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CHEMISTRY
CHAPTER 4
PROPERTIES OF LIGHT (P91-93)
• Originally thought to be a wave
• It is one type of ELECTROMAGNETIC RADIATION
(exhibits wavelike behavior as it travels through
space)
• Ex. X-ray, ultraviolet and infrared light, microwaves, and
radio waves
FEATURES OF WAVES
• Wavelength
• Distance between corresponding points on adjacent waves
• Frequency (v)
• The number of waves that pass a point in a specific time
(usually one second)
PHOTOELECTRIC EFFECT (P93-94)
• Emission of electrons from a metal when light shines
on the metal
• Because the light is needed to be a certain
frequency before it can emit an electron, the wave
theory didn’t fit
• (like a soda machine that takes quarters only)
• Max Planck suggested that light is emitted in small
specific amounts called QUANTA
MAX PLANCK
• Quantum is the minimum
quantity of energy that can be
lost or gained by an atom
E = hv
E = energy in joules
v = frequency
h = Planck’s constant: 6.626 x 10-34 J s
ALBERT EINSTEIN
• Said that electromagnetic radiation has a dual
wave- particle nature
• Each particle of light is called a photon
• Energy of light depends on frequency
• Ephoton = hv
PRACTICE PROBLEMS
Wavelength (m)
Frequency (Hz)
Energy (J)
.001
7.0 x 1013
5.0 x 10-7
2.0 x 10-15
1.2 x 1022
C=ƛv
E = hv
C = speed of light = 3.0 x 108 m/s
H = planck’s constant = 6.626 x 10-34 J s
ANSWERS
Wavelength (m)
Frequency (Hz)
Energy (J)
.001
3.0 x 1011
2.0 x 10-22
4.3 x 10-6
7.0 x 1013
4.6 x 10-20
5.0 x 10-7
6.0 x 1014
4.0 x 10-19
1.0 x 10-10
3.0 x 1018
2.0 x 10-15
2.5 x 10-14
1.2 x 1022
8.0 x 10-12
C=ƛv
E = hv
C = speed of light = 3.0 x 108 m/s
H = planck’s constant = 6.626 x 10-34 J s
NOTICE THE PATTERNS
Wavelength (m)
Frequency (Hz)
Energy (J)
.001
3.0 x 1011
2.0 x 10-22
4.3 x 10-6
7.0 x 1013
4.6 x 10-20
5.0 x 10-7
1.6 x 1014
1.1 x 10-18
1.0 x 10-10
3.0 x 1018
2.0 x 10-15
2.5 x 10-14
1.2 x 1022
8.0 x 10-12
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What do you notice about wavelength?
What do you notice about frequency?
What do you notice about energy?
What happens to wavelength if frequency goes up?
What happens to energy if frequency goes up?
HYDROGEN ATOM LINE EMISSION
(P.94,95)
• Ground state = the lowest energy state of an atom
• Excited state = an atom has a higher potential
energy than it has in its ground state
• Current passes through gas (like neon) and excites
the gas; when it returns to its ground state it gives off
electromagnetic radiation (like the colors of neon
light)
HYDROGEN ATOM LINE EMISSION
• When a current is passed through hydrogen gas it
gave off a pink glow.
• Line emission spectrum = the different wavelengths
of the pink glow when passed through a prism
• This told scientists that each atom has specific
energy levels that electrons jump to and from when
going from excited state to ground state
HYDROGEN ATOM LINE EMISSION
• Every element has a characteristic spectrum
• Spectroscopy = identifying a substance by the
wavelengths that are separated by a prism
BOHR’S MODEL (P. 96-97)
• Earlier models:
• Cubic model
• Plum pudding model
(thomson’s)
• Saturnian model
• Rutherford’s model
BOHR’S MODEL (P. 96-97)
• 1913 (Niels Bohr) proposed a new
model that explained the hydrogen
atom line emission spectrum
• Also called planetary model
• Electrons orbit nucleus at
different levels
• Closer to the nucleus means
lower energy level; farther
away is higher energy level
PROBLEMS WITH BOHR’S MODEL
• This model could NOT explain the spectra of atoms with
more than one electron
• It did not fully explain the chemical behavior of atoms
QUANTUM MODEL
• Electrons behave like waves (De Broglie)
• Electrons are confined to the space around an atomic
nucleus (like the frequency of wave)
• They can be bent (or diffracted)
• They can interfere with each other (overlapping and
reduction: p.99 fig.4-10)
SCHRODINGER WAVE EQUATION
• Erwin Schrodinger used the dual wave-particle
theory to develop an equation that treated
electrons in atoms as waves.
• This with the Heisenberg Uncertainty principle laid
the foundation for modern
quantum theory
QUANTUM THEORY
• Describes mathematically the wave properties of
electrons and other very small particles
• The wave function (schrodinger’s equation) gives
only the probability of finding an electron at a given
place around the nucleus
ORBITALS
• Bohr’s idea about neat orbit’s around the nucleus
was wrong
• An orbital is a 3 dimensional region around the
nucleus that indicates the probable location of an
electron
HEISENBERG UNCERTAINTY
• It is impossible to determine simultaneously both the
position and velocity of an electron
ATOMIC ORBITALS AND
QUANTUM NUMBERS
• Quantum numbers specify the properties of atomic
orbitals and the properties of electrons in orbitals.
• There are four quantum numbers:
PRINCIPAL QUANTUM NUMBER
• Symbolized by n
• Tells the main energy level occupied by the
electron
• Only positive numbers
• as n increases, energy increases
AS N INCREASES,
ENERGY INCREASES
ANGULAR MOMENTUM QUANTUM
NUMBER
• Symbolized by l
• Indicates the shape of the orbital
• Shapes are indicated by
• S spherical
• P dumbell shaped
ANGULAR MOMENTUM
QUANTUM NUMBER (CONTINUED)
• D
pair of dumbells mostly
(p. 103)
ANGULAR MOMENTUM
QUANTUM NUMBER (CONTINUED)
• F
very complicated
MAGNETIC QUANTUM NUMBER
• Symbolized by m
• Indicates the orientation of an orbital around the
nucleus
• In S level m=0
• In P orbital m= -1, 0, or +1
• In D; m=-2, -1, 0, 1, or 2
• In F: there are 7 possible orientations
SPIN QUANTUM NUMBER
• Only 2 possible values (+1/2 or -1/2)
• One orbital can hold 2 electrons but they must
have the opposite spin
ELECTRON CONFIGURATION
• The arrangement of electrons in an atom
• 3 basic rules
AUFBAU PRINCIPLE
• An electron occupies the lowest energy orbital that can
receive it
PAULI’S EXCLUSION PRINCIPLE
• No two electrons in the same atom can have the
same set of four quantum numbers
HUND’S RULE
• Orbitals of equal energy are each occupied by one
electron before any orbital is occupied by a
second electron, and all electrons in singly
occupied orbitals must have the same spin