Electrons

Atomic theory Overview

1) The Humble BeginningsDemocritus (460-370 BC) and Leucippus (~500 BC)

• The atom is an indestructible thing, it is the smallest piece that any substance can be broken in to.

• It is indivisible, that is, it cannot be broken down any further.

2) Thousands of years passed: John Dalton (1808)

• Atoms are the smallest part that any sample of element can be broken into.

• Atoms of the same element have the same atomic mass, atoms of different elements have different atomic mass.

3) Not so much time passed: a Crookes Tube inJ. J. Thomson (1897)

• The atom is a sphere made of a diffuse (thin) positive charge, in which negatively charged electrons are embedded (stuck).

• He called his model the “plum pudding” model, but who eats plum pudding anymore? It’s more like a “chocolate chip cookie dough” model, where the atom is a positively charged cookie dough ball with negative chocolate chip electrons stuck in it.

4) But then Ernest Rutherford discovered the alpha particle and HAD to play with it! (1911) • The atom is made of a small, dense, positively

charged nucleus with electrons orbiting outside the nucleus at a distance with empty space making up the rest of the atom.

• The majority of an atom’s volume is empty space, and the majority of the atom’s mass is in the nucleus.

5) He saw the light! Broken up into bright lines though a spectroscope!

Neils Bohr! (1913) • Bohr observed the light given off when several elements are

heated and give off light. Different elements gave off different colors of light.

• When this light was passed through a prism, the light was broken up into lines of color. Each element’s lines were different.

• Bohr figured that electrons falling from high energy levels to low energy levels were causing the light.

• Each element’s spectrum of colored lines was different, meaning that the energy levels of different elements have a different amount of energy.

• This process, called spectroscopy, is useful for identifying element samples.

6) Werner Heisenberg may have slept here: we’re uncertain! The Quantum-Mechanical Model• The atom contains a small, dense positive nucleus surrounded by

electrons that travel in a wave-like motion around the nucleus. • This motion is modified by mass and charge interactions between

electrons and the nucleus. • The interactions and the fast speed of the electron make it

impossible to know with any certainty both where an electron is and where it is going in any particular instant.

• All we can know is the general area of space in which the electron might be found. They very from the most general location to the most specific.

• Electrons travel in principal energy levels, which are made up of sublevels, which are made up of orbitals that contain up to two electrons each. If two electrons are in the same orbital, they will spin in opposite directions.

Light

1) Electrons (charged –1 each, with a mass of 1/1836 amu each) surround the nucleus of the atom in distinct energy levels.

Electrons occupy the lowest possible energy levels when the atom is in the ground state.

2) When electrons are given energy (in the form of light, heat or electricity), electrons will rise in energy level by the same amount of energy that the electrons were given. The more energy electrons absorb, the higher they

rise. This is called the excited state. This is in accordance with the Law of Conservation of Energy, which states that energy cannot be created

or destroyed by physical or chemical change.

3) Since electrons are negatively charged, and therefore attracted to the positively charged nucleus, they will eventually fall back to the ground state.

As the electrons fall back to the ground state, they release the energy that caused them to

rise in the first place.

4) The energy is released in the form of photons. They travel at the fastest theoretical speed possible, 3.00 X 108 m/sec, otherwise known as the speed of light.

Photons are, in fact, particles of light.

5) The color of the light is determined by the amount of energy lost by the electron when it dropped back to the ground state.

Light particles travel in a wave pattern. The length of each wave is called, strangely enough, a wavelength. The more energy a photon has, the shorter its wavelength is.

An excited lithium atom emitting a photon of red light to drop to

a lower energy state.

An excited H atom returns to a lower energy level.

Electromagnetic

Spectrum

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The Nature of Light

• The electromagnetic spectrum includes many different types of radiation.

• Visible light accounts for only a small part of the spectrum

• Other familiar forms include: radio waves, microwaves, X rays

• All forms of light travel in waves

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Electromagnetic Spectrum

Figure 06.01Figure 06.01

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Wave Characteristics

• Wavelength: (lambda) distance between identical points on successive waves…peaks or troughs

• Frequency: (nu) number of waves that pass a particular point in one second

• Amplitude: the vertical distance from the midline of waves to the top of the peak or the bottom of the trough

Copyright McGraw-Hill 2009 21

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Wave Characteristics

• Wave properties are mathematically related as:

c = where c = 2.99792458 x 108 m/s (speed of light) l = wavelength (in meters, m) = frequency (reciprocal seconds, s1)

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Wave Calculation

• The wavelength of a laser pointer is reported to be 663 nm. What is the frequency of this light?

c =

1147

8

s104.52m106.63

m/s103.00

m106.63nm

m10nm 663 7

9

c

Copyright McGraw-Hill 2009 24

• Calculate the wavelength of light, in nm,of light with a frequency of 3.52 x 1014 s-1.

c =

c

m108.52s103.52

m/s103.00 7114

8

nm 852m

nm10m108.52

97