Download for the p sublevel

ATOMIC THEORY
In 460 B.C., a Greek philosopher, Democritus,
develop the idea of atoms.
He asked this question: If you break a piece of
matter in half, and then break it in half again,
how many breaks will you have to make before
you can break it no further?
Democritus thought that it ended at some point,
a smallest possible bit of matter. He called these
basic matter particles, atoms.
Democritus
(460-370 BC)
• Proposed that matter is made of
tiny particles and empty space
• atoms = smallest part of matter
• different types of atom for each
type of matter
People considered
Aristotle's opinions very
important and if Aristotle
thought the atomic idea
had no merit, then most
other people thought the
same also.
For more than 2000 years nobody did anything
to continue the explorations that the Greeks
had started into the nature of matter.
Lavoisier
(1743-1794)
• Made measurements of a chemical
change in a sealed container
• law of conservation of matter
Proust (1799)
• Observed that water is 11% H and
89% O.
• Law of definite proportions
definitions
• The principle that the elements
that comprise a compound are
always in a certain proportion
by mass.
In the 1800's an English
chemist, John Dalton
performed experiments
with various chemicals
that showed that matter,
indeed, seem to consist of
elementary lumpy
particles (atoms).
http://antoine.fsu.umd.edu/chem/senese/101/ato
ms/dalton.shtml
Postulates
1. All matter consists of tiny particles.
2.Atoms are indestructible and
unchangeable.
3.Elements are characterized by the mass
of their atoms.
4.When elements react, their atoms
combine in simple, whole-number ratios.
J. J. Thomson
(1897)
• Discovered the electron and proton
• Vacuum Tube Experiment
cathode -
anode +
power
source
• The beam starts at the cathode.
• Cathode ray tube.
• CRT
cathode -
anode +
power
source
• The beam is bent away from a
negatively charged plate.
• Like charges repel.
• The beam must be negative particles.
cathode -
anode +
power
source
• Positive particles were deflected from
the anode.
THOMSON’S MODEL
In 1897, the English physicist J.J. Thomson
discovered the electron and proposed a model for
the structure of the atom. Thomson knew that
electrons had a negative charge and thought that
matter must have a positive charge. His model
looked like raisins stuck on the surface of a lump
of pudding.
The Gold Foil Experiment
“You, lab tech, just stay here and count
the flashes of light. I’m going out to
have a drink with the other professors”
- Rutherford to Geiger
“I say! It was like firing a gun at a
tissue and having the bullet bounce
back at you!
“You’re wrong. Do you hear me? You’re
wrong! I say!” -Rutherford to Thomson
Experiment Conclusion
• The atom is mostly empty space.
• The atom contains a small, dense
positive core.
– Nucleus
– If the nucleus was the size of a golf
ball the nearest electron would be one
mile away.
– model
RUTHERFORD’S MODEL
Rutherford knew that atoms consist of a
compact positively charged nucleus, around
which circulate negative electrons at a relatively
large distance.
The nucleus occupies less than
one thousand million millionth
of the atomic volume, but
contains almost all of the
atom's mass. If an atom
had the size of the earth,
the nucleus would have the
size of a football stadium.
• The small, dense, positively
charged central core of an
atom.
smallest
same
mass
Comparing Subatomic Particles
Particle Symbol Charge
Mass (u)
Mass (g)
Proton
P+
+1
1
1.67x10-24
Neutron
n0
0
1
1.67x10-24
Electron
e-
-1
.00055
9.11x10-28
Atomic Number
17
Cl
Chlorine
35.45
• Equals the
number of
protons
• defines the
element
– all chlorine
atoms have 17
protons
Mass number
17
Cl
Chlorine
35.45
• Protons + neutrons
• Isotope have
different masses
– because they have a
different # of
neutrons
• Chart mass is an
average of the
isotopes
Mass number
17
Cl
Chlorine
35.45
• What is the unit for
mass?
• Atomic mass unit
• symbol = u
• based on carbon-12
• 1p+= 1n0 = 1u
Why aren’t the electrons
attracted to the nucleus?
BOHR’S MODEL
In 1912 a Danish physicist, Niels Bohr came up with
a theory that said the electrons do not spiral into the
nucleus and with some rules for what does happen.
RULE 1: Electrons can orbit only at certain allowed
distances from the nucleus.
RULE 2: Atoms radiate energy when an electron
jumps from a higher-energy orbit to a lower-energy
orbit. Also, an atom absorbs energy when an electron
gets boosted from a low-energy orbit to a highenergy orbit.
BOHR-SOMMERFIELD’S MODEL
According to the Bohr-Sommerfeld model, not
only do electrons travel in certain orbits but the
orbits have different shapes and the orbits could
tilt in the presence of a magnetic field. Orbits
can appear circular or elliptical, and they can
even swing back and forth through the nucleus
in a straight line.
QUANTUM MECHANICAL MODEL
The visual concept of the atom now appeared as an
electron "cloud" which surrounds a nucleus. The
cloud consists of a probability distribution map which
determines the most probable location of an electron.
ATOM
divisions
NUCLEUS
ELECTRON CLOUD
particles
particles
ELECTRONS
PROTONS
charge
charge
NEGATIVE
Relative mass
NEUTRONS
NEUTRAL
POSITIVE
is
1/1900
If atom is neutral
ATOMIC NUMBER (Z)
Relative mass
1
is
ATOMIC MASS (A)
N. Bohr
(1885-1962)
• Proposed that the electrons orbited
the nucleus .
• The further away the more energy
was needed.
• Electrons only occupy orbits of
certain energy.
•Carry energy not matter
•Wavelength is the distance between
corresponding points on the wave.
•Wavelengths per second = frequency
•1 wavelength/second = 1 Hertz (Hz)
The Modern Model
• Visible light
can be split
into colored
light.
• Each color has
its own
– frequency
– wavelength
– energy
• Red light has
– low frequency
– long wavelength
– less energy
• Blue light has
– high frequency
– short wavelength
– more energy
Gamma Rays….....Pass through most substances.
X rays…....Pass through soft tissue, but not bone.
UV light…….....causes sunburn, stopped by ozone
Light……………………………………………………..visible
Infrared………………………………………..radiant heat
Radiowaves…..Carry the sounds of radio stations.
The
Electromagnetic
Spectrum
Electrons & Light
• When energy is put into an atom,
electrons absorb the energy and
become “excited”.
• Electrons can absorb only fixed
amounts of energy.
• When excited, electrons move into
a higher energy level (state).
Electrons & Light
• The electrons return to their
original ground state.
• The electrons must give off energy
in the form of light.
• This light has a specific frequency
(and color.)
• This frequency shows up as a line
when seen through a prism.
The Modern Model
Bright Line Spectra
When atoms absorb energy, from fire or
electricity, the electrons give off light.
•When viewed
through a prism H
a bright line
spectra appears. Hg
•This is unique
for every
Ne
element.
hydrogen
The big gaps between spectral lines
represent electron transitions from
one energy level to another.
helium
The groups of fine lines indicate
that electrons are jumping from
energy levels that are close in
energy.
3
ENERGY LEVEL
2
1
ULTRAVIOLET
4
INFRARED
7
6
5
Quantum Theory Model
(Electron Cloud Model)
• Heisenberg Uncertainty
Principle
– You can never know exactly
where an electron is if you
know how fast it is moving.
– If you know its exact
location, you can’t know how
fast it’s moving.
Electron Cloud Model
• Electron paths are not neat orbits.
• The region around a nucleus where
an electron may be found is the
electron cloud or orbital.
• An electron cloud is associated with
a certain energy level.
In 1927 Heisenberg formulated an idea which agreed with test
results, that no experiment can measure the position and
momentum of a quantum particle simultaneously. Scientists call
this the "Heisenberg uncertainty principle." This implies that as
one measures the certainty of the position of a particle, the
uncertainty in the momentum gets correspondingly larger. Or,
with an accurate momentum measurement, the knowledge
about the particle's position gets correspondingly less.
Uncertainty Principle
The position of a particle as well as its attributes cannot be determined with
absolute precision. In association with this fact, a particle may disappear from
one place and reappear elsewhere acausally.
The velocity (causal indeterminism) is coupled with uncertainty of position
(acausal indeterminism). These indeterminate factors are added to the
deterministic velocity to obtain the motion of the electron as the changing
position of its "locus of activity".
• The visual concept of the atom now appeared as
an electron "cloud" which surrounds a nucleus.
The cloud consists of a probability distribution
map which determines the most probable location
of an electron. For example, if one could take a
snap-shot of the location of the electron at
different times and then superimpose all of the
shots into one photo.
• Note: Just as no map can equal a territory, no
concept of an atom can possibly equal its nature.
These models of the atom simply served as a way
of thinking about them, albeit they contained
limitations (all models do).
What should a Model look like?
Scientific models may
not always look like
the actual object. A
model is an attempt
to use familiar ideas
to describe
unfamiliar things in a
visual way.
This is a painting of a young woman
by Pablo Picasso. Does it actually
look like a young woman?
Is this really an Atom?
Many of the models that you have
seen may look like the one below. It
shows the parts and structure of the
atom. Even though we do not know
what an atom looks like, scientific
models must be based on evidence.
The model above represents the
most modern version of the atom.
(Artist drawing)
Atomic Models
Electron Cloud Model
• A certain amount of electrons can
exist in each energy level
–1st level = 2
–2nd level = 8
–3rd level = 18
• Electrons in the outer most levels
are valence electrons.
Sublevels & Orbitals
• Each energy level can be split into
“sublevels”
• Each sublevel can be divided into
“orbitals”.
Quantum Theory Model
Orbitals
One “s” orbital
Three “p” orbitals
Five “d” orbital
Tells the amount of energy
Tells the shape of orbital
Is a specific volume around the
nucleus where two electrons exist.
1st energy level
has 1 sublevel = 1s
2nd energy level
has 2 sublevels = 2s, 2p
3rd energy level
has 3 sublevels = 3s, 3p and 3d
4th energy level
has 4 sublevels = 4s,4p, 4d and 4f
s sublevel holds maximum
p sublevel holds maximum
d sublevel holds maximum
f sublevel holds maximum
2 e6 e10 e14 e-
Placing electrons in orbitals:
•orbital can hold a maximum of 2 e•orbital has the same name as the sublevel
•orbitals in the “s” sublevel are “s” orbitals
y
2s
z
x
Placing electrons in orbitals:
•orbital can hold a maximum of 2 e•orbital has the same name as the sublevel
•orbitals in the “p” sublevel are “p” orbitals
2py
2px
2pz
1, 2, 3, 4, 5, 6 or 7
s, p, d, or f (for the 4th energy level)
x, y, or z (for the p sublevel)
SUBLEVEL ORBITALS ELECTRONS
s
1
2
p
3
6
d
5
10
f
7
14
The most stable arrangement
of electrons in sublevels and
orbitals.
Electron Configurations
For Aluminunum
2
2
6
2
1s 2s 2p 3s 3p
sublevel
1
number of
electrons in
the sublevel
Electron Configurations
For Manganese
2
2
6
2
6
2
1s 2s 2p 3s 3p 4s 3d
sublevel
number of
electrons in
the sublevel
5
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
4f
5f
Sublevel Filling Order
7s
7p
7d
6s
6p
6d
6f
5s
5p
5d
5f
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
Try these electron configurations!
2 2s2 2p6
1s
Ne
=
10
2 2s2 2p6 3s1
Na
=
1s
11
2 2s2 2p6 3s2
Mg
=
1s
12
2 2s2 2p6 3s2 3p6
Ar
=
1s
18
19K = 1s2 2s2 2p6 3s2 3p6 4s1
Déjà vu!
2 2s2 2p6
Look! It’s
1s
Ne
=
10
neon!
2 2s2 2p6 3s1
Na
=
1s
11
2 2s2 2p6 3s2
Mg
=
1s
12
2 2s2 2p6 3s2 3p6
Ar?!
Ar
=
1s
18
19K = 1s2 2s2 2p6 3s2 3p6 4s1
Let’s make this shorter.
2 2s2 2p6
1s
Ne
=
10
2 2s2 2p6 3s1
Na
=
1s
[Ne]
11
2 [Ne]
2 2p6 3s2
Mg
=
1s
2s
12
2 2s2 2p6 3s2 3p6
Ar
=
1s
18
2p6 3s2 3p6 4s1
19K = 1s2 2s2 [Ar]
Noble Gas Notation
2 2s2 2p6
1s
Ne
=
10
1
Na
=
[Ne]
3s
11
2
Mg
=
[Ne]
3s
12
2 2s2 2p6 3s2 3p6
Ar
=
1s
18
19K = [Ar] 4s1
1s2
2p6
3p6
4p6
5p6
6p6
Last Filled Sublevel
for Noble Gases
Try these electron configurations!
26Fe =
47Ag =
79Au =
20Ca =
40Zr =
[Ar] 4s2 3d6
[Kr] 5s2 4d9
[Xe] 6s2 4f14 5d9
[Ar] 4s2
[Kr] 5s2 4d2
Principle and Angular quantum
numbers ( n and l )
• Principle quantum number ( n ) - describes the SIZE of
the orbital.
Since the distance from of an electron from the nucleus is
directly ( l ) - describes the SHAPE of the orbital.
proportional to the energy of the electron (as described in the
Bohr model), the principle quantum number is also a
measure of the orbital.
• Angular quantum number ( l ) - describes the SHAPE of
the orbital.
– The s orbitals are spherical ( l = 0 ).
– The p orbitals are polar ( l = 1 ).
– The d orbitals are clover-leaf shaped (l = 2 ).
l - describes SHAPE
Magnetic quantum number ( m )
• Magnetic quantum number ( m ) - describes
an ORIENTATION of orbital in space.
• For s orbitals (l = 0), there is only one
orientation possible, so m must equal 0.
• For p orbitals (l = 1), there are three possible
orientations, so m can be -1, 0, or 1.
• For d orbitals (l = 2), there are five possible
orientations, so m can be -2, -1, 0, 1, or 2.
m - describes ORIENTATION
Spin quantum number (s)
• Spin quantum number ( s ) - describes the
SPIN or direction (clockwise or counterclockwise) in which an electron spins.
• If there are two electrons in any one orbital,
they will have opposite spins, that is, one
will have a + spin and the other will have a
- spin.
• The maximum number of electron in any
one orbital is two.
Rules for Allowable Combinations of
Quantum Numbers
• The three quantum numbers ( n, l, and m ) that
describe an orbital must be integers.
• “ n " cannot be zero. " n " = 1, 2, 3, 4...
• "l " can be any integer between zero and ( n-1).
• e.g. If n = 4, l can be 0, 1, 2, or 3.
• "m" can be any integer between -l and +l.
• e.g. If l = 2, m can be -2, -1, 0, 1, or 2.
• "s" is arbitrarily assigned as + or - , but for any
one subshell (n , l , m combination),
there can only be one of each.
• SUMMARY
• 1. De Broglie first pointed out the wave-particle duality
of nature. His idea
was that all particles exhibit some wave characteristics,
and vice versa.
• 2. The momentum of an object is the product of its
mass and velocity. Wavelength varies inversely as
momentum.
• 3. Heisenberg's uncertainty principle concerns the
process of observing an electron's position or its
velocity. It is impossible to know accurately both the
position and the momentum of an electron at the same
time.
• 4. Schrodinger developed a mathematical equation
which describes the behavior of the electron as a wave.
The solution set of the wave equation can be used to
calculate the probability of finding an electron at a
particular point.
• 5. Because of the electron's high velocity,
it effectively occupies all volume defined
by the path through which it moves. This
volume called the electron cloud.
• 6. The principal quantum number (n = 1,
2, 3,..) is the number of energy level and
describes the relative electron cloud size.
• 7. Each energy level has as many
sublevels as the principal quantum
number. The second quantum number (1
= s, p, d, f...) describes the of the cloud.
• 8. The third quantum number, m, describes the
orientation in space of orbital. 9:10, 9
• 9. The fourth quantum number, s, describes the spin
direction of an electron.
• 10. Each orbital may contain a maximum of one pair of
electrons. Electrons in the same orbital have opposite
spins. s:12
• 11. Pauli's exclusion principle states that no two
electrons in an atom have the same set of quantum
numbers.
• 12. Electrons occupy first the empty orbital giving the
atom the lowest energy.
• 13. The diagonal rule can be used to provide the
correct electron configuration for most atoms.
• 14. The chemist is primarily concerned with the
electrons in the outer energy level. Electron dot
diagrams are useful in representing the outer level
electrons.