It took me a while to knock up a bit of a not too mathematical explanation
of pH during the thread on the acids wiki.
In that thread it has aroused no comment so I post it here in case anyone is
still interested, and had given up following that thread.

S


"The Medway Handyman" wrote in message
news:eOe9o.12094$Nn4.5764@hurricane...
Huge wrote:
On 2010-08-13, The Medway Handyman
wrote:
Tabby wrote:
Feedback welcome...

NT


Would a simple explanation of the ph scale help?

The negative of log10 molar hydrogen ion concentration? (*) Or perhaps
you meant simpler than that.

Much simpler than that :-)

Most people have no idea of the difference between acid & alkali. Nor
that the scale is logrithmic IIUIC.

When I sold pressure washers & demonstrated with detergents of ph 13
punters would often ask if it was "a acid that burned the dirt off".


--
Dave - The Medway Handyman
www.medwayhandyman.co.uk


Spamlet says:


Some of the confusion comes from the way we tend to label acids as
'corrosive: causes burns', and alkalis as 'caustic: causes burns'...

The negative log gets you a nice round figure for one particular ion instead
of having to quote for H+ *&* OH-, but it is by no means an easy concept to
explain.

Let's have a go...

You could just say that: "Extremes of acidity or alkalinity are highly
reactive and can cause serious burns": but how to indicate just how reactive
they are? A useful way to make a scale of reactivity is to imagine a series
of concentrations where each step is ten times stronger than the last. Such
is the pH scale.

[You can think, similarly when you are heating solutions: a 10C rise in
temperature approximately doubles the rate of reaction.]

The 'H' in pH stands for hydrogen ion (H+). Positive ions like H+, are
reactive atoms from which an electron has been removed: this is what *makes*
them
reactive in solution: they 'want' to get their electrons back, from other
substances, to
become neutral again. The more concentrated these ions are, the more
strongly
they will be able to attract electrons back and the more vigorous the
reaction
will be. The pH scale *only* deals with the concentration of the H+ ion:
the concentration of the chemicals that will produce a solution of a
particular pH varies with the chemical, so speaking of Strong Acid, is not
the same as speaking of concentrated acid.

When the chemist says Strong acid, he means that it is a substance that
forms these H+ ions easily *in water*. Water is what is termed a 'polar'
solvent, because its molecules are arranged such that the O atom has a
bigger attraction for the molecule's electrons than the two H atoms. This
makes each molecule like a little magnet with slightly positive poles at the
H end, and slightly negative poles at the O end. This is what makes water
such a good solvent: its slightly charged molecules can surround ions like
H+ and keep them free in solution to make them available for chemical
reactions.

Each step in the pH scale from 7 down to 1 indicates a tenfold increase in
strength in the *acid* direction, with pH1 being ten million times as H+ ion
concentrated as pH7. Each step in the scale from pH7 up to pH14 indicates a
tenfold increase in the *alkaline* direction, with pH14 being ten million
times stronger than pH7. [This may seem an awkward way of looking at
things, but it saves having to have separate scales for H+ and its
counterpart, OH-. For the DIYer it might have been more intuitive to call
it a pOH scale, so then we would have a scale where the numbers went up with
the concentration instead of down, but we are stuck with it I'm afraid!]

The concentration of the H+ ion is not directly related to the concentration
of the acid that produces it. Weak acids with a capital W do not so readily
form H+ ions in water, and so can still have a moderate pH even when quite
concentrated. Solubility is also important: HCl is a gas that won't
dissolve in water to give highly concentrated solutions, whereas H2SO4 will
concentrate almost to treacle like consistency, and is so attracted to water
that the pure acid is only found in space. In concentrated form it is
extremely dangerous, but, because it is so lacking in the water the H+ ion
needs to become mobilised, the term pH is no longer much use to indicate
just how dangerous it is. Remember, the pH only refers to the concentration
of H+ : not of the acid itself.


We normally make solutions of chemicals with water as the solvent. Everyone
knows that the 'formula' of water is H2O, but if you want to imagine a bit
more about pH, it is better to think of it as 'HOH'. This is a neutral
structure in which the electrons are shared between the three atoms, though,
as above, the O attracts them more strongly than the H, so it could be said
to behave like a little magnet. Despite its general stability, in a volume
of water containing billions of these molecules, a small percentage break up
into ions: the H+ that we have already seen, has its counterpart in the
negative OH- ion. This ion still has the electron attached which has been
'lost' by H to make the H+ ion. So, in the case of the OH- ion, it wants to
'give away' or share its extra electron to become neutral again, and, as in
the case of H+, the more concentrated OH- becomes, the more strongly will be
reactions, as it seeks to share its extra electron to become neutral again.
The fact that water is thus slightly ionised itself allows us to quote a pH
for it even when nothing is dissolved in it.

As, in pure water, there are always the same number of H+ ions as OH- ions,
it is neutral over all, and should have a pH of 7 but will often be a little
lower in practice. Once acids or alkalis are dispersed in the water, the pH
changes dramatically as extra H+ or OH- ions are mobilised. Acids are, thus
chemicals which break up in water to greatly increase the proportion of H+
ions relative to OH- ions. Alkalis have the opposite effect: increasing the
concentration of OH- relative to H+.



Hope this gives a reasonable insight without having to go into the maths.
Wikipedia proper will give the full picture, for the mathematically
inclined:
http://en.wikipedia.org/wiki/PH

S