Why don’t things burn under water?
Friends, when they see me, often remark that many of my posts start
off with me writing ‘just the other day, a friend asked me.…’ Well, to break from that
pattern, just the other day it was in fact an acquaintance who asked me, over a rather
good glass of Belgian beer I might add, why things do not burn under water. It was
actually a fascinating question, and I had much fun answering it, so I’ve decided to
elaborate on my recent back-of-the-beer-mat drawings & hand-waving arguments.
As it turns out, the customary claim is not entirely true; it rather depends what
substance one is trying to burn, as well as how one actually defines burning. Before
getting to that however, we must first make some distinctions: there are two levels at
which water can be said to contain oxygen, which is one of the various substances in
which things can oxidise (i.e. burn): first there is the oxygen which constitutes the water
molecule itself, the ‘O’ part of H2O, and second there is the oxygen which (along with
other gasses like nitrogen and carbon dioxide), is dissolved in water – there are around
40~50 milligrams of oxygen dissolved in one litre of water, and this is the stuff which
fish, etc. breathe. Incidentally, it is the lack of dissolved gasses in boiled water which
gives it its peculiar taste.
Source: http://sharon-taxonomy2009-p2.wikispaces.com/Chordata-Fishes
Lets look at the second type of oxygen (the oxygen which is dissolved in the
water) first. This oxygen, just like the oxygen which is present in the air we breathe, is
free in the sense that it is not chemically bound-up with any other element and is thus
readily available for use in chemical reactions (breathing provides organisms with the
oxygen necessary for the chemical reactions which are an essential part of the lifeprocesses in all animals). This is why fish can breathe underwater.
The problem with using this dissolved oxygen to actually burn things though (as
opposed to using it for breathing) is that water has a very high specific heat capacity; in
other words, it can absorb a lot of heat before warming up. Now since burning requires a
reasonably elevated temperature before an object can catch fire, a substance which can
burn in the air will generally go out once it is plunged into water because the
temperature of the burning substance rapidly drops to well below its ignition point once
it comes into direct contact with water; it is not entirely because water cuts off the
oxygen supply (as people are sometimes told), but rather in large part because it cools
them down too much that water is such a good fire retardant.
Aside from the oxygen it contains in solution though, water also has a lot of
oxygen in its chemical make up (the ‘O’ part of H2O). However this oxygen is not free and
is not generally available to take part in chemical reactions, and this is why this oxygen
component is not usually available for combustion. In brief then, this is why things do
not generally burn underwater.
So much for the simple explanation. There are actually a few exceptions to what I
just wrote, but to understand how it all works, we have to make a little foray into
chemistry, and specifically into the world of chemical bonding, as well as talk about
burning, and what this process really involves.
All burning involves a chemical reaction in which one substance gets oxidised
(the thing being burnt), and another gets reduced (the thing doing the burning): all
common forms of combustion involve what is called a reduction-oxidation (red-ox)
reaction, but not all red-ox reactions actually involve combustion rusting being the
commonest example of a flameless red-ox reaction – when blood gets oxygenated in the
lungs or gills, in a sense it in fact rusts.
Source: http://zendarie.com/2011/oxidation-reduction-reactions-redox/
The terminology is a bit tricky here. The oxidising agent gets reduced, and the reducing
agent gets oxidised. Furthermore, the stuff which gets reduced actually gains electrons,
rather than losing them!
Another point to keep in mind is that oxidation does not necessarily have to
involve oxygen, but since it is one of the most readily available oxidising agents on the
surface of our planet, the bulk of oxidising reactions which most people are familiar with
involve oxygen as the oxidising agent. Thus, because prior to the emergence of modern
chemistry hardly any other oxidising agents were known, and because the intimate
details of the process was not scientifically understood yet, the reaction itself took its
name from this element. Many, many things can also oxidise in fluorine or chlorine gas
for example, and certain very reactive substances can even be oxidised by nitrogen gas
(which is normally quite an un-reactive element).
To reiterate what I wrote under the above diagram, in a red-ox reaction, the
substance being oxidised surrenders some electrons to the oxidising agent, as you can
see in the above diagram, and as will shortly become more clear. So, looking from the
reducing agent’s point of view, when iron rusts, the iron atom gives up some electrons to
the oxygen atom, whilst looking at it from the oxidising agents point of view, when
chlorine and sodium are combined, the chlorine atom captures an electron from the
sodium: remember, the oxidising agent (the stuff which gets reduced) gains electrons,
whilst the reducing agent (the stuff which gets oxidised) loses them. Since fluorine,
chlorine, bromine, oxygen, nitrogen, etc would all ‘like’ to have a few more electrons,
they can oxidise things. However, as I noted earlier, not all of these substances are highly
reactive. Nitrogen is a chemically ‘reluctant’ element, whilst fluorine is a chemically
‘enthusiastic’ element.
The reason that this is so is because of the way electrons are arranged around an
atomic nucleus.
Source: http://sciencespot.net/Pages/kdzchem.html
Although the atom is typically drawn like this, in fact the electrons are not all equidistant
from the nucleus. There is a definite though rather complex but orderly arrangement to the
way electrons are distributed around a nucleus. As a first approximation, the electrons may
be said to be arranged in ‘shells’, rather like the layers of an onion (in fact the situation is
more complex than this because the shells themselves may actually be divided into subshells of differing shapes, but for the moment this is irrelevant to us).
The first shell can contain a maximum of two electrons, and the second shell can
contain a maximum of eight electrons. The third shell can contain a maximum of
eighteen electrons, but two and even more so eight are ‘magic’ numbers in the world of
atomic chemistry, and so something slightly odd happens when the third shell gets to
eight: the forth shell (which can contain a maximum of 32 electrons) takes on its first
two electrons before the third shell (which at this point already has eight electrons)
starts to take on its final load of ten – remember, the third shell can contain a maximum
of eighteen electrons. Once the third shell has taken on all of its final ten electrons, the
forth shell finishes off taking on its intermediate load of eight electrons. After this, the
fifth shell (mirroring the forth shell’s behaviour) takes on its two initial electrons, before
the forth shell starts filling up). After that, things get even more complicated, and I’ll not
be looking at this process in this post, but shall save that for a later date. The filling order
of the shells and sub-shells is reflected in the form of the periodic table.
Source: http://www.homework-help-secrets.com/periodic-table.html
In looking at a periodic table, please note that the numbers at the top of each cell
correspond to the number of protons in the nucleus (and hence to the number of electrons
orbiting that nucleus. In a neutral atom, the number of protons equals exactly the number
of electrons, so Francium (Fr) for example has 87 protons and 87 electrons, whilst chlorine
(Cl) has 17 of each).
The filling up of the first shell, which can contain a maximum of 2 electrons, is reflected in
the top row of the periodic table which contains only hydrogen (H – which only has 1
electron), and helium (He – which has 2 electrons).
The filling up of the second shell, which can contain a maximum of 8 electrons, is reflected
in the second row, starting with lithium (Li – with 3 electrons), and ending with neon (Ne –
with 10 electrons).
The filling up of the third shell, which can contain a maximum of 18 electrons, is partially
reflected in the third row, which starts with sodium (Na – with 11 electrons), and ends with
argon (Ar – with 18 electrons). The completion of the filling up process of the third shell
however does not take place until the forth shell has taken on 2 electrons first.
The filling up of the forth shell, which can contain a maximum of 32 electrons, is partially
reflected in the forth row starting with potassium (K – with 19 electrons) and ending with
krypton (Kr with 36 electrons). As just noted however, the forth row also reflects the filling
up of the remainder of the third shell (the final 10 electrons which can occupy this shell),
starting with scandium (Sc) and ending with zinc (Zn). Only after the remainder of the
third shell is filled does the forth shell take on its intermediate load of 6 more electrons (to
make the total of the ‘magic’ if only intermediate 8). After this, the filling process gets even
more complicated. In a sense, the filling order may be seen as being done in batches, 2 then
8, then 8 again, then 2 more, then 6 more, then 2 then 10, then 6 more, then 2, then 10 then
6, then 2, then 14, then 10, then 6, and so on. I shall write more about this in a subsequent
post.
With the exception of hydrogen and helium, all atoms would either ‘like’ to add
enough electrons so as to have 8 of them in their outermost shell (called the valence
shell), or get rid of a few electrons so as to expose the underlying complete shell. The
elements which would like to shed electrons are generally metals, whereas the elements
which would like to gain electrons are generally non-metals. If an atom only needs to
add or get rid of one or two electrons, then it tends to be chemically ‘enthusiastic’ (i.e.
highly reactive).
Furthermore, an element which is a metal is more reactive the larger the atom is,
whereas an element which is a non-metal is more reactive the smaller the atom is. Thus,
if you look at the first column (headed by hydrogen (H)), the elements get more reactive
as you move down this column. This is why lithium (Li), sodium (Na), potassium (K) etc
can all strip the oxygen part of water away from the hydrogen part of it very
aggressively.
If you look at the second column though, the elements represented here are less
‘aggressive’. So, even though magnesium (Mg), calcium (Ca), etc can do this too, their
reactions with water are much milder than those elements found in the first column.
This is because the elements in the second column all have to get rid of 2 electrons
rather than just 1.
Conversely, if you look at the seventh column, headed by fluorine (F or Fl), as you
work your way down this column, the elements become progressively less reactive. If
you move across to the sixth column though, you will find that oxygen (O) is less reactive
than fluorine because oxygen needs to add 2 electrons to attain the ‘magic’ 8, whereas
fluorine, chlorine, etc only need to add one electron to get to the ‘magic’ 8.
So, highly reactive metals can actually burn underwater (though to be more
accurate, they actually ‘burn’ using the water itself). If you take a look at the images
below you will clearly see the difference in reactivity between a group one element (like
sodium) and a group two element (like calcium). The calcium is reacting, but not
violently. It is only bubbling gently; the sodium on the other hand is reacting so violently
that the heat of the reaction is sufficient to ignite the hydrogen gas which is released
when the water molecule is broken up by the sodium (water is composed of 1 oxygen
atom, an 2 hydrogen atoms, however the sodium is so ‘desperate’ to grab hold of this
oxygen that it forcefully ‘kicks’ the hydrogen out).
Source: http://www.uncp.edu/home/mcclurem/ptable/hydrogen/h.htm
Calcium reacting with water. The bubbles are hydrogen gas, liberated in the reaction as
the oxygen is stripped away from the hydrogen.
Source: http://www2.uni-siegen.de/~pci/versuche/english/v44-1-1.html
Sodium reacting with water. Notice how much more violent this reaction is than the one
involving magnesium (shown above).
So, can things burn underwater, the answer is as I noted at the beginning a
qualified yes. It depends both on what substance is being burnt, and how one defines
burning.
As a little bonus, here is a photograph of a liquid breathing experiment with
mammals. So although you cannot breathe under water, you could certainly breathe
under some oxygenated liquids!
Source: http://marianuniversityscienceblog.wordpress.com/2010/04/07/oxygenatedwater/
A mouse breathing oxygenated perfluorocarbon, which can contain 25 percent more
oxygen than air and can transfer oxygen to the lungs three times more effectively than air.
When mice are submerged in oxygenated perfluorocarbon, they can survive for several
weeks and will make a complete recovery when the perfluorocarbon is drained from their
lungs. If a deep-sea diver were to take a single breath of oxygenated liquid
perfluorocarbon, s/he could remain submerged for several minutes without having to take
another breath.