Transcript ATOMIC STRUCTURE - New York Science Teacher

ATOMIC
STRUCTURE
?
Kenneth E. Schnobrich
A Brief History
• About 460 B.C. - a Greek Philosopher,
Democritus, developed the idea of
atoms (atomos) as small indivisible
particles
• About 400 B.C. - a number of Greek
philosophers said matter consisted
of FIRE, EARTH, WATER, and AIR.
A Brief History
• NEXT 2000 YEARS - Alchemy (a
pseudoscience) dominated - they were
concerned with turning base metals into
gold. During this time Hg, S, and Sb were
discovered. Alchemists also discovered
how to make mineral acids.
• 1754-1826 Joseph Proust showed that a
given compound always had the same
proportions by mass. Law of Definite
Proportions.
A Brief History
• 1766-1844 - John Dalton (following the
work of Robert Boyle) discovered that
atoms can combine in more than one way.
He proposed the Law of Multiple
Proportions. He theorized that the basic
unit was the atom.
• 1808 - Michael Faraday worked on the
electrolysis of molten salts and coined the
word ion (Greek meaning wanderers).
A Brief History
• 1808 - John Dalton published “A New
System of Chemical Philosophy” which
proposed his theory of atoms
– All elements are composed of tiny, discrete,
indivisible and indestructible particles called
“atoms”.
– All atoms of a given element are identical
– Atoms of of different elements are different; they
have different masses and properties.
– Chemical combinations of these “atoms”
compose all matter - different atoms combine
differently to form compounds.
DALTON’S ATOM
JOHN DALTON - he envisioned the atom as a hard
spherical unit of matter (the ultimate unit)
Lithium
LiCl
Chlorine
Sulfur
Li2S
DALTON’S ATOM
JOHN DALTON - he envisioned the atom as a hard
spherical unit of matter
Oxygen
Hydrogen
H 2O
H2O2
HISTORY (cont.)
• 1875 - Eugen Goldstein discovered the
existence of a charged stream from the
cathode using a “Crookes Tube” and called
them “Cathode Rays”
HISTORY (cont.)
• 1886 - Eugen Goldstein discovered the
existence of positively charged particles he
called “Canal Rays”
HISTORY (cont.)
• 1897 - J.J. Thomson using a modified
“Crookes Tube” determined that the
“Cathode Rays” behaved like charged
particles and measured the charge.
HISTORY (cont.)
• 1907 - J.J. Thomson proposed
his “Raisin Pudding Model” of
the atom.
-
Negatively charged
Electrons
-
-
-
-
- -
-
Positive Matrix
*Atoms are neutral
HISTORY (cont.)
1911 - Ernest Rutherford suggested the
atom was “nuclear” based on a
famous experiment - “The Scattering
Experiment”. He also suggested that
the proton was the fundamental unit
of positive charge
HISTORY (cont.)
Metal Foil(Au)
Alpha
Particles
ASSUMPTIONS
RUTHERFORD’S WORK
• Most of the atoms mass is
concentrated in the nucleus.
• All of the positive charge is
concentrated in the nucleus
• Neutral atoms have equal numbers of
protons and electrons.
• The protons and neutrons are located
in the nucleus of the atom.
HISTORY (cont.)
THOMSON MODEL
RUTHERFORD MODEL
HISTORY (cont.)
• 1932 - Chadwick discovered and
determined the properties of the
neutron.
Proton
Neutron
Electrons
Nucleus
SUBATOMIC PARTICLES
PARTICLE
CHARGE
MASS
LOCATION
SYMBOL
PROTON
+1
1 AMU
NUCLEUS
1H
NEUTRON
0
1 AMU
NUCLEUS
1
0n
ELECTRON
-1
1/1836
AMU
OUTSIDE
0
-1e
1
or 1p1
LOOKING AT THE
ATOMS STRUCTURE
• Atomic Number = # protons and
electrons in a neutral atom
• Atomic Mass Number* = sum of the
protons and neutrons
• #Neutrons = Mass# - Atomic#
*ATOMIC MASS MAY VARY (ISOTOPES)
ISOTOPES
AVERAGE ATOMIC
MASS
Most of the elements on the periodic table have several
Isotopes. The Mass that you see is the weighted
average of known isotopes.
Example: Carbon has two stable isotopes
C12 = 98.89% and C13 = 1.108%
12(0.9889) + 13(0.01108) = 12.01
Average Atomic Mass
FORMING AN ION
• METALS - usually like to lose
electrons to form positive ions
called CATIONS.
• NONMETALS - usually like to
gain electrons to form negative
ions called ANIONS.
FORMING IONS
Na
Cl
Na+1
+ electron
+ electron
Cl-1
Atomic Mass
54.94
Mn
Atomic Number
+2 Oxidation
+3 States
+4
+7
25
2-8-13-2
Electron Arrangement
THE KERNEL
AND
VALENCE ELECTRONS
39
K
19
2-8-8-1
Valence Electrons – those in the outer principal energy
level
Kernel – the nucleus and all of the electrons except
those in the valence level
THE KERNEL
AND
VALENCE ELECTRONS
39
K
19
2-8-8-1
Valence Electrons = 1
Kernel – has a charge of +1
*Now Lewis Dot Structures
LEWIS DOT STRUCTURES
39
K
19
-2
S
16
2-8-8-1
2-8-8
K
-2
[ Sxx]
PUTTING IT TOGETHER
PARTICLE
N
N-3
Sn
Sc+3
Na
PROTONS
ELECTRONS
NEUTRONS
CONTINUOUS
SPECTRUM
VISIBLE REGION OF THE SPECTRUM
HYDROGEN AND
HELIUM LINE
SPECTRUM
HYDROGEN
More spectra
HELIUM
BRIGHT-LINE SPECTRA
BRIGHT-LINE SPECTRA ARE LIKE FINGER
PRINTS. EACH ELEMENT HAS ITS OWN
CHARACTERISTIC SET OF BRIGHT LINES IN
THE VISIBLE REGION OF THE SPECTRUM.
Hydrogen
Helium
Carbon
THE BOHR MODEL
Bohr’s Model was based on the simplest atom,
Hydrogen. Bohr based his model on the following:
(1) Electrons do not follow the rules of large
macroscopic bodies.
(2) Electrons in atoms have only specific energies.
(3) Electrons are only in specific orbits outside the
nucleus (ground state).
(4) When an electron moves from one orbit to
another it absorbs or releases energy of a specific
frequency.
(5) When electrons absorb energy they move to an
excited state (higher energy orbit).
THE BOHR MODEL
Hydrogen
2
1
Excited State
Ground State
2
2
1
1
Energy Absorbed
Energy Released
THE QUANTUM MODEL
As the science of spectroscopy grew and the
resolution of the bright-line spectra of an element
improved and the dual nature of the electron was
explored scientists formulated a new picture of the
atom.
This new model of the atom retains some of the
original features but changes the concept of electron
location. The electron, instead of occupying a specific
orbit now is thought to occupy a region of 3-D space
called the orbital.
THE QUANTUM MODEL
Dual Nature of the Electron – the electron to this
point, had been described as being “particle-like” in
nature, but it also exhibits “wave-like behavior.”
DeBroglie – was the first to suggest that, based on its
extremely small size, the electron does have a
measureable wavelength.
Double-click on the You
Tube video
THE QUANTUM MODEL
After viewing the video we see that Erwin
Schrodinger allows us to describe the electrons in an
atom with a set of 4 Quantum Numbers.
• The quantum numbers help us to describe the
relative energies and probable locations of the
electrons.
•The Principal Quantum Number (n) – corresponds
very closely with the energy levels described in the
Bohr Model. The PQN can only have small whole
number values (n = 1, 2, 3, 4, 5, 6 etc). The greater
the value of “n” the greater the energy and distance
from the nucleus for the electron.
THE QUANTUM MODEL
• The Sublevel Quantum Number (l) – describes the
sublevels the electrons can occupy within a Principal
Energy Level.
•The SQN – has values that are determined by the
value of the PQN.
• It can have values from 0 … n-1
• So, if n = 0, l = 0
• If n = 2, l = 0, 1 (which means, in the second
Principal Energy Level, there are two available
sublevels the electron can occupy.
• There are also corresponding letter values for
the sublevels 0(s); 1(p); 2(d); 3(f)
THE QUANTUM MODEL
• In the 2nd PEL there were two sublevels, 0, 1 or s
and p.
• Within a PEL, as the value of l increases the energy
and distance from the nucleus increases.
• In the 3rd PEL, there are three sublevels, 0, 1, and 2
or s, p, and d sublevels.
•The Orbital Quantum Number (m) (also sometimes
called the Magnetic Quantum number) - describes
the number of orbitals (3-D orientations in space)
within a sublevel.
THE QUANTUM MODEL
• The OQN’s are determined by the values for l
• m can have values from 0… +/- l
• So, if l = 0, m = 0, which means that there is only
one possible 3-D description (or orbital) in that
sublevel.
• If l = 1, m = 0, +1, -1, which means in the “p”
sublevel there are three, 3-D descriptions (or
orbitals), in that sublevel.
Along the X axis
Along the Y axis
Along the Z axis
THE QUANTUM MODEL
• If an orbital is located in an s sublevel it is referred
to as an s-orbital and has a spherical distribution
along the X, Y, and Z axes.
• If an orbital is located in an p sublevel it is referred
to as an p-orbital and has a “dumbell”
distribution”along the X, Y, and Z axes.
THE QUANTUM MODEL
• Of course there are other orbital shapes but they are
complicated and for our purposes, our concerns will
be limited to the s and p orbital shapes.
• The fourth quantum number is the Spin Quantum
Number – based on the Stern/Gerlach experiment it
is thought that an electron can have one of two
possible spins, +1/2 and -1/2 (it spins on its axis).
• Since no two electrons can have exactly the same set
of four quantum numbers, only two electrons can
occupy an orbital, provided they have opposite spins.
THE QUANTUM MODEL
• Based on the work of many scientists, including
deBroglie, Shrodinger, and Heisenberg, we now know
that –
• we can only speak in terms of the probable
location of the electrons
• the bright line spectra available for the elements
gives us additional information on the energy
associated with the electrons
THE QUANTUM MODEL
OF
HYDROGEN
Note: the electron is
pictured as a cloud or
region of space where
you will most probably
find the electron.
Nucleus
Quantum Atom
Relationships
Increasing Energy
ORBIT
Principal
Energy Level
SUBLEVELS
s
p
d
f
1 orbital
3 orbitals
5-orbitals
7-orbitals
Electron Filling
When we fill the energy sublevels that are several rules we
must follow –
• The Aufbau Principle – you must always fill from lowest
energy to highest energy
• Hund’s Rule – you must completely half-fill an energy
sublevel before you start pairing electrons
• Pauli Exclusion Principle – no two electrons can have the
same set of four quantum numbers in a given orbital, they
must have opposite spins to exist in the same orbital.
Filling the Sublevels & Orbitals
• When filling the sublevels and orbitals remember the rules
• It is also important to remember that for multi-electron
atoms some of the sublevels do overlap from an energy
standpoint.
• there is a simplified relationship to help us with this
overlap – Sublevel Energy = n + l (n is the PQN and l is
the SQN ). It is why the 4s sublevel fills before the 3d
sublevel (see the diagram on the next slide).
• 4s = 4 + 0 = 4 and 3d = 3 + 2 = 5
• the sublevel energy of 4s is lower than that of 3d,
therefore, the 4s sublevel fills before the 3d sublevel.
General Sublevel
Arrangement
3d
4s
3p
3s
2p
2s
1s
General Sublevel
Arrangement
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
* The idea of sublevel overlap can be much more complicated for larger atoms
General Sublevel
Arrangement
3d
4s
3p
3s
2p
2s
1s
For 19K39 the
sublevel filling
would look like this
General Sublevel
Arrangement
3d
4s
3p
3s
2p
2s
1s
For 7N14 the sublevel
filling would look
like this
Note: the sublevel is
half-filled, the
electrons have
parallel spins (the
same)
Electron Arrangement
Let’s take a sample and show you how the electron
arrangement can be written in three formats.
39
K
19
2-8-8-1
39
K
19
1s22s22p63s23p64s1
39
K
19
[Ar]4s1
Principal Energy
Level
1
4s
# of electrons
Energy sublevel