Acids, Bases, and Salts - NSG 741: Genetics, Chemistry, and Physics of Anesthesia

Acids, Bases, and Salts NSG 741: Genetics, Chemistry, and Physics of Anesthesia School of Nursing Who cares? Overview Many biological and drug molecules have acidic and or basic properties. In the patient, acidic and basic propertie...

Acids, Bases, and Salts NSG 741: Genetics, Chemistry, and Physics of Anesthesia School of Nursing Who cares? Overview Many biological and drug molecules have acidic and or basic properties. In the patient, acidic and basic properties of a drug are important for two reasons. Water solubility (lipophilic-alkanes, benzene; hydrophilic- alcohols, amines) Binding to the site of action An acid is a molecule that donates a proton. A base is a molecule that accepts a proton. An acid that has donated a proton is a conjugate base. A base that has accepted a proton is a conjugateSchool acid. of Nursing Who cares? Overview An acid can accept a pair of electrons. A base can donate a pair of electrons. pH is a value between 0 and 14. A strong acid fully dissociates into H+ and its conjugate base. A weak acid does not fully dissociate. The pKa puts a number to a molecule's acidity or basicity – how weak of an acid or base. Knowing the pH and pKa allows you know if the drug will be more ionized or unionized. School of Nursing Chemical Equilibria Le Châtelier’s (le-SHOT-lee-ay) Principle States that equilibrium is a good thing, and nature strives to attain and/or maintain a state of equilibrium. Changing Concentration If you add products, the equilibrium will shift toward reactants. If you remove products, the equilibrium will shift toward products. Changing Temperature Increasing temperature favors endothermic processes. Changing Volume and Pressure Significantly impacts equilibrium reactions only when at least one of the reactants or products is a gas. Decreasing volume increases pressure. School of Nursing Chemical Equilibria The Equilibrium Constant A system is in a state of equilibrium when there is a balance between reactants and products. This balance is defined by thermodynamic parameters, namely bond strengths and the intermolecular forces between all the molecules in the system. The equilibrium constant (K) is the numerical description of that balance. School of Nursing Chemical Equilibria As K increases, the reaction tends to increasingly favor products and the forward reaction becomes more favorable. As K decreases, the reaction tends to increasingly favor starting materials, and the reverse reaction becomes more favorable. Solids and Liquids Pure solids or liquids comprise a different phase from where reactions in aqueous media occur. The concentration of solids, liquids, and water (as a solvent) do not appear explicitly in the equilibrium constant expression. Reversing a Reaction When you reverse the equation for a chemical reaction, Kforward is the reciprocal of Kreverse. School of Nursing Acids and Bases Definition of Acids and Bases Arrhenius definition (most operational definition) Acid–species that increases the hydronium ion (H3O+) concentration in an aqueous solution. Base–species that increases the hydroxide ion (OH–) concentration in an aqueous solution. Brønsted definition (most generally useful definition) Acid–species that donates a hydrogen ion (H+) to a base. Base–species that accepts a hydrogen ion (H+) from an acid. School of Nursing Acids and Bases Conjugate Acid-Base Pairs The charge on the conjugate acid is always one greater than the charge on its conjugate base. When an acid gives away its hydrogen ion to a base, the acid is converted into its conjugate base. When a base accepts a proton from an acid, the base is converted into its conjugate acid. In generic form, we can express this process as: HA + B → A− + BH+ Amphiprotic Species Can behave as either an acid or a base. School of Nursing Acids and Bases Strong Acids Very determined to foist their proton off onto a base. Essentially 100% ionized when dissolved in water, but this is not an equilibrating process, so all of the starting materials are converted into products. Strong acids are relatively rare. Strong Bases In water, the strongest possible base is the hydroxide ion, OH−. A strong base ionizes essentially 100% to produce the OH − ion, so a strong base is a soluble ionic hydroxide. School of Nursing Acids and Bases Weak Acids Are able to donate hydrogen ions to bases but are less determined to do so than strong acids. When a weak acid dissolves in water, it establishes a dynamic equilibrium between the molecular form of the acid and the ionized form. Weak Bases Can accept hydrogen ions from acids but are less determined to do so than strong bases. Do not completely ionize in water to produce an equivalent concentration of the hydroxide ion, because when a weak base dissolves in water, it establishes a dynamic equilibrium between the molecular form and the ionized form. School of Nursing Acids and Bases Polyprotic Acids A diprotic acid has two hydrogen ions to donate, so a diprotic acid can behave as an acid twice. A triprotic acid has three hydrogen ions to donate. The number of acidic protons is not necessarily the number of hydrogens in the molecular formula. The acidic hydrogen is bonded to a highly electronegative oxygen atom. The O—H bond is polarized toward the oxygen to a point that a base can snatch away the hydrogen as an H+ ion from the acetic acid molecule. The other three hydrogens are bonded to a carbon atom, School of and Nursing carbon and hydrogen have almost identical electronegativities, Acids and Base Strength Acid/Base Strength of Conjugate Acid–Base Pairs The stronger the acid, the weaker its conjugate base. The stronger the base, the weaker its conjugate acid. General guidelines The conjugate base of a really strong acid has no base strength. The conjugate base of a weak acid has base strength. The conjugate acid of a weak base has acid strength. School of Nursing Acid–Base Reactions Involve a transfer of a hydrogen ion from the acid to the base In order to predict the products of an acid–base reaction: Identify which is the acid and which is the base. Move an H+ ion from the acid to the base, converting the acid into its conjugate base and the base into its conjugate acid. Any acid–base reaction has two acids and two bases: One acid and one base on the reactant side. Conjugate acid and conjugate base on the product side. The base almost always has a lower (more negative) charge than the acid. School of Nursing The reaction equilibrium always favors the formation of the Measuring Acidity: The pH Function The p in pH is a mathematical operator that means the negative logarithm of, and the H in pH means hydrogen ion concentration, so the definition of pH is pH = −log[H+]. A logarithm function is a way to map a vast range of values onto a much smaller set of values. pH values have no intrinsic units—logarithms represent “pure numbers”. Each change of 1 pH unit means the hydrogen ion concentration is changing by a factor of 10, so small changes in pH correspond to much larger changes in acidity level. School of Nursing Measuring Acidity: The pH Function Water has some very weak acid–base properties, and therefore sets some limits on the parameters of the pH scale. A tiny fraction of water molecules dissociates or ionizes into a hydrogen ion and a hydroxide ion: H2O ⇄ H+ + OH−. In pure water, the concentrations of the H+ and OH− ions are equal: [H+] = [OH−] in pure water The pH of pure water is 7.00. School of Nursing Measuring Acidity: The pH Function Relationship Between pH and pOH Because pH and pOH are derived from the ionization of water, there is a fixed relationship between them. The pH plus the pOH of any aqueous solution (at 25°C) always adds up to 14.00: 14.00 = pH + pOH pKa and pKb The equilibriums constants for acids and bases are commonly presented as p-functions. The pKa and pKb of a conjugate acid–base pair sum to give School of Nursing 14.00: Other Acidic Species Nonmetal oxides dissolve in water to give acid solutions. The most physiologically important example is carbon dioxide. Buildup of carbon dioxide in the blood results in acidosis. In cellular tissue, where the carbon dioxide concentration is relatively high, the increased acidity slightly alters the structure of hemoglobin and facilitates the release of oxygen. School of Nursing Other Acidic Species Carbon dioxide is a nonmetal oxide because it is a compound composed of a nonmetal and oxygen. Nonmetal oxides are sometimes called acid anhydrides because they are produced by stripping water from an acid. When carbon dioxide dissolves in water, it combines with a water molecule to give carbonic acid: CO2 + H2O ⇄ H2CO3. When the carbonic acid forms, it dissociates according to its acid strength: H2CO3 ⇄ H+ HCO−3. So, when CO2 dissolves in water, the pH drops. School of Nursing Buffers A pH buffer is a solution that resists changes in pH. It contains a weak acid (HA) and its conjugate base (A−) or a weak base and its conjugate acid. If a strong base is added to a buffered solution, the weak acid in the buffer HA reacts with the hydroxide ion to give water and the weak base A−. HA + OH− ⇄ H2O + A− This results in converting a strong base OH− into a weak base A−. If a strong acid is added to a buffered solution, the weak base in the buffer (A− ) reacts with the H+ ion to give HA. School of Nursing A− + H+ ⇄ HA Henderson–Hasselbalch Equation (Estimates pH of the buffer system) This equation predicts that when a basic drug is placed into a relatively more acidic environment, the conjugate acid will predominate. If the molecule is a weak base and pH is > the pKa of the drug, then the unionized fraction predominates. If the molecule is a weak acid and the pH is < the pKa of the drug, then the unionized fraction predominates. If the molecule is a weak base and pH is < the pKa of the drug, then the ionized fraction predominates. If the molecule is a weak acid and the pH is > the pKa of the drug, then the ionized fraction predominates. As pKa get further away from physiologic pH, the degree of ionization increases. The closer the pKa is to the pH of the blood, the faster the onset. School of Nursing Henderson–Hasselbalch Equation School of Nursing School of Nursing Acid-Base Balance & Respiratory System The respiratory system plays an important role in maintaining normal pH balance within the body. It works along with the kidneys and the buffer systems to balance the acids and bases of the blood and other body tissues, allowing them to function normally. Hydrogen ions interact with negatively charged regions of other molecules, such as proteins, altering their structural conformation and in doing so altering their behavior. Blood pH alters the activity of various enzymes, thereby changing metabolic functions in all body tissues. School of Nursing Acid-Base Balance &Respiratory System In addition to the efforts of the respiratory system and kidneys to regulate pH levels, buffers in the human body maintain pH within the physiologic range. The buffers consist mainly of bicarbonate, phosphate, and proteins. From values for PaCO2 and HCO3 are used to determine if the disturbance is respiratory or metabolic. The respiratory system can rapidly compensate for metabolic acidosis or alkalosis by altering alveolar ventilation. The total body bicarbonate deficit is equal to the base deficit (in mEq/ L) that is obtained from the blood gas values. Complete correction of the base deficit is not indicated; only half of the calculated dose of bicarbonate is initially recommended. School of Nursing Human Buffer Systems Blood System: Normal: Bicarbonate buffer is the most important. pH – 7.35-7.45 Hemoglobin is the second PaCO2 – 35-45 most important. HCO3 – 22-26 Respiratory System: Altering ventilation to change PaCO2 is the key. Renal System: Reabsorption of filtered bicarbonate. Removal of titratable acids. Formation of ammonia. School of Nursing Anion Gap – Metabolic acidosis Anion gap helps determine the cause of the metabolic acidosis. Anion gap = Major cations – Major Anions or NA – Cl + HCO3 The value is usually 8-12 mEq/L. Accumulation of acid? > 12? Gap acidosis. Loss of bicarbonate or ECF dilution? Non-gap acidosis. School of Nursing Interpretation of Arterial Blood Gases Analysis of ABGs can provide useful information concerning the relationship of acid production and acid removal by the lungs and kidneys. Acid-base disturbances can be categorized into four major groups: respiratory acidosis, metabolic acidosis, respiratory alkalosis, and metabolic alkalosis. Each can be considered compensated under certain conditions. In spite of observing a compensatory change in CO2 or bicarbonate, an acid-base disorder is considered uncompensated if that mechanism has not been able to bring the pH back into a normal range. School of Nursing Acidosis Any process that leads to an elevation in PaCO2 tends to lower arterial pH, resulting in respiratory acidosis. An acute change in PaCO2 of 10 mm Hg is associated with a change in pH of 0.08 units. An increase in PaCO2 with a normal bicarbonate level is termed uncompensated respiratory acidosis. Metabolic acidosis should more properly be referred to as nonrespiratory acidosis because it does not always involve alterations in metabolism. Causes of this condition include ingestion (poisoning), infusion, production of a fixed acid (lactic acidosis), and decreased excretion of acid by the kidneys. A base change of 10 mEq/L is associated with a pH change of 0.15 unit. School of Nursing Alkalosis As the pH rises, this hyperventilation results in respiratory alkalosis. Metabolic alkalosis occurs when fixed acid loss is increased or when the intake of bases is abnormally high. School of Nursing School of Nursing Treatment of Blood Gas Abnormalities For the patient who is mechanically ventilated, respiratory acidosis and respiratory alkalosis can be treated with a simple increase or decrease in the amount of alveolar ventilation. To restore stable and spontaneous circulation, mild to moderate metabolic acidosis can be treated with hyperventilation and correction of shock. School of Nursing School of Nursing