Chapter 5

Chapter 5
Electrons in Atoms
Actually, the Chemical History powerpoint talked about a lot of the stuff from
Chapter 5 too, specifically part of section 1 and most of section 3.
This powerpoint will talk about what is still left over in Chapter 5.
Quantum Mechanics…where we
left off from Chem History
S Better than any previous model, quantum
mechanics does explain how the atom
S Quantum mechanics treats electrons not
as particles (which they are), but as waves
(like light) which can gain or lose energy.
S But they can’t gain or lose just any
amount of energy. They can only gain or
lose a “quantum” of energy.
A quantum is just an amount of energy that the electron needs to
gain (or lose) to move to the next energy level. Max Planck,
another German Nobel Prize winning scientist first came up with
this idea.
What the heck is a Quantum?
S Think of a quantum as a “packet” of
energy, much like a sugar packet at a
restaurant. A sugar packet contains a
teaspoonful of sugar.
S If the electron absorbs energy, it moves
to a higher energy level. If it emits
(loses) energy, it moves to a lower
energy level.
S But like Bohr suggested in his model,
the electron has to gain or lose exactly
the right amount. That amount is a
quantum of energy.
Neils Bohr: The Planetary Model
& Energy Levels
S You can’t just step anywhere.
S You have to step on the rungs of a ladder.
S An electron has to jump from one level to
S The steps on a ladder are all the same distance
S But in Bohr’s model, the energy levels get closer
and closer the further away you get from the
Quantum Mechanics “borrowed”
the concept of energy level.
S The electron really doesn’t orbit (like a little
planet) around the nucleus.
S Quantum mechanics describes “electron clouds”
and where they are in relation to the nucleus.
S Again, the electron can move from one energy
level to another, IF it absorbs a quantum of
Energy Levels &
Quantum Numbers
S Quantum mechanics has a principal
quantum number.
S It is represented by a little n. It
represents the “energy level” similar
to Bohr’s model.
S n=1 describes the first energy level
S n=2 describes the second energy
level. Etc.
S Each energy level represents a
period or row on the periodic table.
S Isn’t it amazing how all this stuff
just “fits” together?
Atomic Orbitals
S The energy levels in quantum mechanics describe locations
where you are likely to find an electron cloud.
S Schroedinger used calculus to calculate the PROBABILITY
of finding an electron in a particular location.
S These locations are called ORBITALS.
S Orbitals are “geometric shapes” around the nucleus where
electrons are found.
S There must be at least a 90% probability of finding an
electron there.
S The 4 different types of orbitals are s, p, d, and f.
Atomic Orbitals
S Think of orbitals as sort of a "border” for spaces
around the nucleus inside which electrons are allowed.
S No more than 2 electrons can ever be in 1 orbital.
S The orbital just defines an “area” where you can
find 1 or 2 electrons.
S No more than 2 can fit into any one orbital.
S What is the chance of finding an electron in the
S Yes, of course, it’s zero.
S There aren’t any electrons in the nucleus.
S A node = a location where the probability of finding
an electron there = 0.
Atomic Orbitals define an area
where electrons are moving
S Quantum mechanics doesn’t predict a SPECIFIC orbit,
like the Bohr model does.
S We don’t really know how the electron is moving, or if
it follows any particular path as it moves.
Energy Sub-level = Specific
Atomic Orbital
S Each energy level has 1 or more “sub-
levels” which describe the specific
“atomic orbitals” for that level.
Blue = s block (0)
S n = 1 has 1 sub-level (the “s” orbital)
S n = 2 has 2 sub-levels (“s” and “p”)
S n = 3 has 3 sub-levels (“s”, “p” and
S n = 4 has 4 sub-levels (“s”, “p”, “d”
and “f ”)
S s, p, d, f refer to specific areas on the
Periodic Table where those orbitals
are being filled with electrons.
S A second quantum number identifies the
specific orbital.
Shapes of These Orbitals
(the nucleus is ALWAYS at the center of the orbital)
S The s orbital looks like a ball or sphere.
S The p orbital looks like a dumb-bell.
S These orbitals are all perpendicular to each other.
S The d orbitals have two shapes.
S 4 of the 5 look like “4-leaf clovers.”
S The 5th one looks like a “big dumb-bell” with a “hula-hoop”
around the middle.
S The shapes of the f orbitals are complex.
S We have a slide showing them, but you don’t need to remember
them, nor will they be on the test. But s, p and d will be.
Shapes of s, p, and d Orbitals
S In the s block, electrons are going into s orbitals.
S In the p block, the s orbitals are full.
S New electrons are going into the p orbitals.
S In the d block, the s is full but the p orbitals are not full.
S New electrons are going into the d orbitals, because we are in the transition
metals. THIS is characteristic of the d block.
f orbitals
g orbitals = Science Fiction?
S Dr. Seaborg predicted the g orbitals would start with element number
121, which has not been invented yet. The g block will have 18
S Will his hypothesis be proven true?
To Summarize
Total Orbitals
Total Electrons
per Level
1 (1s orbital)
1 (2s orbital)
3 (2p orbitals)
1 (3s orbital)
in your
(3p orbitals)
5 (3d orbitals)
as6 we
discuss this.
S The first level (n=1) has an s orbital. It has only 1.
orbitals in the first
s are no other
1 (4s orbital)
2 energy level.
3 (4p orbitals)
7 (4f orbitals)
S We dcall this orbital
the 1s orbital.
5 (4d orbitals)
Island of Stability
S This is another hypothesis from Dr. Seaborg. His thought was that element 114
would be an “island of stability,” especially if it also had 184 neutrons. It would
aehv a mass number of 298.
S However, other “islands” might be 120 or 126. Detailed and complicated math
calculations are necessary to figure out these numbers.
S Most synthesized elements only last for fractions of seconds. However, in 1998
researchers synthesized element 114 and it lasted for 30 seconds. Perhaps this is
the “shore” of the Island of Stability that Dr. Seaborg hypothesized.
S The element 114 was made using some of the original Pu-244 that Dr. Seaborg
himself made in the early 1940s. They bombarded plutonium with Ca-48 atoms
to form some of the new element 114.
S Element 114 is now know as Flerovium (symbol Fl); it was named in 2012.
S It took 14 years to agree on the name.
S All of the atoms so far have had mass numbers of 285-289. Therefore, the “island”
still remains undiscovered.
Island of Stability
Famous picture of the “Island of Stability” showing the island off in the distance (top
right) with 114 protons and 184 neutrons. An element with Z = 184 is also predicted to
be another “island of stability.”
Timeline = Homework
S Check out the History timeline on page 133 in your book.
S Prepare a timeline listing the major developments listed up to 1932.
S Answer questions 1&2 at the bottom of the page (2 requires a 5-
sentence paragraph as a minimum requirement).
S Add 3 things to your timeline that have happened in Chemistry SINCE
1935 that you think are significant. You might have to do research to
answer this.
Electron Configurations
Section 2
S What do I mean by “electron configuration?”
S The electron configuration is the specific way
in which the atomic orbitals are filled.
S Think of it as being similar to your address.
The electron configuration tells me where all
the electrons “live.”
Rules for Electon Configurations
S In order to write an electron configuration,
we need to know the RULES.
S 3 rules govern electron configurations.
S Aufbau Principle
S Pauli Exclusion Principle
S Hund’s Rule
S Using the orbital filling diagram at the
right will help you figure out HOW to
write them
S Start with the 1s orbital.
S Fill each orbital completely and then go
to the next one, until all of the electrons
have been accounted for.
S FOLLOW the arrows!!!
Fill Lowest Energy Orbitals
Each line represents
ONE orbital.
1 (s), 3 (p), 5 (d), 7 (f)
S The Aufbau Principle states
that electrons enter the lowest
energy orbitals first.
S The lower the principal quantum
number (n) the lower the energy.
S Within an energy level
S s orbitals have the lowest energy
S followed by p then d and then f.
S f orbitals are the highest energy
for that level.
No more than 2 Electrons in Any
S The next rule is the Pauli Exclusion
S The Pauli Exclusion Principle states that
an atomic orbital may only have 1 or 2
electrons and then it is full.
S The spins have to be paired.
S We usually represent this with an up arrow
and a down arrow.
Hund’s Rule (Dog’s Rule?)
S Hund’s Rule states that when you
get to degenerate orbitals, you fill
them all half way first, and then
you start pairing up the electrons.
S Degenerate means they have the
same energy.
S p orbitals are degenerate because
there are 3 of them on EACH
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
S d and f orbitals are also degenerate.
Let’s Try Some…
S NOW that we know the rules, we can try to write some electron
S Remember to use your orbital filling guide to determine WHICH
orbital comes next in the sequence (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p,
S Follow the arrows!!
S Lets write some electron configurations for the first few elements,
and let’s start with hydrogen.
S There are also shorthand electron configurations, but we will look
at those after Chapter 6.
Electron Configurations
H Z=1
He Z=2
1s2 (1s is now full)
Li Z=3
Be Z=4
1s22s2 (2s is now full)
N Z=7
Ne Z=10
(2p is now full)
Na Z=11
Cl Z=17
K Z=19
Sc Z=21
Fe Z=26
Br Z=35
Electron Configurations of
Alkali Metals (and H)
H Z=1
Li Z=3
Na Z=11
K Z=19
Exceptions to the Rules for
Electron Configurations
Cr should be
1s22s22p63s23p64s23d4 (Z=24)
BUT Cr is
1s22s22p63s23p64s13d5 (d half full)
Cu should be
BUT Cu is
1s22s22p63s23p64s13d10 (d is full)
More HW…OMG!
S Chemistry: write full electron
configurations for elements 1-36.
S Also write orbital diagrams for 3-10
S Advanced Chemistry: write full
electron configurations for 1-36 + Rb,
Sr, Y, Ag, I, Kr, Cs, Ba, La, Ce, Hf, Pb.
S Also write orbital diagrams for 11-18
Emission Spectra = Fingerprint of
the Elements (Section 3)
S Atomic emission spectrum is sometimes
called a line spectrum, to distinguish it from
the continuous spectrum.
Emission Spectra = Fingerprint of
the Elements (Section 3)
S The top 3 (H, Hg, Ne) are emission spectra.
S The bottom one is an absorption spectrum of H.
Emission Spectra = Fingerprint of
the Elements
S Atomic emission spectra are “unique.” You can use
the spectrum to identify the element (like a
S Bohr’s model predicted and explained emission
spectra by pointing out how electrons can move from
one energy level to another.
S His model also explained why metals glow red when
they are heated.
S Scientists can look at light from a distant star and
analyze it and determine what types of elements
make up that star.
S Just by looking at the light!
S No element (except H) has those same 4 lines in its
All the Rest of Section 3….
S …was covered in the Chemical History power point.
S Photoelectric effect
S A photon is a quantum of light. It is light behaving as a
particle. A photon has a certain wavelength, frequency
and energy.
S De Broglie equation
S Showed that particles could also act as waves.
S Heisenberg’s uncertainty principle
S Principal = Dr. Gordon
S Principle = a statement that explains how or why
something works scientifically
The End

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