Build your own periodic table at home

Build your own periodic table at home
The first and most important thing to note about the periodic table is that it is
periodic; viz., its structure reflects the recurrent patterns found in the chemical and physical
properties, the periodicity if you like, of the elements. These patterns however were not
apparent to us until rather recently; until the advent of modern chemistry in fact.
In the distant past, we did not know the true chemical composition of the world
around us. As a result, the true nature of the composition of the stuff around us, and of
which we too are made was more a matter of philosophical speculation than of scientific
discovery. At least 2,500 years ago, people in Ancient Greece and in India were wrestling
with the question of what things are made of, and their speculations and arguments
eventually did help to induce the birth of modern science. However, their musings were not
based on anything resembling the modern scientific method (which had not been invented
yet), being rather a combination of unstructured observations and logical deductions. If you
are interested in the thoughts and ideas of these people, please look for the works of the
‘Nature Philosophers’ both on this site (under the rubric ‘Notes for H’) and elsewhere on the
Raphael’s famous painting generally known as Scuola di Atene (School of Athens) in which are
depicted some of the more famous Ancient Greek Philosophers (though not those, like
Leucippus and Democritus, who were of indirect significance to atomic chemistry)
Roughly two millennia after the work of these Nature Philosophers, and once we
managed to get through the terrible interregnum caused by the Medieval Church’s
retarding grip on the development of scientific ideas: a period which, though great for art,
really put a serious damper on both the advancement of scientific thought and the
development of scientific thinking in Europe, however, things finally did change, and we
ended up being able to begin investigating the world in a fundamentally new and different
way; I refer of course to the birth of modern science, a fascinating story in its own right, and
one detailed elsewhere, both on this site and on the net, as well as in a plethora of books on
the history and philosophy of science.
Source: (Democritus on the
left) and (Leucippus on the right).
These two Ancient Greek Philosophers were the founders of European atomism, an idea which
in some ways was surprisingly close to modern atomic theory and chemistry.
Anyway, as we began to discover more and more about chemistry we eventually came
to understand that almost all substances on our planet were made up of a number of basic
building blocks called elements. The elements are basic in a chemical sense. The elements
cannot be further reduced to yet more basic components and still remain as normal matter.
What this means will become more apparent shortly.
As we began isolating more and more of these chemical elements and investigating
their properties, it gradually became clear to various scientists that there was a definite
pattern to the chemical and to some extent to the physical properties of these elements, and
that they could be organised into groups evincing similar behaviours: the inert gasses for
example are all gasses, and are all extremely non-reactive, whilst lithium, sodium,
potassium, etc, (which are amongst the group I elements) are all silvery to silvery-gold, soft
solids or – by the time one gets to caesium and francium – metallic liquids by around 30°C,
and are all hyper-reactive. Their densities also increases as one proceeds down the relevant
column of the periodic table. Similarly, fluorine, chlorine, bromine and iodine are all highly
reactive substances which progress respectively from being gasses to liquids to solids at
room temperature and range in colour from pale yellow, through greenish-yellow to brown
and finally dark bluish-black. Again, in a solid state, their densities increase as one proceeds
down the column of the periodic table.
Source: chlorine (left):
bromine (right):
As an aside, please note that in the early days of modern chemistry, it was far easier to
establish the atomic weight of an element than its atomic number. Thus, in 1865, John
Newlands came up with his ‘Law of octaves’ when he arranged the elements known at the
time in order of increasing atomic weight (which roughly is about twice the atomic number
– we will get into the meaning of both these concepts shortly). Essentially, he noted that the
eighth element (starting from any given one) shared many chemical and physical properties
similar to other ‘eighth’ elements.
Newlands law of eighths in tabular form. Please keep in mind that during his time, only a few
of all the elements were known to us, so this table is a thoroughly incomplete list of all the
elements now known to us. To me, this makes the achievements of these pioneering chemists
all the more remarkable, since they were able, with spectacularly sparse data, to adumberate
a profoundly important feature of the physical world around us.
With the advent of modern atomic chemistry, we discovered that all the elements
were made of atoms, and began isolating and studying these elements systematically, as
well as looking for gaps in the periodic occurrence of the elements. Indeed, in 1869, when
Dmitri Mendeleev first put his periodic table together, he and others noticed gaps in it, and
predicted the existence and physical properties of some of the then unknown elements.
Dmitri Mendeleev, the scientist who put all the varied observations of the chemical and
physical properties of the known elements at the time into some sort of order, and came up
with his periodic table. Shown below is a page of one of his early formulations of this table of
periodicity of the chemical elements.
Mendeleev’s first sketch of his periodic table, which looks nothing like the modern version of it,
but which truly was the start of great things.
It was not until we discovered the structure of the atom, and particularly the role
which electrons played in them however that a fully articulated explanation for all this
emerged. Whilst I am not going to go into the history of the periodic table here, I shall start
this post off by writing about the structure of the atom, as this is the key to the periodic
behaviour of the elements, before moving on to explaining how the periodic table was built
up. To make it easier for yourself, it would be good if you had access to a periodic table
whilst reading this post. If you follow this link, you’ll find a good interactive one:
The structure of the atom:
From a chemical point of view, all atoms, regardless of which element we are talking
about, are composed of electrons which orbit a core of neutrons and protons. Furthermore,
in any neutral atom, the number of protons in the nucleus exactly matches the number of
electrons in orbit around that nucleus. Once we get into the realm of chemical bonding
however, this is not always true, and atoms either share electrons in what is called covalent
or coordinate bonds, or some atoms make an outright donation of one or more of their
electrons to other atoms in what is called ionic bonding, but more on that later. For now,
let’s look at a few atoms to see what exactly this means.
The hydrogen atom, the simplest atom, and the most common one in the universe (though not
on the Earth), generally consists of one proton (which is positively charged) at the centre of
the atom (the nucleus), orbited by one electron (which is negatively charged) at some
minuscule distance from the proton. Please note though that compared to the dimensions of
the proton and even more so to those of the electron, any atom is of a vast size and in fact
mostly contains just empty space: the diameter of the hydrogen atom for example is about 5.3
x 10-11 metres, whilst the diameter of the proton is about 1.75 x 10-15 metres (which is a few
tens of thousands of times smaller than the hydrogen atom); in comparison however, the
diameter of the electron is about a thousand times smaller than that of the proton. This means
that both you, and the floor on which your feet are resting are overwhelmingly just empty
space, as the volume of the matter is about a million, millionth (ca. 10-12) of that of the atom!
Aside from electrons and protons, there is also another particle (which is neutrally charged)
called the neutron, which is about the same size and mass as the proton, and which, though
neutral, is very important in almost all nuclei (except that of hydrogen, which only has one
proton in it), as it helps to keep the positively charged protons confined in the small volume of
space which constitutes the nucleus. Remember that like charges repel each other, and even
though the charge on an individual proton is miniscule, the fact that they are all squashed
together in a micro-miniscule space (the nucleus) means that the repulsive forces between
them becomes quite significant. In fact, without neutrons, there would have been no elements
larger than hydrogen, as any larger nuclei would simply fly apart due to the mutually
repulsive positive charges of the protons in them. So again, without neutrons, there would
have been no one to write or to read this post, and no planet Earth on which to do it! There are
actually 3 types of hydrogen, called isotopes (Greek for ‘in the same place’), and they all
behave in the same way chemically, but have slightly different physical properties: one of them
(called deuterium) has one neutron in addition to one proton in its nucleus, whilst the other
(called tritium) has two neutrons and one proton in its nucleus. Since they both also only
contain one proton though, all hydrogen atoms only have one electron in their neutral state,
so chemically their behaviour is essentially indistinguishable. In nature, 99.985% of hydrogen
is the ‘normal’ form (sometimes called protium), having only one proton and no neutrons in its
nucleus. The fact that deuterium and tritium have more neutrons in their nuclei means that
these two isotopes are respectively about twice and thrice as heavy as protium. This is because
the proton and the neutron not only have almost the same size, but also have almost the same
mass. Again, the electron is the Lilliputian here, as it has a mass only about five tenthousandths of that of either of these two particles (5.49 x 10-4 of that of the proton, to be
precise). The fact that the different isotopes have different masses accounts for their different
physical behaviours. These physical differences though can also sometimes marginally affect
things like their chemical reaction rates, and is part of the reason that radio-carbon dating is
a useful tool.
As an aside, the neutron was first proposed by J. J. Thompson in 1920, and was finally
discovered by James Chadwick in 1932. Please note how recent these dates are.
Source: (left) and (right), the two Englishmen most responsible
for our discovery of the neutron (other key names in this discovery were: Viktor
Ambartsumian & Dmitri Ivanenko (Russian), Walther Bothe & Herbert Becker (German), and
Irène Joliot-Curie & Frédéric Joliot (French). The discovery of the neutron was, as you can see a
truly international enterprise, and depended on a phenomenal degree of international
The carbon atom:
The most commonly occurring carbon atom. Note that the most commonly occurring carbon
isotope has 6 protons, and 6 neutrons in its nucleus, with 6 electrons orbiting it (this
arrangement is true for the most common isotope of carbon, carbon 12. There are actually 2
other common isotopes of carbon, one with 7 neutrons (carbon 13), and the other with 8
neutrons (carbon 14) in their respective nuclei. Some isotopes are stable, but others, like
carbon 14, are radioactive, and will eventually decay into other elements, but that’s another
story, and one already partly covered elsewhere in this site. Note one other feature in this
atom: the electrons are not all arranged in the same orbital shell, but rather in two different
shells. This is very important both for chemistry and for the periodicity one observes in the
chemical behaviour of the elements. We shall shortly see how all of this works, but first let’s
look at a few other elements.
As an aside, carbon 14 is very useful to us for radio-carbon dating studies, and is one of the
ways in which we can establish the age of biogenic samples, a process which depends partly on
the fact that the chemical behaviours of isotopes is not totally identical, and that many lifeforms express a differential take-up rate for carbon 12 versus carbon 14.
The iron atom:
The most commonly occurring iron atom has 26 protons and 30 neutrons in its nucleus, with
26 electrons in orbit around this nucleus. There are a total of 7 relatively common, naturallyoccurring isotopes of iron, with anywhere from 28 neutrons to 34 neutrons in their respective
nuclei. To reiterate, in any neutral atom the number of protons always equals the number of
electrons, in ionised atoms though, this is not true and the number of electrons will either
exceed or fall short of the number of protons. The number of protons however cannot change,
or rather, if this changes (as it does in nuclear reactions), then one no longer has the same
element, since only the number of protons determines what the element is. A nuclear reaction
in other words transmutes the element; it’s the old alchemist’s dream actualised. As you can
see, again there are a number of shells into which the various electrons fit. Also, note that as
the number of protons in the nucleus goes up, the number of neutrons goes up faster (i. e. the
larger the atom, the disproportionately greater is the number of neutrons compared to the
number of protons, so iron has a few more neutrons than protons, whilst uranium (shown
below) with 92 protons and between 141 to 146 neutrons (depending on which of its various
isotopes we are talking about), has about 50% more neutrons than protons. This is significant.
Remember that neutrons act as a kind of glue to bind the protons together in a nucleus. As the
number of protons increases, the mutually repulsive electrical force of all these protons (which
are in very close proximity to each other) also increases dramatically, and so the amount of
‘glue’ needed also increases, hence the proportion of neutrons to protons also has to increase.
As an aside, note that from a radioactive perspective, the iron nucleus is one of the most stable
nuclei in the world. This has no bearing on its chemical behaviour, but is an interesting fact
none-the-less. Much of the iron found on our planet and elsewhere in the universe may be the
end-product of nuclear decays, although this is still a matter of speculation.
The uranium atom:
Uranium has 92 protons and between 141 and 146 neutrons in its nucleus. A neutral uranium
atom has 92 electrons orbiting its nucleus arranged into a variety of shells.
Hydrogen has an atomic number of 1 as it only has 1 proton in its nucleus (atomic
number = number of protons). If you look at a periodic table, you will see that the atomic
number of every element is always a whole number. An individual hydrogen atom may have
an atomic mass of 1, 2 or 3 depending on how many neutrons it has in its nucleus, thus if
you look at a periodic table, you will see that the atomic mass is never a whole number. This
is partly because the atomic mass is a proportionate average of all the isotopes of that
particular element, so the atomic mass of hydrogen is 1.00794. (atomic mass = proportional
average of number of protons + number of neutrons for an element. So, since almost
99.99% of all hydrogen atoms have an atomic mass of 1 (1 proton, and no neutrons), one
would expect its atomic mass to be very close to one, however the few hydrogen atoms
which have either one or two neutrons in their nuclei increase the average atomic mass to
slightly over one. There are two other factors which influence the final atomic mass
number: the binding energy, and the standardisation issue, but I shall not get into that in
detail here Briefly however, in any nucleus aside from protium, some energy is required to
bind the protons together, and this energy makes some contribution to the mass of the
nucleus (remember that E = mc2). Also, because of this slight variance, and since atomic
masses are based on the mass of nuclei, it depends very much what particle one takes as
one’s starting point, as having an atomic mass of 1. A proton is ever so slightly lighter than a
neutron, so if one were to assign to either of these sub-atomic particles an atomic mass of 1,
then the atomic masses calculated for all the elements would end up being non-integers.
Similarly, if we were to assign an atomic mass of one to the protium atom, then we would
have to include the mass of the electron in that unit, so we would again end up with a
decimal point in the atomic masses of the elements. We could also take some other element,
say carbon 12, and simply divide its mass by 12, but the atomic mass would always have a
decimal point in it.
Recall that I wrote earlier that all chemical activity is based upon an exchange of
electrons between atoms, and that in certain types of bonds, electrons are shared between
two or more atoms (covalent or coordinate bonds), whilst in others, some atoms outright
donate or accept electrons from some other atoms (ionic bonds). With these caveats in
mind, let’s now start ‘building’ a few atoms, before seeing how they would start bonding
with each other.
As you already know, the most basic atom is hydrogen with one proton and one
electron (with possibly one or two neutrons as well, but this is irrelevant to us). If we add
one proton (and a couple of neutrons) and one electron to hydrogen, we get helium, add
one more again (and some neutrons of course) and another electron, and we get lithium,
and so on along the periodic table.
Now comes the next set of rules: 1. the first shell, called the ‘K’ shell can contain a
maximum of 2 electrons; 2 atoms like their outermost shells to either be empty or full. Thus,
hydrogen, chemically speaking, is ‘happiest’ either when it’s lost its single electron, or when
it’s gained one, thus either emptying, or completely filling its outermost (and in this case,
only shell). Helium on the other hand (the next element along), already has a full ‘K’ shell,
thus it’s very stable, so hydrogen is quite reactive (as it only has to either add one electron
or get rid of one electron in order to achieve the stable configuration of a completely full
outermost shell), but helium is very reluctant to react because its outermost shell is already
Lithium burning in air. The metallic elements all burn with characteristic colours (a topic for a
future post), and is one of the ways in which one can analyse what elements might be
contained in an unknown sample of certain materials.
When we move on to lithium, by adding another proton (some neutrons), and another
electron, the additional electron takes up residence in the next shell, the ‘L’ shell. This ‘L’
shell can contain a maximum of 8 electrons, so in order to be full, it would have to take on
another 7 electrons, which is just too many electrons (remember, lithium only has 3
protons in its nucleus, and these 3 protons couldn’t possibly hold on to 7 additional
electrons). As a result, lithium whilst it’s quite ‘happy’ to get rid of one electron, cannot
possibly hope to complete its ‘L’ shell by taking on an additional seven electrons.
There’s another rule of sorts here: In filling or emptying outermost shells, a balance
needs to be established between the quantum mechanical needs of having a full outermost
shell, and the electrostatic needs of not having too much of a charge imbalance in a
chemically bound atom. Thus, whilst it is relatively easy to lose or gain one additional
electron to complete your outermost shell, it is somewhat more difficult to lose or gain two
electrons in order to achieve the completed outermost shell profile. Thus, when we move to
the next element in the periodic table, beryllium (atomic number 4), we find that it has 2
electrons in its ‘K’ shell, and 2 in its ‘L’ shell, so it needs to lose 2 electrons in a chemical
bond, which is more difficult than simply losing one, so beryllium is less reactive than
A piece of beryllium. It is nowhere near as reactive as lithium, and is very light (as one would
expect since the atom contains only 4 protons and 5 neutrons). It is used in certain specialised
engineering applications (mirror coatings, injector nozzles for space-craft etc).
The next element is boron (atomic number 5, structure 2 ‘K’ electrons, 3 ‘L’ electrons),
which because it would have to lose 3 electrons, is even more un-reactive than beryllium.
When we get to carbon, things could go either way. It could either lose 4 electrons, or it
could gain 4 of them to achieve the stable configuration. This is a lot of electrons to shed or
to latch on to, and so carbon is again not all that reactive. Moving on to atomic number 7
(nitrogen), it’s the same story as boron, except that nitrogen is not able to lose 5 electrons,
and so needs to add another 3 to fill its outermost ‘L’ shell. This again is difficult, so nitrogen
is not all that reactive. We see plenty of evidence of this. Food is sometimes nitrogenpacked, as this element does not support life, and does not react readily. Also, most
substances will not burn in pure nitrogen.
Oxygen is the next element, and it is considerably more reactive than nitrogen,
because it only needs to add two electrons to complete its ‘L’ shell. When we get to fluorine
(atomic number 9) however, we find an element which is one of the most reactive in the
periodic table. Ending this row is neon with its complete ‘L’ shell, the second of the almost
totally un-reactive, noble / inert elements.
Before moving on to the next row, I would now like to explain something in slightly
more detail. As you know, there are 2 general classes of chemical bonds: ionic bonds
wherein certain atoms donate electrons outright to other atoms which accept them
outright; and covalent and coordinate bonds, wherein atoms share electrons. The elements
which are highly reactive (like lithium or fluorine) readily form ionic bonds, whereas
elements which are not so reactive (like nitrogen and carbon) are more likely to form
covalent or coordinate bonds. This is because whilst electrostatically speaking, it is
relatively easy to latch onto or get rid of one or two electrons, it is not so easy to so do when
three or four electrons are involved, thus since covalent bonds involve sharing rather than
an outright donation, this is the way these latter elements usually form their chemical
Sodium burning in air with its characteristic yellow flame.
Having finished with neon (atomic number 10), let’s move to the next element, sodium
(atomic number 11). Now, sodium already has full ‘K’ and ‘L’ shells (with 2 and 8 electrons
respectively), so the additional electron goes into the next shell, the ‘M’ shell. The ‘M’ shell
can contain a maximum of 18 electrons, but unfortunately things are no longer as simple as
they were. You see, there is another rule. The outermost shell must always contain a
maximum of 8 electrons only. Hence, the filling of this row of the periodic table (the 3rd
row) is identical to the exercise we went through just now, and the rules are also identical:
sodium is far more reactive than magnesium, which is considerably more reactive than
aluminium. Silicon behaves much like carbon, and so on until we come to argon, the next
un-reactive element.
There is another point to be aware of now. Sodium is more reactive than lithium, and
bromine is less reactive than chlorine. This is because of something called electronic
shielding. All the electrons whizzing around a given nucleus tend to shield the outermost
shells form the electrostatic influences of that nucleus. Hence, because the sodium atom has
more shells than the lithium atom, the outermost electron of sodium is better shielded than
that of lithium. This means that it is easier for sodium to get rid of its outermost electron
than it is for sodium to so do. Conversely, because of this shielding, it is easier for say
chlorine or oxygen to grab on to an additional electron or two than for bromine or sulphur
to so do.
This shielding also explains why if one looks at a modern periodic table, one notes a
diagonal pattern which runs from boron / aluminium down to astatine, and which divides
the elements into those with a metallic characteristic, and those which are essentially nonmetallic. The elements which lie along this diagonal are often called metalloids, because
they are semi-metallic in their chemical and physical behaviour. This is also why silicon and
germanium are both good semiconductors, whilst carbon is not, and why tin, lead and
bismuth are all metals even though they appear to the far right of the table, where all the
non-metals are located.
When we get to the next element, potassium, it is now the ‘N’ shell which is going to be
filled, but this happens in a two-step process. This is because the shells I have been writing
about so far actually have some finer structure to them; they have what are called subshells. The ‘K’ shell has only one sub-shell, and it is called the ‘1s’ sub-shell. The ‘L’ shell has
2 sub shells in it, the ‘2s’, and the ‘2p’. The ‘M’ shell has 3 sub-shells in it, the ‘3s’, the ‘3p’
and the ‘3d’. The ‘N’ shell has 4 sub-shells in it, the ‘4s’, the ‘4p’, the ‘4d’ and the ‘4f’. These
sub-shells have different maximum capacities. Any ‘s’ sub-shell can contain a maximum of 2
electrons, any ‘p’ sub-shell can contain a maximum of 6 electrons, any ‘d’ sub-shell can
contain a maximum of 10 electrons, any ‘g’ sub-shell can contain a maximum of 14 electrons
and so on.
So, getting back to potassium (atomic number 19), its electronic structure is 1s2 (i. e. 2
electrons in the ‘1s’ sub-shell), 2s2, 2p6 (2 electrons in the ‘2s’ sub-shell & 6 in the ‘2p’ subshell), 3s2, 3p6, 4s1 (2 electrons in the ‘3s’ sub-shell, 6 in the ‘3p’ sub-shell, and 1 in the ‘4s’
sub-shell). Note that the ‘3d’ shell has nothing in it yet.
Next comes calcium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2 (i. e. the same as potassium except
that calcium has 2 electrons in its ‘4s’ sub-shell). After calcium, the ‘3d’ sub-shell starts
filling, and we have the elements from scandium to zinc inclusive (the first row of transition
elements). Scandium’s electronic structure is 1s2, 2s2, 2p6, 3s2, 3p6, 3d1, 4s2, whilst that of
zinc is 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2. Gallium’s electronic structure is 1s2, 2s2, 2p6, 3s2,
3p6, 3d10, 4s2, 4p1. As one continues along this row (the 4th one in the periodic table), the
‘4p’ sub-shell continues to fill until by the time one gets to krypton, all available shells are
full, and we have another un-reactive element.
When we get to row 5, the story is the same as that of row 4, except that in this case
it’s the ‘4d’ sub-shell which gets filled after the ‘5s’ sub-shell is taken care of. Here again,
there is a skip, as the ‘4f’ sub-shell does not start getting filled until one gets to the 6th row.
This then, is why the modern periodic table looks the way it does, and why this gaptoothed object is really a thing of profound beauty. As C. P. Snow upon learning about the
periodic table put it: “For the first time I saw a medley of haphazard facts fall into line and
order. All the jumbles and recipes and hotchpotch of the inorganic chemistry of my
boyhood seemed to fit themselves into the scheme before my eyes — as though one were
standing beside a jungle and it suddenly transformed itself into a Dutch garden.”
Spectral emissions from different chemicals burning: (left to right) Methane, Calcium Sulfate,
Calcium phosphate, Sodium chloride, Potassiam phosphate, Sodium borate. One of the factors
contributing to these different characteristic colours is the electrons orbiting the various
nuclei involved jumping from sub-shell to sub-shell, driven by the energy they have derived
from the flame.