We use our taste buds to detect each of the five tastes i.e. bitter, sweet, sour, salty and umami (savory). Foods containing acids such as vinegar and many fruits, have a sour taste. These foods have a low pH. In contrast, foods that are bitter generally have a higher pH due to the presence of basic amine groups that are common in bitter compounds.
The pH of foods affect their potential shelf life. For example foods that have a low pH (high-acid foods) resist the growth of bacteria and require less thermal treatment for pasteurization and commercial sterilization. A minimum pH of 4.6 is used as a standard for acidified foods. Foods such as ketchup and pickles must be acidified to a pH less than 4.6 to prevent the growth of Clostridium botulinum and other pathogens. My expectation is that you already have a basic and general understanding of acids, bases and pH. However, let’s review if you have not looked at this for a while. Â
Definition of Acid and Base
Arrhenius Theory: An acid is a substance that produce hydrogen ions (H+) when dissolved in water, and a base is a substance that produce Hydroxide ions (OH–) when dissolved in water
Bronsted-Lowry Theory: An acid is a substance that donates H+ and a base is a substance that accepts H+
Conjugate Acid and Conjugate Base: In the Bronsted-Lowry theory, the conjugate base is the chemical substance formed after an acid loses a proton (H+) and a conjugate acid is the chemical substance formed when a base accepts a proton.
So, remember:
The acid turns to the conjugate base, and
The base turns to the conjugate acid
Understanding pH
The pH of a substance is a measure of its H+ concentration. The higher the H+ concentration, the more acidic it is and hence the lower its pH. PH ranges from 0 to 14. This is a logarithmic scale which means that each interval differ by a factor of 10.
pH is calculated as the negative log of the concentration of hydrogen ions present in the substance, i.e.
pH = -log[H+]
You will also see this equation:
pH = -log[H3O+]
If you are asked to calculate [H3O+], give a certain pH, you can convert the above equation to
[H3O+] = 10-pH
Note that H3O+ (hydronium ion) is equivalent to H+ once hydrogen is in water (aqueous). Therefore H3O+ and H+ are often used interchangeably.
The concentration of H+ ions in acids is very small. For this reason, the number is converted to pH, which is an easier number to work with. For example, if the concentration of H+ = 1.7 X 10-3 M,
pH = -log[1.7 X 10-3] = 2.77
You may find it interesting to know that the pH of pure water at 23oC is 7.0. This is because the concentration of H+ = 1.0 X 10-7 M. Therefore,
pH = -log[1.0 X 10-7] = 7.0
The pH of an acid will not necessarily tell you how strong the acid is. Strong acids are those that dissociate completely in water while weak acids just partially dissociate. To know how much the acid will dissociate, you needed to know the acid dissociation constant (Ka).
A high Ka means strong dissociation and hence strong acid. Alternately, a low Ka means poor dissociation and hence weak acid. Ka is calculated as the concentration of hydrogen multiplied by the concentration of the conjugate base, divided by the concentration of the undissociated acid; i.e.
Note: Based on this equation, the Ka is the ratio of the amount of acid that is dissociated to the amount of acid that is not dissociated.
As we saw with the concentration of H+, the Ka is generally a very small number. Therefore, it can be converted to pKa by taking the negative log of the Ka to give a simpler number, i.e.,
pKa = -log[Ka]
For example the Ka of acetic acid = 1.6 x 10-5,
pKa = -log[1.6 x 10-5] = 4.8
You can convert pKa to Ka using the formula:
Ka = 10-pka
There is an inverse relationship between the pKa and the Ka. Therefore, strong acids (high Ka) will have a low pKa and weak acids will have a high pKa.
Earlier, you saw that pH can be calculated when you have the concentration of H+. You can also calculate the pH when you have the pKa along with the concentration of the conjugate base and and the concentration of the undissociated acid. See equation below, known as the Henderson-Hasselbalch Equation:
Buffer Solutions
The pH of most solutions will change drastically when a small amount of acid or base is added. This can be undesirable especially when working with enzymes in the lab. Drastic pH changes will cause them to denature. Therefore buffers are used to “protect” enzymes from denaturation. Buffers undergo very little change in pH when acids and bases are added. They are able to achieve this by neutralizing any base or acid that is added to the solution. The figure below shows how they work.
Buffers consist of nearly equal concentrations of an acid (in this case acetic acid) and its conjugate base. They will not neutralize each other since they are in equilibrium, but other acids or basses added to the buffer will be neutralized. In the figure, notice that when an acid (H3O+) is added, it is neutralized by the conjugate base and when a base (OH–) is added, it is neutralized by the acid. The result is only a small or no change in pH.
