Dmitri Mendeleev is known as the father of the periodic table. He originally arranged the
elements in order of their increasing atomic mass. If organized as such, the elements show a
distinct periodicity to their properties. Mendeleev even left room for elements that had not yet
been discovered but that he predicted existed! The ordering of the elements by their atomic
mass was done at the time because atomic number was not yet known! Today, the periodic table
is arranged by atomic number – and as previously mentioned a few chapters ago – that mimics
atomic mass with a few exceptions. It is now known that it is the outermost electrons that are
involved in chemical reactions. These electrons are called valence electrons. The vertical
columns, called groups, are families of elements that behave similarly chemically. Each column,
as previously mentioned, is identified by number and a letter (A or B). The horizontal row is
called a period. As we go across the row, we add electrons first into the same subshell (e.g. we
fill the s) and then we begin to fill additional subshells (e.g. the p and d subshells) until
eventually we reach the noble gasses.
Hydrogen only has one electron. We have discussed how the Bohr theory worked realllllly well
to describe the single electron in a hydrogen atom, but when examining multi-electron atoms,
his theory could not account for electron-electron repulsion, electron attraction with the
nucleus, and more importantly, it was determined that electrons do not orbit the nucleus. The
orbitals of many electron atoms differ from hydrogen in their electron energies and in the most
probable distances of the electrons from the nucleus.
Schrödinger’s model for the atom was then adopted. Instead of orbits (like planets around the
sun) he perceived orbitals – 3-dimensional structures within which an electron can travel
around the nucleus.
Previously, we talked about how to calculate the energy of an electron in hydrogen using:
E = -2.179 x 10-18 J
n2
where n = the level the electron is in
However, electron energies in many electron atoms depend on n (the level they are in – their
distance from the nucleus) but they also depend , to a lesser extent, on what orbital they are in:
s, p, d, or f. In a many electron atom, the energy of the electron in a given shell increases with
the l value of its orbital (e.g. an electron in the 3s shell has less energy than the electrons in 3p
which has less energy than an electron in the 3d shell). Also, the energy of an electron in a
given orbital decreases as the atomic number (number of protons) increases. Why? Think
about it, if the nucleus contains 3 protons has 3 electrons in it is can “pull” or attract those
electrons with some force. However, when the nucleus contains 56 protons and there are 56
electrons, that nucleus has more “pull” on those first few electrons and thus can attract them
closer to the nucleus than the atom with only 3 protons.
1
We’ll discuss this a bit later when we talk about shielding . . .
When electrons for multi-electrons atoms are in their lowest energy state it is important to know
that they do NOT all congregate in the 1s orbital! Each electron has its own place in the
electron. Each electron has a level (n), and a subshell (s, p, d or f), and a particular orbital (px,
pz, pz for example) that it resides. An important rule regarding electron organization within the
atom is:
Pauli’s Exclusion Principle: Only two electrons can occupy a given orbital, and those two
electrons within the orbital must have opposite spins. This means that no two electrons in an
atom on the periodic table will have the same set of 4 quantum numbers. This exclusion
principle should not be underestimated, as the resulting electron distributions are largely
responsible for the physical and chemical properties of the element.
The fourth quantum number:
spin quantum number: (ms): indicates the direction of the electron spin. There is a maximum
of 2 electrons contained in a particular orbital (px, py, pz, or dxy, dz2, etc . . .) and each electron is
said to have its own spin. There are only two possible values for spin – one electron is said to
have a spin of +1/2 the other is said to have a spin of -1/2. Arbitrarily, the first electron is given
the +1/2 assignment. Now all electrons on the periodic table can be described by a unique set of
4 quantum numbers – unique meaning that no two electrons in the same atom will share the
same four quantum numbers! Given the quantum numbers for electrons in an atom, you could
determine which elements/element those quantum numbers correspond to. As we journey
through the periodic table, we will see that elements have similar electron configurations: the
electrons fill up orbitals in a specific manner, such that all begin their electron configurations the
same way – but at the end – they have a unique location for the electron and the location for that
electron, thus, if you write quantum numbers for every element’s electrons on the periodic
table, even though the quantum numbers for all electrons in the 1s subshell or 2 s subshell are
the same, the final set written for any element will be unique to that element, making the total
set of all quantum numbers unique to that particular element. Doing electron configurations
will help you see this!!
Given: Write the quantum numbers for each electron in the same 3d subshell (e.g. pretend that
the electrons are both in the dxy orbital)
For electron #1
For electron #2
n=
l=
ml =
ms =
2
Notice that the two electrons share the first three quantum numbers. They are in the same level
(n=3) they are in the same d subshell so they share the same l value , and they are in the same d
orbital, so they share the same ml values. Since each electron is said to have its own direction
within the orbital, the first is assigned the +1/2 value and the second is assigned the -1/2 value.
A hydrogen atom only has one electron, for that reason, all other sublevels above the 1s subshell
have the same energy. There are no electrons in them to repulse one another, and there are no
electrons in them to have any attraction to the nucleus. One consequence of an atom containing
more than one electron is the splitting of the energy levels – e.g. the 2nd level into sublevels as
the electrons repulse one another and attract with the nucleus. Therefore, the 2nd energy level
splits its 2s and 2p from one another creating two “levels” each a different distance away from
the nucleus.
Evidence for the energy levels splitting can be seen in the spectral line data from multi-electron
atoms. The lines get more and more complex as the number of electrons increases, indicating
more available energy levels which the electrons can jump to in the excited state and which the
electrons can then fall from emitting photons of light.
The farther apart opposite charges are the weaker their attraction to one another. When the
nucleus and the electron are far apart the potential energy is high (e.g. less stable), than when
the nucleus and the electron are close together.
The greater the charges are, the stronger the attraction. The more protons, or positive charge
the nucleus contains, the greater the attraction is with the electrons in the atom, the lower the
energy and the more stable the system is.
The energy of the orbital is measured in terms of how much energy is needed to strip that
orbital of its electrons. It takes more energy to remove an electron that is closer to the nucleus
(more stable) than it does to remove an electron that is far away (less stable) from the nucleus.
Orbital energy depends on three factors: the charge of the nucleus, electrons repulsions, and
orbital shape
Nuclear charge: the higher the nuclear charge, the lower the orbital energy, the more stable the
system, the greater the nucleus-electron attraction. Thus, the more protons that the atom has,
the higher the nuclear charge.
Electron repulsions: like charges do NOT want to be anywhere near each other! We have
talked about that previously. Each electron in an orbital “feels” not only the nuclear attraction
but also the electron-electron repulsion with the other electron in the same orbital. This
repulsion counteracts the attraction that each electron feels in the same orbital. Thus, a filled
orbital (one which contains two electrons) raises the orbital’s energy (makes it less stable)
because of this repulsion. There is in effect, some shielding going on within the same orbital.
3
The text gives the example of comparing He – which has 2 protons and 2 electrons, with He+1
which has 2 protons and now only 1 electron (remember, to make a positive ion we lose an
electron. He is in row 1, it is part of the s block, and it is the second element over on the
periodic table. Since each orbital can hold 2 electrons, He’s electron configuration is written: 1s2.
1 indicating the level, s indicating the orbital, and 2 indicating how many electrons are in that
orbital, in this case it is filled.
eep+
p+
p+
p+
eHe+1: 1s1
He: 1s2
Two things need to be considered when comparing the atom to the ion. The first is that the
orbital, in this case the 1s orbital, for He contains 2 electrons. These electrons are going to
repulse one another. This repulsion and subsequent shielding makes is easier to remove the
electron than when compared to removing the electron from He+1 on the right. HOWEVER,
He+1 has already lost an electron from its 1s shell. So there is no electron electron shielding
going on but now we have nuclear charge coming into play. There are still 2 protons, but now
instead of 2 protons interacting with 2 electrons (as in the neutral atom on the left) we have 2
protons attracting 1 electron in the ion. This “pulls” that electron even closer to the nucleus,
forming a greater attraction with the nucleus and making it MUCH more difficult to remove.
Thus, it is not a simple comparison between atom and ion when comparing which electron will
be lost more easily.
A second type of shielding occurs with the electrons in the lower orbitals compared to electrons
in the higher orbitals. The electrons in the inner orbitals spend 90% of their time between the
electrons in the outer shells and the nucleus. Clearly, inner electrons do more shielding than
electrons in the same sublevel!! It is said that these inner electrons reduce the nuclear charge to
some degree, making it seem, to the outer electrons, that the nuclear charge is somewhat less
than the total number of protons present. This is termed the effective nuclear charge.
The
4
effective nuclear charge gets less and less the more electrons there are in the atom. Thus, the
outer electrons in an atom are shielded more so than the inner electrons.
The probability diagrams which result from Schrödinger’s wave equations show that with
increasing l value, there is less and less chance for the electron to be found near the nucleus.
Therefore, not only do the diagrams show that density towards the nucleus decreases as the n
value increases, density also decreases as the l value increases. This is termed penetration. The
lower the l value, the easier it is for the electron to penetrate closer to the nucleus – and the
closer the electron is to the nucleus, the greater its attraction with the nucleus and the harder it
is to remove that electron from the atom.
We have already previously discussed that only 2 electrons can be held per orbital. We have
also discussed that there is 1 s orbital, 3 p orbitals, 5 d orbitals, and 7 f orbitals. Since each
orbital can only hold 2 electrons that means that at most, there are 2 electrons in the s subshell, 6
electrons in the p subshell, 10 electrons in the d subshell, and 14 electrons in the f subshell. We
also examined the periodic table and saw how the s block, p block, d block, and f block contain
the same number of elements as possible electrons that the orbitals can contain (e.g. the s block
is 2 elements wide, the p block is 6 elements wide, the d block is 10 elements wide, and the f
block is 14 elements wide). We write electron configurations straight off the periodic table, and
the electrons fill the subshells according to their distance from the nucleus.
We will be writing the ground state electron configurations for neutral atoms. The atomic
number from the periodic table equals the number of protons and also the number of electrons
for a neutral atom.
Writing the electron configuration we first list the principal quantum number followed by the
orbital that electron is located in, then we list the number of electrons that orbital contains as a
superscript. The sum of the superscripts must add up to the total number of electrons in the
atom!
Use the periodic table, very similar to a type writer, fill the levels and the orbitals keeping in
mind that the s orbital has a maximum of 2 electrons, the p orbital has a maximum of 6 and the
d orbital has a maximum of 10. We will not do f orbital electron configurations at this time.
Given: write the ground state electron configurations for Li, C, Ne, Na, Ca, Ti and Br
Li has 3 electrons. Looking at the periodic table you see that Li is in the second row first
column, and is in the s block. Therefore we KNOW that it will have a 2s in its electron
configuration. We begin, however, at the beginning of the periodic table. First we have to
realize that there are two other electrons in front of that 2s shell. Therefore we write the electron
configuration as:
5
Li3: 1s22s1 read 1 s 2, 2 s 1. When we sum the superscripts we see that 2+1 = 3 which does
indeed indicate the number of electrons in lithium.
Now do the same for C, Ne, Na, Ca, and Ti and Br
C:
Ne:
Na:
Ca:
Ti:
Br:
Write quantum numbers for several of the electrons in the above elements using different n
values.
Remember when you get to the d subshell that the principal quantum number is 1 behind the
row number. Also, notice that you drop down to the lower principal quantum number AFTER
beginning with the higher principal quantum number for the s subshell (e.g. fill the 4s first then
fill the 3d, then jump back to the 4 p). The order of filling shows us that the 3d is lower in
energy than the 4p subshell, but higher in energy than the 4s subshell. Eventually, the farther
away the electron gets from the nucleus, the closer together the subshells get and the lines
between them get a little blurry. With elements with larger numbers of electrons, sometimes the
electron configurations might not be what you would expect given the normal filling patterns
we have discussed to this point.
It is important to pay attention to the fact that within a group, the outer electron configurations
are very similar. Notice when writing electron configurations for the elements, specifically the
noble gasses, whether or not electrons fill up the orbitals. Notice that the noble gasses have
filled s and p orbitals. Remember that noble gasses are also termed inert gasses, so perhaps this
filled orbital situation leads to unreactivity . . . (!!!) Orbitals are filled in the same order over and
over again for the elements, which leads to outer electron configurations that recur periodically,
which leads to chemical properties that recur periodically.
If filled orbitals mean stability, then perhaps the same can be said of 1/2 filled orbitals as well.
Given what we know about ion formation it would seem that electrons want to take their place
in the atom such that they form a more stable species as opposed to a less stable one (we have
talked about stability extensively!!) and thus, when given the chance to form a 1/2 filled or filled
orbital, the electrons take it. This seems to happen with some of the transition metals –
specifically the 4s23d4 and the 4s23d9 elements Chromium, Molybdenum, Copper, Silver, and
Gold. Notice the figure below for the anomalous elements on the periodic table.
6
You should know the NAMES and how to write electron configurations for the elements
previously mentioned in chapters 1-7 AND now know names and how to write electron
configurations from elements 1-54.
Categories of electrons:
1.) inner (core) electrons: those electrons up to the previous noble gas and any completed
transition series. These make up the lower energy levels of the atom
2.) outer electrons: those in the highest energy level (highest n value), these electrons are
the farthest away from the nucleus.
3.) valence electrons: those electrons involved in chemical reactions. For the main group (A)
elements, they are the outer electrons, for the transition metals (B), sometimes inner d
electrons can be involved as well.
7
Group and Period numbers:
1.) for main group elements (A), the number of outer electrons is the group number (1-7)
2.) the period number (row number) is the highest n value for that row of elements
3.) 2n2 gives the total number of electrons possible for a particular period number. For
example, if n=3, the 2*(32) equals 2*9=18. 2 electrons in the s subshell, 6 electrons in the p
subshell and 10 possible electrons in the d subshell shows us that yes indeed, 18 electrons
are possible in the third level.
We have previously determined electron configurations for the elements. And the easiest way
to do so is to begin at the beginning of the periodic table and add one electron per element to
the lowest energy level available. This requires knowing which level is the lowest in energy!!
The periodic table does “tell” us this information but below is a line diagram which might help
you also:
8
You can see that as the n value increases, the discrepancy in energy begins to get smaller and
smaller and smaller. This explains why electron configurations begin to do some crazy, not so
predictable things when you get into larger electron containing atoms (like the f block!)
Orbital box diagrams are just what they sound like – a box which indicates the orbital that the
electron is occupying, and its spin. Traditionally, the first electron is placed in the box arrow up
and that indicates the positive spin. The second electron in the orbital is placed in the box arrow
down, which indicates the equal – but opposite – spin state. Once again, just like the orbital can
only hold two electrons (Pauli Exclusion Principle) so can the box only hold two electrons AND
they must be of opposite spin. The orbital diagrams will show all of the available orbitals even
if they do not contain an electron. This shows us that all the orbitals for a specific n value are
equal in energy.
We have been writing ground state electron configurations following the aufbau principle,
because the configuration for each atom is based on building off the configuration for the
previous atom. We follow the same principle for filling our orbital boxes as well.
The other rule to filling subshells is Hund’s Rule: which tells us “how the hund they fill up”.
Electrons occupy the orbitals of a subshell singly at first, then they pair up. Remember,
electrons want to minimize repulsions between them!! Since the s orbital only has one box, the
second electron will get paired with the first, but remember that there are 3 p orbitals – a px, a
py, and a pz. So the box diagram will look like:
px
py
pz
INCORRECT
px
py
pz
CORRECT
The first organization violates Hund’s rule. Electrons will first fill up 1 in each available orbital
BEFORE pairing. Electrons want to minimize electron electron repulsions!!! Remember also,
that the first electron placed in the orbital is assigned as spin “up” and when paired, the second
electron will be assigned spin “down”.
Draw orbital box diagrams for: Be, B, C, N, O, F, and Ne paying careful attention to how the
orbital boxes fill up.
9
You will not be writing shorthand notation when asked to write electron configurations – you
can use the noble gas shorthand method when drawing orbital box diagrams. Pay attention to
the last column in the figure below:
The noble gas shorthand notation is used to simplify/shorten the electron configuration.
Once again, this will ONLY be used for drawing orbital box diagrams. Noble gas
10
configurations will NOT be accepted for written electron configurations – only for orbital
box diagrams when indicated!!!!
All physical and chemical behaviors of an element are based, ultimately, on their atom’s
electron configuration. Three properties of atoms that are directly influenced by the atom’s
electron configuration are: atomic size, ionization energy, and electron affinity.
Trends in Atomic Size
Schrödinger has shown us that electrons lie in orbitals. And solving the wave function has
given us probability diagrams which show us the probable distance and location of those
electrons from the nucleus. Atoms are spherical in shape, and our orbital diagrams confirm
this. When we measure atomic size, we do not look at a single atom by itself, instead we
examine the atomic radii of that atom next to another atom. Theoretically, we measure the
distance between two identical atoms and divide that distance in half. But atoms are not hard
spheres, so their size can change a little bit depending on what other element they are bonded
to.
Metallic radii: one-half the distance between nuclei of adjacent atoms in a crystal structure
composed of that element (a metal).
Covalent radii:
(non-metals).
one-half the distance between the nuclei of identical covalently bonded atoms
Factors that Influence Atomic Size
1.) As the principal quantum number increases there are more outer electrons and they get
farther and farther away from the nucleus making the atom larger in size.
2.) As the number of protons increases the effective nuclear charge for the outer electrons
increases and this tends to pull the electrons in closer
Net Effect of two factors
1.) Looking at atomic size DOWN a group/column, the principal quantum number
dominates. This means that larger principal quantum numbers mean larger atoms: size
INCREASES
2.) Looking at atomic size ACROSS a period/row effective nuclear charge dominates. This
means that electrons are added, but they are added to the same subshells – e.g. adding
11
more p electrons into the p subshell doesn’t make the p subshell bigger. In fact, for the
main group elements (the A elements) these outer electrons shield each other poorly, so
the electrons feel a greater effective nuclear charge and are better able to attract to the
nucleus making them smaller in size: size DECREASES
Transition metals are not as well behaved as the main group elements. The inner d electrons
are sandwiched between the s and p for those particular elements across the periodic table.
As we move from left to right across the row, initially the size shrinks because of increasing
nuclear charge (more protons). From then on however, once you get some d electrons in the
orbitals, those d electrons shield a little and the size doesn’t change all that much. For
example, Vanadium 4s23d3 has the same atomic radius as Zinc (4s23d10). Just note that there
are always exceptions and these are just general “guidelines” to follow when thinking about
an atom’s size
The ionization energy (IE) is the amount of energy needed to remove 1 mole of electrons from 1
mole of gaseous atoms or ions. Removing an electron from an atom or ion requires energy to
overcome the attraction that the electron feels with the nucleus. Because the energy is used on
the system, ionization energies are always (+) in value.
When we talk about Hydrogen, you can calculate this energy by thinking about the transition
being between n=1
n=, at which point the electron is no longer associated with the
atom, and we have generated an ion. Many electron atoms can lose more than one electron –
we have seen this already when we examined ions such as Ca+2, Cu+2, Cr+3. The first ionization
energy is always less than the second which is always less than the third which is always less
than the fourth etc. . . basically, it gets harder and harder to remove more and more electrons
since the number of protons is now greater than the number of electrons, the electrons left
behind are held onto with a greater force of attraction by the nucleus. It also becomes very
difficult to continue to remove electrons once all the valence electrons are removed. The core
electrons generally do not want to leave the ion.
Atoms with low ionization energies tend to easily lose their electrons and then tend to easily
become positive (e.g. metals known as cations) Atoms with high ionization energies do not
want to lose their electrons, and tend not to. These atoms also tend to have an affinity for
gaining an electron (instead of losing) and therefore tend to be negative species (e.g. nonmetals
known as anions).
The relationship between size and ionization energy is such that the larger the atom is, the
easier it becomes to remove an electron. Those electrons that are far far far away from the
nucleus are feeling a lower effective nuclear charge, they are in orbitals whose principle
quantum number indicates that they are far away from the nucleus, and it becomes quite easy to
strip them away. Quite frankly, it’s the equivalent of me being the nucleus and noticing if you
don’t show up to your physics class later in the day – the nucleus has NO idea what’s going on
out there!
12
Ionization energies therefore generally decrease down a group
Ionization energies tend to increase across a period since the size of the atom decreases
the electrons are closer to the nucleus that you might imagine, this would be the
equivalent of you not showing up for chemistry class, and yes, the nucleus noticed!
Some exceptions to consider: if removing an electron will result in a stable filled or halffilled subshell it is a little easier to remove that electron than you might imagine. Also,
removing the first electron in a new subshell (e.g. the p1) would leave behind a stable ns2
level behind.
Electron affinity (EA) the amount of energy needed to add 1 mole of electrons from 1 mole of
gaseous atoms or ions. Again, there is a first EA energy and a second EA energy associated
with the addition of multiple electrons to an atom or ion. We have seen evidence of this
occurrence in the formation of S-2, N-3, and O-2 etc. . . The first addition of an electron results in
a monovalent (-1) charge on the ion. Energy is released by the system we say that EA energies
are negative in value. The second EA value is positive as energy must be used in order to
overcome electron-electron repulsions that are associated with shoving more electrons into the
ion.
Trends are not always as clear, but there are some general ones to keep in mind:
Reactive nonmetals (VIA, VIIA) do NOT want to lose electrons (have high ionization
energies, and want to gain electrons (have very negative electron affinity energies).
Reactive metals have low ionization energies and slightly negative electron affinities
indicating the desire to lose, not gain electrons.
Noble gases have very high ionization energies and positive electron affinities. This
means that the noble gases want neither to lose nor gain electrons. In fact, only the larger
noble gases, Kr, Xe, and Rn form compounds at all!
Metals are located to the left of our staircase line. They are typically shiny solids that have
moderate to high melting points, they are good thermal and electrical conductors. Metals tend
to lose their electrons in reactions with non-metals.
Nonmetals are located to the right of our staircase line. They are typically not shiny, have
relatively low melting points are poor thermal and electrical conductors, and tend to gain
electrons in reactions with metals.
Metalloids are located in the region between the two classes and have properties of both.
Metallic behavior increases from the right side (the non-metal side) to the left (the metal side)
and increases down a row. Remember that a characteristic of a metal is that it wants to lose an
13
electron (or two) and that it is very easy for larger atoms to lose electrons (remember size
increases down a column).
Once again, there are always exceptions to the classifications. For example, carbon is a nonmetal but in its graphite form it can conduct electricity. Iodine, a non-metal is a shiny solid.
Gallium and cesium are metals, but they melt at temperatures below body temperatures and
mercury is a liquid at room temperature. But since we have learned there are always exceptions
we will try and focus on the generalizations that we can make about metals and their behavior.
Tendency to lose electrons: metals tend to lose their electrons in chemical reactions because they
have lower ionization energies compared to non-metals. Examining Group 5A: N, P, As, Sb,
Bi. Nitrogen is a gaseous nonmetal, Phosphorus is a solid non-metal. Both occur as -3 anions.
Arsenic and Antimony are metalloids, Sb happens to be the more metallic of t he two, but
neither one readily forms ions. Bismuth, the largest atoms of the group behaves as a metal,
forming mostly ionic compounds as a +3 ion.
As we move across a period, it becomes harder and harder to lose an electron (higher IE) and
easier to gain an electron (more negative EA). With regard to monatomic ions (single atoms
become ions) elements on the left hand side of the periodic table lose their electrons to become
cations, while elements on the right hand side tend to gain electrons becoming anions.
Previously we have talked about ion formation for main group elements and I told you that the
ion charge could be determined by looking at the number over the column (e.g. Group 1A all
have a charge of +1 and Group 2A atoms form +2 ions). Now we find out WHY!!
Why is Na+1 and not Na+2, why is F-1 and not F-2? Why are the noble gases so unreactive?
Electron configurations, and writing electrons configurations answers the why.
As was mentioned previously, noble gases have very high ionization energies and positive
electron affinities, they do not form ions, and they are chemically stable due to their ns 2np6
electrons configuration.
Elements in Groups 1A, 2A, 6A, and 7A readily gain or lose electrons in order to attain this
ns2np6 electron configuration. This is termed a noble gas configuration and the ion is said to be
isoelectronic with a noble gas (the same number of electrons).
The valence electrons are the electrons that are involved in chemical reactions and are also the
subshells that contain the electrons that will either be lost, or they are the subshells to which an
element will gain electrons. The valence electrons are located – for the main group elements –
in the HIGHEST principle quantum level. (the highest n value).
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Writing electron configurations for the ions:
1.) Write the electron configuration for the neutral atom
2.) Rewrite the electron configuration grouping the same principal quantum numbers
together. Add or subtract electrons from the highest principle shell
Na+1: 1s22s22p63s1
1s22s22p6 notice the generation of the noble gas like configuration
Br-1: 1s22s22p63s23p64s23d104p5
1s22s22p63s23p64s23d104p6
The larger metals in groups 4A, 5A, 6A, and 7A would have to lose too many electrons to
become isoelectronic with the noble gases. They will actually gain electrons. Transitions metals
would have to lose too many electrons to become isoelectronic with noble gases, therefore, they
will only lose a few electrons in order to try and have a filled, or half-filled subshell.
For example, let’s examine the two ions formed by Tin (Sn)
Sn+2: 1s22s22p63s23p64s23d104p65s24d105p2 rewrite the electron configuration grouping the n
values together.
Sn+2: 1s22s22p63s23p64s23d104p64d105s25p2:
1s22s22p63s23p64s23d104p64d105s2
Sn+4: 1s22s22p63s23p64s23d104p64d105s25p2
1s22s22p63s23p64s23d104p64d10
Transition metals also would have to lose too many electrons in order to reach a noble gas
configuration. Thus, these metals will lose their electrons from the highest principle quantum
shell and also some from the (n-1)d shell at times.
The filling order, and the stability diagram above shows that the ns subshell is lower energy
than the (n-1)d subshell. However, even though the 4s subshell is more stable, the 3d subshell is
still closer to the nucleus, so they are not shielded by the 4s electrons in that subshell. As a
result, as the electrons fill the d subshell, it becomes more stable. In effect, a crossover of orbital
energy occurs. This means that the electrons are instead lost from the 4s subshell instead of the
3d. A quick and easy way to remember this is “first in, first out”. The electrons fill the 4s first,
and they leave the 4s first. Thus, the electrons are lost from the highest principle quantum shell,
which we mentioned earlier.
Rules for ion formation
1.) For main group (#A) elements, s block, remove ALL electrons from the highest n level
2.) For main group (#A) elements, p block, remove np electrons before ns electrons
3.) For transition elements, d block, remove ns electrons before (n-1)d electrons
4.) For non-metals, add electrons to the p orbitals of the highest n value
15
Another way of determining if an electron configuration is correct (the first is atomic spectra
and exciting electrons and paying attention to their emission photons) is to examine the
magnetic properties of the element and its compounds.
The spinning electron generates a tiny magnetic field. Only atoms with one or more unpaired
electrons, paired electrons have balanced spins – equal and opposite to one another and
therefore are not affected. A species with unpaired electrons exhibits paramagnetism, and it is
attracted by an external magnetic field. A species with all electrons paired exhibit
diamagnetism, it is not attracted to, in fact it is slightly repelled by a magnetic field.
Write electron configurations for Zn and Mn+2. Are they diamagnetic or paramagnetic?
Cations are smaller than parent atoms: when electron are removed from the atom, the number
of protons remains the same and those protons in the nucleus are able to pull the electrons
closer due to reduced e-/e- repulsions
Anions are larger than parent atoms: when electrons are added to the atom, this increases the e/e- repulsions in the atom causing them to take up more space. The number of protons is less
than the number of electrons present and they cannot pull the electrons as effectively.
Across a periodic table, the size decreases as the cations are formed - +1 are larger than +2 which
are larger than +3. Then we enter the non-metal side of the periodic table and begin to form
anions. -3 are larger than -2 which are larger than -1 anions. Down a group, ionic size increases
because the number of electrons increases and are placed in higher energy levels. There are
more electrons in Fr than K and they both form +1 ions, but Fr has more electrons, so losing one
is not really a big deal to the atom concerning its overall size.
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The greater the positive ionic charge, the smaller the cation.
The greater the negative ionic charge, the larger the anion.
Which ion is smaller?
Ca+2, Sr+2, Mg+2
K+1, Cl-1, S-2
Au+1, Au+3
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