ATOMIC STRUCTURE

1
ATOMIC STRUCTURE
?
Kenneth E. Schnobrich
2
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.

3
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.

4
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).

5
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.

6
DALTONS ATOM
JOHN DALTON - he envisioned the atom as a hard
spherical unit of matter (the ultimate unit)
LiCl
Lithium
Chlorine
Li2S
Sulfur
7
DALTONS ATOM
JOHN DALTON - he envisioned the atom as a hard
spherical unit of matter
Oxygen
H2O
Hydrogen
H2O2
8
HISTORY (cont.)

• 1875 - Eugen Goldstein discovered the existence
  of a charged stream from the cathode using a
  Crookes Tube and called them Cathode Rays

9
HISTORY (cont.)

• 1886 - Eugen Goldstein discovered the existence
  of positively charged particles he called Canal
  Rays

10
HISTORY (cont.)

• 1897 - J.J. Thomson using a modified Crookes
  Tube determined that the Cathode Rays behaved
  like charged particles and measured the charge.

11
HISTORY (cont.)

• 1907 - J.J. Thomson proposed his Raisin Pudding
  Model of the atom.

Negatively charged Electrons
-
-
-
-
-
-
-
-
Positive Matrix
-
-
-
-
-
-
Atoms are neutral
12
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

13
HISTORY (cont.)
Metal Foil(Au)
Alpha Particles
14
ASSUMPTIONSRUTHERFORDS 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.

15
HISTORY (cont.)
THOMSON MODEL
RUTHERFORD MODEL
16
HISTORY (cont.)

• 1932 - Chadwick discovered and determined the
  properties of the neutron.

Proton
Neutron
Nucleus
Electrons
17
SUBATOMIC PARTICLES
PARTICLE CHARGE MASS LOCATION SYMBOL
PROTON 1 1 AMU NUCLEUS 1H1 or 1p1
NEUTRON 0 1 AMU NUCLEUS 0n1
ELECTRON -1 1/1836 AMU OUTSIDE -1e0
18
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)
19
ISOTOPESAVERAGE 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
20
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.

21
FORMING IONS
Na
Na1
electron
Cl
Cl-1
electron
22
Mn
2 3 4 7
54.94
Atomic Mass
Oxidation States
25
Atomic Number
2-8-13-2
Electron Arrangement
23
THE KERNELANDVALENCE ELECTRONS
19K39
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
24
THE KERNELANDVALENCE ELECTRONS
19K39
2-8-8-1
Valence Electrons 1 Kernel has a charge of
1
Now Lewis Dot Structures
25
LEWIS DOT STRUCTURES
26
PUTTING IT TOGETHER
PARTICLE PROTONS ELECTRONS NEUTRONS
N
N-3
Sn
Sc3
Na
27
CONTINUOUSSPECTRUM
VISIBLE REGION OF THE SPECTRUM
28
HYDROGEN AND HELIUM LINE SPECTRUM
HYDROGEN
More spectra
HELIUM
29
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
30
THE BOHR MODEL

• Bohrs Model was based on the simplest atom,
• Hydrogen. Bohr based his model on the following
• Electrons do not follow the rules of large
  macroscopic bodies.
• Electrons in atoms have only specific energies.
• Electrons are only in specific orbits outside the
  nucleus (ground state).
• When an electron moves from one orbit to another
  it absorbs or releases energy of a specific
  frequency.
• When electrons absorb energy they move to an
  excited state (higher energy orbit).

31
THE BOHR MODELHydrogen
2
Excited State
1
Ground State
2
2
1
1
Energy Released
Energy Absorbed
32
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.
33
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
34
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.

35
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)

36
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.

37
THE QUANTUM MODEL

• The OQNs 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 Y axis
Along the X axis
Along the Z axis
38
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
  distributionalong the X, Y, and Z axes.

39
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.

40
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

41
THE QUANTUM MODEL OFHYDROGEN
Note the electron is pictured as a cloud or
region of space where you will most probably find
the electron.
Nucleus
42
Quantum AtomRelationships
Increasing Energy
43
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
• Hunds 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.

44
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.

45
General SublevelArrangement
3d
4s
3p
Energy Distance from Nucleus
3s
2p
2s
1s
46
General SublevelArrangement
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
47
General SublevelArrangement
3d
4s
3p
For 19K39 the sublevel filling would look like
this
Energy Distance from Nucleus
3s
2p
2s
1s
48
General SublevelArrangement
3d
4s
3p
For 7N14 the sublevel filling would look like this
Energy Distance from Nucleus
3s
2p
Note the sublevel is half-filled, the electrons
have parallel spins (the same)
2s
1s
49
Electron Arrangement
Lets take a sample and show you how the electron
arrangement can be written in three formats.
19K39 2-8-8-1
19K39 1s22s22p63s23p64s1
19K39 Ar4s1
4s1
of electrons
Principal Energy Level
Energy sublevel