Download Chapter 2 – Atoms and Elements

Chapter 8 – Atomic Electron Configurations
and Chemical Periodicity
We have seen that an atomic orbital is described by three
quantum numbers: n, l, and ml. An electron in the atomic orbital
is also described by these three quantum numbers as well as a
fourth, the electron spin magnetic quantum number, ms.
Electron Spin and Magnetism
Electrons behave as if they spin on their axes. Because they are
charged, this generates a _________________________ around
the electron (i.e. the electron acquires a magnetic dipole).
When an electron is exposed to a magnetic field, its dipole can
take on one of two possible orientations. It can line up with or
be opposed to the dipole of the external magnetic field:
Electrons with one of these spin orientations are assigned a spin
quantum number of ms = +½; the rest are ms = -½. There are no
other possible values for ms.
When a beam of hydrogen atoms is passed through a magnetic
field, it will split into two beams – one of atoms in which the
electron has ms = +½ and one of atoms in which the electron has
ms = -½. When a beam of helium atoms is passed through the
same magnetic field, it does not split. This is because every
helium atom contains two electrons, one with ms = +½ and one
with ms = -½. We say that the two electrons in the helium atom
are spin-paired. Effectively, their spins cancel each other out.
Substances in which all of the electrons are paired are called
______________________. When exposed to a magnetic field,
they are weakly repelled by it.
Substances in which one or more electron is unpaired are called
______________________. When exposed to a magnetic field,
they are attracted to it.
Substances which have many unpaired electrons that have all
had their spins aligned are called ________________________.
They can be used to make magnets since they do not require an
external magnetic field to keep the electron spins aligned.
Assigning Electrons to Atomic Orbitals
In 1925, Wolfgang Pauli (1900-1958) stated the Pauli exclusion
principle: No two electrons in an atom can have the same set of
four quantum numbers (n, l, ml, and ms).
Because all electrons in an atomic orbital have the same n, l, and
ml, and there are only two possible values for ms, each atomic
orbital can only contain ______________ electrons.
In order for an atom to be in its ground state, each electron must
have the lowest energy wavefunction permitted by the Pauli
exclusion principle. In other words, we put electrons in the
lowest energy orbitals first. (Orbitals are not actually ‘boxes’ or
‘slots’ to be filled but it’s easiest to think of them that way.)
We can draw orbital box diagrams for the first four elements:
H
(1e-)
He (2e-)
Li (3e-)
Be (4e-)
1s
2s
2p
The orbitals are listed in order of increasing energy (from left to
right). Electrons are filled in from left to right, completely
filling one type of orbital before moving to the next. Arrows
show the electrons’ relative spin (‘up’ or ‘down’).
Electron configuration can also be written using a shorthand
notation which lists how many electrons are in each subshell.
e.g. Be = 1s22s2
e.g. Write electron configurations for H, He and Li.
Recall that the number of orbitals in a subshell is 2l+1. Since
two electrons can be placed in each orbital, 2(2l+1) electrons
can be placed in each subshell. You may find it easier to just
remember: s orbitals can have up to ____ electrons each (l=0),
p orbitals can have up to ____ electrons each (l=1),
d orbitals can have up to ____ electrons each (l=2),
f orbitals can have up to ____ electrons each (l=3).
When we assign electrons to orbitals in a p, d or f subshell, we
have two options: (a) fill one orbital at a time then move to the
next, or (b) assign one electron to each orbital then go back and
add a second electron once all orbitals are half full. Unpaired
electrons have lower energy if their spins are aligned, and this is
only possible using method (b).
Knowing this, we can now draw orbital box diagrams for
elements five to ten:
B
(5e-)
C
(6e-)
N
(7e-)
O
(8e-)
F
(9e-)
Ne (10e-)
1s
2s
2p
While there are three p orbitals, they are not listed separately
when writing electron configuration. e.g. Ne = 1s22s22p6
So far, we have filled orbitals starting at the lowest shell and
working up. This order breaks down when we get up to n = 3.
This is because subshells with more nodes are higher energy
than subshells with fewer nodes.
Instead of filling orbitals in order of increasing n, we should
really be filling them in order of increasing n + l (using n as a
‘tiebreaker’). See Figure 8.4 in Kotz et al.
This is the aufbau order and is often illustrated as follows:
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
This gives the following general order for filling orbitals:
1s, 2s,
2p, 3s,
3p, 4s,
3d, 4p, 5s,
4d, 5p, 6s, …
e.g. Write electron configurations for the following elements:
Ti
Pb
Electron configurations for the larger elements are lengthy to
write out. Because of this, chemists have adopted a short cut
called noble gas notation in which the symbol for a noble gas is
used as an abbreviation for its electrons.
e.g. Ne has an electron configuration of 1s22s22p6.
For Na, we can write either 1s22s22p63s1 or [Ne]3s1.
In this way, the core electrons are represented by the noble gas
while the configuration of the valence electrons is still listed.
Conveniently, it turns out that the periodic table is arranged
according to electron configuration. The properties which were
used to group the elements together into columns were caused
by those elements having similar electron configurations!
Now that we know how to write electron configurations, we can
build the periodic table ourselves by writing the last orbital filled
(and how many electrons went in it):
1s1
2s1
3s1
4s1
5s1
2s2
3s2
4s2 3d1 3d2 3d3 etc…
5s2 4d1 4d2 4d3 etc…
2p1
3p1
4p1
5p1
2p2
3p2
4p2
5p2
2p3
3p3
4p3
5p3
2p4
3p4
4p4
5p4
2p5
3p5
4p5
5p5
1s2
2p6
3p6
4p6
5p6
As such, it is also possible to determine electron configuration
by following along the periodic table, using it to tell us what
type of orbital to fill next.
Identify the elements with the following electron configurations.
a) 1s22s22p3
c) [Ne]3s23p3
b) 1s22s22p63s23p64s23d7
d) [Kr]5s24d5
How many core and valence electrons do these atoms have?
a) ____core, ____valence
c) ____core, ____valence
b) ____core, ____valence
d) ____core, ____valence
Valence electrons are particularly important because they are the
electrons that ‘do the chemistry’. (Electrons in a d or f subshell
are only considered to be valence electrons if it is not full.)
Exceptions to the Aufbau Order
There are a few exceptions to the aufbau order, most of which
stem from the fact that full and half-full subshells are most
stable. As such, if an atom gains a full or half-full subshell by
breaking from the aufbau order, it often will.
e.g. Cr is not [Ar]4s23d4; it is [Ar]4s13d5
Cu is not [Ar]4s23d9; it is [Ar]4s13d10
All of these exceptions are elements from either the ‘d-block’ or
the ‘f-block’, so they are either transition metals (Cr, Cu, Nb,
Mo, Ru, Rh, Pd, Ag, Pt, Au), lanthanides (Ce, Pr, Gd), or
actinides (Th, Pa, U, Np, Pu, Cm).
Electron Configurations of Ions
We use the periodic table to determine how many electrons
reside in each atom. If an atom is neutral, this will be the
__________ number (Z).
For an ion, the number of electrons = Z – charge.
Thus, an anion will have ______ electrons than the neutral atom
while a cation will have ______ electrons than the neutral atom.
Subshells in anions are filled using the ____________________.
Subshells in cations are filled in order of increasing ____.
Write electron configurations for the following atoms and ions:
as a neutral atom
as an ion
+
Li
P3Ti+
Sn2+
Sn4+
We can use electron configurations to predict which ions are
likely to form. As a general rule, ions try to achieve:
1. the same configuration as the closest noble gas,
2. a pseudo noble gas configuration (noble gas configuration
plus a full d or f subshell), or
3. a noble gas configuration for everything but d or f electrons
Which of the following ions are likely to form? For those which
are not, what ion would you expect to form from that element?
a) O2c) Cl+
e) Pb4+
b) Mg6-
d) Ca+
f) Ga3+
Effective Nuclear Charge
Why is the 4s orbital in neutral atoms lower in energy than 3d?
Looking at the radial probability density plot for an s orbital, we
see that there is a small (but significant) fraction of electron
density in the region very close to the nucleus. Thus, an electron
in an 4s orbital spends a small (but significant) amount of time
very close to the positively charged nucleus. This is not the case
for a 3d orbital (which has nodal surfaces but no radial nodes):
Most electrons do not ‘feel’ the full positive charge of the
nucleus. Other electrons in the atom (particularly those in lower
energy orbitals) ‘shield’ some of this charge. The amount of
positive charge ‘felt’ by an electron in a given orbital is called
the effective nuclear charge (Z*).
The following table lists the atomic number (Z) and effective
nuclear charges (Z*) for electrons in the 2s and 2p orbitals of
neutral atoms of the elements in the second period:
Element
Li
B
C
N
O
F
Z
3
5
6
7
8
9
Z* (2s)
1.28
2.58
3.22
3.85
4.49
5.13
Z* (2p)
2.42
3.14
3.83
4.45
5.10
Note that the effective nuclear charge on an s orbital is slightly
higher than on a p orbital in the same shell. Why?
Also, note that Z* does not increase by 1 when Z increases by 1.
Why not?
Effective nuclear charge explains several of the periodic trends
(atomic properties that can be predicted using the periodic table)
including atomic radius, ionization energy and electron affinity.
Atomic Radius
• half the distance between nuclei in an A–A bond
• _____________________ from left to right across a period.
As the effective nuclear charge on the valence electrons
increases, they are attracted more strongly to the nucleus.
• _____________________ from top to bottom down a group
as more shells of electrons are added.
Ionization Energy (aka First Ionization Energy)
• the energy that must be absorbed in order to remove a
valence electron from a neutral atom in the gas phase:
• _____________________ from left to right across a period.
As the effective nuclear charge on the valence electrons
increases, they are attracted more strongly to the nucleus.
• _____________________ from top to bottom down a group
as more shells of electrons are added, and the valence
electrons are more shielded.
• the energy required to remove a valence electron from an
atom that has already had one electron removed is the second
ionization energy (IE2):
Electron Affinity
• the energy released when a neutral atom in the gas phase
acquires an extra electron in the lowest energy orbital
available:
• electron affinity values are negative, so an ‘increase’ in
electron affinity is represented by a more negative number
• _____________________ from left to right across a period
except for the noble gases. As the effective nuclear charge on
the new valence electron increases, it is attracted more
strongly to the nucleus.
• _____________________ from top to bottom down a group
as more shells of electrons are added, and the new valence
electron is more shielded.
Ionic Radius
• periodic trends for ionic radii are the same as for atomic radii
• the radius of an anion is ______________ than that of a
neutral atom of the same element. This is because adding the
extra electron __________________ Z*.
• the radius of a cation is _______________ than that of a
neutral atom of the same element. This is because removing
the electron __________________ Z*. Often the valence
electrons of a cation are in a lower energy shell than the
valence electrons of the corresponding neutral atom. This
also significantly __________________ the ionic radius.
Important Concepts from Chapter 8
• diamagnetic vs. paramagnetic vs. ferromagnetic substances
• electron spin
• Pauli exclusion principle
• drawing orbital box diagrams
• writing electron configurations
• using the aufbau order and predicting exceptions to it
• using electron configuration to predict if an ion will form
• core vs. valence electrons
• effective nuclear charge
• periodic trends (atomic radius, ionization energy, electron
affinity, ionic radius)