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What Is A Buffer Biology Essay

Paper Type: Free Essay Subject: Biology
Wordcount: 3049 words Published: 1st Jan 2015

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A buffer is an aqueous solution that has a highly stable pH. If you add acid or base to a buffered solution, its pH will not change significantly. Similarly, adding water to a buffer or allowing water to evaporate will not change the pH of a buffer.

A buffer is a substance in a solution that can neutralize either an acid or a base. A substance is said to be naturally buffered if it has a buffering action in its natural state .Buffers are used in chemistry to modulate and stabilize the pH of a solution. A buffered solution can be made in a variety of ways, most simply by creating a mixture of a weak acid and its conjugate base.

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Buffers are chemicals that, when added to water, tend to maintain a certain pH. This is due to the buffer’s ability to either accept or donate a proton (hydrogen ion, H+) or hydroxyl (OH-) to keep the pH in a certain range. Buffers are different than strong acids and bases because buffers do not donate all of their protons or hydroxyls within their buffering range.

A buffer is a substance, usually a salt, that can bind with either the positively-charged hydrogen ion of an acid or the negatively-charged hydroxide ion of a base. If a solution is buffered, acids or bases can be added to the solution without changing its pH as drastically as if the acid or base were added to an unbuffered solution

Natural Buffers

A solution is said to be naturally buffered if it contains buffering compounds as it exists in nature. Blood is an example of a naturally buffered solution. Blood must maintain a pH close to 7.4 in order to carry oxygen effectively and contains compounds that react to acids or bases in order to neutralize them.


A compound can buffer the pH of a solution only when its concentration is sufficient and when the pH of the solution is close (within about one pH unit) to its pKa.  To make a buffer you must first pick a compound whose pKa   is close to the pH  you want for the solution, and then decide what the buffer concentration should be.  Typically, buffer concentrations are between 1 mM and 200 mM, depending on the desired ionic strength and the buffering capacity required.  If the pH is expected to decrease during the experiment, choose a buffer with a pKa slightly below the working pH.  Conversely, if the pH is expected to increase during the experiment, select a buffer with a pKa slightly above the working pH.  Having decided on the total buffer concentration, you must adjust the ratio of the protonated and unprotonated forms of the buffer in your solution so as to give the desired pH.  Typically, buffers are composed of  weak acids and their salts, or weak bases and their salts.  If the protonated form is uncharged, it is an acid (like acetic acid), and its unprotonated form is a salt (e.g., sodium acetate).  Conversely, if the unprotonated form is uncharged it is a base (like Tris base), and its protonated form is a salt (e.g., TrisHCl).

Four practical ways to make a buffer are described below:

The Slow and Stupid Method – 

To avoid adding extra salt to a solution, prepare a buffer composed of an acid and its salt by dissolving the acid form of the buffer in about ~60% of the water required for the final solution volume.  Adjust the pH using a strong base, such as NaOH.  When preparing a buffer composed of a base and its salt, start with the base form and adjust the pH with strong acid, such as HCl.  After the pH is correct, dilute to just under the final solution volume.  Check the pH and correct if necessary, then add water to the final volume. 


Advantages: Easy to understand. 

Disadvantages: Slow.  May require lots of base (or acid).  If the base (or acid) is concentrated, it is easy to overshoot the pH. If the base (or acid) is dilute, it is easy to overshoot the volume.  Ionic strength will be unknown. Adding a strong acid or base can result in temperature changes, which will make pH readings inaccurate (due to its dependence to temperature) unless the solution is brought back to its initial temperature. 

(2)The Mentally Taxing Method – Using the buffer pKa , calculate the amounts (in moles) of acid/salt or base/salt present in the buffer at the desired pH.  If both forms (i.e., the acid and the salt) are available, convert the amount required from moles to grams, using the molecular weight of that component, and the weigh out the correct amounts of  both forms.  If only one form is available, you can prepare the buffer by adding all of the buffer as one form, and then adding acid or base to convert some of the added buffer to the other form.  Decide what the total concentration of buffer will be in the solution, and convert the concentration to amount (in moles) using the volume of solution, and then to grams, using the molecular weight of the buffer form available.  Then calculate the amounts (in moles) of each form that will be present in the final solution, using the buffer pKa  and the desired pH.   Then calculate how much strong acid or base must be added to convert enough of the buffer form added to the other form,  to give the correct amounts of each form at the pH of the final solution.  Dissolve the buffer and strong acid or base in slightly less water than is required for the final solution volume.  Check the pH and correct if necessary.  Add water to the final volume.


Advantages: Fast.  Easy to prepare.  Additional pH adjustment is rarely necessary, and when necessary, the adjustment is small. Ionic strength easily calculated.


The Two Solution Method – Make separate solutions of  the acid form and base form of the buffer, both solutions having  the same buffer concentration (and ionic strength, if required) as the concentration of total buffer in the final solution.  To obtain the desired pH, add one solution to the other while monitoring the pH with a pH meter. 


Advantages:  Easy to do. 

Disadvantages: Requires both forms of buffer.  The required solution volumes are proportional to the ratio of buffer components in the final solution at the desired final pH, so making equal amounts of each form may waste a lot of one solution.

The Completely Mindless Method – Find a table of the correct amounts of acid/salt or base/salt required for different pH’s, and dissolve the components in slightly less water than is required for the final solution volume.  Check that the pH and correct if necessary. Add water to the final volume. 


Advantages: Easy to do (with appropriate table).  Convenient for frequently prepared buffers. 

Disadvantages: May be impossible to find table.  Table may be incorrect.  Requires both forms of buffer.  Component amounts from table will need to be adjusted to give the buffer concentration and volume in your solution. Ionic strength is unknown.

 Common Buffer Preparations


Stock Soln


Amount per Liter Soln.

Conc. Stock Soln

Final Conc.


(Phosphate Buffered Saline)  

adj. pH ~7.3






 80 g  

   2 g  

 11.5 g  

   2 g

 1.37 M  

   27 mM  

   43 mM  

   14 mM

137 mM  

    2.7 mM  

    4.3 mM  

    1.4 Mm


adj. pH ~7.0



Sodium citrate

 175 g  

   88 g

 3 M  

 0.3 M

 150 mM  

   15 Mm


(Saline Tris EDTA) 


Tris base  


EDTA (acid)

  1.2 g  

  0.6 g  

  0.29 g

 10 mM  

 10 mM  

   1 mM

 10 mM  

 10 mM  

   1 mM


(Tris acetate EDTA)  

pH ~8.5


Tris base  

Acetic acid (glacial)  


 242 g  

   57.1 mL

  37.2 g 

 2 M (Tris acetate)

0.1 M  

 40 mM (Tris acetate)  

   2 mM


(Tris borate EDTA)  

pH ~8.0


 Tris base  

 Boric acid  


 108 g  

   55 g  

   40 mL 

(0.5 M pH 8) 

 0.89 M  

 0.89 M  

 0.02 M

 89 mM  

 89 mM  

   2 mM


(Tris EDTA)  

pH ~7.5


Tris base  

H2 EDTA (acid)

 1.2 g  

 0.29 g

 10 mM  

   1 mM 


Useful buffer mixtures


pH range

HCl, Sodium citrate

1 – 5

Citric acid, Sodium citrate

2.5 – 5.6

Acetic acid, Sodium acetate

3.7 – 5.6


5.8 – 8 [3]

Na2HPO4, NaH2PO4

6 – 7.5 [4]

Borax, Sodium hydroxide

9.2 – 11

 “Universal” buffer mixtures

By combining substances with pKa values differing by only two or less and adjusting the pH a wide-range of buffers can be obtained. Citric acid is a useful component of a buffer mixture because it has three pKa values, separated by less than two. The buffer range can be extended by adding other buffering agents. The following two-component mixtures (McIlvaine’s buffer solutions) have a buffer range of pH 3 to 8.

0.2M Na2HPO4 /mL

0.1M Citric Acid /Ml




















A mixture containing citric acid, potassium dihydrogen phosphate, boric acid, and diethyl barbituric acid can be made to cover the pH range 2.6 to 12.[6]

Other universal buffers are Carmody buffer and Britton-Robinson buffer, developed in 1931.


An important number for any aqueous solution is its pH. The pH is the negative logarithm of the concentration of hydrogen ions (often represented as a hydronium ion, which is a water molecule with an extra proton attached). Anything with a pH of less than 7 is considered to be acidic, and a solution that has a pH of greater than 7 is basic. Pure water, which has an equal amount of acid and base in it, is defined as having a pH of 7.


Buffered solutions are used to make a solution that exhibits very little change in its pH when small amounts of an acid or base are added to it. A buffer can be made by addition of a weak acid and its conjugate base to a solution. Alternately, a weak base and its conjugate acid can be used. Buffers are most effective when the amounts of the weak acid/base and the conjugate base/acid are used—generally to be effective, neither quantity should be more than 10 times that of the other.

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Control of ph by buffer

Buffers are chemicals that, when added to water, tend to maintain a certain pH. This is due to the buffer’s ability to either accept or donate a proton (hydrogen ion, H+) or hydroxyl (OH-) to keep the pH in a certain range. Buffers are different than strong acids and bases because buffers do not donate all of their protons or hydroxyls within their buffering range.


The measure of pH gives the concentration of hydrogen ions in a solution. Pure water has a pH of 7.0 and has an equal balance of hydrogen ions and hydroxyl ions. The pH will be lower if an acid is added and higher if a base is added. For various purposes, it is useful to add a buffer so that the pH remains constant even if an acid or base is added to the solution. Buffers are used to maintain constant pH.

Buffer Equilibrium:-

Different buffers work at different pHs. A buffer is most efficient at maintaining a constant pH when the pH equals its acid dissociation constant, its pKa. At this pH, the buffer consists of an equal solution of protonated and de-protonated conjugate base. This often denoted by a theoretical buffer molecule, AH, dissociating into a proton, H+, and the conjugate base, A-. The pKa is determined by the pH where the concentration of AH equals the concentration of A-.

Buffering Acids:-

At its pKa, a buffer can maintain a constant pH by accepting free protons. Adding an acid to the solution, such as HCl, will cause free protons to enter the solution. The conjugate base will accept the free protons, causing the pH to be unchanged. The solution will remain at the same pH as long as there is enough conjugate base to accept the additional protons. The reaction can be written as the dissolution of the acid: HCL–H+ + Cl- leading to free H+ in the solution. The further reaction with the buffer occurs: H+ + A- — HA. This removes the free protons and constant pH is maintained.

Buffering Bases:-

Adding a base such as sodium hydroxide to a solution will cause an increase in the hydroxyl concentration. In a buffered solution, the protons attached to the undissociated buffer are donated to the solution, forming H2O with the free hydroxyls. This counters the effect of the base and maintains the pH of the solution as long as there is undissociated buffer available in the solution. In this case, the reaction can be written as the dissolution of the base: NaOH– Na+ + OH- leading to the second reaction of OH- + HA — H2O + A-.

Different Buffers:-

Different buffers have different pKas and can be used to buffer solutions at a wide range of pHs. Chemicals used as buffers often have the ability to donate or accept multiple protons or bases. These buffers hence have two or more different pKas corresponding to how many protons or hydroxyls per molecule they can donate or accept. Different buffer molecules can be combined to form customized buffer ranges. For a list of biological buffers, see Resources.


When hydrogen ions are added to a buffer, they will be neutralized by the base in the buffer. Hydroxide ions will be neutralized by the acid. These neutralization reactions will not have much effect on the overall pH of the buffer solution.

When you select an acid for a buffer solution, try to choose an acid that has a pKa close to your desired pH. This will give your buffer nearly equivalent amounts of acid and conjugate base so it will be able to neutralize as much H+ and OH- as possible.


A buffer solution is used to resist changes in pH when a certain amount of strong acid or base is added to the solution. It is an important part of biological systems in living organisms as well as in the laboratory. The reason a buffer works to maintain a certain pH is that the concentration of weak acid and base is kept in a specific ratio in line with the acid titration curve.

The titration curve

A titration curve is a graph that relates the relative concentrations of a weak acid to its conjugate base by graphing pH versus amount of base added. In the region of the titration curve where the graph is almost flat, the pH changes very little with added base and therefore this would be a good pH for the buffer.

Buffer in cells and blood

The main buffer found in living cells is the H2PO4/HPO4- buffer pair. In blood, the main buffer is the H2CO3/HCO3- pair. This system relies on dissociation of carbonic acid, which has a pKa of 6.37. The pH of human blood needs to remain at around a pH of 7.4 and therefore the system also involves carbon dioxide transported to the lungs.

Phosphate buffer

The phosphate buffer is based on tris (hydroxymethyl) aminomethane or TRIS. This buffer has a pKa of 8.3 and is found both in living organisms as well as used in the laboratory. It is a good buffer because it does not tend to interfere with the system being studied.

Buffers in enzymatic reactions

Laboratory methods to isolate an enzyme use buffered solutions because an enzyme can only function in a narrow pH range. Enzymes are very sensitive to pH as well as salt concentrations. Therefore, it is important to use a buffer with a very good buffering capacity for the specific pH in order for the experiment to be successful.

Physiological consequences

Respiration plays a role in buffering of blood by controlling the rate of respiration depending on the need for hydrogen ions or increased acidity. Increasing the rate of respiration is helpful when there is a buildup of hydrogen ions or acidity in the blood. The H+ ions bind to bicarbonate to form carbonic acid. This raises the level of carbon dioxide in the lungs. Increasing the level of respiration removes the excess carbon dioxide. Therefore, here the buffering system is used to keep pH level of the blood within the required narrow range.


The pH of a buffered solution is defined by the Henderson-Hasselbalch equation, which states that the pH of a solution is equal to the acid dissociation constant of the weak acid plus the logarithm of the ratio of the concentrations of conjugate base to the weak acid. The acid dissociation constant is a number that defines the tendency of an acid to dissociate and form hydrogen ions. A strong acid will have a very low dissociation constant, whereas a weaker one may have a significantly higher one, around 5.


Buffered solutions have many applications in chemical manufacturing for processes that require a specific pH range to work. This is also true for the human body, which contains many enzymes that are only functional at a specific pH. Outside of this range, the enzymes are either unable to catalyze reactions, or in some case will misfold and become broken down; thus, a mixture of carbonic acid and bicarbonate is used by the body to keep pH of the blood between 7.35 and 7.45.


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