The pH scale is something we’re all familiar with; most people will remember it from school chemistry lessons. It’s the scale used to rank how strong an acid (or alkali) a solution is. The colours associated with each number correspond to the colour that universal indicator turns in solutions of that particular pH. A fair proportion of people probably don’t know the chemistry behind the pH scale, though – where exactly do these numbers come from?
The clues are actually partly hidden in the scale’s name. The ‘H’ in pH stands for the element, hydrogen. On a simple level, the pH scale can be thought of as a ranking of the amount of hydrogen ions in a solution: the more hydrogen ions, the lower the pH number. The ‘p’ in pH, to chemists at least, stands for the mathematical operation ‘-log10‘. pH, then, is simply equal to -log10[H+], where [H+] is the hydrogen ion concentration in a particular solution. (Note that, strictly speaking, we’d usually use H3O+ to represent hydrogen ions, as this is the form they take in solutions. However, to keep things simple, we’ll continue to use H+ here).
Looking at the graphic above, you can see that an increase in pH of one point actually involves a tenfold decrease in the concentration of hydrogen ions in a solution. The scale is what’s known as a logarithmic scale. Why do we bother with this mathematical manipulation in the first place? Well, it’s clearly a lot easier to state a single number when referring to the acidity (or alkalinity) of a solution than it is to quote the many-numbered hydrogen ion concentration. A pH of spot on 7 denotes a neutral solution (neither acidic or alkaline). Any pH below 7 is acidic, whilst any pH above 7 is termed alkaline.
Water molecules have the chemical formula H2O. However, these molecules are capable of splitting up slightly in solution, in H+ and OH– (hydroxide) ions. In a neutral solution, the concentrations of these two ions are equal. However, the addition of an acid or alkali can cause them to vary. Acids are a source of hydrogen ions, and adding them to water increases the concentration of hydrogen ions in solution, lowering the concentration of hydroxide ions. For alkalis, the opposite is true: they decrease the concentration of hydrogen ions, whilst increasing the concentration of hydroxide ions.
Something a lot of people don’t realise is that pH is temperature dependent. Strictly speaking, pure water only has a pH of 7 at ‘room temperature’ (25˚C). Above and below this temperature, it can vary: for example, at 100˚C, the pH of pure water is 6.14, whilst at 0˚C, it’s 7.47. This doesn’t mean that the pure water is becoming acidic or alkaline, but that, at these temperatures, those particular pH numbers represent the neutral point. As a side note, it’s also worth pointing out that the pH scale isn’t limited to the usual 0-14 range shown here – some strong acids and alkalis can fall outside this range, into negative pH values or values higher than 14.
Another common misconception about pH concerns the human body. Diets such as the alkaline diet claim that it is possible to affect the pH value of your body by altering your diet to include ‘alkalising’ foodstuffs that make the pH of your body more alkaline. Whilst what’s actually suggested, a diet heavy in fruits and vegetables, is certainly by no means unhealthy, it’s pretty much impossible for what you eat to change the pH of your body. I’m not going to go into too much detail, as Kat from The Chronicle Flask blog has already done an excellent job of debunking alkaline diets, but it’s worth reiterating a few key points.
The pH of the stomach can vary, between 1.5 and 3.5 on the pH scale. However, this has no effect on the pH of our body, or, more specifically, our blood. Human blood has a pH value that’s always slightly alkaline, between 7.35-7.45. If we were able to purposefully change the pH of blood outside of this small range, we could actually cause ourselves a good deal of harm; even a pH change of 0.5 either way could result in irreversible cell damage. Luckily, the food we eat has next to no effect on blood pH.
What we eat can affect the pH of our urine. The pH of urine has an average value of around 6, but can range between 4.5 and 8. However, whilst you might be able to have the slightest effect on your urinary pH with your diet, it’s unrelated to the pH of your blood; this stays in the range previously stated, regardless of any change in urinary pH.
Moving on, and whilst we’re talking about acids, it’s worth talking about the strength of acids. Strictly speaking, it’s hard to place particular chemical compounds at specific points on the pH scale, as their positions vary depending on concentration. Concentration is a measure of how much of a substance is dissolved in a solution. If we have a lot of an acid dissolved in a relatively small amount of water, we’d say we have a solution with a high concentration. Similarly, if we had very little acid dissolved in lot of water, we’d have a solution with a low concentration.
Since concentration of solutions can easily be varied, solutions of varied concentration of the same acid can have different pH values. Some acids, however, are stronger than others. Hydrochloric acid, the same acid found in stomach acid, is a strong acid as it can easily split up into its component ions. On the other hand, acetic acid, the acid found in vinegar, is a comparatively weak acid – it doesn’t split up into its component ions easily. Another example of a weak acid is hydrofluoric acid; contrary to what a certain popular chemistry-themed TV show would have you believe, it’s actually a fairly weak acid, which certainly couldn’t dissolve a body. It’s an unpleasant compound for different reasons though. We can rank acids in terms of their strength (or how easily they split up into their component ions), but that’s a topic for another post!
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References & Further Reading
- Amazing alkaline lemons – K Day, The Chronicle Flask
- A few points on pH – M Leger, Atoms & Numbers
- This post is pH-balanced for all – M Leger, Atoms & Numbers
- pH, pOH, and Ka – Bodner Research Web
14 replies on “Acids, Alkalis, and the pH Scale”
Always nice and instructive graphics as usual !
Just one little point : negative values or values above 14 should be mentioned or at least one should give the impression that pH scale does not end at zero or 14.
You’re right – I actually intended to mention it in the post, then forgot! I’ve added mention of it in now. I might in the future develop a further graphic from this, showing acids/alkalis with pH values outside the ‘standard’ range.
Thank you. This is an excellent resource for high school chemistry and the introduction of acids and bases. I particularly like the columns showing the concentrations of the two ions. You don’t see that much….
Thanks for the great chart. The term Alkaline isn’t used as often in the US (favoring Basic). Are the two terms synonymous or is there some distinction?
In the UK, we use both – a base is used to describe a solid that reacts with an acid, such as a metal oxide. An alkali is a base that is soluble in water. So all alkalis are bases, but not all bases are alkalis!
Acid and alkaline graphics are a staple in many science labs and I predict that this graphic will be useful to many. Another great post.
I believe there is a typo in the 9th paragraph from the top – “Movign on, and whilst we’re talking about acids”
There is, thanks for the spot! I’ll fix that one right away.
Gorgeous and informative as always. One small suggestion: indicate explicitly on the graphic itself that the colour scale is based on the “universal” indicator mixture — it might currently give the impression that “all acids are red” or something. Or perhaps include multiple colour gradients, as per the pH indicators infographic (http://www.compoundchem.com/2014/04/04/the-colours-chemistry-of-ph-indicators/)?
Yeah, you’re right, I’m intending to add in a mention that it’s the colour with universal indicator. I’d add other indicators too, but I don’t think there’d be enough space!
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