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Effect of pH on life? Use buffer or HCl? How about lactic acid?

Effect of pH on life? Use buffer or HCl? How about lactic acid?


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I'm trying to find the effect of pH on living things like E. coli and such. At first I thought of using a buffer solution which maintains its level of pH, but it's possible that the inactive ions of each buffer have their own effect on living things that I can't account for. On the other hand, I could use hydrochloric acid and sodium hydroxide at different concentrations by diluting it. This will be a little bit more effort (I'll be diluting the acid and base myself), but much cheaper than the buffers, and I'll know that the only other ions present beside H+ and OH- are just Na+ and Cl-, the ingredients of table salt. According to this paper, the effect of pH on the population also depended on the acid (HCl vs lactic acid).

I would like a second opinion on this: If I'm investigating the effect of pH on growth of living things, do I use specially made buffers or just properly diluted HCl and NaOH?


That is a variable that you have the pleasure defining yourself. Just know that using a buffer will introduce other atoms that could affect the "living things". Conversely, if you don't use a buffer, it is likely that the pH of your "living things" will be easily altered by the metabolism of your living things (IMHO, this is worse).

I don't know the details of your project, but recommend:

-use a buffer like… Tris, phosphate… -use 10 mM Tris-Cl without any extra electrolyte (aside of the HCl or NaOH that you use to modulate the pH of your buffer to different pH's that you want to test) -e.g. Tris-Cl pH 7, 8, 9… or Phosphate pH 2,3,4,5,6,7,8,9…


What Is Lactic Acid Fermentation & How Does It Preserve Food?

Most people are familiar with Dill Pickles and Sauerkraut as pickled foods. The familiar tang of a pickle many people associate with vinegar and not another form of acid created via lactic acid fermentation.

There are key differences between pickles preserve with vinegar and those pickled with lactic acid so let’s take a look at what lactic acid fermentation is and how people for thousands of years have been using lactic acid fermentation to preserve food.


Acid-Base Homeostasis: Response & Regulation | Excretory System | Biology

In this article we will discuss about:- 1. Definition of Acid-Base Homeostasis 2. Response to an Acid-Base Imbalance 3. Regulation.

Definition of Acid-Base Homeostasis:

Acid-base homeostasis is the part of human homeo­stasis concerning the proper balance between acids and bases, in other words, the pH. The body is very sensitive to its pH level, so strong mechanisms exist to maintain it. Outside the acceptable range of pH, proteins are denatured and digested, enzymes lose their ability to function, and death may occur.

The term pH refers to the negative log of hydrogen ion concentration-

The normal range of blood pH falls between 7.35 and 7.45 and our Acid-base balance has to maintain pH within this normal range.

pH of Some Body Fluids:

Survival range of pH is – 6.8 to 8.0

An acid is a molecule containing hydrogen atom that can release hydrogen ions in solutions, e.g.

There is always a constant production of acid by the body’s metabolic processes and to maintain balance, these acids need to be excreted or metabolized. The various acids produced by the body are classified as respiratory (or volatile) acids and metabolic (or fixed) acids.

The acid is more correctly carbonic acid (H2CO3) but the term ‘respiratory acid’ is usually used to mean carbon dioxide. Carbon dioxide is the end-product of complete oxidation of carbohydrates and fatty acids. It is called a volatile acid meaning in this context it can be excreted via the lungs. Of necessity, considering the amounts involved there must be an efficient system to rapidly excrete CO2.

This term covers all the acids the body produces which are nonvolatile. Because they are not excreted by the lungs they are said to be ‘fixed’ in the body and hence the alternative term fixed acids. All acids other than H2CO3 are fixed acids.

For Acid-base balance, the amount of acid excreted per day must equal the amount produced. The routes of excretion are the lungs (for CO2) and the kidneys (for the fixed acids).

Acid-base imbalance occurs when a significant insult causes the blood pH to shift out of the normal range (7.35 to 7.45). An excess of acid is called acidosis (pH less than 7.35) and an excess of base is called alkalosis (pH greater than 7.45). The process that causes the imbalance is classified based on the etiology of the disturbance (respiratory or metabolic) and the direction of change in pH (acidosis or alkalosis).

There are four basic processes:

(iii) Metabolic alkalosis, and

One or a combination may occur at any given time.

Response to an Acid-Base Imbalance:

The body’s response to a change in Acid-base status has three components:

Respiratory compensation by alteration in arterial PCO2

Renal compensation by alteration in HCO3 – excretion.

A buffer is any substance that can reversibly bind H + . Buffer + H + ↔ H buffer

80 mEq of H + are produced per day.

1. Bicarbonate Buffer System:

The major buffer system in the ECF is the CO2-bicarbonate buffer system. This is responsible for about 80% of extracellular buffering but it cannot buffer respiratory Acid- base disorders.

2. Phosphate Buffer Systems:

The phosphate buffer systems are not important blood buffer as its concentration is too low. It plays an important role in renal tubular system.

Protein buffers in blood include hemoglobin (150 g/1) and plasma proteins (70 g/1). Buffering is by the imidazole group of the histidine residues. Hemoglobin is quantitatively about 6 times more important than the plasma proteins as it is present in about twice the concentration and contains about three times the number of histidine residues per molecule. For example, if blood pH changed from 7.5 to 6.5, hemoglobin would buffer 27.5 mmol/1 of H + and total plasma protein buffering would account for only 4.2 mmol/1 of H + . Deoxyhemoglobin is a more effective buffer than oxyhemoglobin.

‘Whenever, there is a change in H + concentration in the ECF, the balance of all the buffer systems changes at the same time’.

H + = K1 × HA1/A1 = K2 × HA2/A2 = K3 × HA3/A3.

K1, K2 and K3 are dissociation constants of 3 respective acids.

Regulation of Acid-Base Balance:

I. Respiratory Regulation of Acid-Base Balance:

i. Regulates H + concentration through CO2 in venti­lation.

ii. ↑ [H + ] → ↑ alveolar ventilation

iii. Buffering power of respiratory system is 1-2 times greater than chemical buffers.

iv. Lung diseases decrease the efficacy of the buffering power. Respiratory regulation refers to changes in pH due to PCO2 changes by altering the ventilation. This change in ventilation can occur rapidly with significant effects on pH. Carbon dioxide is lipid soluble and crosses cell membranes rapidly, so changes in PCO2 result in rapid changes in [H + ] in all body fluid compartments.

II. Renal Regulation of Acid-Base Balance:

There are three systems that regulate H + concentration in the body fluids to prevent acidosis or alkalosis.

1. The chemical Acid-base buffer systems which combine with acid or base to prevent excessive changes in H + concentration.

2. The respiratory centers which regulates the removal of CO2 from ECF.

3. The kidneys regulate blood pH by three mechanisms.

i. The chemical Acid-base buffer systems which combine with acid or base to prevent excessive changes in H + concentration.

ii. The respiratory centers which regulates the removal of CO2 from ECF.

a. Excretion of acid in the form of titrable acid and ammonium ions.

b. Reabsorption of the filtered HCO3

c. Generation of new NaHCO3

Mechanism of H + Secretion by PT:

i. Formation of carbonic acid

ii. Secretion of H + into the lumen via Na + H + counter-transport in luminal membrane—an example of secondary active transport.

iii. H + secreted in the lumen combines with filtered HCO3 and helps in reabsorption.

iv. HCO3 formed in the cell diffuses into interstitial fluid through basolateral membrane. This is done by Na HCO3 transport and CI HCO3 exchanger. Thus for each H + secreted one Na + and one HCO3 ion enter the interstitial fluid.

Fate of H ion into the Lumen:

1. Nontitrable Acidity:

H + ion combines with HCO3 and NH3 producing non-titrable acids. The reactions are:

The process by which NH3 is secreted into the urine and then changed to NH4 maintaining the concentration gradient for diffusion of NH3 is called nonionic diffusion.

Ammonium Ion Secretion:

Glutamine is metabolized in PCT cells yielding ammonium and bicarbonate. The NH4 + is actively secreted by Na + NH4 + pump and bicarbonate is returned to blood.

Ammonium Ion Secretion in CD:

The CD is permeable to NH3 which diffuses into tubular lumen but less permeable to NH4, therefore, NH4 is trapped in the tubular lumen and excreted in the urine.

The H + ions that combine with dibasic phosphate produce monobasic phosphate which contributes titrable acidity.

Net acid secretion = Titrable + urinary NH4 – urinary HCO3 acidity

Total acid excreted by the kidney = 50 to 100 mEq/day.

Mechanism of H + Secretion by DT and CD is independent of Na + :

i. ATP driven pumps increase H + concentration by 1000 times. Aldosterone acts on this pump to increase H + secretion.

ii. H + K + ATPase is also responsible.

Reabsorption of Filtered HCO3:

PT reabsorbs 80% of filtered HCO3, H + secreted in lumen of PT combines with HCO3 to form H2CO3. It is converted to CO2 and H2O. CO2 diffuses into tubular cells. CO2 combines with H2O to form H2CO3 which dissociates into H + and HCO3.

The H + ion is secreted into tubule and HCO3 ion diffuses into interstitial fluid. When each molecule of HCO3 is reabsorbed into lumen, one molecule of HCO3 diffuses into blood even though it is not the same molecule. The pH of fluid in proximal tubule is very little altered since the H + ion secretion is neutralized by HCO3 ion reabsorption.

LOH reabsorbs 15% of filtered HCO3.

DT and CT reabsorb only 5% of the filtered HCO3.

Generation of New NaHCO3 Ions:

Phosphate and ammonia buffers in the tubule carries excess H + ions to generate new NaHCO3 ions. Therefore, whenever H + ion secreted into the tubules combines with a buffer other than HCO3, the net effect is addition of new bicarbonate to the blood. For examples if H + reacts with NH3 to form NH4, NH4 is trapped in the tubular lumen and eliminated in the urine. For each NH4 excreted, a new HCO3 is generated and added to the blood.

Kidneys filter 4320 mEq of HCO3 – /day.

To reabsorb 4320 mEq of HCO3, equal amount of H + are secreted.

In addition, 80 mEq of H + from nonvolatile acids are also secreted, making the total of 4400 mEq of H + to be secreted/day. Only a small amount of excess H + can be secreted in ionic form in the urine. Minimal urine pH is about 4.5 corresponding to an H + concen­tration of 0.03 mEq/L. For every liter of urine only 0.03 mEq of H+ can be excreted.

To excrete 80 mEq of H + , 2667 liters of urine would have to be formed.

Defined as an increase in H + concentration or decrease in pH (<7.4).

Defined as a decrease in H + concentration or increase in pH (>7.4).

Any disturbance of Acid-base balance resulting from changes in HCO3 – concentration in ECF.

Disturbances in Acid-base balance due to changes in PCO2.

Respiratory Alkalosis:

Either a decrease in tubular secretion of H + or increased excretion of HCO3 – .

Either an increase in excretion of H + or by generation of new HCO3 – .

Any factor that decreases the rate of pulmonary venti­lation increases the PCO2 of ECF → ↑ H2CO3 → ↑ H + .

i. Damage to respiratory center.

ii. Obstruction of passages of the respiratory tract.

Respiratory Alkalosis:

ii. Physiologically at high altitudes.

i. Failure of kidneys to excrete metabolic acids.

ii. Formation of excess quantities of metabolic acids.

iii. Addition of metabolic acids to the body.

iv. Loss of base from the body.

Renal Tubular Acidosis:

v. Vomiting of intestinal contents.

vii. Ingestion of acids (aspirin, methyl alcohol).

i. Excess retention of HCO3 – .

ii. Loss of H + from the body.

iii. Use of diuretics (except carbonic anhydrase inhi­bitor).

v. Vomiting of gastric contents.

vi. Ingestion of alkaline drugs.

vii. Treatment of acidosis.

ix. Infusion of sodium lactate and sodium gluconate.

Treatment of Alkalosis:

ii. Lysine monohydrochloride.

↓ Tubular section of H+ ions.

To diagnose acid-base disorders quickly and to find out the severity. pH, PCO2 and HCO3 values are used. Sufficient time should be given for compensatory res­ponse. 6-12 hours for lungs and 3-5 days for kidneys.


Buffers: Definition, Principles and Uses

In this article we will discuss about Buffers:- 1. Definition of Buffers 2. Principles of Buffers 3. Determining the pH 4. Buffer Mixture 5. Buffer Pairs in the Blood 6. Uses 7. Tissue Fluids and Tissues 8. Role in pH Regulation 9. Acidosis and Alkalosis Acidosis 10. Role of Lungs and Kidneys in pH Regulation 11. Elimination of Free Acids 12. Renal Correction of Acidosis 13. Renal Correction of Alkalosis.

  1. Definition of Buffers
  2. Principles of Buffers
  3. Determining the pH of Buffers
  4. Buffer Mixture
  5. Buffer Pairs in the Blood
  6. Uses of Buffers
  7. Buffers of Tissue Fluids and Tissues
  8. Role of Buffers in pH Regulation
  9. Acidosis and Alkalosis Acidosis
  10. Role of Lungs and Kidneys in pH Regulation by Means of Buffers
  11. Elimination of Free Acids
  12. Renal Correction of Acidosis
  13. Renal Correction of Alkalosis

1. Definition of Buffers:

Buffers are the mixtures of weak acids and their salts of strong bases (or strong acids and their salts of weak bases).

Acetic acid (CH3COOH) + Sodium acetate (CH3COONa).

2. Principles of Buffers:

HAC + NaAC → Na + + H + + 2AC −

where, HAC=Acetic acid NaAC = Sodium acetate.

If alkali (NaOH) is added to this system, it will form salt and no free H + or OH − will be available.

HAC + NaAC + NaOH → 2NaAC + H2O

If acid (HCl) is added to this system, it will also form salt and no free H + or OH − will be available.

HAC + NaAC + HCl NaCl + 2HAC

In either cases there is no change in hydrogen ion concentration. The buffer acts almost as if it were “absorbing” the added free hydrogen or hy­droxyl ions.

3. Determining the pH of Buffers:

The pH of buffers can be determined by the Henderson-Hasselbalch equation:

In case of blood, the ratio between [BHCO3]: [H2CO3] can be found out by applying the above equation to maintain average pH of blood 7.4:

The pH of human blood is 7.4. In normal health, it lies between 7.3 and 7.5 although CO2 (i.e. carbonic acid) is always added. If the pH of human blood becomes 7.0 and 7.6, it alarms dan­ger if not fatal.

The difference in pH between arte­rial and venous blood is rarely more than 0.04. A marked decrease in pH of blood has been observed during severe muscular exercise when the blood lactic acid content rises over 100 mg per 100 ml.

4. Buffer Mixture:

(b) Acid potassium phthalate and HCl.

(c) Acid potassium phosphate and NaOH.

(d) Sodium bicarbonate and Sodium carbon­ate.

5. Buffer Pairs in the Blood:

Actually in blood the buffering of carbonic acid is complicated by the presence of the red cells:

6. Uses of Buffers:

i. Buffers are used for preparing standard so­lutions in which it is always desired to maintain a constant pH. This is required for the colorimetric determination of the pH of unknown solutions.

ii. These are used to maintain H + concentra­tion which is necessary for optimal activ­ity of enzymes.

iii. These are practically important in all physiological systems.

7. Buffers of Tissue Fluids and Tissues:

i. The buffering system of lymph, cerebros­pinal fluid etc., are similar to that of blood although the quantity is much less.

ii. The chief buffering system in these fluids is BHCO3 and H2CO3.

iii. The buffering system in tissues is mainly BHCO, and H2CO3 protein buffers and organic acid salt.

8. Role of Buffers in pH Regulation:

(i) Bicarbonate Buffer:

a. It is the main buffer in blood plasma and consists of bicarbonate (HCO − 3) and car­bonic acid (H2CO3).

b. The bicarbonate buffer neutralizes stronger dietary and metabolic acids (HA) converting them into weak bases (A – ) with the increase in H2CO3. Stronger bases (B) are also changed into weak acids (BH + ) with the rise in HCO − 3.

c. The pH of blood is maintained 7.4 when the buffer ratio becomes 20. If the bicar­bonate buffer neutralizes any acid or base, there may be the change of buffer ratio and the blood pH value. But the buffer ratio remains unchanged by the respira­tory elimination of H2CO3 as CO2 or the urinary elimination of HCO − 3.

d. Since cells contain much lower amounts of HCO − 3 the importance of bicarbonate buffer inside the cell is negligible.

(ii) Phosphate Buffer:

a. Since the concentration of phosphate buffer in the blood plasma is about 8 per cent of that of the bicarbonate buffer, its buffering capacity is much lower than bi­carbonate in the plasma.

b. The phosphate buffer consists of dibasic phosphate (HPO −− 4) and monobasic phosphate (H2PO − 4). Its PKa value is about 6.8. It is more effective in the pH range 5.8 to 7.8. Plasma has a [HP0 — 4]:

c. The concentration of phosphate buffer is much higher in intra-cellular fluid than in extracellular fluids. The pH of intracellu­lar fluids (6.0 – 6.9) is nearer to the PKa of the phosphate buffer. Therefore, the buff­ering capacity of the phosphate buffer is highly elevated inside the cells and the phosphate buffer is also effective in the urine inside the renal distal tubules and collecting ducts.

d. In case the ratio of [HPO −− 4]: [H2PO − 4] tends to be changed by the formation of more H2 PO − 4, there occurs the renal elimi­nation of H2PO4 for which the ratio ulti­mately remains unaltered.

(iii) Protein Buffers:

a. The protein buffers are very important in the plasma and the intracellular fluids but their concentration is very low in C.S.F., lymph, and interstitial fluids.

b. They exist as anions serving as conjugate bases (Pr) at the blood pH 7.4 and form conjugate acids (HPr) accepting H + .

3=c. They have the capacity to buffer some H2CO3 in the blood:

(iv) Hemoglobin Buffers:

a. They are involved in buffering CO2 in­side erythrocytes. The buffering capacity of hemoglobin depends on its oxygena­tion and de-oxygenation. Inside the eryth­rocytes CO2 combines with H2O to form H2CO3 under the action of carbonic anhydrase.

At the blood pH 7.4, H2CO3 dissoci­ates into H + and H2CO3 and needs imme­diate buffering. Oxy-hemoglobin (HBO − 2) on the other side loses O2 to form de-oxy-hemoglobin (Hb − ) which remains un-dissociated (HHb) by accepting H + from the ionization of H2CO3.

Thus, Hb buff­ers H2CO3 in erythrocytes:

Some of the HCO − 3 diffuse out into the plasma to maintain the balance between intracellular and plasma bicarbonates. This causes influx of some CI − into erythrocytes along the electrical gradient produced by the HCO − 3 outflow (chloride shift).

b. HHbo2, produced in lungs by oxygena­tion of HHb, immediately ionizes into H + and HbO − 2. The released hydrogen ions (H + ) are buffered by HCO – 3 inside eryth­rocyte to form H2CO3 which is dissociated into H2O and CO2, by carbonic anhydrase. CO2 diffuses out of erythrocytes and es­capes in the alveolar air. Some HCO − 3 re­turn from the plasma to erythrocytes in exchange of Cl − and are changed to CO2.

9. Acidosis and Alkalosis Acidosis:

a. Accumulation of acid or loss of alkali is called acidosis.

b. It occurs due to the loss or fall of [HCO – 3] : [H2CO3] of blood below 20.

c. There are two types of acidosis: (a) Meta­bolic (b) Respiratory.

(i) The concentration of plasma bicarbonate is decreased in excessive loss of bases.

(ii) It happens in renal failure, diabetic ketosis, severe diarrhoea.

(i) The retention of CO2 is caused in hypoventilation resulting in the rise of H2CO3. This lowers [HCO − 3]: [H2CO3] ra­tio.

(ii) It happens in chronic obstructive airway diseases (asthma, respiratory paralysis), prolonged anesthesia, and unconscious­ness due to any cause.

a. Accumulation of alkali and loss of acid is called alkalosis.

b. There is increase in the ratio of [HCO − 3]: [H2CO3] of blood above 20, resulting the rise in blood pH.

3. There are two types:

(a) Metabolic:

(i) High intake of alkaline substances like NAHCO3 may elevate the plasma HCO − 3.

(ii) It happens in severe vomiting due to any cause, indiscriminate use of antacid.

(b) Respiratory:

(i) The excess CO2 is removed from the blood due to hyperventilation and causes de­crease of H2CO3.

(ii) It happens in high altitude (hyper-ventilation syndrome), hysteria.

10. Role of Lungs and Kidneys in pH Regulation by Means of Buffers:

a. Lungs maintain the normal ratio of [HCO – 3]: [H2CO3] and the pH of blood by altering the rate of respiratory elimination of CO2 from the blood. This lowers the alveolar PCO2 and increases the diffusion of CO2 to the alveolar air. Since the con­centration H2CO3 is in equilibrium with that of dissolved CO2 in the blood, hyper­ventilation increases the ratio of [HCO − 3]. : [H2CO3] with the fall in the CO2 concen­tration. The doubling of ventilation may raise the blood pH by 0.4.

b. Hypoventilation, on the other hand, raises the blood concentration of dissolved CO2 and consequently lowers the buffer ratio. A fall in alveolar ventilation to one-fourth of the normal value may lower the blood pH by 0.46. The pulmonary ventilation is adjusted according to the blood pH. Hypoventilation not only retains CO2 de­crease the ratio of [HCO − 3] : [H2CO3] and the blood pH but also curtails the O2 sup­ply—an undesirable effect.

c. Lungs also have a role in the functioning of hemoglobin buffers through oxygena­tion and de-oxygenation which has been discussed earlier.

a. The pH of the glomerular filtrate is about 7.4. But the pH falls to about 6.9 in the proximal tubule, then to about 6-6.5 in the distal tubule and ultimately to about 4.5-4.7 in the collecting duct.

b. The urinary pH is maintained by a coop­eration between the urinary buffers and the renal ion-exchange mechanism. The major urinary buffers are bicarbonate and phosphate buffers. As the filtrate proceeds along the tubules, the ratio between base member and the acid member of each uri­nary buffer falls progressively with a con­sequent fall in the urinary pH.

c. In the renal ion exchange mechanism, some urinary Na + are actively reabsorbed in the exchange of H H secreted in the tu­bular filtrate. 85 percent of H + are secreted by the proximal tubule and 15 per cent from the distal tubules and collecting ducts. The hydrogen ions (H + ) are mainly formed from the ionization of H2CO3 formed from CO2 and H2O by carbonic anhydrase in the tubule cells.

The bicarbonate formed is returned to blood along with the reabsorption of Na + .

d. Most of the secreted H + are immediately buffered by HCO – 3 filtered from the plasma into the glomerular filtrate.

e. H2CO3 produced by such buffering in the proximal tubules is immediately removed through its cleavage to CO2 and cannot play any role in lowering the pH there.

f. The buffering of the secreted H + is essen­tial for continuing their secretion in the urine. This buffering is held in the tubular filtrate in several ways.

(a) Buffering by Bicarbonate:

(i) Normally about 3.50 mM of H + are secreted in the tubules per minute.

(ii) The major portion of the secreted H + can be buffered by HCO − 3 in the tubular fil­trate to form H2CO3, except a small amount of free H + to pass into the urine.

(iii) H2CO3 is then immediately cleaved into H2O and CO2 in the proximal tubular lu­men by carbonic anhydrase.

(iv) CO2 then diffuses very readily into the proximal tubule cell and therefrom to the blood.

(v) In erythrocytes, carbonic anhydrase con­verts this CO2 into H2CO3 which dissoci­ates to form fresh HCO – 3. This HCO – 3 is restored in the plasma.

(vi) When H + are secreted in excess due to the fall in blood pH, almost all the filtered HCO − 3 changes into H2CO − 3 and is re­turned to the plasma. So the urinary HCO − 3 is negligible so long as the uri­nary pH does not exceed 6. But whenever the blood pH tends to rise, much more HCO − 3 is filtered than the amount of H + secreted. So, some of the filtered HCO3 fails to get H + to combine and fails to re­turn to the plasma. This causes the urinary elimination of bicarbonate.

(vii) Buffering of the secreted H + by the filtered HCO − 3 serves two purposes:

1. It does not allow the pH to fall below 6.9 in the proximal tubule and allows more H + to be eliminated by tubule cells into the urine.

2. It helps to reabsorb the filtered HCO – 3 and to restore it in the blood.

(b) Buffering by Phosphate Buffer:

(i) In the distal tubules, some secreted H + is buffered by the phosphate buffer. HPO −− 4, filtered into the glomerular filtrate, re­ceives the secreted H + to form H2PO − 4. This changes the ratio of [HPO −− 4]: [H2PO − 4] from 4 of Bowman’s capsule to 0.02-0.05 in the final urine. The lower this ratio, the more acidic is the urine.

(ii) H2PO – 4 is eliminated in urine carrying some Na + with it resulting in the urinary loss of Na + .

(iii) The total buffering capacity of urinary phosphate is much less than that of bicar­bonate.

(iv) As the urine gets more concentrated in the distal tubules and collecting ducts, the rise in tubular concentration of HPO −− 4 en­hances the buffing capacity of the phos­phate buffer.

(c) Buffering by Ammonia:

(i) The base NH3, synthesized and secreted by the tubule cells, can buffer some H + in the distal tubule.

(ii) In the tubular lumen, NH3 combines with H + to form NH + 4 which is excreted in the urine in association with CI – and SO −− 4 left behind by the reabsorbed Na + .

(iii) NH + 4 behaves like a weak acid and does not dissociate much. Its formation lowers the tubular concentration of free H + , ena­bling further secretion of H + . The tubular membrane is not permeable to NH + 4 which is retained in the tubular filtrate in­stead of diffusing back into the tubule cells (diffusion trapping).

(iv) The secretion of NH3 rises whenever the free H + concentration is high enough to lower the urinary pH below 6 the more acidic the urine, the higher is the urinary ammonia.

(v) An alkaline urine contains little or no am­monia, kidneys normally excrete about 40 mEq of NH + 4 in 24 hours. The elimina­tion of highly acidic urine may enhance more ammonia secretion.

11. Elimination of Free Acids:

Strong conjugate bases such as lactate, acetoacetate, urate and oxalate anions accept some H + re­placing the Na + reabsorbed from their salts. As a result, the free acids such as lactic acid, acetoacetic acid, uric acid are excreted. Their elimination changes the urinary pH a little only.

12. Renal Correction of Acidosis:

In acidosis, the blood carries a high amount of dis­solved CO2 compared to that of HCO − 3 and the tubule cells secrete far more H + than the HCO − 3 filtered from glomeruli. As a result, all the filtered H2CO − 3 combines with H + to form H2CO3.

Thus, the urine does not contain HCO − 3 while the reabsorbed HCO − 3 is retained in blood increasing the buffer ratio. In addition, tubular secretion of NH3 increases for buffering the H + left in the tubules after all the HCO3 has been reabsorbed.

13. Renal Correction of Alkalosis:

In alkalosis, the blood carries a high amount of HCO − 3, and the glomerular filtrate contains far more HCO − 3, than H + secreted in the tubules. The urinary elimination of unabsorbed HCO − 3, causes a loss of HCO − 3, from blood and finally lowers the buffer ratio and the pH of blood. Alkalosis also reduces the tubular secretion of NH3.


1 Answer 1

I'm not sure I follow your argumentation, but I guess you don't really need $K_mathrm$ at all since addition of strong acid ($ce$) influences the dissociation of the one buffer component - the weak acid ($ce$):

Henderson–Hasselbalch equation applied to this buffer system before the addition of acid allows to find initial $mathrm$ (not required by the problem, I do this solely for demonstration):

Once the strong acid ($ce$, assuming complete dissociation) is added, the equilibrium shifts accordingly:


Results

Lactic acid yield from single sugar sources

To establish the preference for single carbon sources and the overall conversion yield to LA, a defined MRS medium was formulated containing individual sugars and was fermented with each strain (Table 2). AB39 consumed more fructose than glucose, but converted glucose more efficiently into LA than fructose (yield of 89% compared with 71%, respectively). Maltose was consumed to a lesser extent and LA yields were low (12%). A high level of assimilation was seen for FST1.7, and sugar-to-LA conversion rates were very high for all sugars (78–95%). Very high yield values (> 90%) are probably due to the production of LA from sugar impurities in the medium and/or from sources other than carbohydrates, e.g. amino acids. FST2.11 consumed maltose and glucose in higher quantities compared with fructose however, yields from glucose were poorer (40%) than for AB39. Finally, FST2.11 was the only strain producing LA (10.19 ± 0.2 g/L) when incubated in 2% (w/w) soluble starch in MRS. The consumption of starch amounted to 5.07 ± 0.77 g/L, while the rest of the carbon was provided by dextrins and simpler sugars present in the medium.

Fru-MRS Glu-MRS Mal-MRS
Sugar consumption (%) LA (g/L) YLA/Sa a LA yield was calculated as a percentage from the amount (g) of LA produced divided by the amount (g) of sugar consumed.
Sugar consumption (%) LA (g/L) YLA/Sa a LA yield was calculated as a percentage from the amount (g) of LA produced divided by the amount (g) of sugar consumed.
Sugar consumption (%) LA (g/L) YLA/Sa a LA yield was calculated as a percentage from the amount (g) of LA produced divided by the amount (g) of sugar consumed.
AB39 19.2 13.65 ± 0.24 0.71 ± 0.03 14.5 12.85 ± 0.10 0.89 ± 0.04 6.5 0.77 ± 0.04 0.12 ± 0.07
FST1.7 16.9 13.22 ± 0.17 0.78 ± 0.05 18.2 17.36 ± 0.10 0.95 ± 0.03 18.6 17.75 ± 0.12 0.95 ± 0.02
FST2.11 8.2 4.95 ± 0.16 0.60 ± 0.02 15.3 6.1 ± 0.37 0.40 ± 0.01 15.8 10.36 ± 0.14 0.66 ± 0.05
  • a LA yield was calculated as a percentage from the amount (g) of LA produced divided by the amount (g) of sugar consumed.

PH inhibition of LAB

To estimate the effect of end-product accumulation on cell self-inhibition, control wort was pH adjusted by adding either LA or HCl (Fig. 2). Growth gradually declined with decreasing pH values, and dropped abruptly at pH 3.9 for FST1.7 and pH 3.4 for AB39 and FST2.11. The pH adjusted with LA caused a stronger inhibitory effect on the bacterial growth than HCl at pH 4.9 and lower. Growth stopped at pH 3.4 for all strains when pH was corrected with LA. Among the strains, FST2.11 showed higher viability at lower pH values. As an example, in terms of relative growth, FST2.11 was inhibited only by 63% when growing in a LA-acidified wort (pH 3.9 with LA), while it was reduced by 81 and 94% for AB39 and FST1.7, respectively.

Improvement of buffering capacity during mashing

Values of mash BC were quantified from mashing-in until the end of the proteolytic rest (50°C Fig. 3). Starting at a value of 0.94, the BC rose to a maximum of 1.34 after 90 min (+43%) with no significant change over the last 15 min. A further increase in the BC during the remaining mashing process could not be observed.

To further improve the BC of both control wort (CW) and optimized wort (OW), an external protease was added at mashing in. The increase in FAN in the worts amounted to +23% in CW + P and +34% in OW + P. Regression analysis of corresponding BC and FAN showed that wort FAN had a statistically significant (p = 0.000) linear relationship with BC (R 2 = 0.957 Fig. 4).

Impact of buffering on LA production

The four worts were fermented with each bacterial strain over a period of 48 h (Fig. 5). In addition, a citrate-based buffer was added to control wort (CW + B) and diluted (50:50 with water) control wort (CW0.5 + B) to further improve LA production. Trials in diluted wort were performed to study the gradual deficiency of nutrients on LAB activity. The BC values of CW + B and CW0.5 + B were 5.7 and 5.3 times higher than CW, respectively.

Compared with the control, the LA released in worts obtained by extending proteolysis and/or adding protease showed contrasting results amongst the strains tested. AB39 showed no notable increase in LA production, while FST1.7 achieved a significant (p < 0.05) increase of LA in OW + P compared with CW (8.65 ± 0.11 and 7.23 ± 0.29 g/L, respectively). This strain showed a linear correlation (R 2 = 0.990) between the BC of the substrate and LA released. In contrast, FST2.11 reacted positively to the application of external protease, but only when added in the control wort (CW + P). The fermentation of this substrate led to high LA accumulation (11.3 g/L), corresponding to an increase of +24% compared with CW, while reaching the lowest pH values after fermentation (average of 0.25 lower than the other strains). The low final pH values reached during the trials were likely to be responsible for cessation of bacterial growth. Fermentation of CW + B resulted in higher LA concentration for all strains (32–53% compared with CW), with maximal LA released by FST2.11 (12.8 g/L). Lower LA values were found in the diluted, buffered wort (CW0.5 + B). These were comparable with the values found for CW, but the significantly higher final pH values (4.25–4.71 compared with 3.05–3.31 in CW) suggested that, in this case, depletion of an essential nutrient(s) or co-factor(s) could have led to suboptimal fermentations.

Metabolite consumption and cell count in buffered worts

A closer look at the sugar consumption (Table 3) revealed that fructose and glucose were the preferred carbon sources of all strains, with complete assimilation of these monosaccharides in buffered trials, while residual sugars remained during control fermentation in CW for AB39 and FST1.7. The only culture that consumed maltose was FST2.11, while maltotriose was not utilized by any of the strains. Trials in CW + B led to the highest cell counts compared with CW and CW0.5 + B, with the latter trials sharing similar values. Maximal cell growth for FST1.7 corresponded also to the largest decrease in FAN.

Cell count Fructose Glucose Maltose FAN
CW
Unfermented ND 1.89 ± 0.00 10.30 ± 0.14 67.72 ± 0.28 169 ± 6
AB39 8.67 ± 0.05 <LOD 8.25 ± 0.03 65.85 ± 0.25 159 ± 2
FST1.7 9.23 ± 0.13 <LOD 7.52 ± 0.21 65.02 ± 1.23 143 ± 1
FST2.11 8.18 ± 0.10 <LOD <LOD 60.93 ± 0.28 144 ± 6
CW0.5 + B
Unfermented ND 0.94 ± 0.06 4.99 ± 0.27 32.10 ± 0.26 78 ± 7
AB39 8.63 ± 0.08 <LOD <LOD 34.17 ± 0.27 79 ± 1
FST1.7 8.95 ± 0.01 <LOD <LOD 31.60 ± 0.16 57 ± 2
FST2.11 7.91 ± 0.10 <LOD <LOD 29.27 ± 0.36 86 ± 6
CW + B
Unfermented ND 1.82 ± 0.03 9.78 ± 0.12 63.88 ± 0.88 165 ± 4
AB39 8.92 ± 0.06 <LOD <LOD 65.79 ± 1.21 152 ± 3
FST1.7 9.47 ± 0.07 <LOD <LOD 64.61 ± 0.43 126 ± 8
FST2.11 8.79 ± 0.09 <LOD <LOD 59.37 ± 0.86 158 ± 2
  • LOD, Limit of detection for fructose and glucose was 0.10 g/L and 0.25 g/L, respectively.
  • ND, Not detected (< 3 log CFU/mL).

The analysis of 18 FAA was performed in diluted buffered wort (CW0.5 + B) to investigate substrate-specific causes for bacterial growth cessation (Table 4). The results showed a strain-dependent consumption of single amino acids, with glutamine being completely assimilated by all strains, serine by AB39 and FST1.7, and arginine, phenylalanine, tyrosine and tryptophan by AB39 and FST2.11. In this regard, FST1.7 depleted only two of the amino acids tested, while AB39 depleted up to six. The total consumption of FAA correlated neither with the FAN utilization (Table 3) nor with the amount of LA produced.

Amino acid Control AB39 FST1.7 FST2.11
Alanine 36.9 ± 1.5 b 82.3 ± 4.4 a 13.3 ± 3.2 c 75.9 ± 6.1 a
Arginine 47.6 ± 2.9 a < 6 b 44.6 ± 2.9 a < 6 b
Asparagine 32.0 ± 1.4 a 12.1 ± 0.8 c 15.1 ± 1.1 b 10.9 ± 0.3 c
Aspartic acid 27.5 ± 1.1 b 33.9 ± 1.7 a 10.6 ± 1.3 c 22.8 ± 3.5 b
Glutamic acid 22.2 ± 0.9 c 51.0 ± 2.9 b 9.1 ± 1.7 d 61.5 ± 5.3 a
Glutamine 41.5 ± 1.4 a <5 b <5 b <5 b
Glycine 11.3 ± 1.0 b 11.0 ± 0.9 b 4.9 ± 0.9 c 21.5 ± 1.6 a
Histidine 22.0 ± 1.7 a 26.5 ± 1.6 a 20.2 ± 1.2 a 21.1 ± 1.8 a
Isoleucine 23.4 ± 1.2 a 25.3 ± 1.2 a 8.9 ± 0.8 b 25.1 ± 1.7 a
Leucine 50.7 ± 3.0 a 37.6 ± 2.8 b 22.7 ± 2.5 c 40.8 ± 2.5 b
Lysine 30.2 ± 2.7 a 30.1 ± 3.0 a 13.3 ± 2.6 b 26.3 ± 7.3 a
Methionine 10.2 ± 0.3 a <10 a <10 a <10 a
Phenylalanine 41.7 ± 1.7 a <5 c 15.3 ± 2.8 b <5 c
Serine 23.6 ± 2.3 a <7 b <7 b 11.9 ± 1.1 b
Threonine 20.1 ± 1.0 a 10.6 ± 1.8 b 7.0 ± 0.6 c 13.6 ± 1.0 b
Tryptophan 14.0 ± 1.1 a <7 b 12.1 ± 1.1 a <7 b
Tyrosine 30.8 ± 1.6 a <6 c 11.9 ± 2.3 b <6 c
Valine 42.3 ± 1.4 a 44.9 ± 3.0 a 28.5 ± 3.6 b 45.3 ± 3.6 a
Total amino acids 518.5 ± 31.8 a 399.0 ± 36.1 b 271.4 ± 24.7 c 404.0 ± 38.1 b
  • For each amino acid, a different letter in each row denotes a significant difference at p < 0.05.

Effect of pH on life? Use buffer or HCl? How about lactic acid? - Biology

Solutions that contain a weak conjugate acid–base pair, such as those discussed in Section 17.1, resist drastic changes in pH when small amounts of strong acid or strong base are added to them. These solutions are called buffered solutions (or merely buffers). Human blood, for example, is a complex buffered solution that maintains the blood pH at about 7.4 (see the “Chemistry and Life” box on page 713). Much of the chemical behavior of seawater is determined by its pH, buffered at about 8.1 to 8.3 near the surface (see “Chemistry and Life” box on page 728). Buffered solutions find many important applications in the laboratory and in medicine (FIGURE 17.1).

FIGURE 17.1 Buffered solutions. For laboratory work, prepackaged buffered solutions can be purchased.

Composition and Action of Buffered Solutions

A buffer resists changes in pH because it contains both an acid to neutralize added OH – ions and a base to neutralize added H + ions. The acid and base that make up the buffer, however, must not consume each other through a neutralization reaction. These requirements are fulfilled by a weak acid–base conjugate pair, such as CH3COOH–CH3COO – or NH4 + –NH3. Thus, buffers are often prepared by mixing a weak acid or a weak base with a salt of that acid or base. The CH3COOH–CH3COO – buffer can be prepared, for example, by adding CH3COONa to a solution of CH3COOH. The NH4 + –NH3 buffer can be prepared by adding NH4Cl to a solution of NH3. By choosing appropriate components and adjusting their relative concentrations, we can buffer a solution at virtually any pH.

GIVE IT SOME THOUGHT

Which of these conjugate acid–base pairs will not function as a buffer:

To understand how a buffer works, let's consider one composed of a weak acid HX and one of its salts MX, where M + could be Na + , K + , or any other cation that does not react with water. The acid-dissociation equilibrium in this buffered solution involves both the acid and its conjugate base:

The corresponding acid-dissociation-constant expression is

Solving this expression for [H + ], we have

We see from this expression that [H + ] and, thus, the pH are determined by two factors: the value of Ka for the weak-acid component of the buffer and the ratio of the concentrations of the conjugate acid–base pair, [HX]/[X – ].

If OH – ions are added to the buffered solution, they react with the buffer acid component to produce water and X – :

FIGURE 17.2 Buffer action. The pH of an HF/F – buffered solution changes by only a small amount in response to addition of an acid or base.

This reaction causes [HX] to decrease and [X – ] to increase. As long as the amounts of HX and X – in the buffer are large relative to the amount of OH – added, the ratio [HX]/[X – ] does not change much and, thus, the change in pH is small.

If H + ions are added, they react with the base component of the buffer:

This reaction can also be represented using H3O + :

Using either equation, we see that the reaction causes [X – ] to decrease and [HX] to increase. As long as the change in the ratio [HX]/[X – ] is small, the change in pH will be small.

FIGURE 17.2 shows an HX/MX buffer consisting of equal concentrations of hydrofluoric acid and fluoride ion (center). The addition of OH – reduces [HF] and increases [F – ], whereas the addition of [H + ] reduces [F – ] and increases [HF].

GIVE IT SOME THOUGHT

a. What happens when NaOH is added to a buffer composed of CH3COOH and CH3COO – ?

b. What happens when HCl is added to this buffer?

Calculating the pH of a Buffer

Because conjugate acid–base pairs share a common ion, we can use the same procedures to calculate the pH of a buffer that we used to treat the common-ion effect in Sample Exercise 17.1. Alternatively, we can take an approach based on an equation derived from Equation 17.5. Taking the negative logarithm of both sides of Equation 17.5, we have

where [acid] and [base] refer to the equilibrium concentrations of the conjugate acid–base pair. Note that when [base] = [acid], we have pH = pKa.

Equation 17.9 is known as the Henderson–Hasselbalch equation. Biologists, biochemists, and others who work frequently with buffers often use this equation to calculate the pH of buffers. In doing equilibrium calculations, we have seen that we can normally neglect the amounts of the acid and base of the buffer that ionize. Therefore, we can usually use the starting concentrations of the acid and base components of the buffer directly in Equation 17.9.

SAMPLE EXERCISE 17.3 Calculating the pH of a Buffer

What is the pH of a buffer that is 0.12 M in lactic acid [CH3CH(OH)COOH, or HC3H5O3] and 0.10 M in sodium lactate [CH3CH(OH)COONa, or NaC3H5O3]? For lactic acid, Ka = 1.4 × 10 –4 .

Analyze We are asked to calculate the pH of a buffer containing lactic acid (HC3H5O3) and its conjugate base, the lactate ion (C3H5O3 – ).

Plan We will first determine the pH using the method described in Section 17.1. Because HC3H5O3 is a weak electrolyte and NaC3H5O3 is a strong electrolyte, the major species in solution are HC3H5O3, Na + , and C3H5O3 – . The Na + ion is a spectator ion. The HC3H5O3–C3H5O3 – conjugate acid–base pair determines [H + ] and, thus, pH [H + ] can be determined using the acid-dissociation equilibrium of lactic acid.

Solve The initial and equilibrium concentrations of the species involved in this equilibrium are

The equilibrium concentrations are governed by the equilibrium expression:

Because Ka is small and a common ion is present, we expect x to be small relative to either 0.12 or 0.10 M. Thus, our equation can be simplified to give

Solving for x gives a value that justifies our approximation:

Alternatively, we can use the Henderson–Hasselbalch equation to calculate pH directly:

PRACTICE EXERCISE

Calculate the pH of a buffer composed of 0.12 M benzoic acid and 0.20 M sodium benzoate. (Refer to Appendix D.)

Answer: 4.42

In Sample Exercise 17.3 we calculated the pH of a buffered solution. Often we will need to work in the opposite direction by calculating the amounts of the acid and its conjugate base needed to achieve a specific pH. This calculation is illustrated in Sample Exercise 17.4.

SAMPLE EXERCISE 17.4 Preparing a Buffer

How many moles of NH4Cl must be added to 2.0 L of 0.10 M NH3 to form a buffer whose pH is 9.00? (Assume that the addition of NH4Cl does not change the volume of the solution.)

Analyze We are asked to determine the amount of NH4 + ion required to prepare a buffer of a specific pH.

Plan The major species in the solution will be NH4 + , Cl – , and NH3. Of these, the Cl – ion is a spectator (it is the conjugate base of a strong acid). Thus, the NH4 + –NH3 conjugate acid–base pair will determine the pH of the buffer. The equilibrium relationship between NH4 + and NH3 is given by the base-dissociation reaction for NH3:

The key to this exercise is to use this Kb expression to calculate [NH4 + ].

Solve We obtain [OH – ] from the given pH:

Because Kb is small and the common ion [NH4 + ] is present, the equilibrium concentration of NH3 essentially equals its initial concentration:

We now use the expression for Kb to calculate [NH4 + ]:

Thus, for the solution to have pH = 9.00, [NH4 + ] must equal 0.18 M. The number of moles of NH4Cl needed to produce this concentration is given by the product of the volume of the solution and its molarity:

Comment Because NH4 + and NH3 are a conjugate acid–base pair, we could use the Henderson–Hasselbalch equation (Equation 17.9) to solve this problem. To do so requires first using Equation 16.41 to calculate pKa for NH4 + from the value of pKa for NH3. We suggest you try this approach to convince yourself that you can use the Henderson–Hasselbalch equation for buffers for which you are given Kb for the conjugate base rather than Ka for the conjugate acid.

PRACTICE EXERCISE

Calculate the concentration of sodium benzoate that must be present in a 0.20 M solution of benzoic acid (C6H5COOH) to produce a pH of 4.00.

Answer: 0.13 M

Buffer Capacity and pH Range

Two important characteristics of a buffer are its capacity and its effective pH range. Buffer capacity is the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree. The buffer capacity depends on the amount of acid and base used to prepare the buffer. According to Equation 17.5, for example, the pH of a 1-L solution that is 1 M in CH3COOH and 1 M in CH3COONa is the same as the pH of a 1-L solution that is 0.1 M in CH3COOH and 0.1 M in CH3COONa. The first solution has a greater buffering capacity, however, because it contains more CH3COOH and CH3COO – .

The pH range of any buffer is the pH range over which the buffer acts effectively. Buffers most effectively resist a change in pH in either direction when the concentrations of weak acid and conjugate base are about the same. From Equation 17.9 we see that when the concentrations of weak acid and conjugate base are equal, pH = pKa. This relationship gives the optimal pH of any buffer. Thus, we usually try to select a buffer whose acid form has a pKa close to the desired pH. In practice, we find that if the concentration of one component of the buffer is more than 10 times the concentration of the other component, the buffering action is poor. Because log 10 = 1, buffers usually have a usable range within ±1 pH unit of pKa (that is, a range of pH = pKa ± 1).

GIVE IT SOME THOUGHT

The Ka values for nitrous acid (HNO2) and hypochlorous (HClO) acid are 4.5 × 10 –4 and 3.0 × 10 –8 , respectively. Which one would be more suitable for use in a solution buffered at pH = 7.0? What other substance would be needed to make the buffer?

Addition of Strong Acids or Bases to Buffers

Let's now consider in a more quantitative way how a buffered solution responds to addition of a strong acid or base. In this discussion, it is important to understand that reactions between strong acids and weak bases proceed essentially to completion, as do those between strong bases and weak acids. Thus, as long as we do not exceed the buffering capacity of the buffer, we can assume that the strong acid or strong base is completely consumed by reaction with the buffer.

Consider a buffer that contains a weak acid HX and its conjugate base X – . When a strong acid is added to this buffer, the added H + is consumed by X – to produce HX thus, [HX] increases and [X – ] decreases. (See Equation 17.7.) Upon addition of a strong base, the added OH – is consumed by HX to produce X – in this case [HX] decreases and [X – ] increases. (See Equation 17.6.) These two situations are summarized in Figure 17.2.

To calculate how the pH of the buffer responds to the addition of a strong acid or a strong base, we follow the strategy outlined in FIGURE 17.3:

1. Consider the acid–base neutralization reaction and determine its effect on [HX] and [X – ]. This step is a stoichiometry calculation. (Section 3.6)

2. Use the calculated values of [HX] and [X – ] along with Ka to calculate [H + ]. This step is an equilibrium calculation and is most easily done using the Henderson–Hasselbalch equation.

FIGURE 17.3 Calculating the pH of a buffer after addition of an acid or base.

SAMPLE EXERCISE 17.5 Calculating pH Changes in Buffers

A buffer is made by adding 0.300 mol CH3COOH and 0.300 mol CH3COONa to enough water to make 1.000 L of solution. The pH of the buffer is 4.74 (Sample Exercise 17.1). (a) Calculate the pH of this solution after 5.0 mL of 4.0 M NaOH(aq) solution is added. (b) For comparison, calculate the pH of a solution made by adding 5.0 mL of 4.0 M NaOH(aq) solution to 1.000 L of pure water.

Analyze We are asked to determine the pH of a buffer after addition of a small amount of strong base and to compare the pH change with the pH that would result if we were to add the same amount of strong base to pure water.

Plan Solving this problem involves the two steps outlined in Figure 17.3. First we do a stoi-chiometry calculation to determine how the added OH – affects the buffer composition. Then we use the resultant buffer composition and either the Henderson–Hasselbalch equation or the equilibrium-constant expression for the buffer to determine the pH.

Solve (a) Stoichiometry Calculation: The OH – provided by NaOH reacts with CH3COOH, the weak acid component of the buffer. Prior to this neutralization reaction, there are 0.300 mol each of CH3COOH and CH3COO – . The amount of base added is 0.0050 L × 4.0 mol/L = 0.020 mol. Neutralizing the 0.020 mol OH – requires 0.020 mol of CH3COOH. Consequently, the amount of CH3COOH decreases by 0.020 mol, and the amount of the product of the neutralization, CH3COO – , increases by 0.020 mol. We can create a table to see how the composition of the buffer changes as a result of its reaction with OH – :

Equilibrium Calculation: We now turn our attention to the equilibrium for the ionization of acetic acid, the relationship that determines the buffer pH:

Using the quantities of CH3COOH and CH3COO – remaining in the buffer, we determine the pH using the Henderson–Hasselbalch equation. The volume of the solution is now 1.000 L + 0.0050 L = 1.005 L due to addition of the NaOH solution:

(b) To determine the pH of a solution made by adding 0.020 mol of NaOH to 1.000 L of pure water, we first determine the concentration of OH – ions in solution,

[OH – ] = 0.020 mol/(1.005 L) = 0.020 M

We use this value in Equation 16.18 to calculate pOH and then use our calculated pOH value in Equation 16.20 to obtain pH:

Comment Note that the small amount of added NaOH changes the pH of water significantly. In contrast, the pH of the buffer changes very little when the NaOH is added, as summarized in FIGURE 17.4.

FIGURE 17.4 Effect of adding a strong base to a buffered solution and to water.

PRACTICE EXERCISE

Determine (a) the pH of the original buffer described in Sample Exercise 17.5 after the addition of 0.020 mol HCl and (b) the pH of the solution that would result from the addition of 0.020 mol HCl to 1.000 L of pure water.

Answers: (a) 4.68, (b) 1.70

CHEMISTRY AND LIFE
BLOOD AS A BUFFERED SOLUTION

Chemical reactions that occur in living systems are often extremely sensitive to pH. Many of the enzymes that catalyze important biochemical reactions, for example, are effective only within a narrow pH range. For this reason, the human body maintains a remarkably intricate system of buffers, both within cells and in the fluids that transport cells. Blood, the fluid that transports oxygen to all parts of the body, is one of the most prominent examples of the importance of buffers in living beings.

Human blood has a normal pH of 7.35 to 7.45. Any deviation from this range can have extremely disruptive effects on the stability of cell membranes, the structures of proteins, and the activities of enzymes. Death may result if the blood pH falls below 6.8 or rises above 7.8. When the pH falls below 7.35, the condition is called acidosis when it rises above 7.45, the condition is called alkalosis. Acidosis is the more common tendency because metabolism generates several acids in the body.

The major buffer system used to control blood pH is the carbonic acid–bicarbonate buffer system. Carbonic acid (H2CO3) and bicarbonate ion (HCO3 – ) are a conjugate acid–base pair. In addition, carbonic acid decomposes into carbon dioxide gas and water. The important enuilibria in this buffer system are

Several aspects of these equilibria are notable. First, although carbonic acid is diprotic, the carbonate ion (CO3 2– ) is unimportant in this system. Second, one component of this equilibrium, CO2, is a gas, which provides a mechanism for the body to adjust the equilibria. Removal of CO2via exhalation shifts the equilibria to the right, consuming H + ions. Third, the buffer system in blood operates at pH 7.4, which is fairly far removed from the pKa1 value of H2CO3 (6.1 at physiological temperatures). For the buffer to have a pH of 7.4, the ratio [base]/[acid] must be about 20. In normal blood plasma the concentrations of HCO3 – and H2CO3 are about 0.024 M and 0.0012 M, respectively. Consequently, the buffer has a high capacity to neutralize additional acid but only a low capacity to neutralize additional base.

The principal organs that regulate the pH of the carbonic acid–bicarbonate buffer system are the lungs and kidneys. When the concentration of CO2 rises, the equilibria in Equation 17.10 shift to the left, which leads to the formation of more H + and a drop in pH. This change is detected by receptors in the brain that trigger a reflex to breathe faster and deeper, increasing the rate at which CO2 is expelled from the lungs and thereby shifting the equilibria back to the right. When the blood pH becomes too high, the kidneys remove HCO3 – from the blood. This shifts the equilibria to the left, increasing the concentration of H + . As a result, the pH decreases.

Regulation of blood pH relates directly to the effective transport of O2 throughout the body. The protein hemoglobin, found in red blood cells (FIGURE 17.5), carries oxygen. Hemoglobin (Hb) reversibly binds both H + and O2. these two substances compete for the Hb, which can be represented approximately by the equilibrium

Oxygen enters the blood through the lungs, where it passes into the red blood cells and binds to Hb. When the blood reaches tissue in which the concentration of O2 is low, the equilibrium in Equation 17.11 shifts to the left and O2 is released.

During periods of strenuous exertion, three factors work together to ensure delivery of O2 to active tissues. The role of each factor can be understood by applying Le Châtelier's principle to Equation 17.11:

1. O2 is consumed, causing the equilibrium to shift to the left, releasing more O2.

2. Large amounts of CO2 are produced by metabolism, which increases [H + ] and causes the equilibrium to shift to the left, releasing O2.

3. Body temperature rises. Because Equation 17.11 is exothermic, the increase in temperature shifts the equilibrium to the left, releasing O2.

In addition to the factors causing release of O2 to tissues, the decrease in pH stimulates an increase in breathing rate, which furnishes more O2 and eliminates CO2. Without this elaborate series of equilibrium shifts and pH changes, the O2 in tissues would be rapidly depleted, making further activity impossible. Under such conditions the buffering capacity of the blood and the exhalation of CO2 through the lungs are essential to keep the pH from dropping too low, thereby triggering acidosis.

RELATED EXERCISES: 17.29 and 17.95

FIGURE 17.5 Red blood cells. A scanning electromicrograph of red blood cells traveling through a small branch of an artery.


Selection of Suitable Buffer Mixtures

There are two useful rules of thumb for selecting buffer mixtures:

    A good buffer mixture should have about equal concentrations of both of its components. A buffer solution has generally lost its usefulness when one component of the buffer pair is less than about 10% of the other. Figure (PageIndex<4>) shows an acetic acid-acetate ion buffer as base is added. The initial pH is 4.74. A change of 1 pH unit occurs when the acetic acid concentration is reduced to 11% of the acetate ion concentration.

Figure (PageIndex<4>): The graph, an illustration of buffering action, shows change of pH as an increasing amount of a 0.10-M NaOH solution is added to 100 mL of a buffer solution in which, initially, ([ce] = 0.10: M) and (ce<[CH3CO2^<->]>=0.10:M).
  1. Weak acids and their salts are better as buffers for pHs less than 7 weak bases and their salts are better as buffers for pHs greater than 7.

Blood is an important example of a buffered solution, with the principal acid and ion responsible for the buffering action being carbonic acid, H2CO3, and the bicarbonate ion, (ce). When an excess of hydrogen ion enters the blood stream, it is removed primarily by the reaction:

[ce(aq)+ce(aq)⟶ce(aq)+ce(l)] When an excess of the hydroxide ion is present, it is removed by the reaction: [ce(aq)+ce(aq)⟶ce(aq)+ce(l)] The pH of human blood thus remains very near 7.35, that is, slightly basic. Variations are usually less than 0.1 of a pH unit. A change of 0.4 of a pH unit is likely to be fatal. Buffers Maintain the pH of Intracellular and Extracellular Fluids

A growing cell must maintain a constant pH in the cytoplasm of about 7.2 –𠁗.4 despite the production, by metabolism, of many acids, such as lactic acid and CO2, which reacts with water to form carbonic acid (H2CO3). Cells have a reservoir of weak bases and weak acids, called buffers, which ensure that the cell’s pH remains relatively constant. Buffers do this by “soaking up” H + or OH − when these ions are added to the cell or are produced by metabolism.

If additional acid (or base) is added to a solution of an acid (or a base) at its pKa value (a 1:1 mixture of HA and A − ), the pH of the solution changes, but it changes less than it would if the original acid (or base) had not been present. This is because protons released by the added acid are taken up by the original A − form of the acid likewise, hydroxyl ions generated by the added base are neutralized by protons released by the original HA.

This ability of a buffer to minimize changes in pH, its buffering capacity, depends on the relationship between its pKa value and the pH. To understand this point, we need to recognize the effect of pH on the fraction of molecules in the undissociated form (HA). The titration curve for acetic acid shown in Figure 2-21 illustrates these relationships: at one pH unit below the pKa of an acid, 91 percent of the molecules are in the HA form at one pH unit above the pKa, 91 percent are in the A − form. Thus the buffering capacity of weak acids and bases declines rapidly at more than one pH unit from their pKa values. In other words, the addition of the same number of moles of acid to a solution containing a mixture of HA and A − that is at a pH near the pKa of the acid will cause less of a pH change than it would if the HA and A − were not present or if the pH were far from the pKa value.

Figure 2-21

The titration curve of acetic acid (CH3COOH). The pKa for the dissociation of acetic acid to hydrogen and acetate ions is 4.75. At this pH, half the acid molecules are dissociated. Because pH is measured on a logarithmic scale, the solution changes from (more. )

All biological systems contain one or more buffers. Phosphoric acid (H3PO4) is a physiologically important buffer phosphate ions are present in considerable quantities in cells and are an important factor in maintaining, or buffering, the pH of the cytosol. Phosphoric acid

Figure 2-22

The titration curve of phosphoric acid (H3PO4). This biologically ubiquitous molecule has three hydrogen atoms that dissociate at different pH values thus, phosphoric acid has three pKa values, as noted on the graph. The shaded areas denote the pH ranges — within (more. )

In nucleic acids, phosphate is found as a diester linked to two carbon atoms of adjacent ribose sugars:

The pKa for the dissociation of the single —OH proton is about 3, which is similar to the pKa for the dissociation of the first proton from phosphoric acid. Therefore, each phosphate residue in deoxyribonucleic acid (DNA) or ribonucleic acid (RNA) is dissociated and carries a negative charge at neutral pH, which is why DNA and RNA are called nucleic acids:


Discussion

The level of NADP-dependent GDH activity was markedly strain dependent ( Tanous et al. 2002 Williams et al. 2004, 2006 ). Some strains of Lact. plantarum, Lact. rhamnosus DPPMA19 and Lact. parabucknerii B48F3 were the most active. First, this study showed the presence of NADP-GDH activity in species such as Lact. curvatus and Lact. parabucknerii ( Helinck et al. 2004 ). Anyway, the real potential of NADP-GDH activity has to be assayed under temperature, pH and NaCl conditions of cheese ripening ( Kieronczyk et al. 2004 Williams et al. 2006 ). The experimental design used to study the interactive effect of environmental parameters on enzyme activity had been previously adopted by other authors ( Laan et al. 1998 Curtin et al. 2001 ) for other microbial enzymes involved in cheese ripening. Overall, the values of temperature and pH tested negatively affected the NADP-GDH activity of all strains. On the contrary, the effect of NaCl on NADP-GDH activity was positive, negative or neutral depending on the strain. In addition, the level of adaptation of NADP-GDH to temperature, pH and NaCl conditions of cheese ripening markedly varied among the lactobacilli strains. Lactobacillus plantarum DPPMA13 and, especially DPPMA49, were the only strains retaining relatively high NADP-GDH activity (30–100% of the enzyme activity found under optimal conditions) under values of temperature, pH and NaCl found during cheese ripening. On the contrary, the most NADP-GDH active strain, Lact. plantarum DPPMA63, revealed the lowest enzyme adaptability to values of temperature, pH and NaCl found during cheese ripening. The results found in this study could be useful to explain why some strains showing an in vitro GDH activity did not increase the formation of important flavour compounds when used in cheese model system ( Kieronczyk et al. 2004 ). High numbers of Lact. plantarum were found in Italian (e.g., Pecorino), Spanish (e.g., Manchego, Cabrales and Roncal), Portuguese (e.g., Picante), Greek (e.g., Feta), British, Irish and United States (e.g., Cheddar) cheeses ( Gobbetti et al. 2007 ). Nevertheless, during growth under the hostile conditions of cheese ripening, the expression of stress related proteins and enzymes may markedly vary ( Wouters et al. 2000 De Angelis and Gobbetti 2004 ). To the best of our knowledge, no information about the regulation of the expression of GDH gene in LAB under cheese-like conditions (temperature, pH and NaCl) exists in literature. Such studies on the expression of GDH gene were only reported for Prevotella ruminicola ( Wen and Morrison 1996 ), Neisseria meningitides ( Pagliaruolo et al. 2004 ) and Debaryomyces hansenii ( Alba-Lois et al. 2004 ). The partial GDH sequences found in this study showed high identity with GDH sequences from Lact. plantarum WCFS1, Lact. sakei, Lact. casei BL23, Lact. lactis and Strep. thermophilus. Overall, high variability was found between the partial GDH sequences of strains belonging to the species Lact. plantarum especially for DPPMA49. Compared to the other strains, the highest activity shown by DPPMA49 even at low temperature, pH and high concentration of NaCl could be related to the different GDH amino acid sequence. RT-PCR analysis revealed that GDH expression of Lact. plantarum DPPMA49 was down-expressed by low temperature (<13°C) and over-expressed by NaCl (1·87–5·62%). The same effect of NaCl was also reported for D. hansenii ( Alba-Lois et al. 2004 ). GDH was not expressed under conditions of 4°C, pH 6·0 and 3·75% NaCl (Table 4). This could be apparently in contrast with GDH activity (1·30 U) measured using the cytoplasm fraction of Lact. plantarum DPPMA49 under the same conditions (Table 2). However, such a discrepancy could be explained taking into account that the cytoplasm fraction derived from a 24- h-old culture grown under optimal conditions (MRS, 30°C), and therefore it may be hypothesized that the GDH had been expressed under those conditions. In other words, the enzyme activity reported in Table 2 could be a residual activity. On the contrary, the RT-PCR analysis (Table 4) had been performed with cells growing under conditions of 4°C, pH 6·0 and 3·75% NaCl.

When used as adjunct starter in cheese-making, Lact. plantarum DPPMA49 showed a high count of viable cells after 60 days of ripening, thus demonstrating a high capacity to tolerate acidic conditions and starvation. As previously shown ( Tanous et al. 2002 Williams et al. 2004, 2006 ), the addition of α-ketoglutarate or NADP-GDH active LAB strains as adjunct starter increased the FAA catabolism. Overall, catabolic pathways of amino acids by LAB contribute to the production of ATP directly or by proton translocation, thereby reducing the amount of ATP needed for proton balancing ( Konings 2002 Teusink et al. 2006 ). Recently, Teusink et al. (2006) reported that the transamination and degradation of aromatic and branched-chain amino acids can generate ATP by proton-motive force-driven transhydrogenase reaction. The transhydrogenase is a highly favourable reaction, as it regenerates catabolic NADP and produces NADPH for biosynthesis. Anaerobic oxidation of NADH allows less ethanol formation and hence more acetate with concomitant ATP production ( Teusink et al. 2006 ). Other authors ( Higuchi et al. 1997 ) have demonstrated that whole cells of Lactobacillus sp. were capable of synthesizing ATP in the presence of glutamate. According to the aforementioned considerations, Lact. plantarum DPPMA63 showed a lower level of cell density in cheese ripened for 60 days compared to DPPMA49. An increased cell density of DPPMA63 was found in cheese added with α-ketoglutarate. According to previous studies ( Yvon et al. 1998 Banks et al. 2001 ), the amount of VOC of cheeses made with the commercial starter, the commercial starter and DPPMA49 or DPPMA63, with or without the addition of α-ketoglutarate differed. The highest level of alcohols, aldehydes, miscellaneous and carboxylic acids was found in cheeses made with DPPMA49. No further increase in VOC was found when α-ketoglutarate was added to cheeses made with this adjunct. On the contrary, α-ketoglutarate strongly increased the concentration of VOC in cheese made with the commercial starter and, especially, in cheese made with DPPMA63. The highest level of VOC derived from catabolism of branched-chain and aromatic amino acids found in cheese made with DPPMA49 confirms the importance of GDH for flavour formation in cheese. For instance, VOC such as 3-methyl-1-butanal (cheesy, chocolate, malt), 3-methyl-1-butanoic acid (cheesy, sweaty) and phenylacetaldehyde (floral) are some of the key odour compounds in hard and semi-hard cheese varieties (for reviews see Ardö 2006 Gobbetti et al. 2007 ).

The findings of this study contribute to the knowledge about enzymes involved in the catabolism of amino acids, to be used as an important trait for cheese strains selection. From an industrial point of view, the use of strains favouring the catabolism of FAA could represent an economic vantage, because it would increase cheese flavour and/or shorten the period of ripening.



Comments:

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  2. Bruce

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  3. Shaktirr

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