs. Measured in m, cm, nm. Frequency (υ): The number of waves that pass a given point per second. Measured in 1/s (s-1) or Hertz (Hz). Amplitude: The waves height from the origin to the crest or the trough. C = speed of light, is a constant (always known, never changes) = 3.00 x 108 m/s c = λ x υ

So what does c = λ x υ really show?
If you take any wave and multiply the wavelength λ by the frequency υ, you will get the speed of light c. Use this equation to solve for questions asking you about wavelength or frequency. What can you infer about the relationship of a wave’s wavelength and frequency? Inversely related: if one gets bigger, the other gets smaller

Electromagnetic Spectrum (EM Spectrum)
Encompasses all forms of electromagnetic radiation. Visible light is the most familiar example of electromagnetic radiation. Differences in the wavelength of visible light are manifested as different colors.

Visible Light The small range of light in the EM spectrum that humans can see with the naked eye (about 400 to 800 nm) ROY G. BIV: Red, Orange, Yellow, Green, Blue, Indigo, Violet (in order of longest to shortest wavelength) 1 2 3 5 6 7 4

Bohr Model (yes, copy this all down)
Based on the hydrogen atom Key points: Electrons are in specific orbitals around the nucleus Ground state: orbital closest to nucleus, lowest in energy Excited states: Orbitals further away are higher in energy Absorb light – electron goes from ground state to excited state Emit light – electron goes from excited state to ground state

Atomic Emission Spectra
Electrons can move between orbitals of FIXED energies Only SPECIFIC wavelengths of light can be absorbed and emitted Atomic Emission Spectrum: the unique pattern of light emitted by a specific atom

Atomic Emission Spectrum
When an element is heated in a flame, the emission spectrum will appear to be just one color to the naked eye.

Identifying Elements using Atomic Emission Spectra!
Match each of the following atomic emission spectra to the ones you drew to identify the element! Write the name of the element on your sheet! Krypton Helium Hydrogen Xenon

Up In Flames! Model Lesson
Materials Chemical compounds mL beakers mL beaker Cotton swabs Energy source Bunsen burner Matches or strikers

Summary Visible light differs in color based on different wavelengths
Atoms can absorb specific quantities of energy based on how electron arrangement Each atom has an unique atomic emission spectrum (so we can identify it!)

Up In Flames! Model Lesson
Yellow Sodium (Na) Orange Calcium (Ca) Red Lithium (Li) Strontium (Sr) Plain White – Silver Magnesium (Mg) Purple-Violet Potassium (K) Blue Copper (Cu)

Up In Flames! Model Lesson
Sodium Yellow Calcium Orange Potassium Purple