 Originate from the movement of electrical charges
 Do not require a medium to move
 ________________- the distance of one wave cycle
 Often taken from one crest to the next but only has to be two
identical points
 Symbol is 
 Length units
 __________________- # of wave cycle that pass a given point in a
given unit of time
 Symbol is 
 Units: cycles/sec, sometimes written as 1/s, also equals 1 Hertz (Hz)
 _____________________
 Distance between the zero amplitude line and the top of the crest
 Waves with higher amplitudes carry more energy
 _________________
  is the wavelength in meters
  is the frequency in s-1 (or Hz)
 c is the speed of the wave
 For electromagnetic radiation, use the speed of light
2.99792458 x 108 m/s (round this to 3.00 x 108 m/s)
 There is an _____________ relationship between
wavelength and frequency
 Light can be passed through a prism and makes a
rainbow. This is a ________________ spectrum.
 Proposed that electromagnetic energy
comes in discreet amounts (or
_____________________)
 The energy of a quantum is expressed
by
 ____________where h= 6.626 x 10-34 J s
 Energy can only increase by multiples:
1 h, 2 h, 3 h, and so on
 Applied the idea of quanta to explain how light can
cause electricity:
 Tiny packets of light (photons) hit an atom.
 If they have enough energy, they can kick an electron
out of an atom (photoelectric effect)
 So light is a wave and a “packet”
 Remember: Bohr revised Rutherford’s model, putting
the electrons in orbitals like planets around the sun
 Excited electrons could jump orbital levels
 When electrons returned to their original level
(ground state), they emit energy
 Creates a ________ spectrum
 Created when only a few wavelengths of the
electromagnetic spectrum are seen
 Gases of each element will make them when electricity
passes through them or they’re heated up.
 The lines can be seen when the light is passed through a
prism.
 Lines are different for each element
(_____________________________)
BRIGHT LINE SPECTRA
7500
7000
6500
6000
5500
5000
4500
4000 A
6500
6000
5500
5000
4500
4000 A
Mercury
Calcium
Sodium
Helium
Hydrogen
7500
7000
 Electron behave like ____________ (matter) and
____________ (electromagnetic radiation)
 If this is true, electrons can’t be tied to a fixed orbit
 We can’t even be certain of both an electron’s
momentum and position at any one point in time
(Heisenberg uncertainty principle)
 Can only create shaded areas around the nucleus
where there is a high _________________ of finding
electrons
 System of numbers that gives each electron a unique
address in any atom
 Also called ______
 Represents the primary energy levels around the
nucleus (kind of like Bohr’s rings)
 There are 7 energy levels so _____________
 Electrons with a higher n are in higher energy orbital
and spend more time away from the nucleus
 Electrons with the same n are in the same principle
shell
 Also called l or ___________
 Determines the shape of the orbital
 ____ sublevel
 _____ sublevel
 _____ sublevel
 _____ sublevel
 Spherical in shape
 Look like a dumbbell in each axis (x, y, and z)
 Also called ml or the orbital
 If
 s subshell……_____ orbital
 p subshell……_____ orbitals
 d subshell……_____ orbitals
 f subshell…..______ orbitals
 Two electrons can fill each orbital
 Another quantum number is needed so every electron
has its own number
 Electron Spin Quantum Number
 Also called ms
 Electron can spin _________________________________
 An orbital with two electrons has one of each
 Also represent them with ___________
The magnet
splits the beam.
Silver has 47 electrons
(odd number). On
average, 23 electrons will
have one spin, 24 will
have the opposite spin.
These silver atoms each
have 24 +½-spin electrons
and 23 –½-spin electrons.
These silver atoms each
have 23 +½-spin electrons
and 24 –½-spin electrons.
 Description of the placement of all the electrons in an
atom of an element
 Rules
 Electrons occupy the lowest energy orbitals
available (ground state). Filling order is
determined by Aufbau Principle.
 Orbitals can only hold 2 electrons and they must
have opposing spins (Pauli exclusion principle)
 Electrons will fill empty orbitals when possible.
Half-filled orbitals all have parallel spin. (Hund’s
rule)
 Remember this is the real location of the f block
 Regular
 Energy level is written as a
number
 Subshell letter is written next
 Number of electrons in subshell
is given as an exponent
 Extended
 Same as regular, but subshells are
divided into the individual
orbitals
 Regular
 He _________
 O ______________
 Cl ________________________
 Extended
 He __________
 O _____________________
 Cl _______________________________
 Same as spdf configuration but it starts from the noble
gas of the previous period
 Examples
 Be _______________
 K
_______________
 Br ________________
 Same as spdf configuartions except
 Number of electron in subshell is not written
 Each orbital within subshell is given a line
 ↑ and ↓ arrows are drawn on the lines to represent the
electrons (fourth quantum number)
 Make sure to follow Hund’s rule
 Examples:
 N _________________________
 O __________________________
 Electrons in the __________________of an atom
 Electrons in lower energy levels are called ______
electrons
 Can be predicted by the periodic table for
representative elements
D and f-block elements have 2,
but can vary
Helium
only has
2.
 Dots are placed around the element symbol to
represent to valence electrons
 To be stable, atoms want __________ in their outer shell
 Noble gases already have 8 and don’t usually react with
anyone
 Others will lose or gain electrons to get to 8
 Said to be isoelectronic with the Nobel gas from the last period
 Ion will have the same electron configuration as this noble gas
 Atoms that gain or lose electrons to become isoelectronic with
Helium follow the duet rule (2 electrons in valence shell)
 Group 1- 1 valence electron
 Lose 1 electron, become +1 cation
 Group 2- 2 valence electrons
 Lose 2 electrons, become +2 cation
 Group 13- 3 valence electrons
 Lose 3 electrons, become +3 cation
 Group 15- 5 valence electrons
 Gain 3 electrons, become -3 anion
 Group 16- 6 valence electrons
 Gain 2 electrons, become -2 anion
 Group 17- 7 valence electrons
 Gain 1 electron, become -1 anion
 Transition metals form cations, but the charge cannot
be predicted
 Group 14 does not commonly form ions
 Group 18 does not form ions
 The metal properties of the elements increase as you
move to the left and move down on the periodic table
Metals
Nonmetals
Conduct electricity well
Don’t conduct electricity well
Malleable
Brittle
Ductile
Not ductile
Tend to lose electrons
Tend to gain electrons
Low electronegativity
High electronegativity
Slightly negative electron affinity
Very negative electron affinity
Found in bases
Found in acids
Shiny
Dull
 Ionization energy is the amount of energy needed to
remove an electron
 Need more energy as you move up and to the right of
the periodic table
 Electronegativity is the ability of an atom to attract
electrons
 Increases as you move up and to the right of the
periodic table (noble gases sit out)
 Atomic size increases as you move down and to the left
of the periodic table