Ch 4 Waves and Spectra. Objectives SWBAT Label the parts of a wave. Solve wavelength to frequency calculations.

Ch 4

•Waves and Spectra

Objectives

• SWBAT• Label the parts of a wave.• Solve wavelength to frequency

calculations.

Waves

• All waves, whether they are water waves or electromagnetic waves, can be described in terms of four characteristics:

• amplitude frequency• wavelength speed

AMPLITUDE

The amplitude of a wave is the height of the wave measured from the origin to its crest.

www.hyperphysics.phys-astr.gsu.edu

WAVELENGTH

The wavelength of a wave is the distance that the wave travels as it completes one full cycle of upward and downward motion.

•www.hyperphysics.phys-astr.gsu.edu

http://www.800mainstreet.com/spect/emission-flame-exp.html

VISIBLE LIGHT WAVELENGTHS

Visible light has wavelengths in the range of 400 to 750 nm

Remember that a nanometer is 10-9 meter

You may remember the visible light spectrum as ROYGBIV

FREQUENCY

The frequency of a wave tells how fast the wave oscillates up and down.

The frequency of light is measured by the number of times a light wave completes a cycle of upward and downward motion in one second.

Units can be: cycles/sec or Hertz

SPEED OF LIGHT

Light moves through space at a constant speed of 3.00 x 108 m/s

You will use …c = 3.00 x 108 m/s

WAVELENGTH AND FREQUENCY CALCULATION

λ = c/v

If the frequency of radiation is 3 x 1015 cycles/sec, what is the wavelength?

WAVELENGTH AND FREQUENCY CALCULATION

λ = c/v

If the wavelength of radiation is 3 x 10-4 m, what is the frequency?

PRACTICE PROBLEMS

1. If the wavelength of radiation is 9.6 x 10 - 6 m, what is the frequency?

2. If the frequency of radiation is 0.33 x 109 cycles/sec, what is the wavelength?

3. If the wavelength of radiation is 6.22 x 10 -12 m, what is the frequency?

4. If the frequency of radiation is 7.8 x 1010 cycles/sec, what is the wavelength?

ELECTROMAGNETIC SPECTRUM

Electromagnetic radiation can be described in terms of a stream of photons, which are massless particles each traveling in a wave-like pattern and moving at the speed of light. Each photon contains a certain amount (or bundle) of energy, and all electromagnetic radiation consists of

these photons. The only difference between the various types of electromagnetic radiation is the amount of energy found in the photons. Radio waves have photons with low energies, microwaves have a little more energy than radio waves, infrared has still more, then visible, ultraviolet, X-rays, and ... the most energetic of all ... gamma-rays.

ELECTROMAGNETIC SPECTRUM

Retrieved from: http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html

RADIO WAVESRadio:

this is the same kind of energy that radio stations emit into the air for your iPod to capture and turn into your favorite music.

But radio waves are also emitted by other things ... such as stars and gases in space.

You can use the information gathered from stars to learn what they are made of.

Retrieved from: http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html

RADIO WAVES

Long, Slow Waves

http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html

HOW CELL PHONES WORK

A cell phone is actually a radio. To see how a cell phone works, check out: http://www.howstuffworks.com/cell-

phone.htm

MICROWAVES

In space, microwaves are used by astronomers to learn about the structure of nearby galaxies, including our own Milky Way!


INFRARED RADIATION

Infrared radiation: we often think of this as being the same thing as 'heat', because it makes our skin feel warm.

Have you seen IR heat lamps at fast food restaurants?

In space, IR light maps the dust between stars.


VISIBLE LIGHT

This is the part of the electromagnetic spectrum that our eyes see.

Visible radiation is emitted by everything from fireflies to light bulbs to stars ... also by fast-moving particles hitting other particles.

DON’T ALL UNITS WORK THE SAME?

In the older "CGS" version of the metric system, the units used were angstroms.

An Angstrom is equal to 0.0000000001 meters (10-10 m in scientific notation)

In the newer "SI" version of the metric system, we think of visible light in units of nanometers or 0.000000001 meters (10-9

m). In this system, the violet, blue, green,

yellow, orange, and red light has wavelengths between 400 and 700 nanometers.


VISIBLE SPECTRUM

•http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html

http://http://www.800mainstreet.com/spect/emission-flame-exp.htmlwww.800mainstreet.com/spect/emission-flame-exp.html

ULTRAVIOLET RADIATION

Ultraviolet: We know that the Sun is a source of

ultraviolet (or UV) radiation, because it is the UV rays that cause our skin to burn!

Stars and other "hot" objects in space emit UV radiation.


X-RAY RADIATION

X-rays: your doctor uses them to look at your bones and your dentist to look at your teeth.

Hot gases in the Universe also emit X-rays .


GAMMA RADIATION Gamma-rays: radioactive materials

(some natural and others made by man in things like nuclear power plants) can emit gamma-rays.

Big particle accelerators that scientists use to help them understand what matter is made of can sometimes generate gamma-rays.

But the biggest gamma-ray generator of all is the Universe! It creates gamma radiation in all kinds of ways.

GAMMA RADIATION

Short, fast radiation


WHICH RADIATION REACHES THE EARTH?

Electromagnetic radiation from space is unable to reach the surface of the Earth except at a few wavelengths, such as the visible spectrum, radio frequencies, and some ultraviolet wavelengths.

Do you know what keeps the radiation from reaching the Earth?

WHICH TYPES OF RADIATION REACH THE EARTH?


WHAT IS ELECTROMAGNETIC RADIATION?

Astronomers can get above enough of the Earth's atmosphere to observe at some infrared wavelengths from mountain tops or by flying their telescopes in an aircraft.

Experiments can also be taken up to altitudes as high as 35 km by balloons which can operate for months.

Rocket flights can take instruments all the way above the Earth's atmosphere for just a few minutes before they fall back to Earth, but a great many important first results in astronomy and astrophysics came from just those few minutes of observations.

For long-term observations, however, it is best to have your detector on an orbiting satellite ... and get above it all! http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.htmlhttp://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html

PLANCK’S THEORY

Max Planck proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quanta.

E = hvPlanck’s constant is

h= 6.626 x 10-34 Js

PLANCK’S EQUATION

E = hv

E = amount of energy emitted or absorbed

h = Planck’s constant 6.626x10-34 Js

PHOTOELECTRIC EFFECT

Einstein used Planck’s equation E=hv to explain the photoelectric effect.

In the photoelectric effect, electrons are ejected from the surface of a metal when light shines on the metal

Einstein proposed that light consists of quanta of energy that behave like tiny particles of light.

Energy quanta are photons.

PHOTOELECTRIC EFFECT EXPERIMENT

http://zebu.uoregon.edu/2000/ph101/photoexp.gif

PHOTOELECTRIC EFFECT

http://www.uwsp.edu/physastr/kmenning/images/photoelec.gif

ARTHUR COMPTON

Compton demonstrated that a photon could collide with an electron, therefore a photon behaves like a particle.

DEBROGLIE

Louis de Broglie determined that particles exhibit wavelike behavior.

DUAL NATURE OF LIGHT

Light acts like a particle and behaves like a wave.

TYPES OF SPECTRA

Continuous Spectrum – a blend of colors one into the other.

An example of a continuous spectrum is a rainbow.

TYPES OF SPECTRA Emission Spectrum (Bright-line

spectrum) - a spectrum that contains only certain colors, or wavelengths

Energy is added to an element sample. The electrons absorb the energy and jump to a higher energy level. They only stay there for an instant and then fall back to a lower energy level. As the electrons fall back down they emit photons of light. Each photon has a specific wavelength and frequency.

www.cms.k12.nc.us/.../notes/ch04electrons.html

http://www.800mainstreet.com/spect/emission-flame-exp.html

HELIUM SPECTRUM

http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/atspect.htmlhttp://hyperphysics.phy-astr.gsu.edu/hbase/quantum/atspect.html

COLORFUL CHEMICALS

Try this web site to see the colorful spectra that different metals can create.

http://webmineral.com/help/FlameTest.shtml

BOHR’S MODEL OF THE ATOM

Bohr listened to a lecture by Rutherford (about his model of the atom). He realized how Planck’s idea of quantization could be applied to this model to explain line spectra.

He decided that electrons can be found only in specific energy levels with specific amounts of energy.

Each energy level was assigned a quantum number

The ground state is the lowest energy level, n=1, this energy level is closest to the nucleus

Electrons absorb a specific quanta of energy and jump to an excited state, n=2 or above

BOHR’S EXPLANATION OF HYDROGEN’S SPECTRAL LINES Bohr proposed that when radiation is

absorbed, an electron jumps from the ground state to an excited state. Radiation is emitted when the electron falls back from the higher energy level to a lower one. The energy of the absorbed or emitted radiation equals the difference between the two energy levels involved.

BOHR’S EXPLANATION OF HYDROGEN’S SPECTRAL LINES CONTINUED Bohr used his model and Planck’s

equation, E=hv, to calculate the frequencies observed in the line spectrum of hydrogen.

This model worked well for hydrogen with one electron, but not for elements with larger numbers of electrons.

Planck, Einstein and Bohr described light as consisting of photons – quanta of energy that have some of the characteristics of particles.

HEISENBERG’S UNCERTAINTY PRINCIPLE

Heisenberg stated that the position and the momentum of a moving object cannot simultaneously be measured and known exactly.

PROBABILITY OF LOCATING AN ELECTRON IN AN ATOM

think of the electrons as residing in a cloud

more dense areas have a higher probability of finding an electron

Draw a diagram of an atom with a surrounding electron cloud

QUANTUM MECHANICAL MODEL OF THE ATOM Draw a picture of the quantum mechanical

model.

ATOMIC ORBITALS

An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found.

The amount of energy an electron has determines the kind of orbital it occupies.

DRILL

Pick up a copy of the 4-5 Prac Prob WS Try problems # 1, 6, 7, 8 and 9

OBJECTIVES

iWBAT Color a periodic table to enable me to locate

the s, p, d, f sublevels on the periodic table. Determine the elemental composition of a

star by using various emission spectra graphical information.

“TO DO” ACTIVITY

Color the s, p, d, f blocks on the periodic table.

SHAPES OF SUBLEVELS

The next several slides will show the shape of the s, p, d, f sublevels.

S SUBLEVEL

s orbitals are spherical in shape

http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html

P SUBLEVELS

p orbitals are dumbbell shaped

Px Py Pz http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html

D SUBLEVELS d orbitals can be several shapes

dz2 dx2-y2 dxy


There are 2 more on the next slide.

D SUBLEVELS

dxz dyz


F SUBLEVELS

f orbitals are complicated 3D shapes that need to be computer generated

see http://nobel.scas.bcit.ca/chem0010/unit3/3.3.3_QM_econfig.htm#here

DRILL

Pick up a copy of the “Reading an e- Configuration WS” drill on the front desk.

Answer the questions

.

OBJECTIVES

iWBAT Distinguish between principle energy levels

and sublevels Use spectra data for various elements to

determine the composition of Stars

STAR SPECTRA

Pick up a copy of the Star Spectra WS

I will use the document camera to show you how to complete this worksheet.

PRINCIPAL ENERGY LEVELS

Principal energy levels in an atom are designated by the quantum number, n.

“n” must be an integer

look at the left hand margin of the periodic table to find the principal quantum numbers

As “n” increases (i.e. from 1 to 2), the electron energy increases

SUBLEVELS

Each principal energy level is divided into one or more sublevels.

The number of sublevels in each principal energy level equals the quantum number, n , for that energy level.

HOW DO YOU TELL THE DIFFERENCE BETWEEN SUBLEVELS?

Sublevels can be distinguished by their:shapessizesenergies

SUBLEVELS

If “n” = 1 sublevel “s”

If “n” = 2 sublevels “s” and “p”

If “n” = 3 sublevels “s”, “p”, “d”

ORBITALS

Each sublevel consists of one or more orbitals.

There can never be more than 2 electrons in each orbital.

ELECTRONS IN ORBITALS

Electrons behave as if they are spinning on their own axis.

A spinning charge creates an electric and magnetic field.

PAULI’S EXCLUSION PRINCIPLE

1. Each orbital in an atom can hold at most 2 electrons

2. Each of these electrons must have opposite spins.

ELECTRON PAIRING

Two electrons with opposite spins (in the same orbital) are paired.

Sublevel “s” holds 2 e- Sublevel “p” holds 6 e- Sublevel “d” holds 10 e- Sublevel “f” holds 14 e-

YOU KNOW THERE ARE SCIENTISTS

The next three slide list the scientist/proper name for some of the rules that we follow when we fill an orbital diagram.

AUFBAU PRINCIPLE

Electrons are added one at a time to the lowest energy orbitals available until all of the electrons of the atom have been accounted for.

PAULI EXCLUSION PRINCIPLE

An orbital can hold up to 2 electrons

Electrons in the same orbital must have opposite spins

HUND’S RULE

Electrons occupy equal energy orbitals so that a maximum number of unpaired electrons results.

This is commonly known as the “Seat on the Bus” Rule

ORBITAL DIAGRAM “TO DO” ACTIVITY

Acquire a clean orbital diagram. Only use pencil on your diagram. Use an orbital to build an electron

configuration.

Ch 4 Waves and Spectra. Objectives SWBAT Label the parts of a wave. Solve wavelength to frequency calculations.

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