Light, Energy, & Electrons

Discrepant Events/Questions

Chapter 6 Part I

EM Spectrum Light as a wave

v=c Light as a particle

E=hv Line spectra Rydberg Equation Bohr’s Hydrogen Model Hydrogen Equation Wave equation of matter

Dual nature of light…

Light is a wave Light is a particle

Light as a wave….

Light acts as a wave Evidence:

Polarization

Light acts as a wave…

More Evidence Diffraction Grating (Prism)

Light acts as a wave…

More Evidence Laser Laser with Colored Lens Flashlight with 2 colored lenses 3D

Filtering Colored flashlight on other colors

Parts of a “wave”

Wavelength The distance between two adjacent peaks (or

troughs)

Parts of a “wave”

Frequency The number of waves that pass a given point

per second Frequency & Wavelength are related by:

vc V= Frequency C=speed of light (2.998 x 108 m/s) wavelength

SAMPLE EXERCISE 6.2 Calculating Frequency from Wavelength

The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of 589 nm. What is the frequency of this radiation?

Solution  

Analyze: We are given the wavelength, of the radiation and asked to calculate its frequency, .Plan: The relationship between the wavelength (which is given) and the frequency (which is the unknown) is given by Equation 6.1. We can solve this equation for and then use the values of and c to obtain a numerical answer. (The speed of light, c, is a fundamental constant whose value is given in the text or in the table of fundamental constants on the back inside cover.)

Solve: Solving Equation 6.1 for frequency gives = c/ . When we insert the values for c and , we note that the units of length in these two quantities are different. We can convert the wavelength from nanometers to meters, so the units cancel:

Check: The high frequency is reasonable because of the short wavelength. The units are proper because frequency has units of “per second,” or s–1.

PRACTICE EXERCISE

(a) A laser used in eye surgery to fuse detached retinas produces radiation with a wavelength of 640.0 nm. Calculate the frequency of this radiation. (b) An FM radio station broadcasts electromagnetic radiation at a frequency of 103.4 MHz (megahertz; MHz = 106 s–1). Calculate the wavelength of this radiation.

Answers: (a) 4.688 1014 s–1, (b) 2.901 m

EM Spectrum

Electromagnetic Radiation All forms of energy that have “wave-like”

behavior” Electromagnetic Spectrum

A full scale of all the forms of EM radiation

Looking at the spectrum…

Why does Tennent do such a good job of blocking some waves?

Why do microwave windows have a grid?

Light as a Particle

Treating light as a wave accounts for a lot of behaviors, but not all

Examples: Why heated objects act

as they do (Their color changes) Why metals eject electrons

when certain lights shine on them (solar cells)

Energy Relates to Frequency

Absorbing & Emitting Energy Objects can only absorb/emit energy in

certain amounts (packets, quantum) Energy can be determined by:

E = H * frequency of light

Energy Constant Light Emitted or Absorbed

H = 6.626 x 10-34 Js

Quantized Energy

What if energy in a car was “quantized”? This would mean your car can only go at

certain speeds (10, 20, 30mph). Why doesn’t this happen?

Photoelectric Effect Light of certain frequencies can force electrons

out of metals solar cells (Calculators etc.)

Light Intensity does not matter – only frequency Photon: energy packet

Photoelectric Effect

SAMPLE EXERCISE 6.3 continued

Answers: (a) 3.11 10–19 J, (b) 0.16 J, (c) 4.2 1016 photons

PRACTICE EXERCISE(a) A laser emits light with a frequency of 4.69 x 1014 s–1. What is the energy of one photon of the radiation from this laser?(b) If the laser emits a pulse of energy containing 5.0 x 1017 photons of this radiation, what is the total energy of that pulse? (c) If the laser emits 1.3 10–2 J of energy during a pulse, how many photons are emitted during the pulse?

Flame Tests

What was responsible for the different colors?

What can we narrow it down to?

Low-Pressure High Voltage Gas Tubes

What color do you “see”? What color is given off? Are there any other wavelengths given off?

Continuous vs. Line Spectrum

Line Spectrum

Continuous Spectrum

Continuous vs. Line Spectrum

Continuous: The rainbow of colors containing all wavelengths

Line Spectrum: Spectrum containing radiation of only specific wavelengths

Balmer & Rydberg

Mid-1800’s Johann Balmer showed how the wavelengths of the 4

visible lines fit a formula Additional lines found

Ultraviolet & infrared regions

Rydberg Equation Calculation of the spectral lines of Hydrogen

Bohr’s Model

Bohr wanted to describe the hydrogen line spectrum more fully

“Planetary” model of electrons 3 Main Points:

1. Only orbits of certain “radii”, corresponding to certain energies, are allowed for an electron

2. An electron in a “level” has a certain energy 3. Energy is emitted or absorbed only when the

electron changes from one level to another

Bohr’s Model Summarized

Atom has distinct energy levels, starting with n=1 then n=2, n=3…

Ground State: lowest energy level When excited, it jumps to a higher state (excited

state)

When it goes back down, it emits energy (light) ‘Step ladder’

Small orbit = low energy stateLarge orbit = high energy state

Bohr Model ft. Rydberg

Rydberg’s equation showed wavelength

Bohr derived energy from this E=hv and v=c

Bohr’s Line Spectra Energy of light given off is due to how far

the electron is ‘falling’ through levels Not all of it is visible

Different jumps give different wavelengths

Grouped in “series” Lyman series: Emits light in the UV region Balmer series: Emits light in the visible spectrum Paschen series: Emits light in the IR region

Figure 4.16 – Prentice Hall Chemistry

1. It neither emits nor absorbs energy.

2. It both emits and absorbs energy simultaneously.

3. It emits energy.

4. It absorbs energy.

1. It neither emits nor absorbs energy.

2. It both emits and absorbs energy simultaneously.

3. It emits energy.

4. It absorbs energy.

Predict which of the following electronic transitions will produce the longest wavelengthspectral line.

1. n = 4 to n = 22. n = 5 to n = 23. n = 5 to n = 34. n = 6 to n = 4

Correct Answer:

1. n = 4 to n = 22. n = 5 to n = 23. n = 5 to n = 34. n = 6 to n = 4

The wavelength increases as frequency decreases. The lowest frequency (longest wavelength) is associated with the lowest energy, and the smallest energy difference here is between n = 6 and n = 4.

Practice Exercise 6.4

Indicate whether each of the following electronic transitions emits energy or requires the absorption of energy: (a) n = 3 to n = 1; (b) n = 2 to n = 4 .

Answers: (a) emits energy, (b) requires absorption of energy