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Chapter 4
Atomic Structure
Atomic Models
• A model is a working representation of
experimental facts—basically, it’s our best
current guess at why stuff works the way
it does.
• The concept of the atom has been
revised over and over throughout history.
The current model is known as the
quantum model.
Background
• Democritus (465 BC) was the first to
propose that matter was made of tiny
building blocks called atomos.
• The law of definite composition, when
discovered, backed up this idea. All
compounds contain a unique and definite
composition.
• For example, water always contains
hydrogen and oxygen in a 1:8 ratio.
Dalton’s Model
• Elements are made of tiny indestructible
spheres called atoms
• Atoms of different elements have unique sizes
and properties
• An atom of one element cannot become an
atom of another element
• Atoms form compounds by combining with
each other
• Any certain compound always contains the
same relative number and kind of atoms
Thomson’s Model (“Plum
Pudding”)
• Dalton’s model was later revised by J.J.
Thomson and others.
• Thomson identified that all atoms had a
particle in common: the electron
• Thomson’s theory, then—
– Negatively charged electrons are enmeshed
in a positively charged substance
– The charges balance to give electrically
neutral atoms
– Under certain conditions, the electrons can
be removed from the atom
Rutherford and the Proton
• Rutherford did neat stuff: he shot alpha
particles (which are the size of 7350 electrons
and carry a positive charge) at a sheet of gold
foil.
• Most of the particles went through the foil.
However, some were deflected or even
bounced back.
• “It was as if you fired a 15-inch shell at a piece
of paper and it came back to hit you.”
• He reasoned that the gold atoms, then, must
have a small, very dense, region (the nucleus)
made up of positively charged particles
(protons)
Chadwick and the neutron
• Atoms were too big to be made up of
simply protons and electrons.
• The logical explanation, verified by
Chadwick, was an particle of greater
mass with neutral charge.
• We now know that the nucleus consists of
protons and neutrons, which are neutral
and slightly larger than protons.
Atomic sizes
• The average atom is .2 nm across…that’s
.0000000002 meters.
• A 28.5 inch diameter varsity girls’
basketball holds 3,619,500,000 atoms.
• That’s small…but remember, most of the
mass is at the center, with the electrons
orbiting around. If the Rose Bowl were
an atom, and the electrons were orbiting
the upper deck, the nucleus would be the
size of a dime at midfield.
Bohr’s model
• We now knew the basic makeup of the atom.
What we didn’t know was how it was arranged.
• Niels Bohr (1885-1962) couldn’t figure out why
atoms didn’t collapse. If opposite charges
attract, the electrons should “sink” inward
toward the protons in the nucleus.
• The only explanation is that the sheer speed of
the electrons keeps them in orbit—like the earth
around the sun.
Bohr and energy levels
• Light from the sun has a continuous
spectrum; if passed through a prism, it
will break into all the colors of the
rainbow.
• Light coming out of energized electrons,
though, gives off light with unique
(depending on the element) and definite
colored bands: line spectra
continued
• Bohr reasoned that these unique spectra
correlated to the amount of energy used to
excite the electron
• Imagine a dartboard
• Electrons prefer their lowest energy state, close
to the bullseye (remember entropy?). However,
when they take on enough energy, they move
toward the edge of the dartboard. This
transition is quantized and definite, with no
halfway points.
continued
• When the energy ceases to be input, the
electrons immediately return rapidly to
their lowest energy state, giving off light
and heat as they do so.
• It was this light that Bohr observed that
led to his atomic model…electrons
orbiting in concentric rings around the
nucleus
The quantum model
• The problem with Bohr’s model has to do
with uncertainty.
• Heisenberg states that we can know
either a particle’s energy or position, but
we cannot know both simultaneously.
• Think: if we stop a particle to look at it,
we don’t know how it got there. But if we
watch a particle move, we can’t measure
it.
continued
• Since we don’t know exactly how the
electrons move, we can’t state with
certainty that they move along a track.
• We know say that electrons exist in
orbitals, 4-d areas around the nucleus
where we are most likely to find them.
The four dimensions are x, y, z, and the
probability that an electron inhabits that
region.
So where are the
electrons?
• Energy levels>sublevels (s,p,d,f)>orbitals
(x,y,z) > electrons
• There are never more than 2 electrons in
any orbital.
• S orbitals, the simplest, are spherical.
Therefore, any s level can only hold 2
electrons
• See page 80
P orbitals
• The P sublevel contains three separate
dumbbell-shaped orbitals: x, y, and z
• Because it contains three orbitals, the
entire sublevel can hold a total of 6
electrons
The sum of these orbitals can be seen on page 80 of your book
Sublevel d
• D sublevels contain five separate cloverleaf orbitals:
• Because of five orbitals, the d level can
hold 10 electrons.
• The sum of these orbitals can be seen on
page 80 of your book
Sublevel f
• The f sublevel is basically too complex to
depict pictorially.
• It contains seven orbitals, and, therefore,
how many electrons?
• 14…good if you got it.
• Remember…these levels and orbitals are
NOT something we’ve observed. They
are models we’ve guessed at to explain
behavior.
Capacities of energy levels
Principal
energy level
Sublevels
Orbitals in
Electron
Total electron
each sublevel capacity of
capacity
each sublevel
1
S
1
2
2
2
S
P
1
3
2
6
8
3
S
P
D
1
3
5
2
6
10
18
4
S
P
D
F
1
3
5
7
2
6
10
14
32
5
S
P
D
F
1
3
5
7
2
6
10
14
32
Energies of sublevels
• In general…the farther an orbital or level
is from the nucleus, the higher its energy.
• The 1s sublevel has the least energy.
• Everything makes sense until you get to
3p…for some reason, you don’t go to 3d
next…4s comes first.
Sublevel ordering
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
6d
5d
4d
3d
5f
4f
If you’ll draw them out
like this, you’ll never
ever get them wrong.
It’s impossible, unless
you do something
boneheaded
How to build an atom
• The Aufbau principle states that the
electron arrangement of an atom is
determined by adding electrons to a
smaller atom
• Therefore, if you know how many
electrons there are in a mystery atom,
you can draw the electron configuration
perfectly …and if you know the charge,
you can identify the atom.
Hydrogen and Carbon
• Pauli’s exclusion principle states that
each orbital can only contain two
electronw with opposing “spins”
• Hydrogen has 1 electron. Therefore,
hydrogen’s electron configuration is
simple: 1s1
• Carbon, on the other hand, has 6
electrons. Using my diagonal rule,
though, it’s easy to determine its electron
configuration: 1s22s22p2
Filling multiple orbitals
• The p-block has three orbitals: x, y, and z.
How do we know which orbitals get which
electrons?
• Hund’s rule states that, as orbitals fill, each
orbital gets one electron before any get two.
This is because electrons repel each other and
energy must be held to a minimum.
• An atom with four valence electrons in 2porbitals looks like this: 2px22py12pz1 , not this:
2px22py22pz0
Do you get it?
• Give the electron configuration of
manganese (25 electrons).
• 1s22s22p63s23p64s23d5
Quantum numbers:
electron description
• Electron configuration is a description of an
atom. Each type of atom has a unique electron
configuration.
• Quantum numbers work the same way for
electrons. They are a means to quantify and
describe particular electrons.
• N is the first quantum number. It describes the
principal energy level of the electron in
question. As far as you’re concerned, it can go
from 1-7.
More quantum numbers
• L, the second number, identifies the type of
sublevel in which an electron exists—0(s), 1(p),
2 (d), or 3 (f)
• M, the third number, identifies the specific
orbital in which an electron exists. It can range
from –m to +m.
• Ms, the fourth number, identifies the “spin” of
the electron: +1/2 or -1/2
• Remember Pauli: no two electrons can have
the same four quantum numbers in an atom.
Identifying atoms
• Atomic number=number of protons in an
atom. This is the principal way we
identify atoms, because it never changes.
• Mass #=number of protons + number of
neutrons. The number of neutrons is
variable for any atom. More neutrons =
more mass.
• Atoms with numbers of neutrons differing
from the most common are called
isotopes
Isotopes and averages
• Example: 92.58 % of naturally occurring
lithium is Li-7.
• However, that means there’s 7.42 % out
there with some other mass (Li-6).
• How do we know how much lithium
weighs?
• We use a weighted average:
(.0742)(6.015)+(.9258)(7.016)=6.941 amu
• This is similar to grading your work.
Valence electrons
• Valence is a special, fancy, scientific name that
means “outside”
• Valence electrons are important because, as
the electrons feeling the least pull from the
positively charged nucleus, they are the most
likely to be involved in bonding with other
molecules.
• How many valence electrons?
– Hydrogen
– Carbon
– Oxygen
Electron Dot Symbols
• Also called Lewis structures
• A way to visualize an atom’s valence
electrons
• The first two electrons go on as a pair;
any remaining go one at a time:
Mg
S
Ar
Ions: charged atoms
• Anion= A Negative ION
• Cation= (I know it’s cheesy) C A + ION
• The important thing: adding electrons
makes an atom negative; taking them
away makes an atom positive.
• How many electrons would C+2 have?
• What about O-3
HOMEWORK
• Terms: atom, model, electron-neutron,
quantized, orbital, sublevel, electron
configuration, mass number, atomic
number, isotope, valence electron, ion,
cation, anion
• Questions: 2, 5-10, 12-18, 20-24