Transcript Electrons in Atoms - Duplin County Schools

Electrons in Atoms
Models of the Atom – A History

John Dalton
 atom was solid, indivisible mass

J.J. Thomson
 “plum pudding” model
 e- stuck in lump of + charged matter

Ernest Rutherford
 discovered nucleus
 lacked detail about how electrons occupy the space
surrounding the nucleus
 did not address why the charged electrons are not pulled
into the atom’s nucleus

Niels Bohr
 e- in circular paths around nucleus
 “planetary model”
 e- have fixed energy

In the early 1900s, scientists began to unravel
the puzzle of chemical behavior.

They had observed that certain elements emitted
visible light when heated in a flame.

Analysis of the emitted light revealed that an
element’s chemical behavior is related to the
arrangement of the electrons in its atoms.

So…we need to understand a little about light!
Light and Energy
light: viewed as a wave and a particle (Isaac
Newton)
 electromagnetic radiation: any kind of light,
visible or not
 amplitude: height of a wave
 wavelength: distance from crest to crest
 frequency: # of waves that pass a point in a
given time


measured in Hertz (Hz)
Amplitude

All electromagnetic waves, including
visible light, travel at a speed of 3.00 x 108
m/s in a vacuum.

The speed of light is the product of its
wavelength (λ) and its frequency (ν).

Although the speed of all electromagnetic
waves is the same, waves may have
different wavelengths and frequencies

As you can see from the equation,
wavelength and frequency are inversely
related; in other words, as one quantity
increases, the other decreases
Electromagnetic Spectrum
Particle Nature of Light
• The wave model of light cannot explain why
heated objects emit only certain frequencies
of light at a given temperature, or why some
metals emit electrons when colored light of a
specific frequency shines on them.
• Obviously, a totally new model or a revision
of the current model of light was needed to
address these phenomena.
The quantum concept
• In 1900, the German physicist Max
Planck (1858–1947) began searching for
an explanation as he studied the light
emitted from heated objects.
The quantum concept
• His study of the phenomenon led him to
a startling conclusion: matter can gain
or lose energy only in small, specific
amounts called quanta.
• That is, a quantum is the minimum
amount of energy that can be gained or
lost by an atom.
• Matter can have only certain amounts
of energy—quantities of energy
between these values do not exist.
Electrons and Light

ground state: an e- in its lowest energy state

excited state: an e- in a higher than normal
energy level
 energy
level: region around nucleus where
e- is likely to be moving
 e-
can move up or down but cannot exist
between levels
 e-
must gain right amt. of E to move to a
higher level (lose E to go down)
BOHR MODEL ONLY HOLDS TRUE
FOR THE HYDROGEN ATOM!!!!
Atomic Emission Spectra
• The atomic emission spectrum of an element
is the set of frequencies of the electromagnetic
waves emitted by atoms of the element.
• An atomic emission spectrum is characteristic
of the element being examined and can be
used to identify that element.
• The fact that only certain colors appear in an
element’s atomic emission spectrum means
that only certain specific frequencies of light
are emitted.
Emission Spectrum for Hydrogen

In your reference tables

Only ABSORB energy when e- moves to an
excited state

E is EMITTED if an e- moves to a lower level
Emission Spectrum for Hydrogen
Emission Spectra for H, Hg, Ne
Quantum Mechanical Model

Erwin Schrödinger

complex mathematical formula

no definite path for e-

only gives probability of finding
e-

Heisenberg Uncertainty
Principle: it is impossible to
know the exact location and
speed of an electron at any time
Quantum Theory

Four quantum numbers exist – n, l, m, s
I. Principal Quantum Number (n)
 energy level number
 indicates size of electron cloud
 n = whole # > 0
 # of e- in each level = 2n2
 How many e- can exist in levels 1-5?
II. Sublevels (l) - Azimuthal Quantum #
l



= 0 to n-1
indicates shape of the e- cloud
sublevels called s, p, d, and f
energy level # tells # of sublevels
 Ex. Energy level 1 (n=1), 1 sublevel
Energy level 2 (n=2), 2 sublevels
 How
many sublevels exist in energy
levels 3-5?
III. Orbitals (m) – Magnetic Quantum #
 regions where e- are likely to be found
 each orbital can hold 2 e each sublevel has its own specific orbitals

m = -l to +l
sublevel = 1 orbital = 1 pair e- (2)
 p sublevel = 3 orbitals = 3 pair e- (6)
 d sublevel = 5 orbitals = 5 pair e- (10)
 f sublevel = 7 orbitals = 7 pair e- (14)
s
Orbital Shapes
s
– spherical
p
– dumbbell shaped
d
orbital – clover-leaf shaped
f
orbitals are too complex to be visualized
IV. Spin (s)
s
= +1/2 or -1/2
 each e-
in orbital must spin in opposite
direction – WHY?
 one
clockwise, one counterclockwise
Arrangement of Electrons
 Electron
Configurations
in which e- are arranged around the
nucleus
 ways
 high
E is unstable
 unstable
stable
systems lose E to become more

Aufbau Principle
 e- enter orbitals of lowest E first
 s is lowest E, f is highest E

Pauli Exclusion Principle
 no 2 e- can have the same set of quantum
numbers

Hund’s Rule
 e- will occupy empty orbitals of equal E
before pairing up in an orbital
Exceptional Electron Configurations
half-filled and filled sublevels are more stable
than partially filled sublevels
 e- will shift to become more stable
 transition metals are affected
 Examples (you need to know!)
 Copper
 Silver
 Chromium
