Comprehensive Study Notes on the Chemistry of Life

Fundamental Principles of Matter, Atoms, and Molecules

• Chemical Basis of Life: Every living organism is fundamentally a chemical system driven by interactions among basic chemical ingredients. Knowledge of basic chemistry—the scientific study of matter—is required to understand biological processes.

• Definition of Matter: Matter is defined as anything that occupies space and has mass (substance). All physical objects in the universe are composed of matter.

• Three Phases of Matter: On Earth, matter exists in three distinct physical states:

  • Gas: Example includes propane gas used in backyard grills.

  • Liquid: Example includes acetic acid, which gives vinegar its sharp taste.

  • Solid: Example includes acetylsalicylic acid, the active ingredient in aspirin ($100\,\text{mg}$ per tablet).

• Atoms: The smallest units of matter that retain all the chemical and physical properties of their specific type of matter. Atoms rarely exist in isolation in nature.

• Molecules: Structures formed when two or more atoms are chemically bonded together.

  • Propane Gas Molecule: Composed of carbon and hydrogen atoms.

  • Acetic Acid Molecule: Composed of carbon, hydrogen, and oxygen atoms.

  • Acetylsalicylic Acid Molecule: Composed of carbon, hydrogen, and oxygen atoms.

Chemical Elements, Periodic Table, and Essential Life Elements

• Elements: Pure substances that cannot be broken down into simpler substances by chemical reactions. Examples include hydrogen, carbon, uranium, and gold. Elements can occasionally exist as unbonded collections of atoms, such as neon gas lighting up a sign.

• Compounds: Substances consisting of two or more different elements combined in a fixed, definite ratio.

  • Sodium Chloride ($\text{NaCl}$) Example: Sodium ($\text{Na}$) is an explosive solid metal, and chlorine ($\text{Cl}$) is a deadly gas. When chemically combined, they form sodium chloride ($\text{NaCl}$), non-toxic table salt: $\text{Na} + \text{Cl} \rightarrow \text{NaCl}$

• Chemical Reactions: Processes during which molecules are broken down and built up, swapping atoms while keeping the individual atoms whole.

  • Reactants: The starting substances, written on the left side of a chemical equation.

  • Products: The ending substances, written on the right side of a chemical equation.

  • Conservation of Matter: Chemical reactions rearrange atoms, but atoms are never created nor destroyed.

  • Cellular Respiration Reaction: Uses oxygen gas and glucose to yield energy, producing water and carbon dioxide as by-products: $\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{H}_2\text{O} + 6\text{CO}_2 + \text{ENERGY}$

  • Maltose Synthesis Reaction: Two molecules of glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) combine to form one molecule of water ($\text{H}_2\text{O}$) and one molecule of maltose ($\text{C}_{12}\text{H}_{22}\text{O}_{11}$): $2\,\text{C}_6\text{H}_{12}\text{O}_6 \rightarrow \text{H}_2\text{O} + \text{C}_{12}\text{H}_{22}\text{O}_{11}$

• The Periodic Table of the Elements: Organizes all chemical elements in order of their atomic number. There are 92 naturally occurring elements on Earth, alongside several artificial elements created in laboratories.

  • Atomic Number: The number of protons contained within the nucleus of an atom of the element.

  • Element Symbol: One- or two-letter abbreviation for the element (e.g., $\text{Ca}$ for Calcium, $\text{C}$ for Carbon, $\text{He}$ for Helium, $\text{Hg}$ for Mercury, $\text{Ne}$ for Neon).

  • Atomic Weight: Represents the average mass of all naturally occurring isotopes of an element, corresponding to the total number of protons plus neutrons. This is expressed as a decimal value due to isotope variations.

    • Calcium ($\text{Ca}$): Atomic number $20$, atomic weight $40.08$.

    • Carbon ($\text{C}$): Atomic number $6$, atomic weight $12.01$ (commonly found as graphite).

• 25 Elements Essential to Life:

  • Bulk Elements ($96.3\%$ of cell mass): Four elements construct the majority of living cells:

    • Oxygen ($\text{O}$): $65.0\%$

    • Carbon ($\text{C}$): $18.5\%$

    • Hydrogen ($\text{H}$): $9.5\%$

    • Nitrogen ($\text{N}$): $3.3\%$

  • Major Elements ($3.7\%$ of cell mass): Seven elements make up a smaller fraction of cell mass:

    • Calcium ($\text{Ca}$): $1.5\%$

    • Phosphorus ($\text{P}$): $1.0\%$

    • Potassium ($\text{K}$): $0.4\%$

    • Sulfur ($\text{S}$): $0.3\%$

    • Sodium ($\text{Na}$): $0.2\%$

    • Chlorine ($\text{Cl}$): $0.2\%$

    • Magnesium ($\text{Mg}$): $0.1\%$

  • Trace Elements ($0.1\%$ of cell mass): Fourteen elements are required in miniscule quantities, but cells cannot survive without them:

    • Boron, Chromium, Cobalt, Copper, Fluorine, Iodine, Iron, Manganese, Molybdenum, Selenium, Silicon, Tin, Vanadium, and Zinc.

    • Molybdenum Example: The human body contains only $0.0002\,\text{ounces}$ of molybdenum, yet dietary deficiency causes severe neurological damage.

  • Non-Essential Elements: Dietary supplements containing non-essential elements like gold provide no biological benefit, as gold is not among the 25 essential life elements.

Atomic Structure, Subatomic Particles, Isotopes, and Ions

• Subatomic Structure: Atoms consist of three primary subatomic particles:

  • Proton: Relative mass of $1$, electric charge of $+1$, located inside the central nucleus. The number of protons determines the chemical element.

  • Neutron: Relative mass of $1$, electric charge of $0$ (neutral), located inside the central nucleus. The number of neutrons determines the isotope.

  • Electron: Relative mass of $0$ (negligible mass), electric charge of $-1$, orbits the nucleus at high speeds within specific orbital regions called electron shells, forming an electron cloud. Electrons determine ion state and chemical properties.

• Standard Nitrogen Atom Example: Contains 7 protons, 7 neutrons, and 7 electrons (2 in the inner electron shell, 5 in the outer electron shell). Outer shell electrons interact with other atoms to determine chemical reactivity.

• Isotopes: Variants of a single element that contain identical numbers of protons but differ in their number of neutrons. Named according to total atomic weight (protons + neutrons).

  • Nitrogen-15 ($\text{N-15}$ or $^{15}\text{N}$): Contains 7 protons, 8 neutrons, and 7 electrons, giving an atomic weight of 15 (has one extra neutron compared to standard nitrogen).

  • Nitrogen-13: Contains 7 protons and 6 neutrons, giving an atomic weight of 13.

  • Carbon-14 ($^{14}\text{C}$): Contains 6 protons, 8 neutrons ($14 - 6 = 8$), and 6 electrons in its neutral state.

  • Radioactive Isotopes: Useful in medical applications; for example, fluorine-18 ($^{18}\text{F}$) is utilized to detect functional and structural defects in the human brain.

• Ions: Atoms or groups of atoms that have acquired an electrical charge by gaining or losing electrons.

  • Neutral atoms contain an equal count of protons and electrons.

  • Gaining electrons adds negative charge; losing electrons creates a net positive charge.

  • Nitrogen Ion ($\text{N}^{3-}$): Contains 7 protons, 7 neutrons, and 10 electrons (3 extra electrons), giving it a net $-3$ charge.

Chemical Bonds: Ionic, Covalent, and Hydrogen

• Chemical Bond Formation: During chemical reactions, atoms gain, lose, or share electrons. These interactions create attractions that bind atoms together.

• Ionic Bonds:

  • Involve the complete transfer of one or more electrons from one atom to another.

  • The electron-donating atom becomes a positively charged cation; the electron-receiving atom becomes a negatively charged anion.

  • The resulting opposite charges attract one another like magnets to form an ionic bond.

  • Formation of Table Salt ($\text{NaCl}$): A neutral sodium ($\text{Na}$) atom with 11 electrons donates 1 electron to a neutral chlorine ($\text{Cl}$) atom with 17 electrons. This yields a $\text{Na}^+$ ion (10 electrons) and a $\text{Cl}^-$ ion (18 electrons). The electrostatic attraction binds them into $\text{NaCl}$.

• Covalent Bonds:

  • Involve the sharing of one or more pairs of electrons between atoms. Each atom contributes one electron per shared pair.

  • Holds atoms together within individual molecules.

  • Single Bond: Sharing of one electron pair; represented by a single solid line ($\text{-}$).

  • Double Bond: Sharing of two electron pairs; represented by two parallel solid lines ($\text{=}$), as in oxygen gas ($\text{O}_2$ or $\text{O}=\text{O}$).

  • Triple Bond: Sharing of three electron pairs; relatively rare in biological systems.

  • Nonpolar Covalent Bonds: Equal sharing of electrons between bonded atoms. Example: Methane ($\text{CH}_4$), where carbon and hydrogen share electrons equally. Rearranging carbon atoms from coal can alter their structural arrangement to form diamond.

  • Polar Covalent Bonds: Unequal sharing of electrons because one atomic nucleus exerts a stronger pull on the shared electrons than the other.

    • Water ($\text{H}_2\text{O}$): The oxygen nucleus attracts shared electrons more strongly than the hydrogen nuclei. Consequently, oxygen carries a slight negative charge ($\delta-$) and each hydrogen carries a slight positive charge ($\delta+$), while the overall molecule remains neutral.

• Hydrogen Bonds:

  • Weak electrical attractions between partial opposite charges of polar molecules.

  • In water, the slightly negative oxygen atom of one molecule aligns with and attracts the slightly positive hydrogen atom of a neighboring water molecule.

  • Individual hydrogen bonds are weak and short-lived, but extensive interconnected networks of hydrogen bonds create strong cohesion in liquid water.

  • Occur intermolecularly (between separate molecules, such as water) or intramolecularly (between different regions within a single large molecule, such as a protein).

  • Bond Types in Water: Atoms within a single water molecule are held together by polar covalent bonds; separate water molecules are held together by hydrogen bonds.

Unique Properties and Biological Importance of Water

• Biological Importance: Life originated in water and evolved there for billions of years before adapting to land. Most living cells are composed primarily of water by weight.

• Ice Floating and Expansion:

  • In liquid water, hydrogen bonds continuously break and reform, allowing molecules to move in close proximity.

  • When water freezes into solid ice, molecules move further apart, locking into a rigid, spacious crystalline network of long-lasting hydrogen bonds.

  • Unlike almost all other liquids, water expands upon freezing, rendering solid ice less dense than liquid water, so ice floats.

  • Ecological Relevance: In winter, floating ice forms an insulating top layer over bodies of water, maintaining liquid conditions beneath and allowing aquatic organisms to survive until the spring thaw.

• Solvent Properties:

  • Solvent: A liquid dissolving agent in a solution.

  • Solution: A homogeneous liquid mixture of two or more substances.

  • Water is an extremely effective polar solvent capable of dissolving polar and ionic substances. When table salt ($\text{NaCl}$) dissolves in water, polar water molecules surround $\text{Na}^+$ and $\text{Cl}^-$ ions, breaking apart the salt crystal matrix.

  • Extraterrestrial Discovery: In 2015, NASA confirmed the presence of liquid water flowing on the surface of Mars.

• Cohesion and Adhesion:

  • Cohesion: The tendency of identical water molecules to stick to one another via hydrogen bonding. Creates high surface tension, forming a film-like surface capable of supporting light objects or organisms (e.g., water strider insects or spiders walking across a pond).

  • Adhesion: The clinging of water molecules to a different polar surface (e.g., water droplets sticking to a car windshield).

• Temperature Regulation:

  • Liquid water readily absorbs and stores large amounts of heat while resisting dramatic temperature fluctuations.

  • Global Scale: Earth's oceans absorb heat to moderate planetary surface temperatures within livable limits.

  • Personal Scale: Sweating regulates body temperature via evaporative cooling as heat is carried away when surface water evaporates.

Aqueous Solutions, pH Scale, and Biological Buffers

• Aqueous Solutions: Solutions in which water is the dissolving solvent. A small percentage of liquid water molecules naturally dissociate into hydrogen ions ($\text{H}^+$) and hydroxide ions ($\text{OH}^-$).

• pH Scale: Measures the relative concentration of $\text{H}^+$ ions in an aqueous solution, ranging from $0$ (most acidic) to $14$ (most basic).

  • Logarithmic Scale: Each integer change on the pH scale represents a tenfold ($10\times$) change in $\text{H}^+$ ion concentration.

  • Neutrality (pH 7.0): Concentrations of $\text{H}^+$ and $\text{OH}^-$ are equal ($[\text{H}^+] = [\text{OH}^-]$).

• Acids:

  • Chemical compounds that release $\text{H}^+$ ions when dissolved in water, raising $\text{H}^+$ concentration.

  • pH range: $0 \le \text{pH} < 7$.

  • Hydrochloric Acid ($\text{HCl}$): Gastric juice acid in human stomachs; dissociates into $\text{H}^+$ and $\text{Cl}^-$ in solution.

  • Representative pH Values: Battery acid ($1.0$), Lemon juice ($2.4$), Strawberries ($3.5$), Tomatoes ($4.6$), Black coffee ($5.0$), Cow's milk ($6.4$).

  • Comparison: Lemonade has a lower pH (more acidic) and a higher concentration of $\text{H}^+$ ions than pure water.

• Bases:

  • Chemical compounds that accept or remove $\text{H}^+$ ions from solution, typically by releasing $\text{OH}^-$ which combines with $\text{H}^+$ to form $\text{H}_2\text{O}$.

  • pH range: $7 < \text{pH} \le 14$.

  • Sodium Hydroxide ($\text{NaOH}$ / Lye): Dissociates in water into $\text{Na}^+$ and $\text{OH}^-$; $\text{OH}^-$ binds $\text{H}^+$, lowering overall $\text{H}^+$ concentration.

  • Representative pH Values: Human blood ($7.4$), Baking soda ($9.0$), Household ammonia ($11.6$), Household bleach ($12.6$), Lye / Drano ($13.0$).

• Buffers: Substances that minimize drastic changes in pH by accepting $\text{H}^+$ ions when they are present in excess and donating $\text{H}^+$ ions when they are depleted. Blood buffers maintain human physiological pH around neutral ($7.4$), counteracting acidic drops induced during strenuous muscle exercise.

• Environmental Consequences of pH Alterations:

  • Acid Precipitation: Fossil fuel combustion releases compounds that react with atmospheric water vapor to form strong acids. Falling as rain, snow, or fog (sometimes with a pH lower than vinegar), acid precipitation damages aquatic ecosystems, forests, and soils. The U.S. Clean Air Act reduced acid rain levels.

  • Ocean Acidification: The world's oceans absorb approximately $25\%$ of excess atmospheric carbon dioxide ($\text{CO}_2$). Dissolved $\text{CO}_2$ undergoes chemical reactions that lower ocean pH, impairing the ability of marine organisms to build shells and calcified skeletons, disrupting coral reefs.

Carbon Skeletons, Functional Groups, and Biological Macromolecules

• Carbon-Based Life: Aside from water, most biological molecules are organic compounds—molecules containing carbon covalently bonded to other elements.

• Properties of Carbon: Carbon contains 4 valence electrons, allowing it to form up to four covalent bonds with other atoms. Carbon forms large, highly branched, or ringed molecular chains that serve as foundational structural skeletons.

• Carbon Skeletons: Vary in overall length, branching patterns, and ring formations. Every carbon atom consistently forms four total bonds.

  • Examples: 3-carbon straight chain (propane), branched carbon chains, and ring structures.

• Functional Groups: Specific clusters of atoms attached to carbon skeletons that directly participate in chemical reactions and dictate the overall chemical properties of organic compounds.

  • Hydroxyl Group ($\text{-OH}$): Found in ethyl alcohol (ethanol), the alcohol present in alcoholic beverages.

  • Amino Group ($\text{-NH}_2$): Present in amino acids, the primary structural building blocks of proteins.

  • Phosphate Group ($\text{-PO}_4^{2-}$): Present in adenosine triphosphate (ATP), which supplies energy to living cells.

• Four Major Classes of Biological Macromolecules:

  • Carbohydrates: Examples include cellulose (complex carbohydrate forming plant structural cell walls; dietary fiber) and glucose (simple sugar serving as a cellular energy source).

  • Lipids: Examples include coconut oil (fat-rich lipid used as a dietary staple in tropical regions) and cholesterol (circulates in blood, acts as a molecular precursor for steroid hormones).

  • Proteins: Examples include hexokinase (an enzyme driving metabolic reactions) and keratin (a structural protein forming hair, skin, and nails).

  • Nucleic Acids: Examples include DNA (hereditary material of all life on Earth) and RNA (messenger molecule carrying genetic instructions between DNA and cellular machinery across all cell types).

Polymer Synthesis and Breakdown: Hydrolysis and Dehydration Synthesis

• Macromolecules: Large, complex biological molecules constructed from smaller molecular building blocks.

• Polymers and Monomers:

  • Polymer: A large molecule constructed by joining together many smaller repeating sub-units.

  • Monomer: The individual smaller chemical units that serve as the building blocks of polymers.

• Hydrolysis Reactions:

  • Chemical process that breaks down polymers into constituent monomers by adding water.

  • A water molecule ($\text{H}_2\text{O}$) is split; a hydrogen atom ($\text{H}$) attaches to one monomer, and a hydroxyl group ($\text{OH}$) attaches to the adjacent monomer, severing the covalent bond between them.

  • Digestive Application: Digestive enzymes use hydrolysis to break down dietary macromolecules (e.g., proteins in peanut butter) into free amino acid monomers inside the gut. Consuming liquids during meals provides water needed to drive these reactions.

• Dehydration Synthesis Reactions:

  • Chemical process that links monomers together to form larger polymers by removing water.

  • A hydrogen atom ($\text{H}$) from one monomer and a hydroxyl group ($\text{OH}$) from another are removed, creating a new covalent bond between the monomers while releasing one water molecule ($\text{H}_2\text{O}$) per added monomer.

  • Dehydration synthesis is the chemical exact reverse of hydrolysis.

  • Muscular Application: Cells take up absorbed amino acid monomers from the bloodstream and perform dehydration synthesis to build human muscle proteins.

• Metabolism: The sum total of all chemical reactions taking place within an organism's body, including continuous polymer degradation (hydrolysis) and polymer construction (dehydration synthesis).

Structure, Classification, and Roles of Carbohydrates

• Carbohydrates: Molecules composed of one or more simple sugar monomers. Function as a primary source of dietary energy for animals and structural support in plants.

• Monosaccharides (Simple Sugars):

  • The basic monomer building blocks of carbohydrates.

  • Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$): Found in sports drinks; administered intravenously to surgery patients and trauma victims to provide immediate energy.

  • Fructose ($\text{C}_6\text{H}_{12}\text{O}_6$): Fruit sugar found naturally in fruits and honey.

  • Isomers: Glucose and fructose are chemical isomers—molecules sharing identical chemical formulas ($\text{C}_6\text{H}_{12}\text{O}_6$) but possessing different structural arrangements of atoms.

• Disaccharides:

  • Double sugars formed by linking two monosaccharides via a dehydration synthesis reaction.

  • Suffix: Most sugar names end in the suffix "-ose" (e.g., glucose, fructose, sucrose, lactose, maltose).

  • Sucrose ($\text{C}_{12}\text{H}_{22}\text{O}_{11}$): Common table sugar formed by combining glucose and fructose, releasing $\text{H}_2\text{O}$. Most table sugar produced in the United States is extracted from sugar beets rather than sugar cane.

  • Lactose: Milk sugar formed by joining simple sugars.

  • Maltose: Sugar used in brewing and malted milk candy.

• Polysaccharides:

  • Complex carbohydrates created by joining many glucose monomers into long chains.

  • Starch: Long, twisted, unbranched chains of glucose molecules. Plants store excess photosynthesized glucose as starch; animals consume starch as dietary energy.

  • Cellulose: Long, straight, parallel chains of glucose joined by cross-linking bonds to form tough, cable-like fibers. Forms plant cell walls and wood; acts as indigestible dietary "fiber" in human nutrition.

  • Glycogen: Highly branched chains of glucose molecules. Animals store surplus glucose as glycogen granules within liver and muscle cells, providing accessible energy reserves for roughly 24 hours.

  • Chitin: Structural polysaccharide forming the hard exoskeleton of arthropods (insects, spiders) and fungal cell walls. Structurally similar to cellulose, but its glucose monomers feature a nitrogen-containing functional group.

Diversity and Characteristics of Hydrophobic Lipids

• Hydrophobic Nature: Lipids are a structurally diverse group of organic compounds defined by a shared physical property: they are hydrophobic ("water-fearing") and do not mix with or dissolve in water. (e.g., salad oil separating from watery vinegar).

• Phospholipids:

  • Form the structural basis of cell membranes via a double-layered sheet called a phospholipid bilayer.

  • Each phospholipid consists of a hydrophilic ("water-loving") head containing a phosphate group and two long hydrophobic fatty acid tails.

  • Membranes feature floating functional proteins embedded within the phospholipid layers.

• Cholesterol:

  • A lipid embedded within animal cell membranes to regulate membrane fluidity.

  • Acts as a molecular precursor synthesized by animal cells to produce steroid hormones.

  • Synthesized naturally by the human body and ingested through animal-derived foods (e.g., eggs, red meat).

  • Low-Density Lipoprotein (LDL): Known as "bad cholesterol." Elevated by poor diet; high blood levels increase cardiovascular disease risk.

  • High-Density Lipoprotein (HDL): Known as "good cholesterol." Elevated by physical exercise; high blood levels reduce cardiovascular disease risk.

• Triglycerides:

  • Dietary fats consisting of one molecule of glycerol linked to three fatty acid tails via dehydration synthesis.

  • The carbon-hydrogen chains in fatty acid tails store chemical energy. Surplus dietary calories are converted into triglycerides and stored in adipose tissue (body fat).

  • Caloric Density: Body fat is calorie-dense; one pound ($1\,\text{lb}$) of human adipose tissue stores $3,500\,\text{calories}$ of energy.

• Steroid Hormones:

  • Lipids containing a core structure of four fused carbon rings.

  • Cholesterol is modified to form sex hormones, such as estrogen and testosterone.

  • Anabolic Steroids: Synthetic derivatives of testosterone that mimic its tissue-building effects, increasing muscle mass. Usage causes dangerous physiological side effects; professional athletes (e.g., baseball player Alex Rodriguez) have admitted to using synthetic anabolic steroids like THG.

Dietary Fats: Saturated, Unsaturated, Trans, and Essential Fats

• Energy Content: One pound of dietary fat packs more than double ($>2\times$) the energy content of one pound of carbohydrate.

• Saturated Fats:

  • Contain the maximum possible number of hydrogen atoms along the fatty acid tails, with all single covalent carbon-carbon bonds ($\text{C-C}$).

  • Fatty acid tails possess a straight molecular shape, allowing molecules to stack closely together.

  • Tend to remain solid at room temperature.

  • Found in highest proportions in animal products (e.g., beef, butter, cheese).

  • Considered less healthy in human diets.

• Unsaturated Fats:

  • Contain one or more double covalent carbon-carbon bonds ($\text{C=C}$), resulting in fewer than the maximum number of hydrogen atoms.

  • Double bonds cause a bend or kink in the fatty acid tail structure, preventing molecules from stacking tightly.

  • Tend to remain liquid at room temperature.

  • Found in highest proportions in plant products and fish oils (e.g., olive oil, corn oil).

  • Considered healthier in human diets.

  • Proportion Misconception: Animal fats are not exclusively saturated, nor are plant fats exclusively unsaturated; rather, both contain mixtures of saturated and unsaturated fats in varying proportions.

• Trans Fats:

  • Unsaturated fats converted into solid fats through an industrial manufacturing process called hydrogenation.

  • Hydrogenation generates trans fats containing unusual double-bond configurations not typically found in nature.

  • Highly unhealthy; significantly increases risk of heart disease, type 2 diabetes, and high blood pressure. Food labels must explicitly disclose trans fat content.

• Healthy Essential Fats:

  • Fats containing omega-3 fatty acids reduce the overall risk of heart disease.

  • Found in fish, chicken, eggs, peanuts, and beans.

Protein Structure, Diversity, and Biological Functions

• Proteins: The most structurally and functionally diverse class of biological macromolecules. Polymers constructed from combinations of 20 different amino acid monomers.

• Amino Acid Architecture: Every amino acid consists of a central carbon atom bonded to four partners:

  1. A hydrogen atom.

  2. An amino group ($\text{-NH}_2$).

  3. A carboxylic acid group ($\text{-COOH}$).

  4. A unique variable side group (R-group) that confers specific chemical properties.

• Peptide Bonds: Covalent bonds joining amino acids together, formed via dehydration synthesis reactions between the amino group of one amino acid and the carboxylic acid group of another.

• Polypeptides: A long, unbranched chain of amino acids, ranging from hundreds to thousands of amino acids in length.

• Protein Folding and Structure:

  • A polypeptide chain twists and folds into a unique three-dimensional shape held together by chemical bonds formed between side groups of the amino acids.

  • Protein function is dependent upon this precise 3D shape.

  • Some functional proteins consist of multiple individual polypeptide chains joined into a complex assembly (e.g., a single hemoglobin protein molecule contains four individual polypeptide chains).

• Major Biological Functions of Proteins:

  • Structure: Keratin provides structural strength to hair, skin, nails, and animal fur.

  • Transport: Hemoglobin inside red blood cells binds and carries oxygen gas through the bloodstream.

  • Defense: Antibodies produced by the immune system bind to specific foreign pathogens, marking them for destruction.

  • Enzymes: Lactase hydrolyzes the disaccharide milk sugar lactose in the digestive system.

  • Movement: Actin filaments enable muscle fibers to contract.

• Structure-Function Relationships and Pathologies:

  • Altering a protein's amino acid sequence can alter its shape and neutralize its function.

  • Sickle-Cell Disease: Changing a single amino acid out of 146 in one of hemoglobin's polypeptide chains causes the entire protein to misfold. Misfolded hemoglobin distorts red blood cells into a rigid, curved shape, preventing proper oxygen transport.

  • Cystic Fibrosis: The most common underlying cause is a single amino acid mutation among 1,480 amino acids in a critical cellular membrane protein.

Enzymes, Activation Energy, and Mechanisms of Inhibition

• Enzymes: Catalytic proteins that speed up the rate of specific chemical reactions without being consumed or permanently altered by the reaction.

• Substrates and Active Sites:

  • Substrate: The specific target reactant molecule acted upon by an enzyme.

  • Active Site: A specialized pocket on the surface of the enzyme with a shape complementary to the substrate ("hand in a glove").

  • Induced Fit: Upon binding, the active site and substrate undergo slight conformational shape changes to achieve a tighter fit.

  • Catalysis Example: The enzyme lactase binds its substrate lactose, splitting it into two product simple sugars: glucose and galactose. Lactase releases the products and remains unaltered, ready to accept another substrate.

• Activation Energy Reduction:

  • Activation Energy: The threshold quantity of energy required to initiate a chemical reaction.

  • Enzymes accelerate chemical reactions by lowering the activation energy barrier, allowing reactions to proceed rapidly at cellular temperatures.

• Enzyme Inhibitors: Molecules that bind to enzymes and disrupt or prevent their normal catalytic activity.

  • Competitive Inhibitors: Imposter molecules that directly enter and bind to the active site. Block access to the genuine substrate.

  • Noncompetitive Inhibitors: Bind to an alternate site on the enzyme away from the active site. Binding induces a conformational change in the enzyme's shape, altering the active site so the substrate can no longer fit.

  • Theoretical Application: A non-digestible sugar matching the exact shape of lactose would act as a competitive inhibitor, binding to lactase's active site and blocking lactose digestion.

• Pathologies Associated with Enzyme Shape Alteration:

  • Lactose Intolerance: Mutations that alter the structural shape of the active site in lactase render it unable to break down lactose, leading to digestive distress.

  • Tay-Sachs Disease: Genetic mutations disrupting a specific metabolic enzyme cause Tay-Sachs disease, a neurological disorder that is typically fatal by age 4.