Transcript History of the atom

History of the atom
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Dalton
J.J. Thompson
Rutherford
Bohr
Dalton’s Atomic Theory
• All elements are composed of
atoms which are indivisible
• Atoms of the same element are
identical
• Atoms of different elements can
mix together in simple whole
number ratios to form
compounds
• Chemical reactions occur when
atoms are separated, joined or
rearranged. Atoms of one
element can’t be changed into
atoms of another element
Law of Definite
Proportions and Law of
Multiple Proportions
Law of Conservation
of Matter
Early Atomic Models
• J.J. Thompson’s plumpudding model
• Rutherford’s Model of
the atom
Problems with Rutherford Model
• Couldn’t account for
the chemical
properties of the
elements (why do
elements react in the
way that they do?)
The Wave Nature of Light
• Electromagnetic Radiation
Forms of energy that exhibit
wave like characteristics
• Wavelength
– How far apart the waves are
• Frequency
– The # of waves to pass a
point in a certain timeframe
• Wavelength and Frequency
are Inversely Proportional
Wavelength and Frequency
• Not Working very hard
• Long Wavelength
• Low Frequency
Wavelength and Frequency
• Working VERY hard
• Short Wavelength
• High Frequency
What does this tell you about the
relationship between wavelength, frequency
and energy?
The Photo Electric Effect
If light only behaved like a
“wave” ANY frequency of light
would cause the electrons to be
released.
• In the early 1900s an
experiment was done that
COULD NOT be explained
by light being a wave.
• Different colors were
shined on to a metal plate
• Electrons would come off
the metal plate for ONLY
CERTAIN FREQUENCIES
(colors) OF LIGHT
Brick Wall Analogy
No
effect
No
effect
Ping PongSoftball
Ball
Dislodges Brick
and send it
flying
Super Dense
Steel Ball
• Different balls thrown
at wall
• Each one with
different mass
• All at same speed of
90 mile/hour
• Each ball is like a color
or light - each has its
own energy
JB
Einstein and the Photo Electric Effect
• This observation led
Einstein to believe
that light acted like a
particle and a wave
• This is called the
“dual nature” of light
• Light carried packets
of energy called
quanta, or photons
Neils Bohr and the New Model of
the Atom
Mercury Line Spectra
• Bohr hypothesized that
electrons could only be at
certain energy “levels”
around the nucleus
• Electrons could “jump”
from lower to higher
energy states by absorbing
a quantum of energy
• When an electron releases
the energy it gained,
specific colors or
wavelengths of light are
emitted
Electrons and the Atom
Add Energy
Electrons disappear
from one orbit and
reappear at another
without visiting the
space in between!
Release Energy in the
form of colored light
ee-
Nucleus
Just like a jumper has potential
energy at the top of the jump,
the electron has stored potential
energy in the higher orbit.
electron
at “Excited
State”
Quantum
Leap
Electron at
“Ground State”
JB
Energy and the Electron
• Ground State – lowest
energy state for an
electron
• Excited state – high
energy state for e• Quantum – exact
amount of energy to
move an electron from
one energy level to
another
Heisenberg Uncertainty Principle
• It is impossible to
pinpoint the exact
location and velocity of
an electron at any point
in time
• You can estimate where
an electron will be 90% of
the time
• An electron cloud shows
where an electron spends
most of its time
Quantum Mechanical Atomic Model
• Similar to Bohr model
except that e- cannot be
found in distinct orbits
around the nucleus
• Determines how likely it
is for an electron to be
found in various regions
around the nucleus.
Atomic Orbitals
• Region around the
nucleus where an
electron of a given
energy is likely to be
found
• Each orbital has a
characteristic size,
shape, and energy
• There are four different
orbitals: s, p, d, f
Different types of atomic orbitals
Principle Energy Levels
• Symbolized by
n = 1,2,3,4 etc.
• Each energy level
contains sublevels
denoted by a number
and a letter (ex. 1s)
• Each sublevel contains a
certain # of orbitals
• Every orbital can hold 2
e-
Energy Level
# of Sublevels Type of
sublevel and
orbitals
n=1
1
1s (1orbital)
n=2
2
2s (1 orbital)
2p (3 orbitals)
n=3
3
3s (1 orbital)
3p (3 orbitals)
3d (5 orbitals)
n=4
4
4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals)
See table 5.1
on page 131
Electron Configurations
• Aufbau principle: eoccupy the lowest energy
sublevels 1st (see pg. 133
figure 5.7)
• Pauli Exclusion Principle:
every orbital can hold a
maximum of 2 e- (paired espin in the opposite
direction)
• Hund’s Rule: e- fill all
empty orbitals in a sublevel
BEFORE they pair up
Aufbau Diagram
Valence Electrons
• Valence electrons –
electrons found in the
highest energy level of an
atom
• Valence e- determine the
chemical properties of the
element
• A filled outer energy level
(stable octet) of 8 e- makes
an atom stable (or for He
2e- fills its outer energy
level)
Group
Valence e- (and
the orbitals they
occupy)
1A
1 (s1)
2A
2 (s2)
3A
3 (s2p1)
4A
4 (s2p2)
5A
5 (s2p3)
6A
6 (s2p4)
7A
7 (s2p5)
8A
8 (s2p6)
Electron Dot Diagrams
• Shows only number of
valence e• Electrons shown as dots
Carbon has 4 valence e-