Biochemistry: Metabolism and Energy Production

Chemical Reactions and Metabolism

• Chemical reactions are the processes that allow cells to live.
• Most chemical reactions in cells involve taking energy from food and making it usable for bodily functions.
• Two forms of metabolism are anabolism (synthesis) and catabolism (breakdown).

Carbohydrate Metabolism and ATP Formation

• Carbohydrate metabolism and ATP formation occur in chapter 69.
• Fructose and galactose are rapidly converted to glucose.
• Glucose is the final common pathway for transporting carbohydrates into the cell.

Metabolism

• Metabolism refers to the chemical processes allowing cells to live.
• Metabolic reactions take energy from foods and make it usable for bodily functions.
• Anabolism involves building up molecules, and catabolism involves breaking down molecules.
•  Energy production involves proteins, carbohydrates oxidation, and fats.
• Energy utilization includes active ion transport, muscle contraction, and synthesis of molecules.

Reactions Worth Knowing

• Dehydration synthesis is an anabolic reaction where monomers form covalent bonds, releasing water.
• Hydrolysis is a catabolic reaction where polymers break down into monomers by adding water.

Energy Foods

• Fats, carbohydrates, and proteins can be oxidized by the cell to release large amounts of energy.
• The phosphate bond in ATP is a high-energy bond.
• Aerobic metabolism produces more ATP than anaerobic metabolism.

Coupled Reactions

• Chemical energy in foods must be released slowly.
• Energy is linked with cellular processes, such as membrane pumps, protein synthesis, and muscle contraction.

ATP - Adenosine Triphosphate

• ATP is an important intermediate compound in many coupled reactions.
• ATP is present throughout the cell.
• The amount of energy per mole of ATP is 12,000 calories.

Glucose (Dextrose)

• After absorption from the gut, most fructose and galactose convert to glucose in the liver.
• Glucose is the final common pathway for carbohydrate transport into cells.

Transport of Glucose Through Cell Membrane

• Glucose cannot freely diffuse through cell membranes.
• It requires active co-transport or facilitated diffusion with a carrier protein.
• Insulin facilitates this process.

Phosphorylation

• Immediately upon entering cells, glucose combines with a phosphate radical to form glucose-6-phosphate.
• Glucose-6-phosphate can be used immediately for energy or stored as glycogen.

Glycogenesis

• Glycogenesis is the formation of glycogen, a starch.

Glycogenolysis

• Glycogenolysis is the breakdown of glycogen to glucose in the liver to make glucose available.
• Glycogen branches split off through phosphorylation via phosphorylase.
• Hormones like epinephrine and glucagon activate phosphorylase.

Release of Energy from Glucose

• Glycolosis, Mitochondria, Electron transport chain are all parts of ATP production from glucose.
• The ATP is produced is generated through multiple steps.

Glycolytic Pathway

• One gram-mole of glucose releases 686,000 calories of energy in the body.
• Only 12,000 calories of energy are necessary to form 1 gram-mole of ATP.
• Energy from glucose is released in packets yielding 38 moles of ATP per mole of glucose.
• Enzymes assist in the process to maximize ATP production.

Glycolysis

• Glycolysis is an important means of releasing energy from glucose.
• It is initiated by glucose lysing.
• It involves 10 steps.
• The end products of glycolysis are then oxidized.
• Glycolysis occurs in the cytoplasm.
• Glucose is split to form 2 molecules of pyruvic acid.

Glycolysis

• Net gain of ATP from one mole of glucose is 2 pyruvic acid molecules + 2 ATP molecules + 4 H.
• The efficiency of ATP formation is 43%.
• Two pyruvic acid molecules combine with Coenzyme A to form Acetyl-CoA, which will yield up to 6 ATP molecules later.

Citric Acid Cycle (Krebs Cycle)

• Reactions occur in the matrix of mitochondria.
• The cycle starts when acetyl-CoA combines with oxaloacetic acid to form citric acid.
• The cycle ultimately produces oxaloacetic acid.
• 24 hydrogen atoms are released from one glucose.

Citric Acid Cycle (Krebs Cycle)

• 20 hydrogen atoms bind with NAD+ via dehydrogenase.
• Water is added and carbon dioxide and hydrogen atoms released.
• Two molecules of ATP are produced.

Net Reaction per Molecule of Glucose in Citric Acid Cycle

• 2 Acetyl-CoA + 6 H2O + 2 ADP → 4 CO2 + 16 H+ + 2 CoA + 2 ATP

Oxidative Phosphorylation

• Oxygen is necessary in the mitochondria because NADH is split and e- are released from hydrogen oxidation.
• Electrons then enter the electron transport chain.
• The generated energy is stored as a proton gradient that is subsequently converted to ATP.
• This process is called chemiosmosis.
• Electron transport chain + chemiosmosis = oxidative phosphorylation.

Summary of ATP Formation

• Glycolysis makes 4 ATP, using 2, for a net gain of 2 ATP.
• Citric acid cycle produces 2 ATP.
• Electron transport produces 34 ATP.
• Total ATP production is 38 ATPs.
• Complete oxidation of 1 mole of glucose produces 456,000 calories stored as ATP and 686,000 available.
• This process is 66% efficient.
• Feedback control mechanisms control ATP formation.

Anaerobic Glycolysis

• This process is less efficient than aerobic glycolysis and can produce lactic acid.
• Used to monitor cellular oxygen use via lactic acid levels.
• Lactic acid can be converted to glucose when oxygen is available.

Storage of Glucose

• Glycogen stores are filled first.
• Remaining excess stores as fat.

Gluconeogenesis

• Synthesize glucose from fats (glycerol) and proteins (amino acids).
• Low blood glucose triggers this process.
• Cortisol releases to stimulate the adrenal cortex to release amino acids for liver glucose conversion.
• This involves catabolic metabolism.

Lipid Metabolism

• Lipids are the most efficient energy source storage for the cells having the highest energy content.
• Lipids include triglycerides, phospholipids, and cholesterol.

Fatty Acids

• Fatty acids are simple, long-chain hydrocarbon organic (carboxylic) acids.
• They contain a carboxyl group (-COOH).
• Palmitic acid has the generalized formula CH3(CH2)14COOH.

Triglycerides

• Triglycerides are three long-chain fatty acids bonded to one glycerol molecule.
• Glycerol is a triple alcohol with three -OH groups.

Absorption of Fats

• Dietary triglycerides are broken down to monoglycerides and fatty acids in the intestines.
• They are packaged as new triglycerides called chylomicrons.
• Chylomicrons are absorbed into the lymphatic system.
• Blood is slightly cloudy for 1 hour after a high-fat meal.

Uptake Into Cells

• Lipoprotein lipase in the fat and liver cells breaks down triglycerides into fatty acids and glycerol.
• Fatty acids diffuse into the cells.
• Once inside, fatty acids make triglycerides and are stored.

Transport Through Body

• When lipids are needed, fat cells convert triglycerides to fatty acids and glycerol.
• These components enter the blood and combine with albumin to form free fatty acids.

Fat Deposits

• Adipose tissues are called fat deposits or simple tissue fat.
• Fats are also stored in the liver.
• Fat cells in adipose tissue are 80-95% triglycerides.

Use of Triglycerides for Energy

• Up to 50% of the calories in typical American diets come from fats.
• Triglycerides are broken down to fatty acids and glycerol.
• Glycerol converts to glycerol-3-phosphate, which enters the glycolytic pathway for glucose breakdown.
• Fatty acids enter mitochondria and are oxidized.

Beta-Oxidation

• In several steps, fatty acids are converted to Acetyl-CoA, a process called beta-oxidation.
• Acetyl-CoA enters the citric acid cycle.
• Tremendous amounts of ATP are formed via beta-oxidation of fatty acids.
• Complete oxidation of stearic acid produces 148 molecules of ATP (2 used); net gain is 146 ATP molecules.

Ketosis

• When carbohydrates are unavailable for the body, fats are oxidized for energy.
• High concentrations of B-hydroxybutyric acid, acetoacetic acid, and acetone are created and called ketone bodies.
• Starvation, diabetes mellitus, and high-fat, low-carbohydrate diets can cause ketosis.

Regulation of Fat Utilization

• Catecholamines (epinephrine and norepinephrine) activate triglyceride lipase.
• Increased rate of fatty acid utilization is associated with corticotropin release from the anterior pituitary.
• Glucocorticoid release from the adrenal cortex also increases rate of fatty acid use.
• Decreased insulin secretion also increases rate of fatty acid use.

Phospholipids

• Phospholipids contain a fatty acid molecule, a phosphoric acid radical, and a nitrogenous base.
• Three types: lecithins, cephalins, and sphingomyelins.

Uses of Phospholipids

• Cell membranes (important throughout the body)
• Lipoprotein constituents
• Clotting factor (thromboplastin)
• Myelin sheath components (sphingomyelin)
• Phosphate radical donors.

Uses of Cholesterol

• Liver production of cholic acid for fat digestion (bile)
• Precursor for adrenocortical hormones
• Precursor for estrogen, progesterone, and testosterone
• Component of the skin's waterproof outer layer (corneum).

Atherosclerosis

• A disease in large arteries marked by fatty deposits (atheromatous plaques).
• Arteriosclerosis involves similar thickening and stiffening of vessel walls.
• Accumulation of cholesterol is a primary element.
• Connective tissue grows in plaques, causing stiffening (sclerotic).
• Plaque buildup can block vessels (occlusion) and cause rupture.

Factors that Lead to Atherosclerosis

• Sedentary lifestyle, obesity
• Diabetes mellitus
• Hypertension
• Hyperlipidemia
• Smoking

Prevention of Atherosclerosis

• Low-fat diet.
• No smoking
• Exercise
• Control blood pressure
• Control blood glucose
• Oat bran (for bile binding)
• Statins to inhibit HMG-CoA reductase (for cholesterol production inhibition)

Protein Metabolism

• Proteins make up 3/4 of the body's solids.
• Proteins are assembled of 20 amino acids linked by peptide bonds and stabilized through hydrogen bonding.
• Protein functions include structural components, enzymes, oxygen transport, production of nucleoproteins, muscle contraction, and cellular functions.

Regulatory Protein

• Regulatory proteins transmit chemical signals within the cell.
• G protein-coupled receptors transfer messages via the cell membrane.
• Ligands bind to receptors and activate G proteins.
• Downstream events are called second-messenger events, and these messages result in cellular actions.

Transport of Amino Acids

• Protein digestion in the gastrointestinal tract yields amino acids.
• Protein digestion and absorption take 2-3 hours.
• Amino acids enter cells through active transport or facilitated diffusion.
• Amino acids lost in the urine occur when renal threshold limits are exceeded.

Storage of Amino Acids

• Amino acids are immediately incorporated into new proteins in cells.
• When blood amino acid levels are low, amino acids are transported out of cells to replenish the supply.
• A reversible relationship exists between amino acid concentration and protein concentrations.
• Cancer cells actively use amino acids to build proteins, resulting in depletion in other cells.

Major Plasma Proteins

• Albumin regulates colloid osmotic pressure.
• Globulins have functions in enzymes and the immune system.
• Fibrinogen plays a role in blood coagulation.

Dietary Amino Acids

• 10 essential amino acids cannot be synthesized and must be obtained from the diet.
• 10 nonessential amino acids are synthesized in the body but may still be necessary in the diet.

Use of Proteins for Energy

• Protein degradation occurs in the liver.
• Deamination removes the amino group (-NH2) from amino acids.
• Aminotransferases are the enzymes for this process.

Urea

• Ammonia formed through deamination is converted to urea for blood removal.
• Ammonia is a neurotoxin.
• Urea is synthesized in the liver.
• Urea is excreted by the kidneys.

Oxidation of Deaminated Amino Acids

• The resulting keto acid from deamination is converted to a substance that enters the citric acid cycle.
• This substance is oxidized, like acetyl-CoA, to produce ATP.
• Ketogenesis converts amino acids to keto acids or fatty acids.

Obligatory Degradation of Proteins

• 20-30g of protein is degraded to amino acids and oxidized regardless of dietary intake.
• Consuming less protein than this results in deprivation.
• Carbohydrates are known as protein sparers because they reduce protein degradation for energy.

Hormonal Regulation of Protein Metabolism

• Growth hormone increases amino acid uptake and promotes protein synthesis.
• Insulin accelerates the transport of amino acids into cells.
• Glucocorticoids increase the breakdown of extrahepatic proteins to increase available amino acids.
• Testosterone increases protein deposition in muscle tissue.
• Thyroxine increases the rate of metabolism (catabolism or anabolism).

Aqueous Solutions & Concentrations: Solubility, Diffusion, and Osmosis

• A solute is a dissolved substance, and a solvent is the dissolving liquid.
• Solutions are homogenous mixtures where components are uniformly dispersed.
• Phase boundaries separate regions of a mixture with different chemical or physical properties.

Solution Concentrations: Molarity

• Molarity describes moles of solute per liter of solution and is affected by temperature.
• Molality describes moles of solute per kilogram of solvent and isn't affected by temperature.

Solution Concentrations: Molality

• Solutions measured in volume are commonly reported as molarity.
• Molality, when used in chemical analyses, requires calculations as it's based on quantities of solute and solvent.
• Molality is less common than molarity but important in physical chemistry applications.
• Molality and molarity are different, with the difference reducing in dilute solutions.
• Converting them involves determining the solution density.

Solution Concentrations: Percent

• Percentage by weight to volume is expressed in grams of solute per 100 ml of solution for clinical applications such as medicine dosages.
• Percentage by weight is similar to weight-to-weight concentrations, differing in measuring the solute in grams and the solution in grams.

Solution Concentrations

• Percentage by weight is frequently used in topical treatments and is analogous to percent weight to volume.
• Percent-by-volume is rarely used in analytical laboratories due to volume non-additivity.

Solution Concentrations: Normality and Equivalents

• Equivalents (Eq) are analogous to moles, while normality specifies concentration as equivalents of solute per liter of solution.
• Normality relates to the chemical reactivity of a substance, representing how many equivalents of a solute exists in one liter of solution.
• Normality can be challenging because of the ambiguous nature of describing chemical reactivity.

Solution Concentrations: Parts per Million (ppm)

• Parts per million (ppm) is used for extremely dilute solutions, comparing solute quantity to one million grams of solution.
• A safe exposure level for halogenated anesthetics is less than 2 parts per million (ppm) over an hour.
• When nitrous oxide is present with halogenated gases, nitrous oxide should be controlled below 25 parts per million (ppm).

Epinephrine

• Epinephrine concentrations are measured differently; 1 mg ampule = 1000 mcg per ml, making 10 ml of the ampule = 100 mcg per ml.
• Solutions of varying ratios are used for clinical applications.

How do you make up 25 ml of 2% lidocaine with 1:250,000 epinephrine?

• Take 10 ml of a 1:1000 epinephrine solution and add saline to yield a 10,000 dilution.
• Add 1 solution ml to 24 ml lidocaine.

Solubility

• Some solutes are more soluble than others in a solvent.
• Increased water solubility allows a medication to quickly get into the bloodstream.
• The solubility of a substance depends on its intermolecular interactions, temperature, and pressure.
• Polar solutes are typically more soluble in polar solvents
• Nonpolar solutes are generally soluble in nonpolar solvents.

Solubility

• Saturated solutions contain the maximal quantity of a solute.
• Supersaturated solutions contain more solute than they (the solvent) can hold.
• In equilibrium situations, excess solute crystallizes, leaves the solution, or produces a gas.
• Liquids that completely mix are considered miscible.

Effect of Temperature on Solubility

• The solubility of most solids and liquids increases with temperature increase, but gas solubility decreases as temperature increases.
• Increased temperature represents increased kinetic energy for dissolved gas molecules, leading to their escape from the solvent.
• Lower temperatures reduce molecular movement, allowing more gas to dissolve.
• Clinical applications involve hypothermic patients who retain anesthetic gases in solution.

Energy Changes & the Solution Process

• Energy change accompanies dissolution of a solute in a solvent that's often a noticeable temperature change.
• Heat of solution or enthalpy of solution corresponds to the energy change when dissolving a substance.
• Enthalpy of solution is described as the energy change associated with dissolving one mole of solute in a solvent.
• The heat of solution equals the enthalpy change when pressure is constant.

Energy Changes & the Solution Process

• Solution processes are either endothermic or exothermic.
• Exothermic processes result in temperature increase as energy flows from the system into the surroundings.
• Endothermic processes cause a decrease in temperature as energy flows in from the surroundings.
• The relative magnitudes of lattice energy and heat of solution (solvation) determine if the process is endothermic or exothermic.

Energy Changes & the Solution Process

• Energy is absorbed to break bonds. Energy is released when new bonds are formed.
• Whether a reaction is endothermic or exothermic depends on the difference in energy absorbed to break the bonds and the energy released when new bonds are formed.
• More heat release on bond formation equals exothermic reactions. More heat absorbed on bond formation equals endothermic reactions.
• In terms of temperature, heat increase in endothermic reactions equals adding more reactants to the system. In exothermic reactions, heat increase equals adding more products.

Energy Changes & the Solution Process

• Solutions are described as endothermic when the solubility of a substance increases with a rise in temperature.
• Solutions are described as exothermic when the solubility of a substance decreases with a rise in temperature.
• Le Chatelier’s Principle describes how equilibria shift to counter changing conditions, like temperature changes, emphasizing the equilibrium restoring reaction direction.

Factors Affecting Solubility

• Pressure increases the solubility of a gaseous solute in a liquid solvent.
• Solids and liquids are not easily compressed, so pressure change affects gas solubility, not solids and liquid solubility.
• Gas solubility in a liquid is directly proportional to the pressure above the surface; this is known as Henry's Law.
• Solubility (S) = k_H x P_gas (where k_H = Henry's Law constant and P_gas = partial pressure of the gas.)

Henry's Law

• Henry's Law in anesthesia calculations determine the amount of dissolved 02 and CO2 in the blood.
• At a constant temperature, a gas's solubility in a liquid is directly proportional to its partial pressure above the liquid.
• Increasing the gas pressure above the liquid increases the gas concentration and speed of gas uptake in the blood.

Henry's Law

• Blood solubility for 02 is 0.003 ml/100 ml blood/mmHg partial pressure.
• To find the amount of dissolved 02, multiply the partial pressure of 02 by 0.003.

Henry's Law

• The amount of CO2 that dissolves in blood is 0.067 ml/100 ml blood/mmHg.
• To find dissolved CO2, multiply the partial pressure of CO2 by 0.067.

Colligative Properties of Solutions

• Vapor pressure of a liquid is determined by the most energetic molecules at the surface escaping into the gas phase.
• Introducing solute molecules decreases the escape sites for solvent molecules.
• Solution vapor pressure is thus lower than pure solvent vapor pressure.
• Raoult’s law states the vapor pressure of a solution component is equal to the vapor pressure of the pure substance times the mole fraction of the component.

Henry's Law vs. Raoult's Law

• Henry's law describes gas solubility in the presence of a gas above the liquid, as it describes the pressure in the gas phase above the liquid.
• Raoult's law helps to describe nonvolatile solutes in a solvent, such as water, helping to determine the vapor pressure of the solvent itself.

Colligative Properties of Solutions: Boiling Point

• The boiling point is the temperature at which a substance's vapor pressure equals the ambient pressure.
• Solution boiling points increase with increasing solute concentration.
• Boiling point elevation is directly proportional to the molal concentration of solute particles.

Colligative Properties of Solutions: Freezing Point

• The freezing point is the temperature at which the liquid phase of the substance is in equilibrium with its solid phase.
• Solute particles interfere with molecule settling into orderly crystalline lattice structures, requiring lower temperatures to equilibrate.
• The freezing point of a solution is thus less than pure solvent freezing point.
• Increasing solute concentration decreases the freezing point of a substance.

NaCl Example

• Increasing salt concentration to water increases the boiling point and decreases the freezing point.

Colligative Properties of Solutions: Diffusion

• Diffusion is the net movement of molecules from an area of high concentration to an area of low concentration.
• Diffusion is driven by the tendency of a system to minimize concentration differences.
• Rate decreases with molecular weight.

Colligative Properties of Solutions: Diffusion & Membranes

• Diffusion through space or across permeable membranes is determined by factors like concentration gradient, tissue area, fluid tissue solubility, inversely proportional to membrane thickness and molecular weight of the molecules.

Colligative Properties of Solutions: Osmosis

• Osmosis is the movement of water across a semipermeable membrane to reach equilibrium in concentration gradient.
• Tonicity is the relative solute concentration in osmotic systems.
• Solutions with equal solute concentrations are isotonic.
• Osmosis occurs through semipermeable membranes that allow water but not solute passage.

Colligative Properties of Solutions: Osmotic Pressure

• Osmotic pressure is the force needed to prevent osmosis.
• Oncotic pressure is the osmotic pressure exerted in capillaries from plasma proteins and electrolytes, balancing hydrostatic pressure in pushing water out of capillaries.
• Normal oncotic pressure is ~28 mmHg.

Colligative Properties of Solutions: Osmotic Pressure

• Osmotic pressure (π) results from water concentration gradients equalizing.
• Osmotic pressure of a solution increases with solute concentration.
• Most capillary walls are permeable to solutes so don't affect osmotic pressure.
• Albumin (MW=69,000) is an essential determiner of intravascular volume due to its impermeability to capillary walls.

Diffusion and Anesthesia

• Diffusion (gas or liquid movement) is a passive process driven by the tendency to equilibrium (entropy).
• Diffusion rates vary depending on the medium.
• Nitrous oxide diffuses readily into air-filled spaces. Delivery in pneumothorax or situations where cavity expansion is not desirable is contraindicated because nitrous oxide can cause tracheal mucosal damage and bowel distension.

Diffusion and Anesthesia

• Apneic oxygenation demonstrates diffusion benefits.
• Fick's Law describes gases crossing biological tissues, relating diffusion to factors like pressure gradient, membrane solubility, membrane area, membrane thickness, and gas molecular weight.
• When nitrous oxide is delivered and then discontinued, low inspired concentration and diffusion hypoxia can occur.

Fick's Law of Diffusion

• Fick's law of diffusion helps determine diffusion rate, relating it to pressure difference, area, solubility, and membrane thickness.
• Diffusion hypoxia, COPD and slow gas induction, cardiac output calculation, and N20 volume/pressure increase in gas spaces of the body are all related to Fick's Law.
• Graham's law emphasizes how smaller molecules diffuse faster. This concept and others are related to differing gas flow and density.

Need to Know

• Although CO2 is larger than O2, CO2 diffuses 22 times faster across alveolar and capillary membranes due to increased solubility in fluids.
• Equilibration of gases in the body occurs when partial pressure is the same everywhere.
• Oxygen and other medicine delivery across body tissues (such as in the fetus) is described as simple diffusion.
• Simple diffusion requires a difference in gas partial pressure.

Need to Know

• Main factors that affect non-gas diffusion across membranes are concentration gradient, charge concentration, lipid solubility, and molecule size.
• Agents poorly penetrating the blood-brain barrier or placental barrier are mainly lipid-insoluble, and large, ionized agents.
• Ions, such as Na, do not penetrate lipid bilayers because they are hydrophilic, not lipophobic.

Colloids

• Colloids have one phase uniformly dispersed in another but are not true solutions because particles are larger than molecules (10-200 nm).
• Colloids are stable due to controlled conditions, and particles are large enough to scatter light.
• Milk, blood, paints, and jellies are examples.

pH Scale

• The pH scale ranges from 0-14 and measures acidity/basicity levels.
• Values between 0 - 7 describe increasing acid strength, with 7 neutral. Values between 7 - 14 describe increasing base strength.

Who cares? Overview

• Many biological and chemical molecules have acid/base properties.
• Water solubility is a key factor in drug action—hydrophobic drugs diffuse more readily in the body, while hydrophilic drugs are more soluble in aqueous solutions.
• Binding to action sites is a crucial factor for drug action, and ionization status is paramount in drug binding and effect duration.
• Acid molecules release H+, while base molecules accept H+.

Who cares? Overview

• Acidity can be described by an electron pair acceptance. Basicity is described by an electron pair donation.
• When considering acids or bases in aqueous solutions, pH is the negative logarithm of the hydrogen ion concentration.
• pKa is a number relating acid/base strength. It helps determine how ionized a substance will be in a solution.

Chemical Equilibria

• Equilibrium is a state where there's a balance between reactants and products in a chemical system.
• Equilibrium is controlled by thermodynamic parameters, like bond strength, and intermolecular forces within the system.
• The equilibrium constant (K) describes the equilibrium balance numerically.

Chemical Equilibria

• A rise in K indicates that the reaction moves more toward the products, and the forward reaction increases in favorability.
• A decline in K indicates a greater shift toward reactants and the reverse reaction exhibits higher favorability.
• Pure solids or liquids, present in reactions in aqueous solutions, do not influence the equilibrium constant.

Acids and Bases

• Arrhenius definitions for acids and bases focus on hydrogen ion (H⁺) and hydroxide ion (OH⁻) concentration increases in aqueous solutions.
• Brønsted definition details that acids donate H⁺ ions, while bases accept H⁺ ions.

Acids and Bases

• Conjugate acid-base pairs have a charge difference of +1.
• Conjugate pairs are created when an acid loses a proton or a base gains a proton.
• Generic forms illustrate this process (HA + B ⇄ A^- + BH⁺).
• Amphiprotic species can act as either an acid or a base.

Acids and Bases: Strong Acids

• Strong acids readily lose their protons to bases.
• Strong acids are approximately 100% dissociated in solution as they are highly determined to release protons, not necessarily an equilibration process.
• Strong acids are less common than weak acids.

Acids and Bases: Strong Bases

• Strong bases accept protons to form hydroxide ions (OH⁻) freely in water.
• Strong bases completely ionize, converting essentially 100% to OH⁻ ions in water.
• Strong bases include soluble hydroxide salts.

Acids and Bases: Weak Acids

• Weak acids are less determined to give up their hydrogen ions and form equilibria of molecular and ionized forms, resulting in partial ionization in solution.

Acids and Bases: Weak Bases

• Weak bases are less determined to accept hydrogen ions and maintain equilibrium between molecular and ionized forms, resulting in partial ionization in solution.

Acids and Bases: Polyprotic Acids

• Diprotic acids can donate two protons, while triprotic acids can donate three protons.
• The number of acidic protons in a compound may not equal the total number of hydrogens in the molecular formula.
• These protons are typically highly electronegative oxygen atoms due to their polarization toward the oxygen atom.

Acids and Base Strength

• Strong acids have weaker conjugate bases.
• Strong bases have weaker conjugate acids.
• Conjugate bases of strong acids have no base strength. Conjugate bases of weak acids have base strength. Conjugate acids of weak bases have acid strength.

Acid-Base Reactions

• Acid-base reactions involve the transfer of a hydrogen ion (H⁺) to convert acid into its conjugate base or a base into its conjugate acid.
• For predictions on the products of acid-base reactions, identify the acid and base and transfer the H⁺ ion. The acid becomes its conjugate base, the base becomes its conjugate acid.

Measuring Acidity: The pH Function

• pH is the negative logarithm of the hydrogen ion (H⁺) concentration.
• pH = -log[H⁺].
• pH values have no intrinsic units and use logarithms to represent pure numbers, aiding in mapping out a vast range of values.
• pH values change by factors of 10, making minor changes in pH impactful for acidity/basicity level.

Measuring Acidity: The pH Function

• Water has some weak acid-base properties.
• A tiny portion of water molecules (H₂O) dissociate into hydrogen (H⁺) and hydroxyl ions (OH^-).
• Dissociation equilibrium exists; pure water has equal H+ and OH- concentrations.
• Pure water has a pH of 7.00.

Measuring Acidity: The pH Function: Relationship Between pH and pOH

• pH and pOH values of aqueous solutions have a fixed relationship.
• When added together, pH and pOH of any aqueous solution at 25°C always equal 14.00.
• p-functions commonly describe equilibrium constants for acids and bases
• The sum of pKa and pKb for a conjugate acid-base pair is equal to 14.00.

Other Acidic Species

• Nonmetal oxides added to water create acidic solutions.
• Carbon dioxide is a pivotal example of a nonmetal oxide in water dissolution, forming carbonic acid.
• Carbon dioxide buildup in the blood causes acidosis.
• Elevated carbon dioxide levels slightly alter structural properties of hemoglobin for O2 release from the blood.

Other Acidic Species & Carbon Dioxide

• Carbon dioxide is a nonmetal oxide compound combining with oxygen.
• Nonmetal oxides are sometimes called acid anhydrides because they produce acids when combined with water.
• Carbon dioxide combines with water molecules to produce carbonic acid; CO2+H2O → H2CO3
• Carbonic acid dissociates: H2CO3 ⇄ H⁺ + HCO₃^-.
• Dissolving CO2 in water lowers pH.

Buffers

• A buffer solution resists pH changes by containing a weak acid (HA) and its conjugate base (A−) or a weak base and its conjugate acid.
• Adding a strong base converts the weak acid to its conjugate base in a reversible reaction (HA + OH⁻ ⇄ H₂O + A⁻).
• Conversely, adding a strong acid will convert conjugate base to a weak acid. (A⁻ + H⁺ ⇄ HA).

Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation estimates pH of a buffered solution, relating it to the pKa of a drug and its pH.
• The equation predicts which ionization form (unionized or ionized) of a drug will predominate in a certain pH environment.
• The greater the distance between pKa and physiologic pH, the greater the ionization of the drug. Increased ionization rate corresponds with drug effect onset.

Solubility and Pharmacologic Effects

• Ionized molecules have more water solubility and are less likely to be active.
• Unionized molecules have more lipid solubility and are more likely to be active.
• Lipid-insoluble molecules diffuse less well across lipid bilayers and are much less likely to enter the brain or GI tract.
• Both ionized and unionized molecules can cross the placenta effectively.

Acid-Base Balance & Respiratory System

• The respiratory system helps maintain normal pH balance by influencing acid-base balance within body tissues.
• The interplay between the respiratory and renal systems is essential in maintaining blood and tissue acid-base balance.
• This delicate equilibrium can be disturbed in cases such as acidosis and alkalosis, impacting enzyme functions and various metabolic processes.

Acid-Base Balance & Respiratory System & Compensations

• The respiratory and renal systems regulate pH levels utilizing buffers.
• Bicarbonate, phosphate, and proteins are primary buffers, maintaining pH in the physiological range.
• Metabolic acidosis/alkalosis compensations are rapid respiratory changes in alveolar ventilation depending on the underlying metabolic process. The total base deficit, obtained from blood gas results, helps to determine the correct dosage for bicarbonate when indicated.

Human Buffer Systems

• Blood system bicarbonate is the most important buffer.
• Hemoglobin is the second most important buffer in blood.
• Respiratory system influences acid-base balance primarily via PaCO2.
• Kidneys help in reabsorbing filtered bicarbonate and removing titratable acids.

Anion Gap – Metabolic Acidosis

• Anion gap assists in determining the cause of metabolic acidosis.
• Anion gap calculation includes cations and anions to identify acidosis of a metabolic or nonmetabolic origin.
• Acidosis due to an acid accumulation has a gap typically greater than 12 mEq/L, while dilution/bicarbonate loss is classified as non-gap acidosis.

Interpretation of Arterial Blood Gases

• Analysis of arterial blood gases helps clarify the relationship between acid production and removal via a variety of body systems, including the lungs and kidneys.
• Acid-base disturbances are categorized—respiratory and metabolic, each either acidic or alkalotic, with various compensations under certain clinical conditions. A compensatory response isn't enough to normalize pH for a condition to be considered compensated.

Acidosis

• Processes increasing PaCO2 decrease arterial pH, resulting in respiratory acidosis.
• Acute changes in PaCO2 (ten or more mm Hg) change pH by 0.08 units.
• Increased PaCO2 and normal bicarbonate indicate metabolic acidosis is present.
• Causes leading to non-respiratory acidosis include poisons, infusions, lactic acid formation, and reduced kidney acid excretion.
• Changes in base and pH are related—as is a ten mEq/L base change associated with a 0.15 pH change.

Alkalosis

• Hyperventilation decreases PaCO2 resulting in respiratory alkalosis, which is when the pH level increases.
• Metabolic alkalosis is associated with fixed acid loss or high base intake.

pH, CO2, and HCO3 Clinical Examples

• Various conditions (mixed alkalosis, acute hyperventilation, compensated metabolic alkalosis, compensated/normal respiratory alkalosis, compensated/normal respiratory acidosis, renal failure, acute hypoventilation, diabetic ketoacidosis, and respiratory/circulatory arrest) are categorized based on pH, PaCO2, and HCO3 values (often in ranges). The specific values and shifts indicate potential causes and compensations.

This quiz focuses on key concepts in biochemistry related to metabolism and energy production processes, including glycolysis, glycogenolysis, and the citric acid cycle. Test your understanding of the functions and roles of various molecules involved in these metabolic pathways.