Electron Orbital Diagrams
Placing Electrons in Orbitals
• Every orbital may hold up to 2 electrons
• These 2 electrons have opposite spins: an up
spin and a down spin
• You can determine how many electrons are
possible for an energy level by 2n2
• If n=1, only 2 electrons are possible
• If n=2, 8 electrons are possible
• If n=3, 18 electrons are possible
Placing Electrons in Orbitals
• There are rules for placing electrons in
orbitals.
– Hund’s Rule
– Aufbau Principle
– Pauli’s Exclusion Principle
• There is also an order which we follow in
placing electrons in an orbital energy
diagram.
Orbital Diagrams
•The number is the
energy level (n value)
•The letter is the orbital
type or the sublevel
type
•The number of lines
beside the sublevel is
how many orbitals
there are on that
sublevel
•Do you notice any
pattern?
Orbital Diagrams
•For every n value, the
s orbital is the lowest
energy.
•ns < np < nd < nf
•It looks like a pattern
starts, but then it is
broken as 3d is higher
in energy than the 4s.
•So you memorize this!
•How do you place the
electrons in these
orbitals? Let’s try.
Orbital Diagrams
• For ions like F- or Ca2+, their orbital diagram
will be just like a Noble Gas.
• This is an empirical observation called the
Octet Rule:
– Atoms tend to lose, gain, or share electrons in
order to have 8 valence electrons.
• What are valence electrons?
– Valence electrons are the electrons in the
outermost energy level (highest n value) which are
furthest from the nucleus.
Orbital Diagrams
• You can easily count the number of valence
electrons from the Periodic Table for Groups
1-2 and 13-18 (Representative or Main Block
Groups)
• This means that you can predict the ion
charges for these groups as well!
Orbital Diagram Exceptions
• There are 19 exceptions to the rules when
drawing Orbital Diagrams!
• You are only responsible for knowing 5: Cr,
Mo, Cu, Ag, Au.
• These 5 elements all do the same thing: as the
(n-1)d orbital is VERY close in energy to the
ns orbital, a ns electron is “promoted” to a
“higher-energy” (n-1)d orbital.
• This will either fill or half-fill the d orbitals.
Orbital Diagram Exceptions
• Because the d orbitals are filled or half-filled,
the atom gains stability overall.
• These “exceptions” are thus the ground-state
for these atoms.
• If you were to draw other exceptions to the
rules, these would be classified as “excited”
states.
Electron Configurations and
Noble Gas Configurations
Electron Configurations
• Orbital energy diagrams tell us what
electrons/orbitals an element fills in its ground
state (most stable state).
• But they are tedious to draw and they must be in
a figure.
• So how can we get the same info but be able to
write it in a sentence?
• Electron configurations let us do this.
• All we do is translate the vertical orbital diagram
into a horizontal line. (well, almost all)
Electron Configurations
• What about the electron configuration for Sr?
• When you write the correct electron configuration
for the ground state of Sr, notice that you group
the n values together.
• This is different from how you wrote the orbital
diagram which is strictly based on filling order.
• Why do we write the electron configuration this
way?
• Convention and because that way the valence
electrons are at the end of the configuration: and
valence electrons are lost first!
Noble Gas Configurations
• But even electron configurations get long and
tedious to write.
• Write the electron configuration for gold!
• How can we make this shorter?
• Well, the Noble Gases have filled s and p
subshells, so they are good reference elements.
• So we write Noble Gas configurations, which are
based on the LAST Noble Gas BEFORE the
element we are working with.
Periodic Table and Electron
Configurations
• Guess what?
• The Periodic Table is laid out in order of
increasing Atomic Number.
• It is also organized by the order in which the
electrons fill! (Except for the exceptions.)
• Thus, we have the s-block, the p-block, the dblock, and the f-block.
Periodic Properties
Electron Configurations and
Periodic Properties
Periodic Properties & Valence Electrons
•Members of a group have similar chemical
properties. Why?
•They have the same number of valence
electrons, or their outermost electron
configurations are the same.
•And valence electrons are responsible for
chemical properties.
Periodic Properties & Valence Electrons
•Valence electrons are outer shell electrons
and so are furthest from the nucleus
(generally).
•Thus they are not as tightly held to the
nucleus.
•So they are removed more easily than core
electrons (inner shell electrons).
Periodic Properties & Valence Electrons
• Also, the core electrons with their negative
charges block or insulate or "shield" the valence
electrons from feeling the full nuclear charge.
• So the shielded valence electrons don't feel the full
nuclear charge, Z, they instead feel what we call
Zeff, or the effective nuclear charge.
• Think of picking up iron filings or paper clips
with a magnet, eventually the magnet can’t
attract any more filings as it is “shielded”.
Periodic Properties & Valence Electrons
• Members of the same group have the same
number of valence electrons and have similar
properties as those valence electrons behave
similarly.
• For example, all of the Alkali Metals have 1
valence electron, with an outer electron
configuration of ns1.
• This is why Alkali metals are so reactive, their 1
valence electron is easily removed in chemical
rxns.
• And what is their e- configuration then?
Periodic Properties & Valence Electrons
• Alkaline Earth metals all have an ns2 outer
electron configuration, or they have 2 valence
electrons.
• What is the outer electron configuration for
halogens?
Atomic Radii
• But chemical reactivity is not the only periodic
property or trend which we can see in the Periodic
Table.
• Another important Periodic Trend or Property is
the atomic radii, or the radius of an atom, or 1/2 the
diameter of the atom.
• We may measure the radius of an atom.
• It is also common to measure the distance between 2
nuclei in a diatomic element.
• Then the atomic radius would be half of this
distance.
Atomic Radii
There are 2 trends in atomic radii that you can
explain and see from the Periodic Table:
• As you go down a Group, the atomic radius
increases.
• Now you are adding an energy shell which by
definition is further from the nucleus!
Atomic Radii
• As you go across a Period from left to right, the
atomic radius of the elements DECREASES.
• This is because you are adding protons so Zeff is
increasing. As Zeff increases, the outer electrons are
pulled closer to the nucleus, so the radius shrinks.
• This is particularly true as you add s and p
electrons in the same energy level. They don't shield
each other very well, so Zeff increases as you add
protons.
Atomic Radii
• What about the different orbital types, do they have
different attractions, or different Zeff?
• For example, for n = 4, there are 4s, 4p, 4d, and 4f
orbitals. They don’t have the same energy, do they
have different Zeff?
• They do!
• In general, as the subshell energy increases on the
same principal energy level, the Zeff decreases.
• So for the same n value, the s has the highest Zeff.