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Chapter 5. Electrons in Atoms. Wave Nature of Light. Electromagnetic radiation which is a form of energy that exhibits wavelike behavior as it travels through space. Examples: light, radio waves, x-rays, etc. Parts of a Wave. crest. wavelength. amplitude. origin. amplitude.

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Chapter 5

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Chapter 5

Electrons in Atoms

Wave Nature of Light

• Electromagnetic radiation which is a form of energy that exhibits wavelike behavior as it travels through space.

• Examples: light, radio waves, x-rays, etc

crest

wavelength

amplitude

origin

amplitude

wavelength

trough

Wavelength

• Waves have a repetitive nature.

• Wavelength- ( lambda)

• shortest distance between corresponding points on adjacent waves.

• Measured in units like meters, centimeters, or nanometers depending on the size.

• 1 x 10-9 meters = 1 nanometer

Frequency

• # of waves that pass a given point per second.

• Units are waves/sec, cycles/sec or Hertz (Hz)

• Abbreviated n the Greek letter nu or by an f

c = lf

Frequency and wavelength

• Are inversely related

• As one goes up the other goes down.

High frequency, Short Wavelength

Low frequency, Long Wavelength

Wave Formula

• All electromagnetic waves, including visible light, travel at the speed of 3.00 x 10 8 m/s in a vacuum.

• Speed of light = c = 3.00 x 108 m/s

c=f

Speed of light = (wavelength) x (frequency)

Example Problem

• What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz?

Formula: c=f

• = ?

f = 3.44 x 109 Hz

c = 3.00 x 108 m/s

3.00 x 108 m/s =  (3.44 x 109 s-1)

3.00E8 / 3.44E9 = 8.72 x 10-2 m

Practice

• What is the frequency of green light, which has a wavelength of 5.90 x 10-7m?

• A popular radio station broadcast with a frequency of 94.7MHz, what is the wavelength of the broadcast? ( frequency needs to be is Hz)

• Different frequencies produce different types of waves.

• The entire range of frequencies is called the electromagnetic spectrum

• We are only able to see with our eyes a small portion of the spectrum = visible light

• ROY G BIV

• Different colors mean different frequencies/wavelengths

Energy & The Spectrum

• The energy of a wave increases with increasing frequency

• High Frequency = High Energy

• Low Frequency = Low Energy

• Blue light has more energy than Red light

High energy

Low energy

Low Frequency

High Frequency

X-Rays

Microwaves

Ultra-violet

GammaRays

Infrared .

Long Wavelength

Short Wavelength

Visible Light

Quanta

• Max Planck suggested the idea of quanta or packets of energy.

• Quanta is the minimum amount of energy that can be lost or gained by an atom.

• Energy is quantized = it comes in packets (like stairs or pennies only whole numbers)

Planck’s Constant

• h = 6.626 x 10-34 J.s (Joule seconds)

Energy = (Planck’s constant)(frequency)

E = hf

Example: What is the energy in Joules of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 Hz?

E = ?

h = 6.626 x 10-34Js

f = 7.23 x 1014 Hz (or s-1)

E = (6.626 x 10-34Js)(7.23 x 1014 s-1)

E = 4.79 x 10-19 J

Photoelectric Effect

• In the 1900s, scientist studied interactions of light and matter.

• One experiment involved the photoelectric effect, which refers to the emission of electrons from a metal when light shines on the metal.

• This involved the frequency of the light. It was found that light was a form of energy that could knock an electron loose from a metal.

Photon

• Light waves can also be thought of as streams of particle.

• Einstein called these particles photons (He won a Nobel Prize for this)

• A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum energy.

Bohr’s Model

• Why don’t electrons fall into nucleus?

• Bohr suggested that they move like planets around sun.

• Certain amounts of energy separate one level from another.

• Nucleus is found inside a blurry “electron cloud”

Nucleus

Electron

Orbit

Energy Levels

Bohr’s Model

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• Further away from nucleus means more energy.

• There is no “in between” energy

• Energy Levels

Fifth

Fourth

Third

Increasing energy

Second

First

Nucleus

Bohr Model of the Atom

• Ground state- the lowest energy state of an atom.

• Excited state – state in which an atom has a higher potential energy than its ground state.

• Energy is quantized. It comes in chunks.

• quanta - amount of energy needed to move from one energy level to another.

• Since energy of an atom is never “in between” there must be a quantum leap in energy.

Bohr Energy Levels

• K = 2 electrons – 1st

• L = 8 electrons – 2nd

• M = 18 electrons – 3rd

• N = 32 electrons – 4th

Heisenberg Uncertainty Principle

• This is the theory that states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

Quantum Theory

• Schrodinger derived an equation that described energy & position of electrons in atom

• Schrodinger along with other scientists laid the foundation for the modern quantum theory, which describes mathematically the wave properties of electrons and other very small particles.