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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

Chapter 5

Electrons in Atoms

Wave nature of light
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
Parts of a Wave









  • 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


  • # 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
Frequency and wavelength

  • Are inversely related

  • As one goes up the other goes down.

High frequency, Short Wavelength

Low frequency, Long Wavelength

Wave formula
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


    Speed of light = (wavelength) x (frequency)

Example problem
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


  • 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)

Chapter 5

  • 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


  • Different colors mean different frequencies/wavelengths

Energy the spectrum
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

Chapter 5

High energy

Low energy

Low Frequency

High Frequency






Infrared .

Long Wavelength

Short Wavelength

Visible Light


  • 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
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
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.


  • 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
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.

Bohr s model1
Bohr’s Model




Energy Levels

Bohr s model2
Bohr’s Model


  • Further away from nucleus means more energy.

  • There is no “in between” energy

  • Energy Levels




Increasing energy




Bohr model of the atom
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
Bohr Energy Levels

  • K = 2 electrons – 1st

  • L = 8 electrons – 2nd

  • M = 18 electrons – 3rd

  • N = 32 electrons – 4th

Heisenberg uncertainty principle
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
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.