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Chapter 4.

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Presentation on theme: "Chapter 4."— Presentation transcript:

1 Chapter 4

2 The Hook Burning money demo— Draw a picture of what happened. What color did you see? Any idea why? Why didn’t the money burn?

3 Everything you ever wanted to know about where the electrons hang out!
Chapter 4 Everything you ever wanted to know about where the electrons hang out!

4 Building on the Atomic Theory
What did Thompson determine? What did Rutherford’s gold foil experiment prove? Just write the words… we will talk in class!

5 Section 1: Early 1900’s Scientists started doing a lot of experiments looking at the absorption and emission of light by matter. Found that there is a relationship between light and an atom’s electrons.

6 Light can behave as a wave

7 Draw the Wave! Amplitude: height of the wave from the origin to the crest Wavelength ( ) : the distance between the crests (m, cm, nm) Frequency (v): number of waves to pass a given point per unit of time (waves/second = Hz)

8 An Important Relationship
The frequency and wavelength of all waves, including light, are inversely related. As the wavelength of light increases, the frequency decreases.

9 C = v Where: C= speed of light 3.00 x 108 m/sec
 = wavelength (m, cm, nm…) v = frequency (1/sec or sec-1)

10 Electromagnetic Radiation
Includes radio waves, radar, microwaves, visible light, infrared light, ultraviolet light, X-rays, and gamma rays

11 Photoelectric Effect Looks at the emission of electrons from a metal when light shines on the metal. Light causes electrons to be ejected from the metal.

12 Only occurs at certain frequencies!
Photoelectric Effect Only occurs at certain frequencies!

13 Wave Particle Duality Explained by Dr. Quantum Leave some space here to write a reflection on the video clip.

14 Sometimes Light Acts Like Particles!
What would happen if the frequency of the wave increased so much that you could hardly tell where one wave ended and another began? Light would start acting more like a particle than a wave.

15 Max Plank Objects emit small packets of energy- Quanta
Quantum- the minimum quantity of energy that can be lost or gained by an atom. E = hv E = Energy h = x Js (Joule x sec) V = frequency (1/sec)

16 Let the units be your guide!!!!!
Take a look at the WS Let the units be your guide!!!!!

17 The Photon Photon- a particle of electromagnetic radiation having no mass, carrying a quantum of energy.

18 So, what happens when photons hit an atom and eject an electron?
The electron goes from the ground state to an excited state. As the electron returns to the ground state, it gives off the energy that it gained- LIGHT

19 Energy Levels Energy levels are not evenly spaced Energy levels
become more closely spaced the greater the distance from the nucleus

20 Another Cool Illustration

21 The Visible Spectrum From about 400nm to 700nm in wavelength. Blue (400nm) has a shorter wavelength than red (700nm).

22 Spectral Analysis of Emitted Light from Excited Atoms
When emitted light from excited atoms is passed through a prism a spectrum of discrete lines of separate colors (separate energies) is observed rather than a continuous spectrum of ROY G BIV.

23 Each element has a unique line-emission spectra

24 Interpretation of Line Spectrum of Elements
The light atoms give off contain very specific wavelengths called a line spectrum light given off = emission spectrum

25 Continuous Spectrum Atomic Line Spectrum

26 Atomic Spectrum Activity

27 Interpretation of Atomic Spectra
The line spectrum is related to energy transitions of electrons in the atom. Absorption = atom gaining energy Emission = atom releasing energy


29 Energy of an atom is quantized – limited to discrete values
All samples of an element give the exact same pattern of lines because every atom of that element must have certain, identical energy states Energy of an atom is quantized – limited to discrete values If the atom could have all possible energies, then the result would be a continuous spectrum instead of lines

30 Bohr Model Electrons orbit around a nucleus
Each orbit has a fixed energy and because of this cannot lose energy and fall into the nucleus Energy Level of an electron is the region around the nucleus where the electron is likely to be moving

31 This helped explain the spectral lines
Absorption- the electron gains energy and moves to a higher energy level. Emission- when the electron falls to a lower energy level.

32 The Quantum Model Finally– the truth (as we know it!)
Electrons can behave as both waves and particles. Electrons can be considered waves with specific frequencies confined to the space around the nucleus. Electrons can also be considered negatively charged particles.

33 Where are the electrons?
Heisenberg Uncertainty Principle: It is impossible to know the position and the velocity of an electron at the same time.

34 Schrodinger Wave Equation
Developed an equation that treated electrons as waves and described the location of electrons. Helped lay the foundation for modern quantum theory (atomic model).

35 Quantum Theory Estimates the probability of finding an electron in a certain position We denote the position of the electron as a “fuzzy” cloud This volume of space where an electron is most likely to be found is called an orbital. The atomic orbitals have distinct shapes

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