Energyand

Electrons

Rev 11/05/08

Part OnePart One

Pisgah High SchoolPisgah High SchoolM. JonesM. Jones

Goals for this unit – Know about:

1. the lines in the hydrogen spectrum, and Bohr’s atomic theory,

2. the arrangement of electrons in atoms, and the shape of the periodic table,

3. energy diagrams and electron configurations,

4. valence electrons and dot diagrams.

Light and Energy

Some background terms and concepts

Frequency

Symbol: f or f (or – Greek: nu)

Units: Hertz (Hz) or 1/sec

Pronounced “reciprocal seconds”

Historically, frequency was in units of “cycles per second”, but this made too much sense.

WavelengthSymbol: Greek: lambda)Units: meters (m) of radio waves in meters to

micrometers. of light in nanometers (10-9 m).

Wavelengths of light can be measured with a spectroscope.

The sine curve is often used to represent wave motion.

Look at the following graphs.

What is the relationship between frequency and

wavelength?

O n e w a v e l e n g t h

One wavelength

One wavelength

One wave-length

What is the relationship between frequency and

wavelength?

As the frequency increases, the wavelength decreases.

The relationship between frequency and wavelength can

be represented by:

λ

1f

f = frequency (lambda) = wavelength

λ

1f

Frequency is inversely proportional to wavelength

The relationship between frequency and wavelength can

be represented by:

The relationship between frequency and wavelength can

be represented by:

Frequency times wavelength equals a constant

k = f

For electromagnetic energy, the equation is:

c is the speed of light c = 3.00 x 108 m/sec

c = f

The equation can also be written as:

c = speed of light= wavelength = frequency

c =

Light is part of the electromagnetic

spectrum

Refer to your reference tables.

Electromagnetic Spectrum

Longer wavelengthLower frequencyLower energy

Shorter wavelengthHigher frequency

Higher energy

Electromagnetic Spectrum

Longer wavelengthLower frequencyLower energy

Shorter wavelengthHigher frequency

Higher energy

400 nm700 nm

The visible spectrum was discovered by …

Dr. Roy G. Biv

Red Orange Yellow Green Blue Indigo Violet

700 nmLower energy

400 nmHigher energy

Electromagnetic waves carry energy.

…inversely proportional to the wavelength.

…directly proportional to the frequency,

The energy in an electromagnetic wave is …

fhE

hc E

Energy in an EM wave

E = hc E =

hc

Small wavelength – large energy

Large wavelength – small energy

Next, a demonstration …Look at sunlight and fluorescent light

through the spectroscope.

Use the spectroscope to look at the light coming from various gas discharge tubes.

Record the colors and their order in the spectrum of hydrogen.

StopComplete the observation of

atomic spectra, then continue.

What did you see?

Hydrogen Spectrum

And now something completely different,

some history.

Thomson suggested the plum pudding model with many,

many electrons throughout

the atom.

J. J. Thomson discovered the electron in 1897.

Thomson suggested the plum pudding model with many,

many electrons throughout

the atom.

J. J. Thomson discovered the electron in 1897.

Rutherford suggested in 1911 that electrons might exist outside the nucleus in a “planetary” arrangement.

Ernest Rutherford explained that atoms had a small, dense,

positively nucleus.

Hantaro Nagaoka Japan, 1904

Enter Niels Bohr

Bohr also knew about Rutherford’s “planetary” model of the electrons.

Bohr was there right after the “gold foil” explanation was published.

Niels Bohr, with a brand new PhD in physics from U. Copenhagen, went first to Thompson, then to Rutherford to study in 1912.

Niels Bohr

In 1913 he proposed a structure of the atom with Rutherford’s nucleus and

electrons in discrete energy levels.

While at Manchester, Niels Bohr studied the spectra of elements, and how these might relate to the internal structure of atoms.

Niels Bohr

He used this new “quantum theory” to help explain how energy could be absorbed and emitted.

Niels Bohr knew of the work done by Max Planck and incorporated it into

his atomic theory.

Max PlanckIn 1900 Max Planck introduced an

unusual idea.Energy exists in “packets”, or quanta.The beginnings of Quantum Theory.A quantum of energy is the smallest

amount of energy possible.Energy exists only in multiples of

these quanta.

Cadmium Sulfide

Albert EinsteinIn 1905 Einstein used Max Planck’s

idea of quanta to explain the photoelectric effect.

Photon of light

Electron knocked loose

Gave credibility to quantum theory.

Albert Einstein

A photon of just the right energy can knock an electron out of an atom.

A photon is a “packet of light”, or quantum of energy.

Einstein explained with the Quantum Theory what could not be explained with classical physics.

Niels Bohr

… that electrons can only change energy levels when

they absorb or give off a certain amount of energy.

(1913)

Bohr said that electrons could exist only in certain discrete energy levels, and …

Hydrogen atom

nucleus

Discrete energy levels for electrons

electron

Hydrogen atom

Electrons can exist at this level,

Hydrogen atom

…or in this level,

Hydrogen atom

…or in this level,

Hydrogen atom

…but not in between the levels.

Hydrogen atom

Unless the electron is absorbing energy, or …

Giving off energy

Niels Bohr

Why ???

When observed through a diffraction grating, specific lines of color are observed.

When high voltage is connected to the hydrogen discharge tube, a

bluish light is given off.

Niels Bohr

… the electron moves up to a higher energy level.

The electron in a hydrogen atom gains energy from the electricity passing thru the tube and …

Niels Bohr

The electron in the excited state is “unstable”.

… gives off light of a certain energy and wavelength.

The electron drops to a lower energy level, and …

Don’t forget:

E = hc

= hcE

andThe wavelength is inversely proportional to the energy

Hydrogen atom

Hydrogen atom

energy

The electron absorbs energy and …

Hydrogen atom…the

electron is elevated

to the next energy level

The electron is unstable and

“wants” to return to a lower

energy level

Hydrogen atom

Hydrogen atom

Light of a particular wavelength is given off

Each line has a wavelength and color that corresponds to the difference in energy

between the two levels.

A line in the hydrogen spectrum is produced when an electron moves from higher energy level to a lower one.

Don’t forget:

E = hc

= hcE

andThe wavelength is inversely proportional to the energy

Energy and

E1

E2

E1

E2

E is large

E is small Longer

Shorter

54321

Hydrogen atom

The energy levels are

numbered …

54321

Suppose an electron is in level 5 and drops to level 2.

Hydrogen atom

The energy levels are

numbered …

54321

Then purple light with a wavelength of 434 nm will be emitted.

Hydrogen atom

The energy levels are

numbered …

54321

Hydrogen atom

Suppose an electron is in level 4 and drops to level 2.

The energy levels are

numbered …

54321

Hydrogen atom

Then you get blue-green light with a wavelength of 486 nm

The energy levels are

numbered …

54321

Hydrogen atom

Suppose an electron is in level 3 and drops to level 2.

The energy levels are

numbered …

54321

Hydrogen atom

Then red light with a wavelength of 656 nm will be emitted.

The energy levels are

numbered …

54321

The colors of the visible lines come from the energy given off by electrons moving from higher energy levels down to level 2.

Hydrogen atom

The energy levels are

numbered …

Don’t forget:

E = hc

= hcE

andThe wavelength is inversely proportional to the energy

The Hydrogen Spectrum

Each color represents the transitions of gazillions of electrons in gazillions of H atoms going from higher energy levels to the second energy level.

Why can’t we see them?

There can be transitions among seven energy levels.

There must be a lot more lines in the spectrum of

hydrogen.

Transitions in the H-spectrum7654

3

2

1

Transitions to the first energy level produce ultraviolet lines.

Transitions in the H-spectrum7654

3

2

1

Transitions to the second energy level produce visible lines.

Transitions in the H-spectrum7654

3

2

1

Transitions to the third and higher energy levels produce infrared lines.

Bohr successfully calculated the

wavelengths of all the transitions in the

hydrogen spectrum.

But only the hydrogen spectrum.

The spectra of elements with more than one

electron could not be accurately predicted.

“Regardless of it’s shortcomings and the

modifications that were later applied, Bohr’s model of the atom was the first successful attempt to make the internal structure of the atom agree with spectroscopic data.”

Asimov, 1964

The Arrangement of Electrons in the

Quantum Mechanical Model of the Atom

The Modern View of the Atom:1. A small, dense positively

charged nucleus which contains protons and neutrons.

The Modern View of the Atom:2. Electrons which exist outside of

the nucleus at …

1. various distances from the nucleus, and at …

2. various energy levels.

The Electrons3. The electrons can have both

a mass, as does matter, and a wavelength, as does light energy.

The Electrons4. The electrons themselves are

not little solid spheres in orbit around the nucleus, but exist as a “fog” of half-energy, half-matter. The electrons can behave as either matter or energy, depending on the experiment.

Energy Levels5. Based on the ideas of Bohr, the

electrons are located …a)… in major energy levels,

b)… in energy sublevels within major energy levels,

c)… in orbitals within each sublevel.

The energy levels are like an organizational chart for a

business:

e lectron e lectron

O rbita l

e lectron e lectron

O rbita l

Sublevel

e lectron e lectron

O rbita l

e lectron e lectron

O rbita l

Sublevel

M ajor energy level

Quantum Numbers1. Each electron in an atom has a set

of four (4) quantum numbers.

2. The quantum numbers are like an address: name, street, city, state.

3. Pauli exclusion principle: no two electrons in the same atom can have the same set of quantum numbers.

Quantum NumbersAddress Quntm. # Sym. What it tells

State Principal n Major energy level

City Azmuthal L Sublevel

Street Magnetic ML Orbital

Name Spin MS Which e- in orbital

Luckily, we will only deal with the Principal Quantum Number

What’s coming next?1. 2n2, and the shape of the periodic

table

2. Energy levels and sublevels

3. s, p, d, f and the periodic table

4. Orbitals, spin & energy diagrams

5. e- config., valence e-, dot diagrams

Click here to go to Part Two.