Download Chapter 9 - "Atomic Structure"

•Atomic Structure
• This electron microscope
high-resolution image
shows magnification of the
thin edges of a piece of
mica. The white dots are
"empty tunnels" between
layers of silicon-oxygen
tetrahedrons, and the black
dots are potassium atoms
that bond the tetrahedrons
together. Note the 10
Angstrom width, which is
0.000001 mm.
• First Definition of the Atom
• The Early Greeks thought of everything as being
made up of four basic elements.
–
–
–
–
Earth.
Air.
Fire.
Water.
• Democritus and Leucippus thought that matter was
discontinuous or made up of individual particles.
– Democritus called these fundamental particles atoms.
• Atomic Structure Discovered
• Introduction
– In 1661 Robert Boyle defined an element as a simple
substance which could not be broken down into simpler
substances.
– We now define an element as a pure substance that can
not be broken down into simpler things by either
chemical or physical methods.
– Since elements always combine in fixed ratios, this lends
support to the idea of elements being made of discrete
particles.
• (A) Oxygen and
lead combine to
form yellow lead
oxide in a ratio of
1:13. (B) If 1 atom
of oxygen
combines with 1
atom of lead, the
fixed ratio in
which oxygen and
lead combine
must mean that 1
atom of lead is 13
times more
massive than 1
atom of oxygen.
Reasoning
the existence of
atoms from the
way elements
combine in fixedweight rations.
(A) If matter were
a continuous,
infinitely
divisible material,
there would be no
reason for one amount to go with another amount. (B) If
matter is made up of discontinuous, discrete units (atoms),
then the units would combine in a fixed-weight ratio. Since
discrete units combine in a fixed-piece ratio, they must
also combine in a fixed-weight-based ratio.
• Discovery of the Electron
– A cathode ray is a beam of electrons that moves between
metal plates in an evacuated tube from a negative to a
positive terminal.
• The electron beam is seen as a green beam.
– These rays can be deflected by a magnet.
• A vacuum tube with metal plates attached to a high
voltage source produces a greenish beam called
cathode rays. These rays move from the cathode
(negative charge) to the anode (positive charge).
– In 1897 JJ Thompson place a positively charges plate on
one side of the tube and a negatively charged plate on the
other side of the tube.
• The beam was deflected away from the negative plate toward
the positive plate.
• Thompson realized that the particles that made up the beam
must be negatively charged, since like charges repel and
opposite charges attract.
• By balancing the deflections made by the magnet with that
made by the electrical field, Thompson was able to calculate the
ratio of the charge to mass of an electron as 1.7584 X 1011
coulomb/kilogram
• These particles were later named electrons.
• What appears to be visible light coming through the slit in this
vacuum tube is produced by cathode ray particles striking a
detecting screen. You know it is not light, however, since the beam
can be pulled or pushed away by a magnet and since it is attracted
to a positively charged metal plate. These are not the properties of
light, so cathode rays must be something other than light.
A cathode ray passed between two charged plates is
deflected toward the positively charged plate.
– The ray is also deflected by a magnetic field.
– By measuring the deflection by both, J.J. Thomson was
able to calculate the ratio of charge to mass.
– He was able to measure the deflection because the
detecting screen was coated with zinc sulfide, a substance
that produces a visible light when struck by a charged
particle.
– In 1906 Robert Millikan passed mineral oil through a
vaporized into an apparatus where he could observe the
drops with a magnifier and make measurement on them
as they drifted downward.
• he found that the least charge on any of the droplets was 1.60 X
10-19 coulombs and that larger droplets always had a charge that
was some multiple of this value.
• Knowing Thompson’s work of charge to mass ratio and the
charge on an individual electron, it was possible to calculate the
mass of the electron as 9.11 X 10-31 kg.
• Thompson proposed that an atom was a blob of positively
charged matter in which electrons were stuck like raisins in
plum pudding.
• Millikan measured the charge of an electron by
balancing the pull of gravity on oil droplets with an
upward electrical force.
– Knowing the charge-to-mass ratio that Thomson had
calculated, Millikan was able to calculate the charge on
each droplet.
– He found that all droplets had a charge of 1.60 x 10-19
coulombs or multiples of that charge.
– The conclusion was that this had to be the charge of an
electron
• The Nucleus
– Ernst Rutherford determined that there was a positively
charge nucleus associated with the atom, that was
surrounded by electrons.
• Rutherford calculated that the radius of the nucleus to be about
10-13 cm and the radius of the atom to be about 10-8 cm.
• Electrons therefore took up about 100,000 times the radius of
the nucleus.
• Rutherford and his co-workers studied alpha particle
scattering from a thin metal foil.
– The alpha particles struck the detecting screen, producing
a flash of visible light.
– Measurements of the angles between the flashes, the
metal foil, and the source of the alpha particles showed
that the particles were scattered in all directions,
including straight back toward the source
– In 1917 Rutherford broke up the nucleus of the nitrogen
atom by bombarding it with alpha particles and was able
to identify a particle with a positive charge called a
proton.
• He also thought that there were neutral particles in the nucleus
called neutrons.
• The atom has a tiny, massive nucleus made up of protons and
neutrons.
• Negatively charged electrons, whose charge balances the charge
on the protons, move around the nucleus at a distance of about
100,000 times the radius of the nucleus.
• Rutherfords's nuclear model of the atom explained
the alpha scattering results as positive alpha
particles experiencing a repulsive forced from the
positive nucleus
– Measurements of the percent of alpha particles
passing straight through and of the various angles
of scattering of those coming close to the nuclei
gave Rutherford a means of estimating the size of
the nucleus.
• From measurements of alpha particle scattering,
Rutherford estimated the radius of an atom to be
100,000 times greater than the radius of the nucleus.
This ratio is comparable to that of the (A) thickness
of a dime to the (B) length of football field.
• The Bohr Model
• The Quantum Concept.
– In 1900 Max Plank introduced the idea that matter emits
and absorbs energy in discrete units called quanta.
– In 1905 Albert Einstein extended the quantum concept
to include light and that light consist of discrete units
called photons.
– The energy of a photon is directly proportional to the
frequency of vibration.
• E=hf
– where E = energy
– h = Plank’s constant = 6.63 X 10-34 Js
– f = frequency
• (A) Light from incandescent solids, liquids, or dense
gases, produces a continuous spectrum as atoms
interact to emit all frequencies of visible light (B)
Light from an incandescent gas produces a line
spectrum as atom emit certain frequencies that are
characteristic of each element.
• Atomic hydrogen produces a series of characteristic line
spectra in the ultraviolet, visible, and infrared parts of the
total spectrum. The visible light spectra always consist of
two violet lines, a blue-green line, and a bright red one.
• Bohr’s Theory
– Allowed Orbitals
• An electron can only orbit around an atom in specific orbits
– Radiationless Orbits
• An electron in an allowed orbit does not emit radiant energy as
long as it remains in the orbit.
– Quantum Leaps
• An electron gains or loses energy only by moving from one
allowed orbit to another.
• The lowest energy state is known as the ground state
• Higher states are known as excited states
• Each time an electron males a "quantum leap," moving from
a higher energy orbit to a lower energy orbit, it emits a
photon of a specific frequency and energy value.
• An energy level
diagram for a
hydrogen atom, not
drawn to scale. The
energy levels (n)
are listed on the left
side, followed by
the energies of each
level in J and eV.
The color and
frequency of the
visible light
photons emitted are
listed on the right
side, with the arrow
showing the orbit
moved from and to.
• These fluorescent lights emit light as electrons of mercury
atoms inside the tube gain energy from the electric current.
As soon as they can, the electrons drop back to their lowerenergy orbit, emitting photons with ultraviolet frequencies.
Ultraviolet radiation strikes the fluorescent chemical coating
inside the tube, stimulating the emission of visible light.
• Quantum Mechanics
• Quantum mechanics states that light and matter,
including electrons, have a dual nature of both
particles and waves.
• Matter Waves.
– Louis de Broglie reasoned that particles must also have a
dual nature.
– He reasoned that the electron should have a certain
wavelength that would fit into its orbit around the
nucleus.
•
•
•
•
•
=h/mv
where  is the wavelength
h is Plank’s constant
m is the mass
v is the velocity
• (A) A schematic of de Broglie wave, where the standing
wave pattern will just fit in the circumference of an orbit.
This is an allowed orbit. (B) This orbit does not have a
circumference that will match a whole number of
wavelengths; it is not an allowed orbit.
• Wave Mechanics
– Electrons do emit light in certain wavelengths based on
their energy levels (orbital radius)
– Since waves spread out from the electron, the wave
mechanic model predicts an area where an electron
would be found, and not a specific place where it would
be found.
• The Quantum Mechanics Model
– Quantum mechanics describes the energy levels of an
electron wave with four quantum numbers.
•
•
•
•
distance from nucleus
energy sublevel
orientation in space.
direction of spin
– Principal quantum number (n)
• Describes main energy level of the electron in terms of its
distance from the nucleus.
• n = 1, 2, 3, 4, 5, 6, 7
– Angular momentum quantum number
• Defines energy sublevels within the main energy
levels
• s, p, d, or f designating the type of orbital and also the
orbital shape.
• The Heisenberg Uncertainty Principle states that
you cannot measure the momentum and exact position
of an electron at the same time.
– What you can measure is the probability that an
electron will be found in a certain area, called an
orbital.
• (A)An electron distribution sketch representing probability
regions where an electron is most likely to be found. (B) A
boundary surface, or contour, that encloses about 90 percent
of the electron distribution shown in (A). This threedimensional space around the nucleus, where there is the
greatest probability of finding an electron, is called an
orbital.
– Magnetic quantum number
• Defines the orientation in space of the orbitals relative
to an electrical field.
• The s orbital has one orientation
• The p sublevel can have 3 orientations
• The d sublevel can have 5 orientations
• The f sublevel can have 7 orientations.
• (A) A contour representation of an s orbital. (B) A
contour representation of a p orbital.
– Spin quantum number
• Describes the direction of spin of an electron in its
orbit.
• Electrons occur in pairs and each of the orientations
for a sublevel can have one electron pair.
• Experimental
evidence
supports the
concept that
electrons can
be considered
to spin one
way or the
other as they
move about an
orbital under
an external
magnetic field.
– Pauli Exclusion Principle
• No two electrons can have the same set of quantum numbers.
• At least one of the quantum numbers must differ.
• Electron Configuration
– This is a shorthand designation for electron orientation.
– The lowest possible energy level is n=1.
• If one electron already occupies this energy level, a second can
only occupy it if it has a different spin quantum number.
– Electron configurations tells you the quantum numbers of
the electron.
–
Energy sublevel

Principle quantum number 1s2  two electrons
• There are three
possible
orientations of the
p orbital, and
these are called
px, py, and pz.
Each orbital can
hold two
electrons, so a
total of six
electrons are
possible in the
three orientations;
thus the notation
p6.
• A matrix showing
the order in which
the orbitals are
filled. Start at the
top left, then move
from the head of
each arrow to the
tail of the one
immediately below
it. This sequence
moves from the
lowest-energy level
to the next higher
level for each
orbital.