Acid-Base Balance

Acid-Base Balance

Acidosis - Alkalosis

Serkan SAYINER, DVM PhD. Assist. Prof.

Near East University, Faculty of Veterinary Medicine, Department of Biochemistry [email protected]

Definition of pH

 It is the negative logarithm of [H⁺] ion concentration in a solution.

• pH = -log [ H⁺ ]

• Therefore a decrease in pH value indicates an increase in free hydrogen ion concentration (relatively acidic) and an increase in pH indicates a decrease in free hydrogen ion concentration

(relatively alkaline). pH is very tightly regulated in the body by a number of buffering systems.

 If the pH of a solution is less than 7, it is an acid. If it is

greater than 7, it is a base, if it is 7, it is a neutral solution.

Definition of pH

 Plasma [H +] is kept in narrow limits. Cells have defensive mechanisms against pH changes in the extracellular

environment.

Body Fluids pH Values

Plasma 7,38 – 7,44

Pancreatic liquid 7.5 – 8.00

Saliva 6.35 – 6.85

Gastric juice 0.9 – 1.6

Milk 6.6 – 6.9

Urine 4.8 – 7.5

Definition of pH

[H

] balance, in other words acid-base balance;

• The amount of H⁺ taken in the diet + endogenous metabolism is preserved as a result of mutual

balancing of the amount obtained and the amount taken from the body.

• Thus, the balance of the extracellular fluid (ECF) is kept within the physiological limits and viability is maintained.

• To ensure equilibrium;

• Volatile acids are removed by respiration (such as CO₂)

• H⁺ and HCO₃^- are removed or retained by the kidneys.

• It forms complex with non-volatile H⁺ chemical buffers and is discarded.

Biological Importance of pH

 Acidosis is an increased acidity in the blood and other body tissue (i.e. an increased hydrogen ion concentration.

 Alkalosis is the result of a process reducing hydrogen ion concentration of arterial blood plasma (alkalemia).

 A number of functions in the body are affected by the pH value.

1. Blood pH changes cause intracellular pH changes.

2. The ability of hemoglobin to bind oxygen is affected.

3. Intracellular pH changes alter the activities of enzymes.

4. Hydrogen bond between protein charges and protein molecules.

5. In acidosis, H⁺ enters the cell and K⁺ exits the cell.

6. Bone making and destruction

Causes of Blood pH Changes

 Blood pH is extremely important and varies within a very narrow range (7.38-7.44).

 Minor changes can be compensated. If not, it results in acidosis or alkalosis.

 The reasons for change are diverse.

• Nutrition

• Digestive secretions

• Vomiting

• Diarrhea

Source: Ecy.Wa

Buffer Systems and pH Regulation

A buffer is a solution whose pH changes very

little when acid or base is added. Most buffers are solutions composed of approximately equal amounts of a weak acid and the salt of its conjugate base.

Buffers are aqueous systems that tend to resist pH

changes when small amounts of acid (H

) or base

(OH

) are added.

Buffer Systems and pH Regulation

 Dilution

 Respiration

 Renal Mechanism

Buffer Systems

• In blood plasma: Bicarbonate-Carbonic acid, Phosphate- Phosphoric acid, Protein-Proteinate Buffer

• In erythrocytes: Hemoglobinate-Hemoglobin, Oxyhemoglobinate-Oxyhemoglobin

• Lymph, CSF, transudates: Bicarbonate, phosphate buffers

Buffer Systems and pH Regulation

 After respiration and metabolic activities, acids are added to extracellular enviroment.

• Respiration:

• Metabolism: Approximately 50-90 mEq of acid per day is added.

• The metabolism of neutral structures to organic acids:

Lactic, pyruvate, acetoacetic acid, beta-hydroxybutyric acid.

• Oxidation of S and P Compounds: Methionine, cysteine, hydrolysis of phosphoesters, degradation of nucleic acids.

• Dietary or drugs: Mineral or organic acids.

(CH₂O)_n + O₂ CO₂ + H₂O H₂CO₃ H⁺ + HCO₃^-

Buffer Systems and pH Regulation

 CARBONIC ACID-BICARBONATE (HCO₃^- : H₂CO₃)

• It's an important and notable buffer in extracellular fluids.

• Normally the ratio is 20:1 and pH is 7.4.

• High levels of CO₂ are produced every day as a result of oxidative metabolism. Despite not being an acid, it forms a carbonic acid as a result of reaction with

water.

• Carbonic anhydrase enzyme catalyzes this reaction.

• It is mainly found in red blood cells.

Buffer Systems and pH Regulation

 Protein Buffer System

• COOH or NH₂ groups,

• The biggest part of the buffers in the body,

• Albumin, globulins such as hemoglobin (Hb).

 HCO₃^- Buffer System

• Available in large quantities,

• Open system,

• The respiratory and kidney systems act on these buffer systems.

• The most important buffer of extracellular fluids.

Phosphate Buffer System

• Low in the extracellular medium, significant (especially muscle tissue) in the intracellular environment,

• Best buffer in kidney and bone.

Buffer Systems and pH Regulation

Defense againts [H

]

• Primary line is Buffers;

• Secondary defense line; Respiratory system (Lungs)

• Hyper- or hypoventilation in response to ΔpH or ΔpCO₂

• Tertiary defense line; Kidneys

• Reabsoprtion of HCO₃^- and exrection of H⁺

[HCO -]₃ [H₂CO₃]

; [HPO₄^-] [H₂PO₄^-]

; [Prot^-] [H×Prot]

Buffer Systems and pH Regulation

Respiratory System

• It is the second line of defense.

• It acts rapidly, but lasts 12-24 hours.

• H₂CO₃ is produced, converted to CO₂ and discarded by the lungs.

• Alveolar ventilation increases the pH, which has fallen.

• The respiratory system can not remove the bound acids.

• In metabolic acidosis, the rate of respiration increases, hyperventilation.

• In metabolic alkalosis, the rate of respiration decreases, hypoventilation.

• Bicarbonate buffer system interactions. At the lung level, hemoglobin binds oxygen; this creates a molecular change that favors dissociation of H⁺ from

hemoglobin. This will push the equilibrium to the left, yielding CO₂ and water that is expired. At the tissue level, metabolism yields CO₂ and considerable acid.

Hemoglobin releases O₂ to tissues and deoxygenated hemoglobin then binds H⁺. Thus, equilibrium is driven to the right.

Buffer Systems and pH Regulation

Regulation of acid-base balance in kidneys

• The kidneys help to regulate blood pH by removing H⁺ and reabsorbing HCO₃^-.

• It's the third line of defense.

• H⁺ is released through the exchange of Na⁺ in the tubular epithelial cells.

• Antiport mechanism

• Na⁺ and H⁺ move in opposite directions.

• Urine is normally mildly acidic, because the kidney

absorbs almost all of the HCO₃^- ions and removes H⁺.

• The pH of the blood returns to its normal value.

Buffer Systems and pH Regulation

Regulation of acid-base balance in kidneys

• HCO₃^- reabsorption

• The apical membranes of tubular epithelial cells are not permeable to HCO₃^- .

• It is indirectly reabsorbed.

• When the urine is acidic, HCO₃^- combines with H⁺ and H₂CO₃ is formed, which is catalyzed by carbonic

anhydrase located on the tubular cell membrane.

• When [CO₂] increases in the filtrate, the CO₂ diffuses into the tubule cell and forms H₂CO₃.

• H₂CO₃ then dissociates to HCO₃^- and H⁺ .

• The HCO₃^- than diffuses to peritubular capillary.

Acid-Base

Disturbances

Acidosis - Alkalosis

Acid-Base Disorders

ACIDOSIS

•Metabolic Acidosis

•Respiratory Acidosis

ALKALOSIS

•Metabolic Alkalosis

•Respiratory Alkalosis

Acidosis

Metabolic Acidosis

• Decrease in pH and HCO₃^- are seen.

• This is due to pathologic metabolic production of acid in the form of hydrogen ions or bicarbonate loss.

• In ECF, The increased hydrogen ions are buffered by combining with bicarbonate to form carbonic acid that then dissociates to CO₂ gas and water. The CO₂ is then rapidly eliminated from the system via respiration.

This is bicarbonate-carbonic acid buffer system.

• Protein and phosphate buffers serve at the ICF level.

• In the ECF, a cation shift is formed to prevent the increase in the H ion concentration. H enters, K exits.

• This can cause hyperkalemia. Even if the body store with renal or gastrointestinal losses is exhausted.

Acidosis

• Common examples of pathologic metabolism resulting in

metabolic acidosis include lactic acidosis, ketoacidosis, renal failure, and acid toxicities (e.g., ethylene glycol toxicity).

• Alternatively, bicarbonate may be lost from the system such as may occur with severe diarrhoea. By any of these

mechanisms, depletion of bicarbonate establishes metabolic acidosis.

• Metabolic acidosis is the most common acid-base

disturbance. This is attributed to the fact that dehydration and

poor tissue perfusion leading to lactic acid production is a process common to many primary internal medical disorders. Renal failure and diabetes mellitus, relatively common disorders in veterinary patients, also contribute to the incidence of metabolic acidosis.

• Anion gap increases in metabolic acidosis.

Acidosis

• Compensation

• The first, fast, and short-term effective response is to increase respiration (hyperventilation) and decrease pCO

.

• The kidney goes into effect for long time action.

The bicarbonate retention and H

excretion occur (in the form of ammonium).

• It is difficult to fully compensate in patients sufferd

from renal problems.

Acidosis

Respiratory Acidosis

• Characterized by a decrease in pH and an increase in pCO

.

• This is due to acute respiratory failure with accumulation of CO

.

• The first buffering usually occures intracellular.

• The main buffer of the ICF which is bicarbonate- carbonic acid can not buffer respiratory acidosis.

• There is a decrease in pO

before the increase of

pCO

.

Acidosis

• Causes include hypoventilation during anaesthesia or any pathologic cause of acute spontaneous

hypoventilation or severe impairment of gas exchange at the blood-lung interface.

• Most often; Acute upper respiratory tract obstruction, pneumonia, pneumothorax, chronic obstructive

pulmonary disease.

• Diseases or medicines affecting the respiratory center in brain.

• Volatile anesthetics used with closed systems.

• Positive pressure ventilator systems should be used and arterial gases should be monitored during anesthesia.

Acidosis

• Compensation

• Renal bicarbonate retention, increased hydrogen ion excretion.

• This effect can take several days.

• Increase in plasma bicarbonate level can be detected by increased renal hydrogen ion excretion.

Alkalosis

Metabolic Alkalosis

• Characterized by a increase in pH and HCO

.

• It can be seen at certain frequency in pets.

• It can be seen especially in ruminants suffered from digestive disorders.

• Excess H ions can be lost, due to bicarbonate retention.

• Contraction alkalosis: Decrease in the volume of

ECF results in a disproportionate loss of Na and Cl

ions compared to bicarbonate (such as vomiting).

Alkalosis

• Causes of H ion loss

• The most common cause of metabolic alkalosis is upper gastrointestinal obstruction. The hydrogen ions are

secreted into the stomach and are lost in the obstructive process, while bicarbonate is retained.

• Mineralocorticoid increase, excess diuretic use, low Cl diet.

• Na and Cl deficit are also seen in the circulation.

• Excessive use of bicarbonate.

• Most animals can tolerate it. Excessive use may not be tolerated because the effective circulation volume is reduced or K, Cl is insufficient.

Alkalosis

• In order to speak of alkalosis, it is necessary not only to form but also the factors that play a role in occurance must be participated.

• These factors disrupt renal bicarbonate excretion.

• Increase of bicarbonate reabsorption from the kidney triggers reduction of effective circulation volume, K and Cl depletion.

• Plasma bicarbonate level increases, Cl and K levels decrease.

• Maintenance of circulating effective volume is very important.

Alkalosis

• Compensation

• As the chemoreceptors in the respiratory center are activated, hypoventilation is formed as a

respiratory response. Thus, the amount of pCO

is increased.

Alkalosis

Respiratory Alkalosis

• Characterized by a increase in pH and decrease in pCO

.

• Causes

• Hyperventilation

• Hypoxia-related pulmonary diseases, congestive heart failure, severe anemia, neurological disorders,

salicylate intoxication, gram negative sepsis

• Physiological stress or pain can be a cause.

• It can also happen during evaporation in dogs.

Alkalosis

• Compensation

• The first way (in acute cases) is cellular buffering;

It is moderately reducing the amount of bicarbonate in ECF.

• Subsequently the renal bicarbonate reabsorption is reduced (response is completed within 2-3

days).

• Decrease in bicarbonate is compensated by Cl retention (for electrical neutrality).

• Hyperchloremia and decreased pCO₂ may be seen in compensated respiratory alcohol.

Alkalosis

• It may take several weeks for pH to return to normal levels in chronic cases.

• In dogs, the plasma bicarbonate concentration may be greatly reduced due to renal

compensation in chronic cases.

Approach to Acid-Base Disturbances

pH

<7.35 >7.45

Normal

Decreased, Acidemia Increased, Alkalemia

Evaluate HCO₃

Step 1

Step 2

Step 3, 4

Normal Low Low High High Normal

Evaluate pCO₂

High CO₂ Low pCO₂

Respiratory Acidosis

Metabolic Acidosis

Metabolic Alkalosis

Respiratory Alkalosis

ALKALOSIS ACIDOSIS

pH compensated, by increasing pCO₂ pH compensated,

by decreasing pCO₂

Mixed acid-base disturbances

 Due to different acid-base imbalances, mixed acid-base disturbances can be seen together.

 Metabolic acidosis and alkalosis can be seen together and in some cases they can be seen together with respiratory

acidosis-alkalosis.

 In order to be able to evaluate these situations, it is necessary to understand the anion gap.

• The relationship between serum Na and Cl changes and the compensation limits of primary acid-base disturbances.

 Clinical findings and anamnesis should be assessed to determine these factors.

Mixed acid-base disturbances

 Considerations to evaluate possible mixed acid-base disturbance

• Is the compensatory response to primer acid-base disturbance excessive?

• Normal pH, a demonstration of the compensation? Or is it a mixed acid-base disturbance indicator?

• Primer response may normalize the pH, but this does not necessarily mean that equilibrium is achieved.

• A pH change in the anticipated opposite direction to a known primer disorder refers to a mixed disorder.

• In primer acid-base disturbances, bicarbonate and pCO₂ diverge in the same direction. If there is an opposite condition, there is a mixed disorder.

Mixed acid-base disturbances

 Although the acid-base imbalance is present in animals and is indicated in veterinary literature, it is often overlooked.

 Accurate evaluation of clinical and clinicopathological data is crucial to diagnose these conditions.

 In suspicious situations, attention should be paid, adequate time should be monitored and expected compensatory

responses should be assessed.

Anion Gap

Anion Gap = {[Na

] + [K

]} - {[Cl

] + [HCO

]}

 Normal values = 10-25 mEq/L – 8-25 mmol/L

 Adding potassium to the equation has a very low diagnostic effect. If calculation is made without adding potassium, the 4 mEq/L should be added to reference value.

 In many animals this value is very close.

 In horses, values may vary due to age.

• It is higher in foals.

Anion Gap

Anion deficit is particularly useful when P, total protein (albumin) is at normal levels.

• The anions (phosphate, sulphate, organic ions) not

involved in the calculation are balanced by cations (Ca, Mg and K) which are not included in the calculation.

Anion deficit can be used as a prognostic guide in the categorization of acid-base disturbances, together

with potential factors.

Anion Gap

 Causes of decrease in Anion Gap

• Increased cationic proteins (Polyclonal gammopathy, IgG)

• Hypoalbuminaemia

• Hyperchloremia acidosis (altered protein anionic equivalents)

• Serum K levels can be correlated with the cause of hyperchloremic acidosis in normal or low anion status.

• If gastrointestinal fluid loss is formed due to diarrhea or renal tubuler acidosis, hypokalemia is accompanied.

• Decrease in mineralocorticoid secretion or activity (Addison's) or renal dysfunction is accompanied by hyperkalemia.

• Laboratory error

Anion Gap

 Causes of increase in Anion Gap

• Metabolic Acidosis: It may rarely increase due to dehydration and alkalosis.

• The main cause of the increase is due to the accumulation of

metabolic acids due to an increase in the number of non-routinely measured anions.

• Seen in uremic acidosis due to accumulation of non-metabolize acids, lactic acidosis due to anaerobic exercise, excessive grain consumption, hypovolemic shock and ketoacidosis (with or without diabetes mellitus).

• Anion gap can also be used to evaluate mixed acid-base disturbances.

• This may be suspected if there is no change in the level of bicarbonate although the anion gap changed.

Anion Gap

 Laboratory error

• Due to excessive grain consumption in herbivores, the anion gap is increased due to accumulation of high D-lactic acid in the ECF.

• D-lactic acid can not be determined by the commonly used lactic acid method. Because the general methods determine the level of L-lactic acid produced in metabolism.

Bicarbonate and Total CO

 If respiratory disturbances are excluded, the metabolic component of the acid-base balance is indicated by

bicarbonate concentration.

 Bicarbonate is generally determined by plasma or serum Total CO₂ assay.

 Bicarbonate is about 95% of the total carbon dioxide

measured, and is a measure of metabolic changes in acid- base balance.

 Decreases in metabolic acidosis, increases in metabolic alkalosis.

 A full blood gas measurement in acid-base disturbances should be considered.

Base excess (BE)

 Calculated mathematically.

Henderson-Hasselbalch Equation

Base excess (BE)

 This value is calculated to account for the combined

bicarbonate and hemoglobin buffering capacity of blood.

• In general, reference value is between -2,0 and + 2,0 mmol/L.

• A positive abnormal value indicates an excess of base, or alkalosis. A negative value indicates the. magnitude of HCO₃ deficit in mmol/L in metabolic acidosis.

• The BE utility is for calculation of the amount of

bicarbonate replacement in fluid therapy formulae.

These calculations result in the target amount of

bicarbonate that would be administered to normalize the bicarbonate concentration and pH.

Bicarbonate dosage (mEq/L) = 0.3 x body weight (kg) x BE (mmol/L)

Kaneko et al., 2008

Your Questions?

Send to [email protected]

References

 Kaneko JJ, Harvey JW, Bruss ML, 2008. Clinical Biochemistry of Domestic Animals, 6th edi. Academic Press-Elsevier

 Karagül H, Altıntaş A, Fidancı UR, Sel T, 2000. Klinik Biyokimya.

Medisan, Ankara

 Prof. Dr. Arif ALTINTAŞ, Ders notları.

 Thrall MA, Weiser G, Allison RW, Campbell TW, 2012. Veterinary Hematology and Clinical Biochemistry, 2nd edi. Wiley-Blackwell

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