pH is a measure of the acidity or basicity (alkalinity) of a solution. It is defined as the negative logarithm (base 10) of the concentration of hydrogen ions (H+) in the solution.
pH values range from 0 to 14, with a pH of 7 being neutral, meaning the concentration of H+ ions is equal to the concentration of hydroxide (OH-) ions. A pH below 7 is acidic, indicating a higher concentration of H+ ions, while a pH above 7 is basic or alkaline, indicating a higher concentration of OH- ions.
pH is an important parameter in many chemical and biological processes, including agriculture, environmental monitoring, and medicine. It’s definitely possible to calculate a negative pH value. But on the other hand, whether or not an acid actually has a negative pH value isn’t something you can verify very well in the lab.
“In practice, any acid that yields a concentration of hydrogen ions with a molarity greater than 1 will be calculated to have a negative pH.” That statement is not entirely accurate. A solution with a hydrogen ion concentration of 1 M (molarity) would have a pH of 0, not a negative pH. In fact, pH values below 0 are not physically possible, as a pH of 0 corresponds to a hydrogen ion concentration of 1 M.
It is true that acidic solutions have a higher concentration of hydrogen ions than basic solutions, and therefore have a lower pH value. However, it is important to note that pH is a logarithmic scale, meaning that each pH unit represents a tenfold difference in hydrogen ion concentration. So, a solution with a pH of 1 is 10 times more acidic than a solution with a pH of 2, and 100 times more acidic than a solution with a pH of 3.
It’s possible. If the molarity of hydrogen ions is greater than 1, you’ll have a negative value of pH. For example, you might expect a 12 M HCl solution to have a pH of -log(12) = -1.08.
Why don’t you hear more about negative pH? There are some complications in high molarity acid solutions that make pH calculations from acid molarity inaccurate and difficult to verify experimentally:
Even strong acids don’t dissociate completely at high concentrations. Some of the hydrogen remains bound to the chlorine, making the pH higher than you’d expect from the acid molarity.SATISH KUMAR – Scientist-D/Joint Director at Bureau of Indian Standards, Ghaziabad, Uttar Pradesh, India
Because there are so few waters per acid formula unit, the influence of hydrogen ions in the solution is enhanced. We say that the effective concentration of hydrogen ions (or the activity) is much higher than the actual concentration. The usual general chemistry text definition of pH as -log [H+] (negative the logarithm of the hydrogen ion molarity) is better written as pH = – log a (negative the logarithm of the hydrogen ion activity). This effect is very strong, and makes the pH much lower than you’d expect from the acid molarity.
If you were to dip a glass pH electrode into the 12 M HCl solution to actually measure the pH, you would get a pH that was higher than the true pH. This well-known defect in glass pH electrode measurements is called the “acid error”; it is sensitive to experimental conditions and difficult to correct for.
It’s possible. If the molarity of hydrogen ions is greater than 1, you’ll have a negative value of pH. For example, you might expect a 12 M HCl solution to have a pH of -log(12) = -1.08.Noor Aljammal10th – Ghent University, Feb, 2017
The definition of pH is the negative log of the hydrogen ion (or hydronium ion) activity. For example, a 1N solution of a strong acid would not have a pH of zero because the hydrogen ion activity is less than one because the activity coefficient is less than one. Also, since there are no standards for such low pH values, the pH meter readings are not very accurate. A significant part of the inaccuracy comes from the large potential developed at the liquid junction of the reference electrode. So while negative pH values are real, the meaning in terms of the true hydrogen ion activity is at best semi-quantitative.Richard A Durst – Cornell University, 17th Feb, 2017
- Sørensen, S.P.L. (1909). “Enzymstudien. II. Über die Messung und die Bedeutung der Wasserstoffionenkonzentration bei enzymatischen Prozessen” [Studies on enzymes. II. On the measurement and significance of the hydrogen ion concentration in enzyme processes]. Biochem. Z. 21: 131–304.