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Packet 8
Periodic Properties of the Elements
last updated: 5/24/2017
CHE 170 Packet 8 - 1
Concept Area I: Terminology
atomic orbital energy
diagram
Aufbau principle
degenerate
diamagnetic
effective nuclear charge, Z*
electron affinity
electronic configuration
ferromagnetic
filled subshell
ground state
half-filled subshell
Hund’s rule
inner-shell/core electrons
ionization energy
Noble gas notation
octet rule
orbital
paramagnetic
shell
shields, shielding
subshell
valence electrons
CHE 170 Packet 8 - 2
Concept Area II: Magnetism and Atomic Structure
a. You should know how to determine if a
substance is paramagnetic or diamagnetic.
b. You should understand how magnets can be
formed from ferromagnetic elements.
CHE 170 Packet 8 - 3
Paramagnetic and Diamagnetic
paramagnetic:
Liquid oxygen is attracted to
strong magnets because oxygen
is paramagnetic – it has
unpaired electrons.
similar to Tro page 379
CHE 170 Packet 8 - 4
Paramagnetic and Diamagnetic
diamagnetic:
Liquid nitrogen is not attracted
to strong magnets because
nitrogen is diamagnetic – it has
no unpaired electrons.
CHE 170 Packet 8 - 5
Paramagnetic and Diamagnetic
How can we show that a sample is paramagnetic?
iron(III) oxide
Nifty mnemonic devise for para and diamagnetic?
CHE 170 Packet 8 - 6
How does paramagnetism allow the formation
of magnets for ferromagnetic elements?
Well, if all of those unpaired
electrons can be aligned……
all of their little magnetic
fields would add up to one
larger field – a magnet!
We can align all the
little electrons by
exposing them to a
very strong magnetic
field.
CHE 170 Packet 8 - 7
Concept Area III:
Atomic Orbital Energy Level Diagrams
a. You need to remember what the four quantum numbers tell
you (names of orbitals, how many degenerate orbitals).
b. You need to know how many orbitals of each type are
present in a shell or subshell of specific energy.
c. You need to know how many electrons can go into each
orbital, subshell and shell.
d. You should know that the different energies of atomic
orbitals control the order in which they fill up with electrons.
e. You should be able to use the Aufbau Principle and know
when and why to break it for certain elements.
f. You must be able to draw an atomic orbital energy level
diagram or write an electron configuration for any atom or
ion
CHE 170 Packet 8 - 8
Lots of Things to Remember!
shell: atomic orbitals of similar energy are grouped
together by their principal quantum number, n.
The energy level is shown by counting numbers:
.
subshell: atomic orbitals of the same type/shape are
grouped together, this is given by the angular
momentum quantum number, ℓ.
What are the names of the different orbitals?
How many degenerate orbitals for each?
.
.
atomic orbital: this is where the electrons live, this is
given by the magnetic quantum number, mℓ.
How many electrons can live in each kind of orbital?
.
.
valence electrons: those electrons that are in the
outermost shell of the atom
CHE 170 Packet 8 - 9
Let’s review everything!
CHE 170 Packet 8 - 10
Generic Atomic Orbital Energy Diagram Handout
Generic Atomic Orbital Energy Diagram
6d
5f
7
6
7s
6p
5d
5
4f
6s
5p
4d
4
5s
4p
3d
Energy
3
4s
3p
3s
2
1
2p
2s
1s
Note that all the
degenerate orbitals are
written on the same line
to indicate that they have
the same energy and are
in the same subshell.
Note that the subshells
have a number by them
to indicate what shell they
are in.
Note that the energies of
the different shells
overlap.
Note: These energies are not plotted to scale. In fact, the further
up the diagram we go, the orbitals get closer in energy.
CHE 170 Packet 8 - 11
How do we draw one of those things?!
First, we have to remember the order that we
fill the orbitals:
Second, easy as counting!
One subshell in the first shell:
Two subshells in the second shell:
Three subshells in the third shell:
Four subshells in the fourth shell:
Finally, we have to remember how many
degenerate orbitals for each kind:
Now, we use this mnemonic device…
CHE 170 Packet 8 - 12
How to Easily Create an Atomic Orbital Energy Diagram Handout
How do we know the energy order to list the orbitals?
This handy mnemonic device!
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f 7g 7h 7i
6f 6g 6h
5f 5g
4f
CHE 170 Packet 8 - 13
Looking at the periodic table, when does a
new subshell start filling?
CHE 170 Packet 8 - 14
The Aufbau Principle
1. Electrons are placed in the lowest
energetically available (i.e. unfilled) orbital.
2. An orbital can hold at most two electrons.
3. If two or more degenerate orbitals are
available (for example: p, d etc.) then
electrons should be spread out before they
are paired up (Hund's rule).
CHE 170 Packet 8 - 15
Hund’s Rule
Just like most folks prefer to sit alone on a city bus,
electrons prefer to be alone (fewer repulsions) in
an orbital. So, if they have the choice, they stay
alone!
CHE 170 Packet 8 - 16
Time to use Hund’s Rule!
1. So, put 3 electrons in a p subshell:
2. Now, 4 electrons in a p subshell:
3. Now, try 7 electrons in a d subshell:
CHE 170 Packet 8 - 17
Putting it all together….
Writing atomic orbital energy diagrams and
electronic configurations!
1. Determine how many electrons the atom has
(same as atomic number on periodic table)
2. Use an energy level diagram and put in
electrons one by one using the Aufbau
principle
spdf notation for electron configurations
3. If requested, con- for H, atomic number = 1
vert diagram to an
1
# of
s
1
electron configuraelectrons
tion or write it
value of ℓ
value of n
directly.
CHE 170 Packet 8 - 18
Let’s write the atomic orbital energy diagram
and then write the electron configuration!
Let’s use Chlorine atom as our example:
How many electrons?
Diagram:
3d
4s
3
3p
3s
2p
2
Energy 1
2s
1s
So, it would be:
Or with noble gas notation:
CHE 170 Packet 8 - 19
Now try a few with a neighbor…
Write atomic orbital energy diagram (use the handout) and the
electronic configuration for the following atoms:
1. Carbon
C
2. Fluorine
F
3. Sodium
4. Iron
Na
5. Zirconium
CHE 170 Packet 8 - 20
What is the correct orbital order for
electronic configurations?
Let’s use krypton to answer this.
Its complete electronic configuration listed on
the previous slide was…
1s22s22p63s23p63d104s24p64d25s2
The orbitals were listed from innermost to
outermost shell.
Instead, we could list them in the order that
they energetically fill…
1s22s22p63s23p64s23d104p65s24d2
Thus, it really doesn’t matter which of these two
ways that the orbitals are listed!
However, wait a few slides and we’ll see when
we do cations that the shell order is better.
CHE 170 Packet 8 - 21
Exceptions to the Rule!
Remember the Aufbau principle?
Guess what, some elements break that rule!
There’s another competing rule:
full subshells and half-full subshells are more
stable than other configurations.
So, for d or higher subshells, we can fill one electron
out of order to get to a “full” or “half-full” subshell.
Note: Cr, Cu, Mo, Ag, etc.
Some of the other transition metals do further
rule-breaking stuff (like Pd and Pt). Don’t worry
about why but do realize there are some more
details we aren’t learning!
CHE 170 Packet 8 - 22
Valence Electrons
Where are the valence electrons?
Why are valence electrons important?
Because they are located on the outside of the atom.
These are the electrons that are removed to make cations –
need to know how many valence electrons so you know the
maximum amount that you can remove.
To make anions, electrons are added to the valence shell –
need to know how many valence electrons so you know
how much room you’ve got!
Is there an easy way to determine number of valence
electrons?
Yes! Let’s go to the next slide to see how…
CHE 170 Packet 8 - 23
How many valence electrons?
The periodic table is very useful to tell how
many valence electrons we have in an atom!
For main group elements, simply count from
left to right: 1,2,3,4,5,6,7,8. (Note, use the group
numbers with “A” after them.)
Transition metals are a bit trickier because of
the d orbitals. They should all be two, but
because of the rule breaking they do, they have
to be determined in context of a compound.
Notes page has the labeled periodic table.
CHE 170 Packet 8 - 24
Electron Configurations of Ions
How do we write an electron configuration for an
anion?
Simple, really!
Anions are just like atoms. Write the atom’s
configuration, and then, add in the additional
electron(s) to the next available spot(s).
Let’s try one!
CHE 170 Packet 8 - 25
What is the electron configuration for a chloride ion?
How many electrons in the atom?17
Diagram for the atom:
3d
4s
3
3p
3s
2p
2
2s
Energy 1
1s
So, e.c. for atom would be:
We actually have a
chloride ion. What is
the symbol & charge?
So, we need to add
how many electrons?
So, e.c. for ion is:
Or in noble gas
notation:
Or with noble gas notation:
CHE 170 Packet 8 - 26
Electron Configurations of Ions
How do we write an electron configuration for an
cation?
This is a bit more complicated!
In a cation we remove electrons from an atom.
We have to remove the outermost electrons
on the atom. Now, the outermost electrons
sometimes are at lower energy. Therefore,
we’re not always removing the last electrons we
added!
Let’s try one!
CHE 170 Packet 8 - 27
What is the electron configuration for a iron(II) ion?
How many electrons in the atom?
Diagram for the atom:
3d
4s
3
3p
3s
2p
2
2s
Energy 1
1s
So, e.c. for atom would be:
Or with noble gas notation:
We actually have a
iron(II) ion. What is
the symbol & charge?
So, we need to
remove how many
electrons?
Now, what is the
outermost subshell?
So, e.c. for ion is:
Or noble gas notation:
CHE 170 Packet 8 - 28
Why don’t you try a few…
Write the electronic configurations for the following.
1. fluoride ion
2. oxide ion
3. calcium ion
4. germanium(II) ion
5. vanadium(II) ion
Now that we’ve done those. Are any of these
ions paramagnetic?
CHE 170 Packet 8 - 29
Concept Area IV: Applications
a. You should remember that the periodic table is arranged so
that elements with similar properties are by each other.
b. You need to be able to calculate an approximation of the
effective nuclear charge of an atom or ion and use it along
with atomic theory to explain some of trends on the periodic
table.
c. You need to recognize the role that ionization energy and
electron affinity play in the chemistry of the elements and be
able to explain the observed trends.
d. You should be able to predict whether an atom becomes
larger or smaller when it becomes an ion.
e. You should be able to explain why we can use the periodic
table to predict the most commonly formed ion for main
group elements.
CHE 170 Packet 8 - 30
Mendeleev (from Russia, 1834-1907)
and the Periodic Table
He ordered elements by atomic mass.
When he did so, he saw a repeating
pattern of properties.
Periodic law: When the elements are arranged in order of
increasing atomic mass, certain sets of properties recur
periodically.
So, he arranged elements with
similar properties in the same column.
Then he used the patterns he saw to predict
properties of undiscovered elements (Ga and Ge).
When atomic mass order did not fit observed
properties, he reordered (Te and I)
CHE 170 Packet 8 - 31
Periodic Trend for the Reactivity of Alkali
Metals with Water
Lithium
Sodium
Potassium
And, for more alkali metal fun: http://youtu.be/uixxJtJPVXk http://youtu.be/uixxJtJPVXk
And, for more sodium fun: http://youtu.be/HY7mTCMvpEM http://youtu.be/HY7mTCMvpEM
CHE 170 Packet 8 - 32
What versus Why
Mendeleev’s periodic law allows us to predict
what the properties of an element will be based
on its position on the table.
But, it doesn’t explain why the pattern exists.
Side note – true of all laws. Tells what we observe
without explaining why it is that way.
Quantum mechanics is a theory that explains
why the periodic trends in the properties exist.
electron configurations/atomic theory
effective nuclear charge, Z*
CHE 170 Packet 8 - 33
Effective Nuclear Charge
Effective nuclear charge, Z*,
of an electron is how much
of a “pull” from the nucleus
it feels.
The greater the pull from the
nucleus, the less repulsive
force it feels from the other
electrons. Electrons with less
repulsive force will be at
lower energy.
The text goes into great detail
about how electrons in s
orbitals shield better than
those in p orbitals shield
better than d orbitals and so
on.
Tro page 274
CHE 170 Packet 8 - 34
Calculating Effective Nuclear Charge
We’re going to ignore that fact that different
electrons shield differently for our calculations
and discussions.
However, whenever an atom seems to buck the trend,
if we account for different electrons shielding
differently, those atoms aren’t so strange after all.
Now, the effective nuclear charge, Z*, of a valence
electron in an atom can be approximated by
subtracting the number of inner-shell or core
electrons from the number of protons:
Z*=Z – S
(where Z is actual nuclear charge and S is the charge screened by other electrons)
CHE 170 Packet 8 - 35
Calculating Effective Nuclear Charge
What’s an inner-shell or core electron?
Remember, the atomic number tells us the total
number of protons and the total number of
electrons in an atom.
Remember to determine number of valence
electrons, we count the main group elements left
to right.
CHE 170 Packet 8 - 36
Effective Nuclear Charge explains why
atoms get smaller across a period.
Z* on
Li’s atomic mass = 6.941amu.
C’s atomic mass = 12.011 amu.
Ne’s atomic mass = 20.179 amu.
CHE 170 Packet 8 - 37
Why do atoms get
smaller from left to
right across the
periodic table?
So, size decreases across a period owing to increase
in Z*. Notice that neon had an effective nuclear
charge of +8 versus lithium at +1.
Therefore, the valence electrons on Ne “feel” more
of a pull towards the nucleus, and the valence
electrons on Li are more shielded from the nucleus
and “feel” less of a pull.
Since the valence electron for Li is more shielded,
Li is larger than Ne, even though Ne has more
electrons and more mass than Li!
Tro page 285
CHE 170 Packet 8 - 38
Atomic Radii Trends
The atomic radii
are given here in
picometers for
main group
elements.
So, atomic radii
.
across a period.
This is because the
effective nuclear
charge, Z*,
.
across a period.
CHE 170 Packet 8 - 39
Atomic Radii Trends - plotted
Notice the periodic trend in atomic radius, starting at a peak with
each alkali metal and falling to a minimum with each noble gas.
Tro page 285
CHE 170 Packet 8 - 40
Smaller atoms hang on to
their electrons more, why?
If two atoms have the same effective
nuclear charge, why is it more
difficult to remove an electron
from the smaller atom?
Well, opposites
. Plus, the closer they are
together, the
that
is!
Thus, smaller atoms hang on to their electrons
than larger atoms of the same effective
nuclear charge because the valence electrons are
to the nucleus in the smaller atoms. This is
also manifested by smaller atoms having
ionization energies.
Note: although ionization energies are determined in the gas phase, we can use those
values to decide whether an atom in the solid/liquid phase would be more likely to
gain or lose an electron. We just don’t have a specific value for solid/liquid.
Tro page 295
CHE 170 Packet 8 - 41
Smaller atoms hang on to their electrons more.
CHE 170 Packet 8 - 42
So, why do we have dips in our ionization energies?
Remember how
completely full
and half-full
sublevels are more
stable?
So, each of those
dips occurs where
if we remove an
electron, we’ll
have a more stable
configuration!
Similar image on Tro page 294 that includes transition metals
CHE 170 Packet 8 - 43
Why do atoms on the top-right side of the table
(ignoring noble gases) have a higher electron affinity?
Well, the more positive the Z*, the more
likely the atom can capture a wayward electron
if it is not energetically bad for it to do so.
Again, even though these values are measured in the gas state, we can generalize that
information to other states.
Tro page 297
CHE 170 Packet 8 - 44
CHE 170 Packet 8 - 45
Ionic Size
How about if we add valence
electrons; would the ionic radii
for an anion be larger or
smaller than the atomic radii?
Why?
CHE 170 Packet 8 - 46
Ionic Size
How about if we take away
valence electrons; would the
ionic radii for a cation be larger
or smaller than the atomic radii? Why?
CHE 170 Packet 8 - 47
Ion size compared to atom size.
Similar images on Tro pages 291 and 292
CHE 170 Packet 8 - 48
What charge doth an ion have?
Using the periodic table to predict the charge on an ion
formed from an atom.
Be2+
B3+ C4±
Si4+
Ga3+ Ge4+
Here are some charges on some common monatomic cations and anions.
Metals usually form cations and nonmetals usually form anions.
Learn how to get the periodic table to tell you the most commonly
formed ion. The transition metals are trickier because of the d shells;
their charges must be determined in context.
CHE 170 Packet 8 - 49
Noble Gases and the Octet Rule
The noble gases have eight valence electrons.
Except for He, which has only two electrons.
We know that the noble gases are especially
nonreactive.
He and Ne are practically inert.
The noble gases are so nonreactive because
the electron configuration of the noble gases is
especially stable which means low energy.
Other atoms would like to be at lower energy
too, so…
CHE 170 Packet 8 - 50
They try and look like one:
Noble Gases and the Octet Rule
Once we have the atomic orbital diagram (or
electronic configuration), simply decide if it would
be easier to add or take away electrons to look like
a noble gas and reach stability/lower energy.
We call this the octet rule because most noble gas
configurations have eight valence electrons.
Eight, two or zero electrons in the valence shell is
usually the magic number, why?
Elements in the 3rd period can sometimes act like
they have more than eight valence electrons, why?
CHE 170 Packet 8 - 51
The End of Packet 8
Any Questions?
Homework assignment!
The Lewis Dot Handout is due when
we get to Lewis Dot structures in the
next chapter!
CHE 170 Packet 8 - 52