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37 Cards in this Set

  • Front
  • Back

List 5 examples of electromagnetic radiation

Radiowaves, microwaves, infared, visible light, ultraviolet light

Photoelectric effect

When Light or a certain frequency shines on a metal, the metal glows

Wavelength and frequency are _________ related.


Energy and frequency are _________ related.


Energy and Wavelength are ________ related.


Wave-particle nature

It behaves like a wave but can be measured like a particle.

2 differences between red and green light on the electromagnetic spectrum

Green has a greater frequency

Red has a greater wavelength

Put in order of increasing wavelength: ultraviolet light, microwaves, radio waves, X-rays

X-rays, ultraviolet light, microwaves, radiowaves

Electromagnetic wave relationship formula

c = λv

c = λv




c= constant speed of a wave =3.00x10^8

λ= wavelength in meters (m)

v= Frequency in Hertz (Hz) or 1/s

Formula for energy of a photon






E= Energy in Joules (J)

h= Planck's constant, =6.626x10^-34Js (Joule seconds)

v= frequency in Hertz (Hz) or 1/s


the distance between 2 corresponding points on a wave

Frequency (v)

the occurrance of a wave

Electromagnetic radiation

A form of energy that travels as a wave

Electromagnetic spectrum

Includes all forms of EM radiation (radiowaves, microwaves, infared, visible light, ultraviolet, X-ray, gamma rays)

The speed of light

3.0x10^8 m/s


a minimum amount of energy that can be lost or gained by an atom


Massless particles that carry a quantum of energy

Planck's constant

relates energy with frequency


Electron configuration

the use of orbitals to tell the exact location of every electron of an atom (assume ground state), arrangement of electrons

Aufbau Principle

electrons fill an atom by lowest energy levels and orbital first

Diagonal rule


2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

Hund's rule

orbitals of equal energy fill with 1 electrons before 2 (doesn't apply to s)

electron capacity





Pauli Exclusion Principle

only 2 electrons can occupy each orbital variation and must be opposite in their spin

highest occupied level

electrons in the outermost ring (from config)

valence electrons

electrons in the highest occupied level used for chemical bonding


-Presented the "Planetary Model" of the atom

-Electrons stay in orbits as close as possible to the nucleus with minimum energy (ground state)

-Electrons could jump orbits when energy is absorbed by the atom (excited state)

-When electrons absorb this energy they quickly release it in the form of light (back to ground state)

Planetary model

electrons travel in specific orbits around the nucleus

Problems with Bohr's Planetary Model

-Only explained hydrogen atom (with 1 electron)

-Didn't account for other chemical behavior of atoms

-Conclusion: there really are no circular orbits of electrons


Tried and used methods to detect electrons

Idea: to detect electrons which behave like light using light (photons)

Impossible to take a measurement of an object without disturbing the object (light interferes with electron movement)

Heisenberg Uncertainty Principle

It is impossible to determine the location and speed of an electron simultaneously

-the more you know about one quantity, the less you can know about the other


Derived a mathematical equation that treats electrons as waves; the "Wave Mechanical Model"

-Plots 3-D regions around the nucleus where electrons are probably located at any given point in time

-3-D region called an Orbital (a section of empty space)


Section of empty space

Main Energy Level (Ring or shell)

Theoretically levels go from 1-infinity

Actuall levels 1-7

Orbital shape (subshell)

Sphere (s), propeller (p), double propeller (d), flower petals (f)