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Basics of pH - Lena Rockstein
The pH is a logarithmic scale used to specify the acidity (or basicity) of aqueous solutions. It was first introduced by the Danish chemist Søren Sørensen in 1909 as the “power of hydrogen” (= pH). You might have come across a pH before – many hand cremes or shower gels claim to be “pH skin neutral” or foods are described as being “basic” and thus claim to be healthier than “acidic” options
The exact (and rather abstract!) definition of pH is “approximately the negative of the base 10 logarithm of the molar concentration of hydrogen ions”. Before we dive into this definition further, it might be useful to catch up on the term “molar concentration” before. Concentrations describe how many particles of one species we can find in a defined volume. In science, we use the so called “molar concentration” to talk about defined portions of particles – one mole is set as 6 ∙ (10)^23 particles. If we look at a 2M NaCl solution for instance, we learn that the molar concentration is 2 moles per litre (= M) and hence 2 ∙ 6 ∙ (10)^23 particles sodium chloride has been dissolved per litre water. The pH is connected to the molar concentration of hydrogen ions, but why do we have to take the negative of the base 10 logarithm? The concentration of hydrogen ions is normally rather small, for example approximately (10)^(-7)M for pure water
To make the value of a pH more relatable for everyone, we convert this abstract value into pH = 7 for water, using the logarithm.
The pH scale reaches from 0 – 14, even though we can still find solutions with a pH outside of that span. The so-called neutral point is pH=7, pH<7 is described as acidic and pH>7 as basic. There are various definitions of what exactly an acid or a base are, but the most famous one is the definition of Brønsted, describing acids by their capability of donating (splitting off) a hydrogen ion and bases by their capability of accepting (binding) a hydrogen ion. .
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