Lecture 9_Acids and Bases (Nov 2024) PDF

Lecture 9: Acids and Bases Senior Lecturer Division of Education, School of Biomedical Sciences, Faculty of Medicine, CUHK Email: [email protected] Email: [email protected] Email: [email protected] Office number: 3943 6795 Important Notice...

Lecture 9: Acids and Bases Senior Lecturer Division of Education, School of Biomedical Sciences, Faculty of Medicine, CUHK Email: [email protected] Email: [email protected] Email: [email protected] Office number: 3943 6795 Important Notice: These slides contain copyright materials. Access is limited to CUHK students of MEDF1011, unless otherwise specified. 1 Copyright © 2018 The Chinese University of Hong Kong  LECTURE OUTLINE The importance of a stable physiological pH range The CO2/bicarbonate buffer system is the most important chemical buffer system in the extracellular (ECF) fluid Conceptual difference between pH and pKa How to choose a suitable chemical buffer The Henderson-Hasselbalch (H-H) equation and its significance to homeostatic control of ECF pH Types of acid-base imbalances and their general compensatory responses Copyright © 2024 The Chinese University of Hong Kong 2 Learning outcomes To recall the normal physiological range of ECF pH To give an account on why humans are called acid generators To summarise various sources of acids in the body To name the three methods that help undergo acid base balance To reproduce the modified Henderson-Hasselbalch (H-H) equation and indicate the normal ratio of HCO3− : H2CO3 To write down the equation for the equilibrium between carbon dioxide, carbonic acid and the bicarbonate/hydrogen ion To describe how a suitable buffer can be chosen based on the understanding of 3 important criteria To construct a table that summarises the four simple types of acid-base imbalance showing the changes in pH, concentration of bicarbonate ion and partial pressure of carbon dioxide Copyright © 2024 The Chinese University of Hong Kong 3 Pre-class micro-module 8 on Blackboard 4 Scenario: A male subject has his blood analyzed for bicarbonate (HCO3−), oxygen, and carbon dioxide. Which of these molecules is a chemical buffer? ✓ Dietary intake affects blood pH ✓ 3 bodily methods to regulate blood pH ✓ Excessive acid can be removed 5  pH is a measure of hydrogen ion concentration [H+] It is defined as the negative logarithm to base 10 of free (unbound) hydrogen ion (H+) concentration – pH decreases as acidity increases pH = − log10 [H+] pH scale pH: The lowercase letter ‘p’ refers to the power the uppercase letter ‘H’ refers to the hydrogen ion (also referred to as a proton). Log10: The pH is the logarithmic scale in which any change in an integer value will change the H+ concentration by the factor of 10. [H+]: The square bracket designates concentration 6 The pH varies between the values ranging between 0-14 Higher [H+] = more acidic Lower [H+] = less acidic = lower pH = higher pH 7 The importance of a stable physiological pH Normal pH of plasma arterial blood: 7.35 –7.45 with the average of 7.4. i.e. slightly alkaline pH affects all protein functions and biochemical reactions, so it is tightly regulated within a narrow range Complications with acid-base disturbance (only *will be explained): *1. Conformational change in protein structure (denaturation) 2. Changes in potassium (K+) balance 3. Changes in excitability of neurons and muscles 4. Cardiac arrhythmia 5. Vasodilation of arterioles 8 Consequence of abnormal pH: conformational change in protein structure (denaturation) Slide recall: Lecture 1 (The ingredients of life) Normal protein Denatured protein Acids and bases can disrupt bonding (e.g. hydrogen bonding) formed between the side chains of amino acids, leading to denaturation which is the loss of structural or conformational changes of protein from the native structure without alteration of the amino acid sequence. – Loss of structure means loss of function 9 Disorders of acid-base balance A person who has a blood pH below 7.35 is considered to be in acidosis A person who has a blood pH above 7.45 is considered to be in alkalosis 10 Our body is an acid generator Most H+ are produced by daily metabolism occurring in body A general rule of thumb: 1 mmol H+ generated/kg body weight/day Sources of H+ gain Sources of H+ loss Generation of H+ from CO2 (volatile acid) Loss of H+ in urine through the kidneys Primarily in the form of H2PO4− and NH4+ HPO42− + H+ H2PO4− NH3 + H+ NH4+ Production of non-volatile (fixed) acids from the Loss of CO through the lungs 2 metabolism of proteins and other organic molecules 11 Definitions of acids and bases Acid is a substance that releases free H+ in solution (forms acidic solution) Base is a substance that accepts free H+ in solution (forms basic solution) Most acids we encounter in physiology (except gastric acid) are weak acids. Examples of strong acid is hydrochloric acid, and weak acids are carbonic acid and lactic acid 12  Dissociation of a weak acid is a *reversible reaction When a reaction is in its chemical **equilibrium, the ratio between [reactants] and [products] is constant – For example, the general equation for the dissociation of a weak acid in water, where HA is the parent acid and A− is its conjugate base HA H+ + A− Parent acid Hydrogen Conjugate (reactant) ion base (product) (product) *A reversible reaction is a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously. **An equilibrium is the state of a reversible chemical reaction when the rate of the reaction in the reverse direction equals that in the forward direction. 13 Acid dissociation constant (Ka) is an equilibrium constant The acid dissociation constant (Ka) Ka measures the extent to which an acid HA H + + A− dissociates in an aqueous solution, and thus a quantitative measure of the strength of an acid in a solution. [H+] [A−] A large Ka means it is a strong acid Ka = [HA] A small Ka means it is a a weak acid The larger the pKa, the weaker the acid. pKa = − log [Ka] The smaller the pKa, the stronger the acid. 14 Summary: Conceptual difference between pH and pKa pH pKa pH = − log [H+] pKa = − log [Ka] Describes the strength of an acid Describes the acidity of a (i.e. its ability to donate H+). solution. A low pKa means it is a strong A low pH means it is acidic and acid and a high pKa means it is a a high pH means it is alkaline. weaker acid. 15 Buffering of hydrogen ion in the body Any substance that can reversibly bind to H+ ions is called a buffer. Buffer + H+ HBuffer Buffering does not eliminate H+ from the body or add it to the body, it only keeps H+ “locked up” until balance can be restored and thus the ultimate homeostatic regulators are done by the lungs and the kidneys The lungs regulate CO2 The kidneys regulate total body H+ balance 16 There are 3 important criteria to be called a good chemical buffer: 1. To minimise change in pH 2. During acid loading, it can accept the H+ 3. During acid removal, it can donate the H+ The major extracellular buffering system is CO2 / HCO3− system The major intracellular buffers are proteins and phosphates 17 Examples of chemical buffer: Protein buffer Proteins are made up of amino acids Some amino acids have side chains which are either 1. A weak acid group (carboxyl group- COOH) 2. A weak base group (amino group- NH2) Aspartic acid containing carboxyl Lysine containing side chain amino side chain Proteins serve as zwitterion / dipolar ion that bear two types of charged groups (i.e. cations and anions) 18 Examples of chemical buffer: Protein buffer Intracellular proteins are most plentiful and powerful buffers; plasma proteins also important Protein molecules are amphoteric meaning that they can function as both weak acid and weak base at the same time Hemoglobin functions as intracellular buffer in red blood cell. 19 Examples of chemical buffer: Phosphate buffer Action nearly identical to bicarbonate buffer Components are sodium salts of: – Dihydrogen phosphate (H2PO4–), a weak acid H2PO4− HPO42− + H+ – Monohydrogen phosphate (HPO42–), a weak base HPO42− + H+ H2PO4− Unimportant in buffering plasma as concentration is not high enough Effective buffer in urine and intracellular fluid (ICF), where phosphate concentrations are high 20 Examples of chemical buffer: the CO2 /bicarbonate buffer system The most important extracellular buffer system The rest step is rate limiting and is very slow unless catalysed in both directions by the enzyme carbonic anhydrase Carbonic acid (H2CO3) dissociates rapidly into H+ and HCO3− under the catalysis of the enzyme. Carbonic anhydrases forms dissociates Increased CO2 drives the reaction to the right, which generates H+ (handled by the kidneys), whereas increased H+ drives the reaction to the left, generating CO2 (which is exhaled from the lungs) 21 There are 3 bodily ways to undergo acid base balance Concentration of H+ ions regulated sequentially by: seconds minutes hours 1. Chemical buffer system 2. Respiratory system 3. Renal system Gives moderate changes in Acts within minutes Most potent but require pH through the coordination hours to days to take effect Rapid and first line of of the brainstem (i.e. slowest acting) defense (in seconds) Monitor through CO2 Directly excreting or Combines with or releases Accounts for most acid/ reabsorbing H+ through H+ base disturbances indirect change in the rate at Cellular proteins, which HCO3– buffer is phosphate ions, and reabsorbed or excreted hemoglobin 22 The Henderson-Hasselbalch (H-H) equation [H+] [A−] Ka = [HA] [A−] log10Ka = log10 [H+] + log10 [HA] [A−] − log10 [H+] = − log10 Ka + log10 [HA] pH = pKa + log10 [A−] [HA] 23 The modified Henderson-Hasselbalch (H-H) equation The modified H-H equation describes the relationship of pH, bicarbonate (HCO3−), and PCO2 The pKa for the CO2 /HCO3− buffer system in human body is 6.1 Partial pressure of CO2 N.B. A value of 0.03 is the conversion factor of PCO2 (mmHg) to [CO2] (mM) 24 Partial pressure of gas molecules Partial pressures (Px) is the pressure exerted by each gas in a container. Total pressure in a system is the sum of all partial pressures exerted by different gases Patm = atmospheric pressure at sea level = PN2 + PO2 + PH20 + PCO2 = 760mmHg 25 Arterial blood gas (ABG) analysis An arterial blood gas (ABG) tests blood taken from an artery and it assesses the patient's partial pressures of oxygen (PaO2) and carbon dioxide (PaCO2) PaO2 gives information on the oxygenation status, and PaCO2 on the ventilation status Example: PaCO2 can be affected by hyperventilation or hypoventilation, and acid-base status 26 Homeostatic regulation of CO2 and HCO3− Both the respiratory and renal systems work together to maintain a normal physiological pH of 7.35 to 7.45 with a ratio of: Controlled by kidneys Controlled by lungs Ratio of HCO3− : H2CO3 HCO3− H2CO3 20 1 This ratio gives us a pH value of 7.4 27 Important terms to indicate plasma [H+] The arterial H+ concentrations can be abnormal in a variety of clinical conditions An increase in [H+] or a decrease in plasma pH is called acidemia A decrease in [H+] or an increase in plasma pH is called alkalemia Processes that tend to raise [H+] is called acidosis Processes that tend to lower [H+] is called alkalosis 28 Types of acid base imbalance A pH  7.35 is acidemia. A pH  7.45 is alkalemia. The human body experiences 4 main types of acid-based disorders – An abnormal change in PCO2 will cause a respiratory disorder – An abnormal change in HCO3− will cause a metabolic disorder Respiratory acidosis Metabolic acidosis Respiratory alkalosis Metabolic alkalosis If a disorder occurs, the body induces a counterbalance (called compensatory mechanism/response) in the form of an opposite condition (a negative feedback). 29 Types of acid-base imbalance Kidneys Lungs Types pH [HCO3− ] PCO2 Metabolic acidosis Metabolic alkalosis Respiratory acidosis Respiratory alkalosis = primary change = compensation 30 Respiratory acidosis Hypoventilation Begin (Chronic Obstructive Pulmonary Disease- bronchitis, emphysema) Occurs in hypoventilation Arterial blood – decreased ratio of alveolar ventilation to  PCO2,  [H+],  pH CO2 production Begin Respiratory acidosis leads to increase in PCO2 Compensatory response Plasma pH drops Primary change Renal compensation will stimulate kidneys to reabsorb more (and sometimes also produce) HCO3− 31 Respiratory alkalosis Hyperventilation Begin (anxiety, attack) panic Occurs in hyperventilation Arterial blood – increased ratio of alveolar ventilation to CO2  PCO2,  [H+],  pH production Begin Respiratory alkalosis leads to decrease in PCO2 Compensatory response Plasma pH increases Primary change Renal compensation stimulate kidneys to excrete more HCO3− 32 Metabolic acidosis Acid (non-volatile) load to the body and buffered by HCO3− – Anaerobic respiration during exercise produces lactic acid begin Metabolic acidosis leads to decrease in plasma [HCO3−] Primary change Plasma pH drops Compensatory response Lowered pH (i.e. acidic pH) stimulates rate of ventilation and expire more CO2 PCO2 drops and this is called a respiratory compensation 33 Metabolic alkalosis Loss of acid (e.g. vomiting) or alkali load to body and buffered by H2CO3 start Metabolic alkalosis leads to increase in Loss of gastric juice rich in H+ plasma [HCO3−] Primary change Plasma pH increases Compensatory response Increased pH (i.e. alkaline pH) depresses ventilation rate and retain CO2 PCO2 increases and this is called a respiratory compensation 34 Post-class micro-modules: Case study 1 and 2 35 Required reading: Basic Concepts in biomedical sciences I Chapter 5, page 99-120 36 36  The end of this lecture Copyright © 2024 The Chinese University of Hong Kong 37