Writing Electron Configurations
Ok...let’s simplify this.
• Every atom has a nucleus.
• In that nucleus we have protons (positive
charge) and neutrons (no charge)
• Surrounding the nucleus we have electrons.
• Electrons have a negative charge, however
they do not “fall” into the nucleus (despite
being attracted to the positive proton), why?
Ok...let’s simplify this
• Because electrons exist in “orbitals”
• An orbital is the region of space in which you
will find an electron. (For our purposes)
• This is where it gets confusing...
• There are different levels of orbitals.
• We represent these different levels of orbitals
with the letter “n”.
• So if an electron is in the 1st orbital
surrounding the nucleus it is n=1.
• If an electron is in the 3rd orbital surrounding
the nucleus it is n=3.
• Easy enough, right?
There’s more.
• We also have the letter “l” this represents
what type of orbital the electron is in.
• The different types of orbitals are as follows...
• S, p, d, f, and g
• These letters refer to the shape of the orbital.
• So now, we have two pieces of information to
help us determine where an electron is.
• We know the number of orbital it is in and we
know the shape of the orbital.
Review: Quantum Numbers
• Principle Quantum Number (n) represents the
energy orbital number (how far the orbital is
from the nucleus).
• Orbital-Shape Quantum Number Orbital (l)
represents the shape of the orbital sublevel (s
is round, p is like 2 balloons, 2 d orbitals put
together)
• Magnetic Quantum Number represents the
direction of the orbital sublevel (l)
• n can be any number from 1 on.
• l can be any number from 0 to (n-1)
• ml is –l to +l
Ok...back to new stuff
• The Fourth Quantum Number (ms) is used to
represent the direction of the spin on the
electron.
• There are only two possible values for this
number, + ½ or – ½ .
• Electrons spin like a top.
Electron Spin
Pauli Exclusion Principle
• We have a new principle to learn!
• The Pauli Exclusion Principle states that no
two electrons in any atom will have the same
4 quantum numbers.
• For example you would not find two electrons
with n=4, l=3, ml=-3 and ms= - ½
• You could find n=4, l=3, ml=-3, ms= - ½ and
n=4, l=3, ml= -3, ms= + ½
Example
Electron Configurations
• An electron configuration can be written for
each atom.
• That is, we can write out where we can find
each electron in an atom.
• And remember, no two electrons will have the
same set of 4 quantum numbers in the same
atom.
• Generally when we write electron
configurations we write the configuration of
the atom in it’s ground state.
• Ground state means that the electron is
neutral and is not an isotope.
Example
• Hydrogen, in it’s ground state, has one
electron (atomic number 1, mass of 1).
• The electron configuration for Hydrogen is 1s1.
• Remember, n=1, l=0 and the small, raised 1
symbolizes that there is one electron in the s
sublevel.
• Helium is He 1s2 symbolizing there are 2
electrons in the s sublevel.
• Lithium is 1s22s1 symbolizing that there are 2
electrons in the 1 s sublevel and 1 electron in
the 2 s level.
• Fluorine is written as 1s22s22p5
• Why do you think that there are only 2
electrons in the s sublevel, but there can be 5
electrons in the p sublevel?
Remember the Pauli Exclusion
Principle?
• What did the Pauli Exclusion Principle tell us?
• That no two electrons can have the same 4
quantum numbers in an atom.
• Let’s write out the possible values of the 4
quantum numbers in the s sublevel.
• How many electrons do you think is possible
in the p sublevel?
• How about the d sublevel?
• There is a maximum of 2 electrons in the s
sublevel.
• There is a maximum of 6 electrons in the p
sublevel.
• There is a maximum of 10 electrons in the d
sublevel.
Aufbau Principle
• Each electron occupies the lowest energy
orbital available
• Atom’s are “built-up”
• Do you think that the direction of the arrow
symbolizes anything?
• Why do you think that in the Carbon electron
configuration the 2 p level has two electrons
in the same direction?
• We can tell the electron configuration of an
element just by looking at the periodic table!
• The “long form periodic table” shows us which
sublevel contains the atom’s valence
electrons.
A helpful guide to what goes where
Let’s try it.
• How would we write the configuration for:
– Hydrogen
– Helium
– Boron
– Carbon
– Nitrogen
Noble Gas Notation
• The last column of the periodic table is called
the Noble Gases.
• The Noble Gases have full valence electron
shells.
• So we when we out electron configurations
for our elements we can shorten it like this...
Shorten Electron Configuration
Notation
• Ne
– 1s2 2s2 2p6
• Na
– 1s 2 2s2 2p6 3s1
– [Ne]3s1
Exceptions!
• There are almost always exceptions to the
rules!
• In this case, Chromium and Copper do not fill
their energy shells as we would expect them
to.
• How so and why?
• Evidence shows that these two elements
achieve a ground state in ways that do not
follow the rules.
Hidden Information
• Recall that in the periodic table a group runs
vertically (up and down) and a period rums
horizontally (left to right).
• Recall that the last group number is also the
number of valence electrons in that atom.
Hidden Information
• The n value of the highest energy level shell
(where the valence electrons are) is also the
period number where you find the element.
s-block
• The s-block includes elements in group 1 and
2 (and hydrogen and helium)
• In the s-block all elements have their valence
electrons in an s block (be it 3s, 4s, etc.)
p-block
• The p-block includes elements in groups 13-18
• The elements in this block have their valence
electrons in the p level.
d-block
• The d-block includes all of the transition
elements (groups 3-12)
• These elements have valence electrons in the
d shell
f-block
• We (in our class) have not looked at any
electron configurations in the f-block, but we
know that it exists.
• The f-block consists of the “inner transition”
metals.
Recall
• We said that the last number of the group
number (i.e. 15....5) is the number of valence
electrons in an atom.
• We also said that the last (highest) n value is
the period number.
• So, using the periodic table what element is
represented by [He]2s2 ?
• How about this one?
• [Kr]5s24d105p5 ?
Questions
• What would the valence shell be for Arsenic
(As)?
• 4p
• Let’s see: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
• OR
• [Ar] 4s2 3d10 4p3
• How about Cesium (Cs)?
• 6s
• Let’s see:
1s2 2s22p6 3s23p63d10 4s24p64d10 5s25p6 6s1
• OR
• [Xe] what? How can we determine what
comes next? Xe valence shell is 5p
• How many more electrons does Cs have more
than Xe?
• What sell do we fill after 5p?
• [Xe]6s1
Atomic Radius
• Atomic radius is a measurable property of an
atom.
• What is radius?
Atomic Radii
• decreases across a period (left to right)
• Increases down a group (up and down)
• There are two factors that affect atomic radii
• When n increases, that means we are moving
farther from the nucleus and thus giving an
larger area of finding an electron and thus a
larger atomic radii.
• So it makes sense that as you increase n
(remember n corresponds to the period
number) that the atomic radii would increase
down a group.
• Another factor that affects atomic radii is
called Zeff, this stands for the attraction
between the nucleus and the electrons.
• The further away from the nucleus a valence
shell is, the more it will be shielded from the
positive attraction of the nucleus.
• HOWEVER, as you travel across a period, the n
value stays the same (same distance away) but
each time an electron is added and the
attraction to the nucleus gets stronger.
Ionization Energy
• Ionization Energy is the energy needed to
completely remove an electron from an atom.
• Hint: an ion is an atom with a charge, meaning
it has lost an electron or has gained an
electron. Remember that ionization energy
has to do with removing an electron, thus
creating an ion.
• The trends of ionization energy are the exact
opposite as atomic radii.
• Ionization energy decreases down a group and
increases across a period.
Trends in Ionization Energy and
Chemical Reactivity
• Summarize the table on page 155 and create a
chart in your notebook.
Electron Affinity
• Change in energy that occurs when an
electron is added to a gaseous atom.
• What does gaseous atom mean?
• Read and summarize the rules of electron
affinity on page 156.
Homework
• Ionization Energy and chemical reactivity chart
• Electron affinity summary
• Section summary (pg. 157) questions #1, 3
and 4.