term paper on mechanism of buffer system

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16.7 Buffers

Learning objectives.

By the end of this section, you will be able to:

• Define what a buffer is and describe how it reacts with an acid or a base

Weak acids are relatively common, even in the foods we eat. But we occasionally come across a strong acid or base, such as stomach acid, that has a strongly acidic pH of 1–2. By definition, strong acids and bases can produce a relatively large amount of hydrogen or hydroxide ions and, as a consequence, have marked chemical activity. In addition, very small amounts of strong acids and bases can change the pH of a solution very quickly. If 1 mL of stomach acid [which we will approximate as 0.05 M HCl( aq )] is added to the bloodstream, and if no correcting mechanism is present, the pH of the blood would go from about 7.4 to about 4.9 — a pH that is not conducive to life. Fortunately, the body has a mechanism for minimizing such dramatic pH changes. This mechanism involves a buffer, a solution that resists dramatic changes in pH.

Watch Buffers, the Acid Rain Slayer: Crash Course Chemistry #31 (11min 40s) .

Video Source: Crash Course. (2013, September 16). Buffers, the acid rain slayer: Crash course Chemistry #31 [Video]. YouTube.

Buffers resist dramatic changes in pH by being composed of certain pairs of solutes: either a weak acid plus a salt derived from that weak acid, or a weak base plus a salt of that weak base. For example, a buffer can be composed of dissolved acetic acid (HC 2 H 3 O 2 , a weak acid) and sodium acetate (NaC 2 H 3 O 2 , a salt derived from that acid). Another example of a buffer is a solution containing ammonia (NH 3 , a weak base) and ammonium chloride (NH 4 Cl, a salt derived from that base).

Characteristics of a Good Buffer

Good buffering systems have the following characteristics:

• The solution contains a weak acid and its conjugate base OR a weak base and its conjugate acid
• The buffer resists changes in pH by reacting with added acid or base, so these ions do not accumulate.
• Any added acid reacts with the conjugate base to resist pH changes
• Any added base reacts with the conjugate acid to resist pH changes

Buffers cannot be made from a strong acid (or strong base) and its conjugate since these solutions ionize completely in water. Also take note, water is not a buffer.

Source: “Characteristics of a Good Buffer” by Jackie MacDonald, CC BY-NC-4.0

How Buffers Work

Let’s consider an acetic acid – sodium acetate buffer to demonstrate how buffers work. If a strong base – a source of OH − ( aq ) ions – is added to the buffer solution, those hydroxide ions will react with the acetic acid in an acid-base reaction:

HC 2 H 3 O 2 ( aq ) + OH − ( aq ) → H 2 O (l) + C 2 H 3 O 2 – ( aq )

Rather than changing the pH dramatically by making the solution basic, the added hydroxide ions react to make water, and the pH does not change much.

If a strong acid – a source of H + ions – is added to the buffer solution, the H + ions will react with the anion from the salt. Because HC 2 H 3 O 2  is a weak acid, it is not ionized much. This means that if lots of hydrogen ions and acetate ions (from sodium acetate) are present in the same solution, they will come together to make acetic acid:

H + (aq) + C 2 H 3 O 2 − (aq) → HC 2 H 3 O 2 (aq)

Rather than changing the pH dramatically and making the solution acidic, the added hydrogen ions react to make molecules of a weak acid.

In chemistry texts and sources, you may have noticed that H + and H 3 O + are used interchangeably in contexts when the proton donor-acceptor mechanism does not need to be emphasized. Since it is easier to write the H + proton, chemists often use it to represent acid-base reactions or to explain general concepts in buffering systems. Thus, it is permissible to talk about “hydrogen ions” and use the formula H + in writing chemical equations as long as you remember that they are not to be taken literally in the context of aqueous solutions.

Source: “ The Hydronium Ion ” by Stephen Lower & Avneet Kahlon In Acids and Bases in Aqueous Solutions , licensed under CC BY 3.0 .

Figure 16.7a illustrates both actions of the acetic acid – sodium acetate buffer.

Figure 16.7b illustrates a basic summary of the action of buffers when small amounts of strong base and acid is added.

Buffers made from weak bases and salts of weak bases act similarly. For example, in a buffer containing NH 3  and NH 4 Cl, ammonia molecules can react with any excess hydrogen ions introduced by strong acids:

NH 3 ( aq ) + H + ( aq ) → NH 4 + ( aq )

while the ammonium ion, NH 4 + (aq) can react with any hydroxide ions introduced by strong bases:

NH 4 + ( aq ) + OH − ( aq ) → NH 3 ( aq ) + H 2 O (l) 

Example 16.7a

Which solute combinations can make a buffer solution? Assume that all are aqueous solutions.

• HCHO 2  and NaCHO 2
• HCl and NaCl
• CH 3 NH 2  and CH 3 NH 3 Cl
• NH 3  and NaOH
• Formic acid (HCHO 2 ) is a weak acid, while NaCHO 2 is the salt made from the anion of the weak acid – the formate ion (CHO 2 − ). The combination of these two solutes would make a buffer solution.
• Hydrochloric acid (HCl) is a strong acid, not a weak acid, so the combination of these two solutes would not make a buffer solution.
• Methylamine (CH 3 NH 2 ) is like ammonia with one of its hydrogen atoms substituted with a CH 3  (methyl) group. Because it is not on our list of strong bases, we can assume that it is a weak base. The compound CH 3 NH 3 Cl is a salt made from that weak base, so the combination of these two solutes would make a buffer solution.
• Ammonia (NH 3 ) is a weak base, but NaOH is a strong base. The combination of these two solutes would not make a buffer solution.

Exercise 16.7a

• NaHCO 3  and NaCl
• H 3 PO 4  and NaH 2 PO 4
• NH 3  and (NH 4 ) 3 PO 4
• NaOH and NaCl

Many people are aware of the concept of buffers from buffered aspirin . Aspirin is well known as a pain reliever and fever reducer. Buffered aspirin contains aspirin (acetylsalicylic acid) and also has magnesium carbonate, calcium carbonate, magnesium oxide, or some other salt. The salt regulates the acidity of the aspirin to minimize its acidic side effects in the stomach. The salt acts like a base, while aspirin is itself a weak acid due to its carboxylic acid group. The H atom in that group can be donated, and therefore, aspirin can act as a Brønsted-Lowry acid. Figure 16.7c and 16.7d show the molecular structure of aspirin in 3D and 2D.

Links to Interactive Learning Tools

Explore Learn the Basics about Buffers from eCampusOntario H5P Studio .

Attribution & References

Except where otherwise noted, this page is adapted by Jackie MacDonald from a section in “ 14.10: Buffers- Solutions that Resist pH Change ” In Map: Introductory Chemistry (Tro) by Marisa Alviar-Agnew & Henry Agnew, shared under a CK-12 license.

Stephen Lower & Avneet Kahlon. (2022, August 10). The Hydronium Ion . Chemistry LibreTexts.

• No; NaHCO 3 and NaCl are not acid/base conjugate pairs;
• Yes; H 3 PO 4 is a weak acid and NaH 2 PO 4 is a salt of its conjugate base;
• No - NaOH is a strong base, a buffer requires a weak base or acid and its conjugate. ↵

Enhanced Introductory College Chemistry Copyright © 2023 by Gregory Anderson; Caryn Fahey; Jackie MacDonald; Adrienne Richards; Samantha Sullivan Sauer; J.R. van Haarlem; and David Wegman is licensed under a Creative Commons Attribution 4.0 International License , except where otherwise noted.

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Describe the chemistry of buffer mechanisms and explain their relevant roles in the body

A buffer is a solution which consists of a weak acid and its conjugate base , that can resist a change in pH when a stronger acid or base is added.

• Is a key part of acid-base homeostasis
• In one experiment, dogs were infused with 14,000,000 nmol.L -1 of H + , with a corresponding rise in H + of only 36 nmol.L -1

Efficacy of a buffer system is determined by:

• pKa of the buffer 80% of buffering occurs within 1 pH unit of the pKa of the system.
• pH of the solution
• Amount of buffer
• This alters the concentration of reactants at either end of the equation, thus altering the speed of the reaction via the Law of Mass Action

Buffer Systems

Important buffer systems include:

• Bicarbonate buffer system
• Haemoglobin buffer system
• Phosphate buffer system

All buffer systems are in equilibrium with the same amount of H + . This is known as the isohydric principle .

Bicarbonate Buffer System

The bicarbonate buffer system is:

• Bicarbonate is formed in the erythrocyte and then secreted into plasma
• Bicarbonate diffuses into the interstitium and is also the dominant fluid buffer in interstitial space
• Formed in the erythrocyte
• A buffer pair consisting of bicarbonate and carbonic acid Carbonic acid is exceedingly short lived in any environment even remotely compatible with life and it rapidly dissociates to HCO 3 - and H + .

Hydrogen ions are consumed or released by the following reaction:

• Carbonic anhydrase (present in erythrocytes) is an enzyme which allows rapid conversion of H 2 O and CO 2 to H 2 CO 3 (and back again)
• The pKa for the second stage of the reaction is 9.3 and so essentially no CO 3 2- exists in blood Clincically this reaction can be ignored.
• CO 2 is then able to be exhaled, which prevents equilibration and allows the system to buffer more acid

Bicarbonate is an effective buffer because it is:

• Present in large amounts
• CO 2 can be adjusted by changing ventilation
• Bicarbonate can be adjusted by changing renal elimination
• This prevents the bicarbonate buffer system from equilibrating and allows it to resist large changes in pH despite its low pKa However, because it relies heavily on changes in pulmonary ventilation it is unable to effectively buffer respiratory acid-base disturbances .

Protein Buffer System

• All proteins contain potential buffer groups However, the useful one at physiological pH is the imidazole groups of the histidine residues .
• Extracellularly, proteins have a small contribution which is entirely due to their low pKa
• Intracellular protein concentration is much greater than extracellular concentration
• Intracellular pH is much lower (~6.8) and closer to their pKa

Haemoglobin Buffer System

Haemoglobin is:

• A protein buffer system
• Exists in greater amounts than plasma proteins (150g.L -1 compared to 70g.L -1 )
• Each molecule contains 38 histidine residues This results in 1g of Hb ~3x the buffering capacity of 1g of plasma protein.

In the cell:

• Additional H + ions are bound to Hb molecules
• HCO 3 - diffuses down its concentration gradient into plasma Electroneutrality is maintained through the inwards movement of Cl - .
• Dissolved CO 2 will also form carbamino compounds by binding to the terminal amino groups
• Deoxyhaemoglobin has a pKa of 8.2 Because of its higher pKa, deoxyhaemoglobin will more readily accept H + ions which makes it a better buffer of acidic solutions.
• Oxyhaemoglobin has a pKa of 6.6
• Both are essentially equidistant from normal pH, and are equally effective buffers
• This is the mechanism behind the Haldane effect , and why venous blood is only slightly more acidic than arterial blood

Phosphate Buffer System

Phosphoric acid is:

• Tribasic and can therefore potentially donate three hydrogen ions
• The quantitative effect is low despite the optimal pKa due to the low plasma concentration of phosphate
• At higher concentrations, such as intracellularly and in urine, it is a significant contributor
• In prolonged acidosis, CaPO 4 can be mobilised from bones and can be considered as an alkali reserve
• Alex Yartsev offers an excellent discussion on buffering in his excellent trademark prose at Deranged Physiology
• Brandis's anaesthesia MCQ is required reading
• Barrett KE, Barman SM, Boitano S, Brooks HL. Ganong's Review of Medical Physiology. 24th Ed. McGraw Hill. 2012.
• Kam P, Power I. Principles of Physiology for the Anaesthetist. 3rd Ed. Hodder Education. 2012.

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26.4 Acid-Base Balance

Learning objectives.

By the end of this section, you will be able to:

• Identify the most powerful buffer system in the body
• Identify the most rapid buffer system in the body
• Describe the protein buffer systems.
• Explain the way in which the respiratory system affects blood pH
• Describe how the kidney affects acid-base balance

Proper physiological functioning depends on a very tight balance between the concentrations of acids and bases in the blood. Acid-balance balance is measured using the pH scale, as shown in Figure 26.4.1 . A variety of buffering systems permits blood and other bodily fluids to maintain a narrow pH range, even in the face of perturbations. A buffer is a chemical system that prevents a radical change in fluid pH by dampening the change in hydrogen ion concentrations in the case of excess acid or base. Most commonly, the substance that absorbs the ions is either a weak acid, which takes up hydroxyl ions, or a weak base, which takes up hydrogen ions.

Buffer Systems in the Body

The buffer systems in the human body are extremely efficient, and different systems work at different rates. It takes only seconds for the chemical buffers in the blood to make adjustments to pH. The respiratory tract can adjust the blood pH upward in minutes by exhaling CO 2 from the body. The renal system can also adjust blood pH through the excretion of hydrogen ions (H + ) and the conservation of bicarbonate, but this process takes hours to days to have an effect.

The buffer systems functioning in blood plasma include plasma proteins, phosphate, and bicarbonate and carbonic acid buffers. The kidneys help control acid-base balance by excreting hydrogen ions and generating bicarbonate that helps maintain blood plasma pH within a normal range. Protein buffer systems work predominantly inside cells.

Protein Buffers in Blood Plasma and Cells

Nearly all proteins can function as buffers. Proteins are made up of amino acids, which contain positively charged amino groups and negatively charged carboxyl groups. The charged regions of these molecules can bind hydrogen and hydroxyl ions, and thus function as buffers. Buffering by proteins accounts for two-thirds of the buffering power of the blood and most of the buffering within cells.

Hemoglobin as a Buffer

Hemoglobin is the principal protein inside of red blood cells and accounts for one-third of the mass of the cell. During the conversion of CO 2 into bicarbonate, hydrogen ions liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen. This buffering helps maintain normal pH. The process is reversed in the pulmonary capillaries to re-form CO 2 , which then can diffuse into the air sacs to be exhaled into the atmosphere. This process is discussed in detail in the chapter on the respiratory system.

Phosphate Buffer

Phosphates are found in the blood in two forms: sodium dihydrogen phosphate (Na 2 H 2 PO 4 − ), which is a weak acid, and sodium monohydrogen phosphate (Na 2 HPO4 2- ), which is a weak base. When Na 2 HPO4 2-  comes into contact with a strong acid, such as HCl, the base picks up a second hydrogen ion to form the weak acid Na 2 H 2 PO 4 − and sodium chloride, NaCl. When Na 2 HPO4 2− (the weak acid) comes into contact with a strong base, such as sodium hydroxide (NaOH), the weak acid reverts back to the weak base and produces water. Acids and bases are still present, but they hold onto the ions.

Bicarbonate-Carbonic Acid Buffer

The bicarbonate-carbonic acid buffer works in a fashion similar to phosphate buffers. The bicarbonate is regulated in the blood by sodium, as are the phosphate ions. When sodium bicarbonate (NaHCO 3 ), comes into contact with a strong acid, such as HCl, carbonic acid (H 2 CO 3 ), which is a weak acid, and NaCl are formed. When carbonic acid comes into contact with a strong base, such as NaOH, bicarbonate and water are formed.

As with the phosphate buffer, a weak acid or weak base captures the free ions, and a significant change in pH is prevented. Bicarbonate ions and carbonic acid are present in the blood in a 20:1 ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic. This is useful because most of the body’s metabolic wastes, such as lactic acid and ketones, are acids. Carbonic acid levels in the blood are controlled by the expiration of CO 2 through the lungs. In red blood cells, carbonic anhydrase forces the dissociation of the acid, rendering the blood less acidic. Because of this acid dissociation, CO 2 is exhaled (see equations above). The level of bicarbonate in the blood is controlled through the renal system, where bicarbonate ions in the renal filtrate are conserved and passed back into the blood. However, the bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body.

CO 2 + H 2 O ↔ H 2 CO 3 ↔ H + + HCO 3 –

Respiratory Regulation of Acid-Base Balance

The chemical reactions that regulate the levels of CO 2 and carbonic acid occur in the lungs when blood travels through the lung’s pulmonary capillaries. Minor adjustments in breathing are usually sufficient to adjust the pH of the blood by changing how much CO 2 is exhaled. In fact, doubling the respiratory rate for less than 1 minute, removing “extra” CO 2 , would increase the blood pH by 0.2. This situation is common if you are exercising strenuously over a period of time. To keep up the necessary energy production, you would produce excess CO 2 (and lactic acid if exercising beyond your aerobic threshold). In order to balance the increased acid production, the respiration rate goes up to remove the CO 2 . This helps to keep you from developing acidosis.

The body regulates the respiratory rate by the use of chemoreceptors, which primarily use CO 2 as a signal. Peripheral blood sensors are found in the walls of the aorta and carotid arteries. These sensors signal the brain to provide immediate adjustments to the respiratory rate if CO 2 levels rise or fall. Yet other sensors are found in the brain itself. Changes in the pH of CSF affect the respiratory center in the medulla oblongata, which can directly modulate breathing rate to bring the pH back into the normal range.

Hypercapnia, or abnormally elevated blood levels of CO 2 , occurs in any situation that impairs respiratory functions, including pneumonia and congestive heart failure. Reduced breathing (hypoventilation) due to drugs such as morphine, barbiturates, or ethanol (or even just holding one’s breath) can also result in hypercapnia. Hypocapnia, or abnormally low blood levels of CO 2 , occurs with any cause of hyperventilation that drives off the CO 2 , such as salicylate toxicity, elevated room temperatures, fever, or hysteria.

Renal Regulation of Acid-Base Balance

The renal regulation of the body’s acid-base balance addresses the metabolic component of the buffering system. Whereas the respiratory system (together with breathing centers in the brain) controls the blood levels of carbonic acid by controlling the exhalation of CO 2 , the renal system controls the blood levels of bicarbonate. A decrease of blood bicarbonate can result from the inhibition of carbonic anhydrase by certain diuretics or from excessive bicarbonate loss due to diarrhea. Blood bicarbonate levels are also typically lower in people who have Addison’s disease (chronic adrenal insufficiency), in which aldosterone levels are reduced, and in people who have renal damage, such as chronic nephritis. Finally, low bicarbonate blood levels can result from elevated levels of ketones (common in unmanaged diabetes mellitus), which bind bicarbonate in the filtrate and prevent its conservation.

Bicarbonate ions, HCO 3 – , found in the filtrate, are essential to the bicarbonate buffer system, yet the cells of the tubule are not permeable to bicarbonate ions. The steps involved in supplying bicarbonate ions to the system are seen in Figure 26.4.3 and are summarized below:

• Step 1: Sodium ions are reabsorbed from the filtrate in exchange for H + by an antiport mechanism in the apical membranes of cells lining the renal tubule.
• Step 2: The cells produce bicarbonate ions that can be shunted to peritubular capillaries.
• Step 3: When CO 2 is available, the reaction is driven to the formation of carbonic acid, which dissociates to form a bicarbonate ion and a hydrogen ion.
• Step 4: The bicarbonate ion passes into the peritubular capillaries and returns to the blood. The hydrogen ion is secreted into the filtrate, where it can become part of new water molecules and be reabsorbed as such, or removed in the urine.

It is also possible that salts in the filtrate, such as sulfates, phosphates, or ammonia, will capture hydrogen ions. If this occurs, the hydrogen ions will not be available to combine with bicarbonate ions and produce CO 2 . In such cases, bicarbonate ions are not conserved from the filtrate to the blood, which will also contribute to a pH imbalance and acidosis.

The hydrogen ions also compete with potassium to exchange with sodium in the renal tubules. If more potassium is present than normal, potassium, rather than the hydrogen ions, will be exchanged, and increased potassium enters the filtrate. When this occurs, fewer hydrogen ions in the filtrate participate in the conversion of bicarbonate into CO 2 and less bicarbonate is conserved. If there is less potassium, more hydrogen ions enter the filtrate to be exchanged with sodium and more bicarbonate is conserved.

Chloride ions are important in neutralizing positive ion charges in the body. If chloride is lost, the body uses bicarbonate ions in place of the lost chloride ions. Thus, lost chloride results in an increased reabsorption of bicarbonate by the renal system.

Disorders of the… Fluid Balance: Acid-Base Balance: Ketoacidosis

Diabetic acidosis, or ketoacidosis, occurs most frequently in people with poorly controlled diabetes mellitus. When certain tissues in the body cannot get adequate amounts of glucose, they depend on the breakdown of fatty acids for energy. When acetyl groups break off the fatty acid chains, the acetyl groups then non-enzymatically combine to form ketone bodies, acetoacetic acid, beta-hydroxybutyric acid, and acetone, all of which increase the acidity of the blood. In this condition, the brain isn’t supplied with enough of its fuel—glucose—to produce all of the ATP it requires to function.

Ketoacidosis can be severe and, if not detected and treated properly, can lead to diabetic coma, which can be fatal. A common early symptom of ketoacidosis is deep, rapid breathing as the body attempts to drive off CO 2 and compensate for the acidosis. Another common symptom is fruity-smelling breath, due to the exhalation of acetone. Other symptoms include dry skin and mouth, a flushed face, nausea, vomiting, and stomach pain. Treatment for diabetic coma is ingestion or injection of sugar; its prevention is the proper daily administration of insulin.

A person who is diabetic and uses insulin can initiate ketoacidosis if a dose of insulin is missed. Among people with type 2 diabetes, those of Hispanic and African-American descent are more likely to go into ketoacidosis than those of other ethnic backgrounds, although the reason for this is unknown.

Chapter Review

A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range—between pH 7.35 and 7.45. A buffer is a substance that prevents a radical change in fluid pH by absorbing excess hydrogen or hydroxyl ions. Most commonly, the substance that absorbs the ion is either a weak acid, which takes up a hydroxyl ion (OH – ), or a weak base, which takes up a hydrogen ion (H + ). Several substances serve as buffers in the body, including cell and plasma proteins, hemoglobin, phosphates, bicarbonate ions, and carbonic acid. The bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory and renal systems also play major roles in acid-base homeostasis by removing CO 2 and hydrogen ions, respectively, from the body.

Review Questions

Critical thinking questions.

1. Describe the conservation of bicarbonate ions in the renal system.

2. Describe the control of blood carbonic acid levels through the respiratory system.

Answers for Critical Thinking Questions

• Bicarbonate ions are freely filtered through the glomerulus. They cannot pass freely into the renal tubular cells and must be converted into CO 2 in the filtrate, which can pass through the cell membrane. Sodium ions are reabsorbed at the membrane, and hydrogen ions are expelled into the filtrate. The hydrogen ions combine with bicarbonate, forming carbonic acid, which dissociates into CO 2 gas and water. The gas diffuses into the renal cells where carbonic anhydrase catalyzes its conversion back into a bicarbonate ion, which enters the blood.
• Carbonic acid blood levels are controlled through the respiratory system by the expulsion of CO 2 from the lungs. The formula for the production of bicarbonate ions is reversible if the concentration of CO 2 decreases. As this happens in the lungs, carbonic acid is converted into a gas, and the concentration of the acid decreases. The rate of respiration determines the amount of CO 2 exhaled. If the rate increases, less acid is in the blood; if the rate decreases, the blood can become more acidic.

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Images, from Anatomy & Physiology by OpenStax , are licensed under CC BY except where otherwise noted.

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General Chemistry/Buffer Systems

← Titration and pH · Reactions of Acids and Bases →

Introduction

Buffer systems are systems in which there is a significant (and nearly equivalent) amount of a weak acid and its conjugate base—or a weak base and its conjugate acid—present in solution. This coupling provides a resistance to change in the solution's pH. When strong acid is added, it is neutralized by the conjugate base. When strong base is added, it is neutralized by the weak acid. However, too much acid or base will exceed the buffer's capacity , resulting in significant pH changes.

Buffers are useful when a solution must maintain a specific pH. For example, blood is a buffer system because the life processes in a human only function within a specific pH range of 7.35 to 7.45. When, for example, lactic acid is released by the muscles during exercise, buffers within the blood neutralize it to maintain a healthy pH.

Making a Buffer

Once again, let's consider an arbitrary weak acid, HA, which is present in a solution. If we introduce a salt of the acid's conjugate base, say NaA (which will provide the A - ion), we now have a buffer solution. Ideally, the buffer would contain equal amounts of the weak acid and conjugate base.

Instead of adding NaA, what if a strong base were added, such as NaOH? In that case, the hydroxide ions would neutralize the weak acid and create water and A - ions. If the solution contained only A - ions, then a strong acid like HCl were added, they would neutralize and create HA.

As you can see, there are three ways to create a buffer:

All six of the combinations will create equal amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid.

Buffers and pH

To determine the pH of a buffer system, you must know the acid's dissociation constant . This value, Failed to parse (MathML with SVG or PNG fallback (recommended for modern browsers and accessibility tools): Invalid response ("Math extension cannot connect to Restbase.") from server "https://en.wikipedia.org/api/rest_v1/":): {\displaystyle K_a} (or Failed to parse (MathML with SVG or PNG fallback (recommended for modern browsers and accessibility tools): Invalid response ("Math extension cannot connect to Restbase.") from server "https://en.wikipedia.org/api/rest_v1/":): {\displaystyle K_b} for a base) determines the strength of an acid (or base). It is explored more thoroughly in the Equilibrium unit, but for now it suffices to say that this value is simply a measure of strength for acids and bases. The dissociation constants for acids and bases are determined experimentally.

The Henderson-Hasselbalch equation allows the calculation of a buffer's pH. It is:

For a buffer created from a base, the equation is:

Using these equations requires determining the ratio of base to acid in the solution.

• Book:General Chemistry
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Chapter 12. Acids and Bases

Learning Objectives

• Define buffer .
• Correctly identify the two components of a buffer.

As indicated in the section “Strong and Weak Acids and Bases and Their Salts” , weak acids are relatively common, even in the foods we eat. But we occasionally encounter a strong acid or base, such as stomach acid, which has a strongly acidic pH of 1.7. By definition, strong acids and bases can produce a relatively large amount of H + or OH − ions and consequently have marked chemical activities. In addition, very small amounts of strong acids and bases can change the pH of a solution very quickly. If 1 mL of stomach acid [approximated as 0.1 M HCl(aq)] were added to the bloodstream and no correcting mechanism were present, the pH of the blood would decrease from about 7.4 to about 4.7 — a pH that is not conducive to continued living. Fortunately, the body has a mechanism for minimizing such dramatic pH changes.

The mechanism involves a buffer , a solution that resists dramatic changes in pH. Buffers do so by being composed of certain pairs of solutes: either a weak acid plus a salt derived from that weak acid or a weak base plus a salt of that weak base. For example, a buffer can be composed of dissolved HC 2 H 3 O 2 (a weak acid) and NaC 2 H 3 O 2 (the salt derived from that weak acid). Another example of a buffer is a solution containing NH 3 (a weak base) and NH 4 Cl (a salt derived from that weak base).

Let us use an HC 2 H 3 O 2 /NaC 2 H 3 O 2 buffer to demonstrate how buffers work. If a strong base — a source of OH − (aq) ions — is added to the buffer solution, those OH − ions will react with the HC 2 H 3 O 2 in an acid-base reaction:

Rather than changing the pH dramatically by making the solution basic, the added OH − ions react to make H 2 O, so the pH does not change much.

If a strong acid — a source of H + ions — is added to the buffer solution, the H + ions will react with the anion from the salt. Because HC 2 H 3 O 2 is a weak acid, it is not ionized much. This means that if lots of H + ions and C 2 H 3 O 2 − ions are present in the same solution, they will come together to make HC 2 H 3 O 2 :

Rather than changing the pH dramatically and making the solution acidic, the added H + ions react to make molecules of a weak acid. Figure 12.2 “The Actions of Buffers” illustrates both actions of a buffer.

Buffers made from weak bases and salts of weak bases act similarly. For example, in a buffer containing NH 3 and NH 4 Cl, NH 3 molecules can react with any excess H + ions introduced by strong acids:

while the NH 4 + (aq) ion can react with any OH − ions introduced by strong bases:

Example 12.12

Which combinations of compounds can make a buffer solution?

• HCHO 2 and NaCHO 2
• HCl and NaCl
• CH 3 NH 2 and CH 3 NH 3 Cl
• NH 3 and NaOH
• HCHO 2 is formic acid, a weak acid, while NaCHO 2 is the salt made from the anion of the weak acid (the formate ion [CHO 2 − ]). The combination of these two solutes would make a buffer solution.
• HCl is a strong acid, not a weak acid, so the combination of these two solutes would not make a buffer solution.
• CH 3 NH 2 is methylamine, which is like NH 3 with one of its H atoms substituted with a CH 3 group. Because it is not listed in Table 12.1 “Strong Acids and Bases” , we can assume that it is a weak base. The compound CH 3 NH 3 Cl is a salt made from that weak base, so the combination of these two solutes would make a buffer solution.
• NH 3 is a weak base, but NaOH is a strong base. The combination of these two solutes would not make a buffer solution.

Test Yourself Which combinations of compounds can make a buffer solution?

• NaHCO 3 and NaCl
• H 3 PO 4 and NaH 2 PO 4
• NH 3 and (NH 4 ) 3 PO 4
• NaOH and NaCl

Buffers work well only for limited amounts of added strong acid or base. Once either solute is completely reacted, the solution is no longer a buffer, and rapid changes in pH may occur. We say that a buffer has a certain capacity . Buffers that have more solute dissolved in them to start with have larger capacities, as might be expected.

Human blood has a buffering system to minimize extreme changes in pH. One buffer in blood is based on the presence of HCO 3 − and H 2 CO 3 [the second compound is another way to write CO 2 (aq)]. With this buffer present, even if some stomach acid were to find its way directly into the bloodstream, the change in the pH of blood would be minimal. Inside many of the body’s cells, there is a buffering system based on phosphate ions.

Food and Drink App: The Acid That Eases Pain

Although medicines are not exactly “food and drink,” we do ingest them, so let’s take a look at an acid that is probably the most common medicine: acetylsalicylic acid, also known as aspirin. Aspirin is well known as a pain reliever and antipyretic (fever reducer).

The structure of aspirin is shown in the accompanying figure. The acid part is circled; it is the H atom in that part that can be donated as aspirin acts as a Brønsted-Lowry acid. Because it is not given in Table 12.1 “Strong Acids and Bases” , acetylsalicylic acid is a weak acid. However, it is still an acid, and given that some people consume relatively large amounts of aspirin daily, its acidic nature can cause problems in the stomach lining, despite the stomach’s defenses against its own stomach acid.

Because the acid properties of aspirin may be problematic, many aspirin brands offer a “buffered aspirin” form of the medicine. In these cases, the aspirin also contains a buffering agent — usually MgO — that regulates the acidity of the aspirin to minimize its acidic side effects.

As useful and common as aspirin is, it was formally marketed as a drug starting in 1899. The US Food and Drug Administration (FDA), the governmental agency charged with overseeing and approving drugs in the United States, wasn’t formed until 1906. Some have argued that if the FDA had been formed before aspirin was introduced, aspirin may never have gotten approval due to its potential for side effects — gastrointestinal bleeding, ringing in the ears, Reye’s syndrome (a liver problem), and some allergic reactions. However, recently aspirin has been touted for its effects in lessening heart attacks and strokes, so it is likely that aspirin is here to stay.

Key Takeaways

• A buffer is a solution that resists sudden changes in pH.
• Define buffer . What two related chemical components are required to make a buffer?
• Can a buffer be made by combining a strong acid with a strong base? Why or why not?
• HNO 2 and NaNO 2
• NH 4 NO 3 and HNO 3
• NH 4 NO 3 and NH 3
• H 3 PO 4 and Na 3 PO 4
• NaHCO 3 and Na 2 CO 3
• NaNO 3 and Ca(NO 3 ) 2
• HN 3 and NH 3
• For each combination in Exercise 3 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
• For each combination in Exercise 4 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
• The complete phosphate buffer system is based on four substances: H 3 PO 4 , H 2 PO 4 − , HPO 4 2− , and PO 4 3− . What different buffer solutions can be made from these substances?
• Explain why NaBr cannot be a component in either an acidic or a basic buffer.
• Two solutions are made containing the same concentrations of solutes. One solution is composed of H 3 PO 4 and Na 3 PO 4 , while the other is composed of HCN and NaCN. Which solution should have the larger capacity as a buffer?
• Two solutions are made containing the same concentrations of solutes. One solution is composed of NH 3 and NH 4 NO 3 , while the other is composed of H 2 SO 4 and Na 2 SO 4 . Which solution should have the larger capacity as a buffer?
• A buffer is the combination of a weak acid or base and a salt of that weak acid or base.
• strong acid: NO 2 − + H + → HNO 2 ; strong base: HNO 2 + OH − → NO 2 − + H 2 O
• strong base: NH 4 + + OH − → NH 3 + H 2 O; strong acid: NH 3 + H + → NH 4 +
• Buffers can be made from three combinations: (1) H 3 PO 4 and H 2 PO 4 − ; (2) H 2 PO 4 − and HPO 4 2− ; and (3) HPO 4 2− and PO 4 3− . (Technically, a buffer can be made from any two components.)
• The phosphate buffer should have the larger capacity.

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LEARNING OBJECTIVES

• Define  buffer .
• Correctly identify the two components of a buffer.

Weak acids are relatively common, even in the foods we eat. But we occasionally encounter a strong acid or base, such as stomach acid, which has a strongly acidic pH of 1.7. By definition, strong acids and bases can produce a relatively large amount of [latex]\text{H}^+[/latex] or [latex]\text{OH}^-[/latex] ions and consequently have marked chemical activities. In addition, very small amounts of strong acids and bases can change the pH of a solution very quickly. If 1 mL of stomach acid [approximated as 0.1 M [latex]\text{HCl} (aq)[/latex]] were added to the bloodstream and no correcting mechanism were present, the pH of the blood would decrease from about 7.4 to about 4.7—a pH that is not conducive to continued living. Fortunately, the body has a mechanism for minimizing such dramatic pH changes.

The mechanism involves a buffer , a solution that resists dramatic changes in pH. Buffers do so by being composed of certain pairs of solutes: either a weak acid plus a salt derived from that weak acid or a weak base plus a salt of that weak base. For example, a buffer can be composed of dissolved [latex]\text{HC}_2\text{H}_3\text{O}_2[/latex] (a weak acid) and [latex]\text{NaC}_2\text{H}_3\text{O}_2[/latex] (the salt derived from that weak acid). Another example of a buffer is a solution containing [latex]\text{NH}_3[/latex] (a weak base) and [latex]\text{NH}_4\text{Cl}[/latex] (a salt derived from that weak base).

Let us use an [latex]\text{HC}_2\text{H}_3\text{O}_2/\text{NaC}_2\text{H}_3\text{O}_2[/latex] buffer to demonstrate how buffers work. If a strong base—a source of [latex]\text{OH}^- (aq)[/latex] ions—is added to the buffer solution, those [latex]\text{OH}^- (aq)[/latex] ions will react with the [latex]\text{HC}_2\text{H}_3\text{O}_2[/latex] in an acid-base reaction: [latex]\text{HC}_2\text{H}_3\text{O}_2 + \text{OH}^- (aq) \rightarrow \text{H}_2\text{O} (\ell) + \text{C}_2\text{H}_3\text{O}_2^- (aq)[/latex] Rather than changing the pH dramatically by making the solution basic, the added [latex]\text{OH}^-[/latex] ions react to make [latex]\text{H}_2\text{O}[/latex], so the pH does not change much. If a strong acid—a source of [latex]\text{H}^+[/latex] ions—is added to the buffer solution, the [latex]\text{H}^+[/latex] ions will react with the anion from the salt. Because [latex]\text{HC}_2\text{H}_3\text{O}_2[/latex] is a weak acid, it is not ionized much. This means that if lots of [latex]\text{H}^+[/latex] ions and [latex]\text{C}_2\text{H}_3\text{O}_2 ^-[/latex] ions are present in the same solution, they will come together to make [latex]\text{HC}_2\text{H}_3\text{O}_2[/latex]: [latex]\text{H}^+ (aq) + \text{C}_2\text{H}_3\text{O}_2 ^- (aq) \rightarrow \text{HC}_2\text{H}_3\text{O}_2 (aq)[/latex] Rather than changing the pH dramatically and making the solution acidic, the added [latex]\text{H}^+[/latex] ions react to make molecules of a weak acid. Figure 12.2 "The Actions of Buffers" illustrates both actions of a buffer.

Figure 12.2 The Actions of Buffers

Buffers made from weak bases and salts of weak bases act similarly. For example, in a buffer containing [latex]\text{NH}_3[/latex] and [latex]\text{NH}_4\text{Cl}[/latex], [latex]\text{NH}_3[/latex] molecules can react with any excess [latex]\text{H}^+[/latex] ions introduced by strong acids: [latex]\text{NH}_3 (aq) + \text{H}^+ (aq) \rightarrow \text{NH}_4 ^+ (aq)[/latex] while the [latex]\text{NH}_4 ^+ (aq)[/latex] ion can react with any [latex]\text{OH}^-[/latex] ions introduced by strong bases: [latex]\text{NH}_4^+ (aq) + \text{OH}^- (aq) \rightarrow \text{NH}_3 (aq) + \text{H}_2\text{O} (\ell)[/latex]

Which combinations of compounds can make a buffer solution?

• [latex]\text{HCHO}_2[/latex] and [latex]\text{NaCHO}_2[/latex]
• [latex]\text{HCl}[/latex] and [latex]\text{NaCl}[/latex]
• [latex]\text{CH}_3\text{NH}_2[/latex] and [latex]\text{CH}_3\text{NH}_3\text{Cl}[/latex]
• [latex]\text{NH}_3[/latex] and [latex]\text{NaOH}[/latex]
• [latex]\text{HCHO}_2[/latex] is formic acid, a weak acid, while [latex]\text{NaCHO}_2[/latex] is the salt made from the anion of the weak acid (the formate ion [\text{CHO}_2 ^-]). The combination of these two solutes would make a buffer solution.
• [latex]\text{HCl}[/latex] is a strong acid, not a weak acid, so the combination of these two solutes would not make a buffer solution.
• [latex]\text{CH}_3\text{NH}_2[/latex] is methylamine, which is like [latex]\text{NH}_3[/latex] with one of its H atoms substituted with a [latex]\text{CH}_3[/latex] group. Because it is not listed in Table 12.2 "Strong Acids and Bases", we can assume that it is a weak base. The compound [latex]\text{CH}_3\text{NH}_3\text{Cl}[/latex] is a salt made from that weak base, so the combination of these two solutes would make a buffer solution.
• [latex]\text{NH}_3[/latex] is a weak base, but [latex]\text{NaOH}[/latex] is a strong base. The combination of these two solutes would not make a buffer solution.

Test Yourself Which combinations of compounds can make a buffer solution?

• [latex]\text{NaHCO}_3[/latex] and [latex]\text{NaCl}[/latex]
• [latex]\text{H}_3\text{PO}_4[/latex] and [latex]\text{NaH}_2\text{PO}_4[/latex]
• [latex]\text{NH}_3[/latex] and [latex]\text{(NH}_4)_3\text{PO}_4[/latex]
• [latex]\text{NaOH}[/latex] and [latex]\text{NaCl}[/latex]

Buffers work well only for limited amounts of added strong acid or base. Once either solute is completely reacted, the solution is no longer a buffer, and rapid changes in pH may occur. We say that a buffer has a certain capacity . Buffers that have more solute dissolved in them to start with have larger capacities, as might be expected.

Human blood has a buffering system to minimize extreme changes in pH. One buffer in blood is based on the presence of [latex]\text{HCO}_3 ^-[/latex] and [latex]\text{H}_2\text{CO}_3[/latex] [the second compound is another way to write [latex]\text{CO}_2 (aq)[/latex]]. With this buffer present, even if some stomach acid were to find its way directly into the bloodstream, the change in the pH of blood would be minimal. Inside many of the body’s cells, there is a buffering system based on phosphate ions.

Food and Drink App: The Acid That Eases Pain

Although medicines are not exactly “food and drink,” we do ingest them, so let’s take a look at an acid that is probably the most common medicine: acetylsalicylic acid, also known as aspirin. Aspirin is well known as a pain reliever and antipyretic (fever reducer).

The structure of aspirin is shown in the accompanying figure. The acid part is circled; it is the H atom in that part that can be donated as aspirin acts as a Brønsted-Lowry acid. Because it is not given in Table 12.2 "Strong Acids and Bases", acetylsalicylic acid is a weak acid. However, it is still an acid, and given that some people consume relatively large amounts of aspirin daily, its acidic nature can cause problems in the stomach lining, despite the stomach’s defenses against its own stomach acid.

Figure 12.3 The Molecular Structure of Aspirin

Because the acid properties of aspirin may be problematic, many aspirin brands offer a “buffered aspirin” form of the medicine. In these cases, the aspirin also contains a buffering agent—usually [latex]\text{MgO}[/latex]—that regulates the acidity of the aspirin to minimize its acidic side effects.

As useful and common as aspirin is, it was formally marketed as a drug starting in 1899. The US Food and Drug Administration (FDA), the governmental agency charged with overseeing and approving drugs in the United States, wasn’t formed until 1906. Some have argued that if the FDA had been formed before aspirin was introduced, aspirin may never have gotten approval due to its potential for side effects—gastrointestinal bleeding, ringing in the ears, Reye’s syndrome (a liver problem), and some allergic reactions. However, recently aspirin has been touted for its effects in lessening heart attacks and strokes, so it is likely that aspirin is here to stay.

KEY TAKEAWAYS

• A buffer is a solution that resists sudden changes in pH.
• Define  buffer . What two related chemical components are required to make a buffer?
• Can a buffer be made by combining a strong acid with a strong base? Why or why not?
• [latex]\text{HNO}_2[/latex] and [latex]\text{NaNO}_2[/latex]
• [latex]\text{NH}_4\text{NO}_3[/latex] and [latex]\text{HNO}_3[/latex]
• [latex]\text{NH}_4\text{NO}_3[/latex] and [latex]\text{NH}_3[/latex]
• [latex]\text{H}_3\text{PO}_4[/latex] and [latex]\text{Na}_3\text{PO}_4[/latex]
• [latex]\text{NaHCO}_3[/latex] and [latex]\text{Na}_2\text{CO}_3[/latex]
• [latex]\text{NaNO}_3[/latex] and [latex]\text{Ca(NO}_3)_2[/latex]
• [latex]\text{HN}_3[/latex] HN 3 and [latex]\text{NH}_3[/latex]
• For each combination in Exercise 3 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
• For each combination in Exercise 4 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
• The complete phosphate buffer system is based on four substances: [latex]\text{H}_3\text{PO}_4[/latex], [latex]\text{H}_2\text{PO}_4 ^-[/latex], [latex]\text{HPO}_4 ^{2-}[/latex], and [latex]\text{PO}_4 ^{3-}[/latex]. What different buffer solutions can be made from these substances?
• Explain why [latex]\text{NaBr}[/latex] cannot be a component in either an acidic or a basic buffer.
• Two solutions are made containing the same concentrations of solutes. One solution is composed of [latex]\text{H}_3\text{PO}_4[/latex] and [latex]\text{Na}_3\text{PO}_4[/latex], while the other is composed of [latex]\text{HCN}[/latex] and [latex]\text{NaCN}[/latex]. Which solution should have the larger capacity as a buffer?
• Two solutions are made containing the same concentrations of solutes. One solution is composed of [latex]\text{NH}_3[/latex] and [latex]\text{NH}_4\text{NO}_3[/latex], while the other is composed of [latex]\text{H}_2\text{SO}_4[/latex] and [latex]\text{Na}_2\text{SO}_4[/latex]. Which solution should have the larger capacity as a buffer?

5.   3b: strong acid: [latex]\text{NO}_2^- + \text{H}^+  \rightarrow \text{HNO}_2[/latex]; strong base: [latex]\text{HNO}_2 +  \text{OH}^- \rightarrow \text{NO}_2^- + \text{H}_2\text{O}[/latex]; 3d: strong base: [latex]\text{NH}_4^+ +  \text{OH}^- \rightarrow \text{NH}_3 + \text{H}_2\text{O}[/latex]; strong acid: [latex]\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+[/latex]

7. Buffers can be made from three combinations: (1) [latex]\text{H}_3\text{PO}_4[/latex] and [latex]\text{H}_2\text{PO}_4 ^-[/latex], (2) [latex]\text{H}_2\text{PO}_4 ^-[/latex] and [latex]\text{HPO}_4 ^{2-}[/latex], and (3) [latex]\text{HPO}_4 ^{2-}[/latex] and [latex]\text{PO}_4 ^{3-}[/latex]. (Technically, a buffer can be made from any two components.)

9. The phosphate buffer should have the larger capacity.

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Buffer Systems

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• Buffer System
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Buffer Definition and Examples in Chemistry

A buffer is a solution that maintains the stability of a system’s pH level when adding small quantities of acids or bases . This characteristic makes buffers important in biological and chemical applications where pH stability is crucial.

Composition of Buffer Solutions

A buffer solution typically consists of a weak acid and its conjugate base , or a weak base and its conjugate acid. These components work in tandem to neutralize any added acid or base.

Examples of Buffers

For example, the following acid and base pairs work together and form buffer solutions. A salt supplies the conjugate acid or base as it dissolves in solution .

• The weak acid acetic acid (CH 3 COOH) and a salt containing the conjugate base (the acetate ion CH 3 COO – ), such as sodium acetate (CH 3 COONa)
• The weak base ammonia (NH 3 ) and a salt containing its conjugate acid (the ammonium cation NH 4 + ), such as ammonium hydroxide (NH 4 OH)

How a Buffer Works

Buffers function through a process of chemical equilibrium. When you add an acid to a buffer, the conjugate base present in the buffer neutralizes it. Conversely, when you add a base, the weak acid in the buffer neutralizes the base.

For example, consider a buffer made of acetic acid (CH₃COOH) and sodium acetate (CH₃COONa). When hydrochloric acid (HCl) is added, the acetate ion (CH₃COO⁻) reacts with the H⁺ ions from HCl, forming more acetic acid and mitigating the pH change. Similarly, adding a base like sodium hydroxide (NaOH) results in the acetic acid reacting with OH⁻ to produce acetate and water, again stabilizing the pH.

There are limits to how much acid or base a buffer solution handles before it can no longer maintain the pH. Also, selecting the right buffer system depends on the desired pH.

Selecting the Best Buffer for a Desired pH

The choice of an appropriate buffer depends on the desired pH and the buffer’s pKa , the dissociation constant of the acid (or conjugate acid). A buffer is most effective when the pH is close to the pKa of the acid in the buffer system.

For example, the K a of hydrofluoric acid is 6.6 x 10 -4 . Therefore its pKa value is -log(6.6 x 10 -4 ) = 3.18. This means a hydrofluoric acid buffer works best around a pH of 3.18.

Meanwhile, the K b value of the weak base ammonia (NH 3 ) is 1.8 x 10 -5 . This means the K a for its conjugate acid (NH 4 + ) is K w /K b = 10 -14 / 1.8×10 -5 = 5.6 x 10 -10 . The pKa of NH 4 + is -log(5.6×10 -5 ) = 9.25. The optimal pH for the NH 4 + /NH 3 buffer system is around a pH of 9.25.

Choose a weak acid and its salt for pH values that are lower than 7. Select a weak base and its salt for pH values that are above 7.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation estimates the pH of the solution when the amounts of conjugate acid and conjugate base are approximately equal (within a factor of 10). This equation relates pH, pKa, and the ratio of the concentration of the conjugate base to that of the weak acid:

pH = pKa + log([Base]/[Acid]​)

Rearrange the equation to calculate the required concentrations of the buffer components to achieve the desired pH.

Checking pH After Adding Strong Acids or Bases

After adding a strong acid or base to a buffer, the pH often shifts. The Henderson-Hasselbalch equation estimate the new pH, considering the change in the concentration of the acid/base pair. However, check the pH using a pH meter or indicators.

Universal Buffers

Universal buffers maintain a stable pH over a wide range. These are mixtures of several different buffers. A classic example is the Britton-Robinson buffer, which is a mixture of phosphoric acid, boric acid, and acetic acid. The Britton Robinson buffer maintains the pH over the range from 2 to 12. Another example of a universal buffer is citric acid, which has three pKa values. Universal buffers are especially useful in applications requiring a stable pH across a broad spectrum.

Biological Buffers

Biological buffers play a critical role in maintaining the pH balance in living organisms, ensuring the proper functioning of biological processes. Here are some notable examples:

• Bicarbonate Buffer System : One of the most important buffering systems in human blood, the bicarbonate buffer consists of carbonic acid (H₂CO₃) and its conjugate base, bicarbonate (HCO₃⁻). This buffer helps maintain the blood pH around 7.4, with carbonic acid acting as the weak acid and bicarbonate as the weak base.
• Phosphate Buffer System : Widely present in biological fluids, this buffer consists of dihydrogen phosphate (H₂PO₄⁻) and hydrogen phosphate (HPO₄²⁻). It plays a significant role in buffering pH changes in cells and tissues, and is particularly effective in the pH range of 6.8 to 7.4.
• Proteins as Buffers : Many proteins, including hemoglobin in red blood cells, act as effective buffers. They contain amino acids with acidic and basic side chains that donate or accept protons, helping to buffer changes in pH. Hemoglobin, for instance, buffers the blood by binding to hydrogen ions and carbon dioxide.
• Amino Acids : Free amino acids in cells also contribute to buffering. The amino group (–NH₂) accepts a proton, while the carboxyl group (–COOH) donates a proton, making amino acids capable of buffering in different pH ranges.
• Citrate Buffer : Citrate, a key intermediate in the citric acid cycle, also functions as a buffer in metabolic pathways. It buffers pH changes in the pH range of 3.0 to 6.2.
• Tris Buffer : Tris (tris(hydroxymethyl)aminomethane) is a common buffer in biochemical and molecular biology labs. It maintains a stable pH in various types of solutions, including cell culture media and buffer solutions for DNA/RNA extraction and PCR.
• HEPES Buffer : Widely used in cell culture, HEPES (4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid) is a zwitterionic buffer that is effective in the pH range of 6.8 to 8.2. HEPES is popular for its minimal interference with biological processes.
• Butler, J. N. (1998). Ionic Equilibrium: Solubility and pH Calculations . Wiley. ISBN 978-0-471-58526-8.
• Carmody, Walter R. (1961). “Easily prepared wide range buffer series”. J. Chem. Educ . 38 (11): 559–560. doi: 10.1021/ed038p559
• Scorpio, R. (2000). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems . ISBN 978-0-7872-7374-3.
• Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2014). Fundamentals of Analytical Chemistry (9th ed.). Brooks/Cole. ISBN 978-0-495-55828-6.
• Urbansky, Edward T.; Schock, Michael R. (2000). “Understanding, Deriving and Computing Buffer Capacity”. Journal of Chemical Education . 77 (12): 1640–1644. doi: 10.1021/ed077p1640

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Mechanism of Buffer Action

Mechanism of buffer action

Buffer Action: The property of the solution to resist the changes in its pH value on the addition of small amounts of strong acid or base is known as buffer action”.

A buffer is a solution containing a weak acid and its conjugate base in similar amounts. The way in which a buffer solution acts may be illustrated by considering a buffer solution made up of CH 3 COOH and CH 3 COONa. A buffer solution may be prepared by mixing comparable molar amounts of ethanoic acid and sodium ethanoate in water. In the solution we have the following situation:

CH 3 COOH (aq) ↔ CH 3 COO – (aq) + H + (aq)

CH 3 COONa (aq) → CH 3 COO – (aq) + Na + (aq)

The mechanism of buffer action can be understood by considering an acidic buffer made of a weak acid like Acetic acid and its sodium salt Sodium acetate. Sodium acetate, however, exists as acetate ions and Na+ ions.

The salt exists completely as ions.

Thus the buffer contains both acid (CH 3 COOH) and its conjugate base (CH 3 COO – ). When a small quantity of an acid is added the hydrogen ions are removed by the conjugate base (CH 3 COO – ) as follows:

H + (aq) + CH 3 COO – (aq) ↔ CH 3 COOH (aq)

Since ethanoic acid is only slightly dissociated and in the form CH 3 COOH, it does not contribute any H+ ion, the pH of the resulting solution will remain practically constant. Due to the removal of most of the added H+, there is no appreciable decrease in pH. The reaction nearly goes to completion because CH 3 COOH is a weak acid and its ions have a strong tendency to form non-ionized CH 3 COOH molecules. If, however, a strong base is added, the added OH – ion is neutralized by the reaction with the acid in the buffer,

CH 3 COOH (aq) + OH – (aq) → CH 3 COO – (aq) + H 2 O (l)

One may also consider that the added OH- ion reacts with the H + ion to produce water. The added OH- ions are removed and the acid equilibrium shifts to the right to replace the H+ ions used up. Thus, there is a very slight change in the pH value. Now, if you add a drop of NaOH, the OH –  ions react with the free acid to give undissociated water molecules.

Since additional OH – ions of the base are consumed or neutralized, the pH of the solution remains unchanged. This resistance to change in pH on addition of base is called as reserve acidity and is due to CH 3 COOH. The effect is the neutralization of most of the added OH – ions by CH 3 COOH. Therefore, there is no appreciable change in pH. When a strong base is added, the acid present in the buffer neutralizes the hydroxide ions (OH-)

In either case, the effect of the added acid or base is effectively balanced so that the pH of the solution remains approximately constant. Similarly, the mechanism of basic buffers can be explained.

So, the weak acid and weak base remain in the solution with high concentrations since they only rarely react with the water. However, they are very likely to react with any added strong base or strong acid.

Characteristics of Buffer: A buffer resists changes in pH by having a relatively high concentration of acid or base, or, most often, both an acid and a base.

• It has a definite pH value.
• Its pH value doesn’t change on keeping for a long time.
• Its pH value doesn’t change on dilution.
• Its pH value doesn’t change even with the addition of a small amount of a strong acid or a base.

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Buffer Action

Buffer action, in general, is defined as the ability of the buffer solution to resist the changes in pH value when a small amount of an acid or a base is added to it.

Mechanism of Buffering Action

To understand the mechanism of buffer action, we can take the example of an acidic buffer that is made up of a weak acid like acetic acid and its sodium salt, sodium acetate. In this acidic buffer, the solution contains equimolar amounts of acetic acid and sodium acetate. Usually, a large number of sodium ions (Na + ), acetate ions (CH 3 COO – ) and undissociated acetic acid molecules are present.

The salt exists completely as ions.

Here, the buffer will consist of both acid (CH 3 COOH) and its conjugate base (CH 3 COO – ). If we add a small quantity of acid, the hydrogen ions will be removed by the conjugate base (CH 3 COO – ). It is represented as follows:

H+ (aq) + CH3COO– (aq) ↔ CH3COOH (aq)

Here, the ethanoic acid will only be slightly dissociated in the form of CH 3 COOH, which means that it will not contribute any H+ ions. Therefore, the pH of the resulting solution will remain more or less constant. The added H+ ions are also removed, due to which there is no appreciable decrease in pH.

Also Read: Buffer Solutions

The reaction comes to a completion as CH 3 COOH is a weak acid whose ions have a strong tendency to form non-ionised CH 3 COOH molecules. On the other hand, if we add a strong base, the OH – ion gets neutralised by the reaction with the acid in the buffer,

CH 3 COOH (aq) + OH– (aq) → CH 3 COO– (aq) + H 2 O (l).

We can also consider that the OH – ion can react with the H+ ion in order to form water. The OH – ions that are added are removed, wherein the acid equilibrium shifts to the right to replace the H+ ions that are exhausted. This results in a minor change in the pH value.

Alternatively, if we add a drop of NaOH, the OH – ions react with the free acid to give undissociated water molecules. The extra OH – ions of the base are neutralised. As a result, the pH of the solution remains the same. This condition, or the resistance offered by the pH when a base is added, is known as reserve acidity. This is mainly due to CH 3 COOH.

If we add a strong base, the acid present in the buffer neutralises the hydroxide ions (OH – ).

So, we can say that when acid or base is added, its effect is practically balanced, and the pH of the solution is always constant.

Key Points to Remember

If we take a solution, the salt will be completely ionised, and the weak acid will be partly ionised.

• CH 3 COONa ⇌ Na+ + CH 3 COO–
• CH 3 COOH ⇌ H+ + CH 3 COO–

On Addition of Acid and Base

1. When acid is added, the protons of acid that are released will be removed by the acetate ions to form an acetic acid molecule.

H + + CH 3 COO – (from added acid) ⇌ CH 3 COOH (from buffer solution)

2. When a base is added, the hydroxide that is released by the base will be removed by the hydrogen ions to form water.

HO– + H+ (from added base) ⇌ H 2 O (from buffer solution)

Buffer Action – Video Lesson

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7.4: Buffers

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  Learning Objectives

• To define buffer and describe how it reacts with an acid or a base.

Weak acids are relatively common, even in the foods we eat. But we occasionally come across a strong acid or base, such as stomach acid, that has a strongly acidic pH of 1–2. By definition, strong acids and bases can produce a relatively large amount of hydrogen or hydroxide ions and, as a consequence, have a marked chemical activity. In addition, very small amounts of strong acids and bases can change the pH of a solution very quickly. If 1 mL of stomach acid [which we will approximate as 0.05 M HCl(aq)] is added to the bloodstream, and if no correcting mechanism is present, the pH of the blood would go from about 7.4 to about 4.9—a pH that is not conducive to continued living. Fortunately, the body has a mechanism for minimizing such dramatic pH changes.

The mechanism involves a buffer , a solution that resists dramatic changes in pH. A buffer (or buffered ) solution is one that resists a change in its pH when H + or OH – ions are added or removed owing to some other reaction taking place in the same solution. Buffers do so by being composed of certain pairs of solutes: either a weak acid plus its conjugate base or a weak base plus its conjugate acid .

For example, a buffer can be composed of dissolved acetic acid (HC 2 H 3 O 2 , a weak acid) and sodium acetate ( NaC 2 H 3 O 2 ) . Sodium acetate is a salt that dissociates into sodium ions and acetate ions in solution. For as long as acetic acid and acetate ions are present in significant amounts a solution, this can resist dramatic pH changes. Another example of a buffer is a solution containing ammonia (NH 3 , a weak base) and ammonium chloride ( NH 4 Cl). Ammonium acetate is also a salt that dissociates into ammonium ions and chloride ions in solution. The presence of ammonium ions with ammonia molecules satisfies the requisite condition for a buffer solution.

How Buffers Work

The essential component of a buffer system is a conjugate acid-base pair whose concentration is fairly high in relation to the concentrations of added H + or OH – it is expected to buffer against. Let us use an acetic acid–sodium acetate buffer to demonstrate how buffers work. If a strong base—a source of OH − (aq) ions—is added to the buffer solution, those hydroxide ions will react with the acetic acid in an acid-base reaction:

\[HC_2H_3O_{2(aq)} + OH^−_{(aq)} \rightarrow H_2O_{(ℓ)} + C_2H_3O^−_{2(aq)} \label{Eq1} \]

Rather than changing the pH dramatically by making the solution basic, the added hydroxide ions react to make water, and the pH does not change much.

Many people are aware of the concept of buffers from buffered aspirin , which is aspirin that also has magnesium carbonate, calcium carbonate, magnesium oxide, or some other salt. The salt acts like a base, while aspirin is itself a weak acid.

If a strong acid—a source of H + ions—is added to the buffer solution, the H + ions will react with the anion from the salt. Because HC 2 H 3 O 2 is a weak acid, it is not ionized much. This means that if lots of hydrogen ions and acetate ions (from sodium acetate) are present in the same solution, they will come together to make acetic acid:

\[H^+_{(aq)} + C_2H_3O^−_{2(aq)} \rightarrow HC_2H_3O_{2(aq)} \label{Eq2} \]

Rather than changing the pH dramatically and making the solution acidic, the added hydrogen ions react to make molecules of a weak acid. Figure \(\PageIndex{1}\) illustrates both actions of a buffer.

A simple buffer system might be a 0.2 M solution of sodium acetate; the conjugate pair here is acetic acid HAc and its conjugate base, the acetate ion Ac – . The idea is that this conjugate pair "pool" will be available to gobble up any small (≤ 10 –3 M ) addition of H+ or OH – that may result from other processes going on in the solution.

Buffers work well only for limited amounts of added strong acid or base. Once either solute is all reacted, the solution is no longer a buffer, and rapid changes in pH may occur. We say that a buffer has a certain capacity. Buffers that have more solute dissolved in them to start with have larger capacities, as might be expected.

Buffers made from weak bases and salts of weak bases act similarly. For example, in a buffer containing NH 3 and NH 4 Cl, ammonia molecules can react with any excess hydrogen ions introduced by strong acids:

\[NH_{3(aq)} + H^+_{(aq)} \rightarrow NH^+_{4(aq)} \label{Eq3} \]

while the ammonium ion [NH 4 + (aq)] can react with any hydroxide ions introduced by strong bases:

\[NH^+_{4(aq)} + OH^−_{(aq)} \rightarrow NH_{3(aq)} + H_2O_{(ℓ)} \label{Eq4} \]

Example \(\PageIndex{1}\)

Which solute combinations can make a buffer solution? Assume all are aqueous solutions.

• HCHO 2 and NaCHO 2
• HCl and NaCl
• CH 3 NH 2 and CH 3 NH 3 Cl
• NH 3 and NaOH
• Formic acid (HCHO 2 ) is a weak acid, while NaCHO 2 is a salt supplying —formate ion (CHO 2 − ), the conjugate base of HCHO 2 . The combination of these two solutes would make a buffer solution.
• Hydrochloric acid (HCl) is a strong acid, not a weak acid, so the combination of these two solutes would not make a buffer solution.
• Methylamine (CH 3 NH 2 ) is like ammonia, a weak base. The compound CH 3 NH 3 Cl is a salt supplying CH 3 NH 3 + , the conjugate acid of CH 3 NH 2 . The combination of these two solutes would make a buffer solution.
• Ammonia (NH 3 ) is a weak base, but NaOH is a strong base. The combination of these two solutes would not make a buffer solution.

Exercise \(\PageIndex{1}\)

• NaHCO 3 and NaCl
• H 3 PO 4 and NaH 2 PO 4
• NH 3 and (NH 4 ) 3 PO 4
• NaOH and NaCl

b. H 3 PO 4 (weak acid) and H 2 PO 4 - (conjugate base of H 3 PO 4 ) make a buffer.

c. NH 3 (weak base) and NH 4 + (conjugate acid of NH 3 ) make a buffer

Food and Drink App: The Acid That Eases Pain

Although medicines are not exactly "food and drink," we do ingest them, so let's take a look at an acid that is probably the most common medicine: acetylsalicylic acid, also known as aspirin. Aspirin is well known as a pain reliever and antipyretic (fever reducer).

The structure of aspirin is shown in the accompanying figure. The acid part is circled; it is the H atom in that part that can be donated as aspirin acts as a Brønsted-Lowry acid. Because it is not given in Table 10.5.1 , acetylsalicylic acid is a weak acid. However, it is still an acid, and given that some people consume relatively large amounts of aspirin daily, its acidic nature can cause problems in the stomach lining, despite the stomach's defenses against its own stomach acid.

Because the acid properties of aspirin may be problematic, many aspirin brands offer a "buffered aspirin" form of the medicine. In these cases, the aspirin also contains a buffering agent-usually MgO-that regulates the acidity of the aspirin to minimize its acidic side effects.

As useful and common as aspirin is, it was formally marketed as a drug starting in 1899. The US Food and Drug Administration (FDA), the governmental agency charged with overseeing and approving drugs in the United States, wasn't formed until 1906. Some have argued that if the FDA had been formed before aspirin was introduced, aspirin may never have gotten approval due to its potential for side effects-gastrointestinal bleeding, ringing in the ears, Reye's syndrome (a liver problem), and some allergic reactions. However, recently aspirin has been touted for its effects in lessening heart attacks and strokes, so it is likely that aspirin is here to stay.

Buffer solutions are essential components of all living organisms.

• Our blood is buffered to maintain a pH of 7.4 that must remain unchanged as metabolically-generated CO 2 (carbonic acid) is added and then removed by our lungs.
• Buffers in the oceans, in natural waters such as lakes and streams, and within soils help maintain their environmental stability against acid rain and increases in atmospheric CO 2 .
• Many industrial processes, such as brewing, require buffer control, as do research studies in biochemistry and physiology that involve enzymes, are active only within certain pH ranges.

The pH in living systems (Figure \(\PageIndex{1}) is maintained by buffer systems.

Human blood has a buffering system to minimize extreme changes in pH. One buffer in blood is based on the presence of HCO 3 − and H 2 CO 3 [H 2 CO 3 is another way to write CO 2 (aq)]. With this buffer present, even if some stomach acid were to find its way directly into the bloodstream, the change in the pH of blood would be minimal. Inside many of the body’s cells, there is a buffering system based on phosphate ions.

Medicine: The Buffer System in Blood

The normal pH of human blood is about 7.4. The carbonate buffer system in the blood uses the following equilibrium reaction:

\[\ce{CO2}(g)+\ce{2H2O}(l)⇌\ce{H2CO3}(aq)⇌\ce{HCO3-}(aq)+\ce{H3O+}(aq) \nonumber \]

The concentration of carbonic acid, H 2 CO 3 is approximately 0.0012 M , and the concentration of the hydrogen carbonate ion, \(\ce{HCO3-}\), is around 0.024 M . Using the Henderson-Hasselbalch equation and the p K a of carbonic acid at body temperature, we can calculate the pH of blood:

\[\mathrm{pH=p\mathit{K}_a+\log\dfrac{[base]}{[acid]}=6.1+\log\dfrac{0.024}{0.0012}=7.4} \nonumber \]

The fact that the H 2 CO 3 concentration is significantly lower than that of the \(\ce{HCO3-}\) ion may seem unusual, but this imbalance is due to the fact that most of the by-products of our metabolism that enter our bloodstream are acidic. Therefore, there must be a larger proportion of base than acid, so that the capacity of the buffer will not be exceeded.

Lactic acid is produced in our muscles when we exercise. As the lactic acid enters the bloodstream, it is neutralized by the \(\ce{HCO3-}\) ion, producing H 2 CO 3 . An enzyme then accelerates the breakdown of the excess carbonic acid to carbon dioxide and water, which can be eliminated by breathing. In fact, in addition to the regulating effects of the carbonate buffering system on the pH of blood, the body uses breathing to regulate blood pH. If the pH of the blood decreases too far, an increase in breathing removes CO 2 from the blood through the lungs driving the equilibrium reaction such that [H 3 O + ] is lowered. If the blood is too alkaline, a lower breath rate increases CO 2 concentration in the blood, driving the equilibrium reaction the other way, increasing [H + ] and restoring an appropriate pH.

Career Focus: Blood Bank Technology Specialist

At this point in this text, you should have the idea that the chemistry of blood is fairly complex. Because of this, people who work with blood must be specially trained to work with it properly.

A blood bank technology specialist is trained to perform routine and special tests on blood samples from blood banks or transfusion centers. This specialist measures the pH of blood, types it (according to the blood’s ABO+/− type, Rh factors, and other typing schemes), tests it for the presence or absence of various diseases, and uses the blood to determine if a patient has any of several medical problems, such as anemia. A blood bank technology specialist may also interview and prepare donors to give blood and may actually collect the blood donation.

Blood bank technology specialists are well trained. Typically, they require a college degree with at least a year of special training in blood biology and chemistry. In the United States, training must conform to standards established by the American Association of Blood Banks.

Key Takeaway

• A buffer is a solution that resists sudden changes in pH.

Concept Review Exercise

• Explain how a buffer prevents large changes in pH.
• A buffer has components that react with both strong acids and strong bases to resist sudden changes in pH.
• Describe a buffer. What two related chemical components are required to make a buffer?
• Can a buffer be made by combining a strong acid with a strong base? Why or why not?
• HNO 2 and NaNO 2
• NH 4 NO 3 and HNO 3
• NH 4 NO 3 and NH 3
• H 3 PO 4 and Na 3 PO 4
• NaHCO 3 and Na 2 CO 3
• NaNO 3 and Ca(NO 3 ) 2
• HN 3 and NH 3
• For each combination in Exercise 3 that is a buffer, write the chemical equations for the reactions of the buffer components when a strong acid and a strong base is added.
• For each combination in Exercise 4 that is a buffer, write the chemical equations for the reaction of the buffer components when a strong acid and a strong base is added.
• The complete phosphate buffer system is based on four substances: H 3 PO 4 , H 2 PO 4 − , HPO 4 2 − , and PO 4 3 − . What different buffer solutions can be made from these substances?
• Explain why NaBr cannot be a component in either an acidic or a basic buffer.
• Explain why Mg(NO 3 ) 2 cannot be a component in either an acidic or a basic buffer.
• A buffer resists sudden changes in pH. It has a weak acid or base and a salt of that weak acid or base.
• No. Combining a strong acid and a strong base will produce salt and water. Excess strong acid or strong base will not act as a buffer.
• not a buffer

4. 1. not a buffer

3. not a buffer

4. not buffer

• 3b: strong acid: H + + NO 2 − → HNO 2 ; strong base: OH − + HNO 2 → H 2 O + NO 2 − ; 3d: strong acid: H + + NH 3 → NH 4 + ; strong base: OH − + NH 4 + → H 2 O + NH 3
• 4b: strong acid: H + + CO 3 2 − → HCO3 − ; strong base: OH − + HCO 3 − → H 2 O + CO 3 2 − ;
• Buffers can be made by combining H 3 PO 4 and H 2 PO 4 − , H 2 PO 4 − and HPO 4 2 − , and HPO 4 2 − and PO 4 3 − .
• NaBr splits up into two ions in solution, Na + and Br − . Na + will not react with any added base knowing that NaOH is a strong base. Br- will not react with any added acid knowing that HBr is a strong acid. Because NaBr will not react with any added base or acid, it does not resist change in pH and is not a buffer.
• Mg(NO 3 ) 2 includes two types of ions, Mg 2 + and NO 3 − . Mg(OH) 2 is strong base and completely dissociates (100% falls apart), so Mg 2 + will not react with any added base (0% combines with OH − ). HNO 3 is strong acid and completely dissociates (100% falls apart), so NO 3 − will not react with any added acid (0% combines with H + ). Because Mg(NO 3 ) 2 will not react with any added base or acid, it does not resist change in pH and is not a buffer.

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16.2: Buffers: Solutions That Resist pH Change

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Learning Objective

• Define buffer and describe how it reacts with an acid or a base.

Weak acids are relatively common, even in the foods we eat. But we occasionally come across a strong acid or base, such as stomach acid, that has a strongly acidic pH of 1–2. By definition, strong acids and bases can produce a relatively large amount of hydrogen or hydroxide ions and, as a consequence, have marked chemical activity. In addition, very small amounts of strong acids and bases can change the pH of a solution very quickly. If 1 mL of stomach acid [which we will approximate as 0.05 M HCl(aq)] is added to the bloodstream, and if no correcting mechanism is present, the pH of the blood would go from about 7.4 to about 4.9—a pH that is not conducive to life. Fortunately, the body has a mechanism for minimizing such dramatic pH changes.

This mechanism involves a buffer, a solution that resists dramatic changes in pH. Buffers do so by being composed of certain pairs of solutes: either a weak acid plus a salt derived from that weak acid, or a weak base plus a salt of that weak base. For example, a buffer can be composed of dissolved acetic acid (HC 2 H 3 O 2 , a weak acid) and sodium acetate (NaC 2 H 3 O 2 , a salt derived from that acid). Another example of a buffer is a solution containing ammonia (NH 3 , a weak base) and ammonium chloride (NH 4 Cl, a salt derived from that base).

Let us use an acetic acid–sodium acetate buffer to demonstrate how buffers work. If a strong base—a source of \(\ce{OH^{-}(aq)}\) ions—is added to the buffer solution, those hydroxide ions will react with the acetic acid in an acid-base reaction:

\[\ce{HC2H3O2(aq) + OH^{-}(aq) \rightarrow H2O(ℓ) + C2H3O^{-}2(aq)} \label{Eq1} \]

Rather than changing the pH dramatically by making the solution basic, the added hydroxide ions react to make water, and the pH does not change much.

Many people are aware of the concept of buffers from buffered aspirin , which is aspirin that also has magnesium carbonate, calcium carbonate, magnesium oxide, or some other salt. The salt acts like a base, while aspirin is itself a weak acid.

If a strong acid—a source of H + ions—is added to the buffer solution, the H + ions will react with the anion from the salt. Because HC 2 H 3 O 2 is a weak acid, it is not ionized much. This means that if lots of hydrogen ions and acetate ions (from sodium acetate) are present in the same solution, they will come together to make acetic acid:

\[\ce{H^{+}(aq) + C2H3O^{−}2(aq) \rightarrow HC2H3O2(aq)} \label{Eq2} \]

Rather than changing the pH dramatically and making the solution acidic, the added hydrogen ions react to make molecules of a weak acid. Figure \(\PageIndex{1}\) illustrates both actions of a buffer.

Buffers made from weak bases and salts of weak bases act similarly. For example, in a buffer containing NH 3 and NH 4 Cl, ammonia molecules can react with any excess hydrogen ions introduced by strong acids:

\[\ce{NH3(aq) + H^{+}(aq) \rightarrow NH^{+}4(aq)} \label{Eq3} \]

while the ammonium ion (\(\ce{NH4^{+}(aq)}\)) can react with any hydroxide ions introduced by strong bases:

\[\ce{NH^{+}4(aq) + OH^{-}(aq) \rightarrow NH3(aq) + H2O(ℓ)} \label{Eq4} \]

Example \(\PageIndex{1}\): Making Buffer Solutions

Which solute combinations can make a buffer solution? Assume that all are aqueous solutions.

• HCHO 2 and NaCHO 2
• HCl and NaCl
• CH 3 NH 2 and CH 3 NH 3 Cl
• NH 3 and NaOH
• Formic acid (HCHO 2 ) is a weak acid, while NaCHO 2 is the salt made from the anion of the weak acid—the formate ion (CHO 2 − ). The combination of these two solutes would make a buffer solution.
• Hydrochloric acid (HCl) is a strong acid, not a weak acid, so the combination of these two solutes would not make a buffer solution.
• Methylamine (CH 3 NH 2 ) is like ammonia with one of its hydrogen atoms substituted with a CH 3 (methyl) group. Because it is not on our list of strong bases, we can assume that it is a weak base. The compound CH 3 NH 3 Cl is a salt made from that weak base, so the combination of these two solutes would make a buffer solution.
• Ammonia (NH 3 ) is a weak base, but NaOH is a strong base. The combination of these two solutes would not make a buffer solution.

Exercise \(\PageIndex{1}\)

• NaHCO 3 and NaCl
• H 3 PO 4 and NaH 2 PO 4
• NH 3 and (NH 4 ) 3 PO 4
• NaOH and NaCl

Buffers work well only for limited amounts of added strong acid or base. Once either solute is all reacted, the solution is no longer a buffer, and rapid changes in pH may occur. We say that a buffer has a certain capacity . Buffers that have more solute dissolved in them to start with have larger capacities, as might be expected.

Human blood has a buffering system to minimize extreme changes in pH. One buffer in blood is based on the presence of HCO 3 − and H 2 CO 3 [H 2 CO 3 is another way to write CO 2 (aq)]. With this buffer present, even if some stomach acid were to find its way directly into the bloodstream, the change in the pH of blood would be minimal. Inside many of the body’s cells, there is a buffering system based on phosphate ions.

Career Focus: Blood Bank Technology Specialist

At this point in this text, you should have the idea that the chemistry of blood is fairly complex. Because of this, people who work with blood must be specially trained to work with it properly.

A blood bank technology specialist is trained to perform routine and special tests on blood samples from blood banks or transfusion centers. This specialist measures the pH of blood, types it (according to the blood’s ABO+/− type, Rh factors, and other typing schemes), tests it for the presence or absence of various diseases, and uses the blood to determine if a patient has any of several medical problems, such as anemia. A blood bank technology specialist may also interview and prepare donors to give blood and may actually collect the blood donation.

Blood bank technology specialists are well trained. Typically, they require a college degree with at least a year of special training in blood biology and chemistry. In the United States, training must conform to standards established by the American Association of Blood Banks.

Key Takeaway

• A buffer is a solution that resists sudden changes in pH.

Contributions & Attributions