Astonishing Astronomy 101With Doctor Bones (Don R. Mueller,

Ph.D.)

EducatorEntertainer

JU

G G LE

RPLANETARY

Scientist

ScienceExplorer

Chapter 4 - Spectroscopy

The Nature of Light

The electromagnetic (EM) wave

A small disturbance in an electric field creates a small magnetic field, which in turn creates a small electric field and so on…

EM waves transport energy.

Particles of light (photons) travel through space.

These photons have very specific energies, that is, light is quantized.

• Light is radiant energy.• Travels very fast –

300,000 km/sec!• Light can be described

as either a wave or a particle.

The Effect of Distance on Light

• Light from distant objects seems very dim

• Why? Because the photons are losing energy?

• No – because the light is simply spreading out as it travels from its source to its destination.

• The farther from the source you are, the dimmer the light seems.

• We say that the object’s brightness or amount of light received from a source is decreasing.

2d4OutputLight TotalBrightness

This is an inverse-square law – the brightness decreases as the square of the distance (d) from the source increases.

The Nature of Matter

• The atom has a nucleus at its center containing protons and neutrons.

• Outside of the nucleus, electrons whiz around in clouds called orbitals– Electrons can also be

described using wave or particle models

– Electron orbitals are quantized – that is, they exist only at particular energies.

– The lowest energy orbital is called the ground state, one electron wave long.

• To move an electron from one orbital to the next higher, a specific amount of energy must be added. Likewise, a specific amount of energy must be released for an electron to move to a lower orbital.

• These are called electronic transitions.

Please insert figure 21.3

The Chemical Elements

• The number of protons (atomic number) in the nucleus determines which element a substance is.

• Each element has a number of electrons equal to the number of protons.

• The electron orbitals are different for each element, and the energy differences between the orbitals are unique as well.

• This means that if we can detect the energy emitted or absorbed by an atom during an electronic transition, we can tell which element the atom belongs to, even from millions of light years away!

Absorption

• If a photon of exactly the right energy (corresponding to the energy difference between orbitals) strikes an electron, that electron will absorb the photon and move into the next higher orbital– The atom is now in an

excited state• If the photon is of

higher or lower energies, it will not be absorbed – it will pass through as if the atom were not there.

• This process is called absorption• If the electron gains enough energy

to leave the atom entirely, we say the atom is now ionized. It is an ion.

Please insert figure 21.6

Emission

• If an atom drops from one orbital to the next lower one, it must first emit a photon with the same amount of energy as the orbital energy difference.

• This is called emission.

Please insert figure 21.4

Seeing Spectrahttp://www.youtube.com/watch?v=QI50GBUJ48s

• Seeing the Sun’s spectrum requires a few special tools, but it is not difficult.

• A narrow slit only lets a little light into the experiment.

• Either a diffraction grating or a prism splits the light into its component colors.

• If we look closely at the spectrum, we can see lines, corresponding to wavelengths of light that were absorbed.

The Bohr Model: http://www.youtube.com/watch?v=45KGS1Ro-sc

• The Danish physicist Niels Bohr developed a model of the hydrogen in the early 1900s that describes its possible energy levels.

• The Bohr model of hydrogen describes the electron as orbiting at a set of discrete radii.

• If we number the energy levels starting from the ground state as n = 1, 2, 3 and so on, then the electron’s orbital radius is r = 0.053 × n2 nanometers (nm). (1 nm = 10−9 m).

• The energy En of level n can be described by the formula: E1 (Hydrogen) = - 2.18 x 10-18 J/n2 = -13.6 eV/n2

n = 1 is the ground state quantum number and E1 is the ground state energy.

1 eV = 1.60 x 10-19 J

Hydrogen Energy Levels

E1 = -13.6 eV/12 = -13.6 eV

E2 = -13.6 eV/22 = -3.4 eV

E3 = -13.6 eV/32 = -1.5 eV

E4 = …

.

.

E∞ = 0

Emission Spectrahttp://www.youtube.com/watch?v=8TJ2GlWSPxI&feature=related

• Imagine that we heat up a gas of atoms.

• Collisions among the atoms cause electrons to jump to higher orbitals, hence higher energy levels.

• Collisions can also cause the electrons to fall to lower levels, emitting photons of energy:

E = hυ = hc/.• If the electron falls from orbitals

3 to 2 in hydrogen, the emitted photon will have a wavelength of = 656 nm. Red

• If the electron falls from orbital 4 to orbital 2, the emitted photon will have a wavelength of = 486 nm. Blue

• We can monitor the gas, counting the number of photons at each wavelength. If we graph this data, we’ll see an emission spectrum.

Emission spectrum of hydrogenhttp://csep10.phys.utk.edu/astr162/lect/light/absorption.html

http://www.youtube.com/watch?v=OJzW2RoZq1Y&feature=related

Atomic Spectra: Elemental Fingerprint

• Every element has its own spectrum. A fingerprint so to speak. Note the differences between hydrogen and helium spectra below.

Absorption Spectra

• What if, instead of hot hydrogen gas, we had a cloud of cool gas between us and a star?

• Photons of an energy that corresponds to the electronics transitions in hydrogen will be absorbed by electrons in the gas.

• The light from these photons is effectively removed from the spectrum.

• The spectrum will have dark lines where the missing light would be:

An Absorption Spectrum: The elemental barcode!

Types of Spectra

• Kirchoff’s Laws:– If the source emits light that is

continuous and all colors are present, we say that this is a continuous spectrum.

– If the gas molecules are moving rapidly (i.e., high temperature), the atoms will emit characteristic frequencies of light. This is an emission-line spectrum.

– If the molecules of gas are cool, they will absorb light of a characteristic frequency as it passes through. This is an absorption line spectrum.

Spectra of Astronomical Objects

Please insert figure 24.7

Composition of the Sun, from SpectroscopyElement Atomic

NumberRelative Number of Atoms

Percent by Mass Energy to Remove Outermost Electron

Hydrogen 1 1012 70.6% 2.179 × 10−18 J

Helium 2 9.77×1010 27.4% 3.940 × 10−18 J

Carbon 6 3.63×108 0.31% 1.804 × 10−18 J

Nitrogen 7 1.12×108 0.11% 2.328 × 10−18 J

Oxygen 8 8.51×108 0.96% 2.182 × 10−18 J

Neon 10 1.23×108 0.18% 3.454 × 10−18 J

Silicon 14 3.55×107 0.07% 1.306 × 10−18 J

Iron 26 4.68×107 0.18% 1.266 × 10−18 J

Gold 79 10 0.000000014 1.478 × 10−18 J

Uranium 92 <0.3 <0.0000000057 0.992 × 10−18 J