pH and pOH

Difference between pH and pOH

Before diving into the world of pH and pOH, we need to recall the definition of acids and bases. Acids and bases have different definitions, depending on who you ask!

Svante Arrhenius was the first chemist to classify acids and bases. Arrhenius acids are acids that dissociate in water and form H⁺ ions. Arrhenius bases, on the other hand, dissociate in water and form OH^- ions.

Other chemists who came up with acid and base definitions were Johannes Brønsted and Thomas Lowry. Together, they defined acids and bases by their ability to donate and accept protons (hydrogen ions). Brønsted-Lowry acids are proton (H⁺) donors, whereas Brønsted-Lowry bases are proton (H⁺) acceptors.

Now that we know a bit more about acids and bases, let's define what pH and pOH are!

pH is measured using a pH scale that ranges from 0 to 14. The smaller the pH, the more acidic a solution is, and the higher the concentration of H⁺ ions in the solution. So, acidic solutions have a very small concentration of OH^- ions compared to the high amount of H⁺ ions in the solution.

• Acidic solutions have a pH of smaller than 7
• Basic solutions have a pH of greater than 7
• Solutions containing pH 7 are considered neutral.

Fig.1-pH scale

pOH scale

The sum of the pH and pOH of an aqueous solution equals 14. The relationship between pH and pOH is as follows:

$\mathrm{pH}+\mathrm{pOH}=14$

So, if you know either the pH or pOH, you can easily calculate the missing value! For example, if you have a pOH of 3, then the pH would be equal to 11, because 11 + 3 = 14. Remember:

• The lower the pH value, the more acidic a solution is, the higher the H⁺ concentration and the lower the OH^- concentration.
• The lower the pOH value, the more basic a solution is, the higher the OH^- concentration and the lower the H⁺ concentration.

Fig.2-Scale showing pH and pOH value

pH and pOH Formula

The formula to calculate pH and pOH is:

$$\begin{gathered} \mathrm{pH}=-{\log }_{10}\left[{\mathrm{H}}^{+}\right] \\ \mathrm{pOH}=-{\log }_{10}\left[{\mathrm{OH}}^{-}\right] \end{gathered}$$

Those formulas might look a bit confusing, but don't worry, we will put them to practice!

Now, let's look at an example involving pOH.

Using pH and pOH to calculate Ion Concentration

We can rearrange the pH and pOH formulas to find the concentrations of hydrogen and hydroxide ions!

$$\begin{gathered} \left[{\mathrm{H}}^{+}\right]={10}^{-\mathrm{pH}} \\ \left[{\mathrm{OH}}^{-}\right]={10}^{-\mathrm{pOH}} \end{gathered}$$

Let's solve another example!

pOH to pH

During your AP chemistry exam, you will most likely be asked to convert pOH to pH. This can be accomplished by using the different formulas that we learned so far. The simplest way is by using the formula pH + pOH = 14.

However, sometimes you will have to use more than one formula to find pH. Let's look at an example.

pH of Strong Acids

Strong acids are acids that completely dissociate in aqueous solutions to form positive hydrogen ions.

Strong acids have very low pH, mostly between 0 and 1.

Some common strong acids you should know include:

• HCl
• HBr
• HI
• HNO₃
• HClO₃
• HClO₄
• H₂SO₄

Since strong acids completely dissociate in the solution, we can assume that the concentration of the strong acid will be equal to the concentration of the hydrogen ions in the solution!

pH of Strong Bases

Strong bases are bases that dissociate completely in aqueous solutions, forming OH^- ions. Contrary to strong acids, strong bases have very high pH, usually around 12-14

Some strong bases you need to be familiar with are:

• NaOH
• LiOH
• KOH
• RbOH
• CsOH
• Ca(OH)₂
• Mg(OH)₂
• Ba(OH)₂
• Sr(OH)₂

The same principle applies to strong bases. Since strong bases also completely dissociate, we assume that the concentration of the strong base will be equal to the concentration of the hydroxide ions in the solution.

Now, you should be familiar with pH and pOH and how to apply their formulas to different problems!

_References:

_{Malone, L. J., Dolter, T. O., & Gentemann, S. (2013). Basic concepts of Chemistry (8th ed.). Hoboken, NJ: John Wiley & Sons.}

_{Swanson, J. W. (2020). Everything you need to Ace Chemistry in one big fat notebook. Workman Pub.}

_{Saunders, N. (2020). Supersimple Chemistry: The Ultimate Bitesize Study Guide. London: Dorling Kindersley.}