This all-in-one online Buffer pH Calculator computes the pH of any given buffer solution, as well as the parameters of the buffer solution to achieve the required pH value. First choose whether your buffer is based on a weak acid and its conjugate base or on a weak base and its conjugate acid. Then enter the known parameters and press ‘Calculate’ to get the missing one.
Buffer Solutions
In many practical cases it is necessary to have solutions with a certain pH value, which remains almost constant when such a solution is diluted or when an acid or base is added. A solution that retain an almost constant pH value (within a certain range) when acid or base is added or when diluted is called buffer solution. In nature, many solutions have this property. And in chemical laboratories or in industrial production such solutions are often prepared on purpose.
A buffer solution (also called pH buffer or hydrogen ion buffer) is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. Usually buffer solutions are prepared from solutions of a weak acid and its salt or a weak base and its salt. For example, an acetate buffer solution is a solution of acetic acid CH3COOH and sodium acetate NaCH3COO. And the ammonia buffer solution is prepared from solutions of ammonium hydroxide NH4OH and ammonium chloride NH4Cl.
Acidic Buffers
The buffering effect of acidic buffers is based on the following reversible reactions:
where HA is a weak acid and A− is its conjugate base.
When analyzing these equations, it is important to remember that we are dealing with a weak acid, so that most of the acid molecules are in an undissociated state. The source of almost all conjugate base anions on the right side of the first equation is the salt included in the buffer mixture.
Therefore, when a strong acid is added to this mixture, hydrogen ions (H+) are added to the solution, and the equilibrium in the first of the above equations shifts to the left in accordance with Le Chatelier’s principle. In other words, there is an active binding of hydrogen ions. Because of this, the concentration of H+ increases less than expected for the amount of strong acid added.
Once most of the conjugate base anions are consumed, the pH drops rapidly because most of the added hydrogen ions are not bound by the conjugate base.
Similarly, if a strong alkali is added to this buffer mixture, the concentration of hydrogen ions will decrease less than expected for the amount of alkali added. This is because most of the added hydroxide ions (OH–) are consumed in the second of the above reactions and only a little are consumed in the neutralization reaction (it is the reaction that leads to an increase in pH):
Once most of the acid molecules are stripped of H+, the pH rises rapidly because most of the added alkali is consumed in the neutralization reaction.
Example of Acidic Buffer
As an example of an acidic buffer solution, consider an acetate buffer solution which is a solution of acetic acid CH3COOH and its salt sodium acetate NaCH3COO.
When a strong acid is added to an acetate buffer solution, hydrogen ions bind to the anions of a weak acid formed during the dissociation of the salt:
If a strong alkali is added to an acetate buffer solution, such as a solution of sodium hydroxide NaOH, then the hydroxide ions will be neutralized by a weak acid:
Thus, as a result of the binding of hydrogen ions or hydroxide ions from the added strong acid or strong base, the pH of the buffer solution practically does not change.
Basic Buffers
The buffering effect of basic buffers is based on the following reversible reactions:
where B is a weak base and BH+ is its conjugate acid.
These equations can be considered in a similar way, bearing in mind that the source of most of the cations of the conjugate acid (BH+) in the first equation is the corresponding salt included in the buffer mixture.
So when a strong acid is added to a basic buffer, the equilibrium in the second of these equations will shift to the left. That is hydrogen ions will actively bind to the base molecules and the pH value of the solution will decrease less than expected.
And when a strong alkali is added to this buffer solution, according to the first equation, the hydroxide ions (OH–) will be actively neutralized by the conjugate acid cations. In this case, the pH value of the solution will increase less than expected.
Example of Basic Buffer
As an example of a basic buffer solution, consider an ammonia buffer solution which is a solution of ammonium hydroxide NH4OH and its salt ammonium chloride NH4Cl.
When a strong acid is added to an ammonia buffer solution, hydrogen ions will be neutralized by a weak base:
If a strong alkali is added to an ammonia buffer solution, then the hydroxide ions bind to the cations NH4+ formed during the dissociation of the salt NH4Cl:
And again, as a result of the binding of hydrogen ions or hydroxide ions from the added strong acid or strong base, the pH of the buffer solution practically does not change.
Principles of Buffers
The above analysis gives us an explanation of why a weak acid (or weak base) is used to prepare a buffer solution. The answer is that only when the degree of dissociation of hydrogen ions (or hydroxide ions) is relatively low, the above reversible reactions are possible even when a sufficiently large amount of strong acid (or alkali) is added to the buffer solution.
In turn, the salt that is part of the buffer solution should be well dissociated and present in an amount commensurate with the amount of the weak acid (or base). This is necessary to ensure resistance to pH change when sufficient amounts of both a strong acid and a strong base are added.
Another important condition for the conjugate acid-base pair to have the property of a buffer system is that its concentration must be high enough compared to the concentration of H+ or OH– added.
Once any component of the buffer has fully reacted, the solution is no longer a buffer and rapid changes in pH can occur. There is a concept of buffer capacity. The buffer capacity is determined by the number of moles of strong acid or strong base that must be added to 1 liter of a buffer solution so that its pH changes by 1.
Obviously, a buffer solution that initially contains a greater amount of buffer constituents has a greater capacity. In addition, the buffer capacity reaches its highest value when both components of the buffer solution are present in approximately equal amounts.
Calculating pH of Buffer Solutions
To calculate the pH of a buffer solution, the Henderson-Hasselbach equation is commonly used. This equation can be derived quite simply from the behavior of weak acids (and bases) in solution, which is described by the kinetics of reversible reactions. This very equation is used in our Buffer pH Calculator.
For the case of an acidic buffer, this equation has the following form:
where
[HA] is the acid concentration,
[A–] is the concentration of conjugate base,
Ka is the acid dissociation constant,
pH = – log ( [H+] ),
pKa = – log ( Ka ).
Similarly, in the case of a basic buffer, the corresponding equation has the following form:
where
[B] is the base concentration,
[BH+] is the concentration of conjugate acid,
Kb is the base dissociation constant,
pKb = – log ( Kb ).
In deriving this formula, we took into account that the equation describing the dissociation of water at 25 °C has the following form: [H+][OH–] = 10-14. Therefore, taking the decimal logarithm from this expression, we arrive at the following relation:
It should be emphasized that the following assumptions were made when deriving the Henderson-Hasselbach equation.
• The acid included in the acidic buffer solution is monobasic (contains only one replaceable hydrogen atom per molecule), and the base included in the basic buffer solution is monoacidic (produces only one hydroxide ion).
• We can completely neglect the self-dissociation of water molecules.
• The salt is completely dissociated in the buffer solution.
Below is a table with pKa values for some common buffers:
Examples of Buffer pH Calculations
Example 1. Demonstration of the properties of acetate buffer solution.
Consider 1 liter of 0.1 M acetate buffer solution (consisting of 0.1 M CH3COOH and 0.1 M NaCH3COO). What is the pH of such a solution? How it changes when 0.01 mole of hydrochloric acid or sodium hydroxide is added? Compare these changes with those that occur when these substances are added to the same volume of pure water.
The pH of the above buffer solution can be easily found from the Henderson-Hasselbach equation given the value of pKa = 4.756 and the fact that the concentration of the buffer acid is equal to the concentration of its conjugate base. Since the decimal logarithm of the ratio of these concentrations is zero, then pH = pKa = 4.756.
When a strong acid is added to our buffer solution, almost all of its molecules dissociate and the released hydrogen ions bind to the acetate anions. This increases the concentration of acetic acid and decreases the concentration of sodium acetate (the main source of acetate anions) by the same amount.
So when 0.01 mole of HCl is added to 1 liter of 0.1 M acetate buffer solution we will have as a result 0.11 M CH3COOH and 0.09 M NaCH3COO. Plugging this figures and the known value of pKa = 4.756 into our Buffer pH Calculator we will get pH = 4.6688.
When a strong alkali is added to our buffer solution, almost all of its molecules dissociate and the released hydroxyl ions bind to the hydrogen ions taken away from the acetic acid molecules. This increases the concentration of sodium acetate and decreases the concentration of acetic acid by the same amount.
So when 0.01 mole of NaOH is added to 1 liter of 0.1 M acetate buffer solution we will have as a result 0.09 M CH3COOH and 0.11 M NaCH3COO. Plugging all the figures into our calculator we will get pH = 4.8432.
We know that the pH of our original buffer solution equals to 4.756. It is easy to calculate that adding the indicated amount of a strong acid or alkali to our buffer solution leads to a pH change of less than 0.09 units.
If now we add the indicated amount of acid or alkali to 1 liter of pure water, we would get a pH value of 2 and 12, respectively, which is 5 units different from the pH value of water equal to 7.
Thus, we have a clear manifestation of the buffer effect!
Example 2. Calculation of the amount of salt for an acidic buffer solution.
It is necessary to prepare a buffer solution with pH = 4.6 containing 0.03 mol of propionic acid (CH3CH2COOH) and some amount of sodium propionate (NaCH3CH2COO) dissolved in water, sufficient to obtain 1 L of solution. How much sodium propionate is needed to prepare such a solution?
We know that for propionic acid the pKa = 4.873. This value along with given values of pH = 4.6 and 0.03 M of propionic acid concentration we can plug into our calculator. As the result we get 0.016 mol of sodium propionate.
Buffer Systems in Nature
Buffer systems are often found in biology. They create a constant pH of biological fluids (blood, lymph, etc.). Buffer solutions such as phosphate and carbonate are of great importance for the survival of organisms, as they maintain a constant pH.
Human blood is a weak base, pH = 7.35-7.47. Substances in the blood, including hemoglobin, create a buffering capacity that the body needs, since a change in pH leads to a loss of oxygen and carbon dioxide transport. When the pH drops to 7.0, a person loses consciousness, and when it rises to 7.8, convulsions begin.
In other body fluids, pH is also kept constant by proteins whose solutions exhibit buffering properties. Proteins have acidic and basic groups that are part of the amino groups and carboxyl groups.
Buffer systems are also essential for sustaining life and biodiversity on our planet.
Some plants can thrive at certain pH values and ion concentrations, which is achieved by various soil buffering properties. Soil microorganisms have the ability to maintain the pH at the level necessary for their existence, releasing the appropriate substances.
The maintenance of a constant concentration of oxygen and carbon dioxide in the atmosphere and in the oceans is due to the most powerful buffer system, which includes green plants (photosynthesis).
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