Ch. 2 The Chemical Level of Organization — Vocabulary Flashcards

Matter

  • Matter has mass and occupies space.
  • Exists in three forms: gas, liquid, and solid.
  • Fundamental components include: Elements, Atoms, and Subatomic particles.

The Periodic Table

  • The periodic table organizes elements by increasing atomic number and by recurring chemical properties.
  • Major groupings include:
    • Alkali metals (Group 1)
    • Alkaline earth metals (Group 2)
    • Transition metals (Groups 3–12)
    • Halogens (Group 17)
    • Noble gases (Group 18)
    • Lanthanide series and Actinide series (bottom rows)
  • Common examples shown in the transcript: H, He, Li, Be, B, C, N, O, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Fe, Co, Ni, Cu, Zn, Ag, Au, Pb, etc.
  • Periodic table emphasizes:
    • Each element is represented by a unique symbol (one or two letters; first letter capitalized).
    • The mass number (A) is the sum of protons and neutrons; the atomic number (Z) is the count of protons.
    • Isotopes differ in neutron number but have the same proton number; they may have different chemical properties if neutron number alters behavior in some contexts.

Elements vs Compounds

  • Elements: Pure substances consisting of only one type of atom.
  • Compounds: Substances composed of two or more elements bonded together.
  • The body contains elements whose numbers and types come from diet and inhaled air.

Elements in the Human Body (Major Elements by Mass)

  • Oxygen (O): ≈ 65.0% of body mass; ~3% of body composition by other measures in some slides; role: major component of water and organic molecules.
  • Carbon (C): ≈ 18.5%
  • Hydrogen (H): ≈ 9.5–10%
  • Nitrogen (N): ≈ 3.2%
  • Calcium (Ca): ≈ 1.5%
  • Phosphorus (P): ≈ 1.0%
  • Potassium (K): ≈ 0.4%
  • Sulfur (S): ≈ 0.3%
  • Sodium (Na): ≈ 0.2%
  • Chlorine (Cl): ≈ 0.2%
  • Magnesium (Mg): ≈ 0.1%
  • Trace elements (examples): boron (B), chromium (Cr), cobalt (Co), copper (Cu), fluorine (F), iodine (I), iron (Fe), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vanadium (V), zinc (Zn).

Structure of the Atom

  • Subatomic particles:
    • Protons (+) and neutrons (no charge) reside in the nucleus.
    • Electrons (-) orbit the nucleus and have negligible mass.
  • Proton number = atomic number (Z).
  • Neutron number = mass number − atomic number (N = A − Z).
  • Electron number in a neutral atom = atomic number (Z).
  • Protons are invariant in most chemical contexts; neutrons can vary in isotopes; electrons can change in charged species.

Isotopes and Radioactivity

  • Isotopes are atoms of the same element with the same number of protons and electrons but different numbers of neutrons.
  • Isotopes have identical chemical properties but different atomic masses.
  • Examples for carbon: Carbon-12 (6 neutrons), Carbon-13 (7 neutrons), Carbon-14 (8 neutrons).
  • Some isotopes are radioactive and undergo radioactive decay, releasing energy and particles.
  • Uses of radioisotopes: imaging (low-level radiation), sterilization (high-level radiation), tracers in biochemical pathways.

The Nuclear Model and Electron Shells

  • Electron shells surround the nucleus; electrons occupy orbitals within shells.
  • The first shell holds up to 2 electrons; all subsequent shells typically hold up to 8 electrons for biologically relevant elements.
  • Atoms are most stable when their outermost shell is full; chemically active electrons are those in the outer shell.
  • Bohr-like shell model: energy levels determine reactivity; outer-shell electrons govern bonding.

Atomic Symbols and Notation

  • Atomic symbol is a one- or two-letter abbreviation (e.g., C for carbon).
  • Isotope notation can be written as a superscript for mass number and a subscript for atomic number before the symbol, e.g., $^{12}_{6} ext{C}$ for carbon-12.
  • In a neutral atom, proton number equals electron number; neutron number is A − Z.

Subatomic Particle Counts and Relationships

  • Proton number Z = atomic number (identity of the element).
  • Neutron number N = A − Z.
  • Electron number in neutral atoms = Z; electrons can change in ions.
  • Protons do not change during normal chemical reactions; neutrons and electrons can vary by isotope formation or ionization.

Chemical Bonds: Ionic, Covalent, and Polar Characteristics

  • Ionic bonds form via transfer of one or more electrons from a donor atom to an acceptor atom.
  • Ions are charged atoms; anions are negatively charged (gained electrons); cations are positively charged (lost electrons).
  • Ionic compounds form crystals rather than discrete molecules.
  • Covalent bonds form when atoms share electrons; this leads to molecular compounds.
  • Bond type depends on electronegativity differences; polarity arises from unequal sharing.
  • Amphipathic molecules have both polar (hydrophilic) and nonpolar (hydrophobic) regions; common in lipids like phospholipids.

Common Ions in the Human Body (Table-like References)

  • The body contains many common ions with physiological roles (e.g., Na+, K+, Ca2+, Cl−, HCO3−, H+, OH−). Tables summarize their concentrations and functions in body fluids and cells.

Covalent Bonding, Molecules, and Molecular Compounds

  • Covalently bonded molecules share electrons between atoms of two or more elements.
  • Resulting substances are called molecular compounds.
  • Examples of covalent bonds:
    • Single bond: H–H (hydrogen gas, H2)
    • Single bond: O=O (oxygen gas, O2)
    • Double bond: N=N (nitrogen gas, N2)
  • Bond strength and bond order increase with more shared electron pairs (single < double < triple).

Polar vs Nonpolar Covalent Bonds

  • Nonpolar covalent bond: electrons shared equally; no partial charges on atoms. Example: O2, N2, H2.
  • Polar covalent bond: electrons shared unequally; partial charges on atoms. Example: H2O shows partial charges on H and O.
  • Ionic bonds: complete transfer of electrons; formation of ions with charges; resulting in ionic compounds (crystalline).

Amphipathic Molecules

  • Large molecules with both polar and nonpolar regions.
  • Examples include phospholipids, which form cell membranes with a hydrophilic (polar) head and hydrophobic (nonpolar) tails.

Lipids: Non-Polymers with Diverse Roles

  • Lipids are diverse, hydrophobic or amphipathic, and not polymers in the sense of repeating monomers.
  • Four primary classes:
    • Triglycerides
    • Phospholipids
    • Steroids
    • Eicosanoids

Lipids: Triglycerides

  • Most common dietary/lipid storage form.
  • Structure: glycerol + three fatty acids.
  • Formed by dehydration synthesis (ester bonds connect glycerol to fatty acids; water is released).
  • Functions: long-term energy storage in adipose tissue; cushions, insulation, and structural support.
  • Fatty acids vary by length and number of double bonds:
    • Saturated: no double bonds
    • Unsaturated: one or more double bonds (monounsaturated/polyunsaturated)
  • Lipogenesis: synthesis of triglycerides from excess nutrients
  • Lipolysis: breakdown of triglycerides to release fatty acids when needed

Lipids: Phospholipids and Steroids

  • Phospholipids: amphipathic; glycerol backbone with a phosphate group and two fatty acid tails; form lipid bilayers of cell membranes; polar (hydrophilic) head and nonpolar (hydrophobic) tails create a bilayer that constitutes membranes.
  • Steroids: four fused carbon rings; cholesterol as a membrane component and precursor to steroid hormones and bile salts.
  • Prostaglandins (a type of eicosanoid): derived from unsaturated fatty acids; regulate inflammation, blood flow, clotting, and labor induction.
  • Other lipids: glycolipids (lipids with carbohydrate groups), fat-soluble vitamins (A, E, K), etc.

Carbohydrates: Structure and Function

  • General formula: CH2O (empirical formula) for carbohydrates; often written as (CH2O)n.
  • Monosaccharides: simple sugar monomers (e.g., glucose).
  • Disaccharides: two monosaccharides joined; examples include sucrose, lactose, maltose.
  • Polysaccharides: many monosaccharides linked; examples include glycogen, starch, cellulose.
  • Glucose: a six-carbon (hexose) sugar; primary energy source for cells.
  • Glycogen: storage form of glucose in liver and skeletal muscle; glycogenesis (formation) and glycogenolysis (breakdown).
  • Gluconeogenesis: formation of glucose from non-carbohydrate sources in the liver.
  • Isomers: same molecular formula but different structures; e.g., glucose, galactose, fructose (all C6H12O6) with different arrangements leading to different properties.

Nucleic Acids

  • Store and transfer genetic information.
  • Two main classes: DNA (deoxyribonucleic acid) and RNA (ribonucleic acid).
  • Both are polymers of nucleotide monomers linked by phosphodiester bonds.
  • Three components of a nucleotide:
    • Five-carbon sugar: pentose (DNA uses deoxyribose; RNA uses ribose)
    • Phosphate group
    • Nitrogenous base (purines A, G; pyrimidines C, T, U)
  • Five nitrogenous bases are grouped as:
    • Purines: Adenine (A), Guanine (G)
    • Pyrimidines: Cytosine (C), Uracil (U) in RNA, Thymine (T) in DNA
  • Structure of nucleic acids:
    • RNA: single-stranded, contains ribose and uracil
    • DNA: double-stranded, contains deoxyribose and thymine; hydrogen bonds between complementary bases hold the double helix together
  • Adenosine triphosphate (ATP) is a nucleotide central to energy transfer in cells; structure includes adenine base, ribose sugar, and three phosphate groups; energy is stored in the bonds between the last two phosphates and released when hydrolyzed.
  • Chemical energy in nucleotides is fundamental for cellular work and energy transfer.

Proteins

  • Functions: catalysts (enzymes) in metabolism, defense, transport, structural support, movement, regulation, and storage.
  • General protein structure:
    • Made of amino acid monomers; 20 different amino acids in living organisms.
    • Each amino acid contains: an amino group (–NH2), a carboxyl group (–COOH), a hydrogen atom, and a distinctive side chain (R group) attached to a central carbon.
    • Amino acids are joined by peptide bonds to form polypeptide chains.
  • Protein architecture:
    • Primary structure: linear sequence of amino acids joined by peptide bonds.
    • Secondary structure: recurring patterns such as alpha helices and beta pleated sheets.
    • Tertiary structure: three-dimensional conformation of a single polypeptide; influenced by hydrophobic interactions, hydrogen bonds, ionic bonds, disulfide bridges, and other attractions.
    • Quaternary structure: arrangement of two or more polypeptide chains (subunits) in a multi-subunit protein (e.g., hemoglobin with four subunits).
  • Prosthetic groups: nonprotein components covalently bonded to a protein (e.g., heme in hemoglobin).
  • Denaturation: loss of protein structure and function due to factors like heat or pH changes; usually irreversible.
  • Factors influencing protein structure:
    • pH changes can disrupt electrostatic interactions and other bonds, potentially lethal if in critical environments (e.g., blood pH changes).
  • Amino acid properties:
    • Side chains can be charged (positive or negative), polar, or nonpolar, influencing folding and interactions.
  • The secondary structure and higher-order structures are essential for protein function; misfolding can lead to dysfunction.

Energy and Metabolism: ATP and Cellular Respiration

  • Energy is the capacity to do work; forms:
    • Kinetic energy: energy of motion.
    • Potential energy: stored energy.
    • Chemical energy: a form of potential energy stored in chemical bonds.
  • ATP (Adenosine Triphosphate): central energy currency in cells; ATP cycling involves:
    • ATP formation (endergonic): ADP + Pi → ATP (energy input required)
    • Splitting ATP (exergonic): ATP → ADP + Pi (energy release)
  • Cellular respiration overview:
    • Organic molecules (e.g., glucose) are oxidized in a multistep process to produce ATP.
    • Overall equation for aerobic respiration (per glucose):
      extC<em>6extH</em>12extO<em>6+6extO</em>2<br/>ightarrow6extCO<em>2+6extH</em>2extO+extenergy(ATP).ext{C}<em>6 ext{H}</em>{12} ext{O}<em>6 + 6 ext{O}</em>2 <br /> ightarrow 6 ext{CO}<em>2 + 6 ext{H}</em>2 ext{O} + ext{energy (ATP)}.
    • Major stages: glycolysis (cytosol), intermediate stage (mitochondrial), citric acid cycle (Krebs), and electron transport chain with ATP synthase.
    • Typical ATP yield per glucose under aerobic conditions is about 30–32 ATP (values vary with cell type and shuttle systems).
  • Electron carriers involved include NADH and FADH2; their oxidation drives ATP synthesis via oxidative phosphorylation.

Water as the Universal Solvent and Its Properties

  • Water (H2O) is a polar molecule; polarity leads to hydrogen bonding between molecules.
  • Hydrogen bonds form between polar molecules; individually weak but collectively strong, influencing water’s unique properties.
  • Water has three phases (gas, liquid, solid) depending on temperature; body fluids are largely aqueous at body temperature.
  • Functions of liquid water in the body:
    • Transports dissolved substances; acts as a solvent for solutes.
    • Lubricates and cushions joints and tissues.
    • Excretes wastes; substances dissolve and are excreted.
    • Provides high surface tension, enabling certain biological structures to function properly.
  • Water’s high specific heat and high heat of vaporization enable it to regulate body temperature effectively.
  • Water as the universal solvent: polar molecules and ions dissolve readily (hydrophilic); nonpolar substances resist dissolution (hydrophobic).
  • Hydration shells form around dissolved ions or polar molecules; electrolytes dissolve and dissociate to form ions that conduct electricity; nonelectrolytes dissolve but do not conduct.
  • Amphipathic molecules can form micelles and membranes (bilayers) in water; polar heads interact with water while nonpolar tails are excluded.
  • pH and buffers interact with aqueous solutions; buffers mitigate pH changes by binding or releasing H+ ions.

Acids, Bases, and pH

  • Acids dissociate in water to release H+ (proton donors); increases free H+ concentration; stronger acids dissociate more completely.
    • Example: HCl in stomach.
  • Bases accept H+ (proton acceptors); decreases free H+ concentration.
    • Example: Ammonia.
  • pH scale ranges from 0 to 14; neutral pH is 7.0. Lower values indicate acidity; higher values indicate basicity.
  • Relationship:
    • pH = −log[H+], therefore higher [H+] means lower pH and more acidic.
  • Buffers help resist changes in pH by binding excess H+ or OH−; example: carbonic acid (H2CO3) and bicarbonate (HCO3−) buffer blood pH around 7.35–7.45.
  • Neutralization occurs when an acid and base react to form a salt and water, restoring neutral pH under certain conditions.
  • pH values for common substances (illustrative): lemon juice ~2–3, stomach acid ~1–3, milk ~6.3–6.6, blood ~7.35–7.45, household ammonia ~11, household bleach ~12–13.

Water in Chemical Reactions: Dehydration and Hydrolysis

  • Dehydration synthesis (condensation): monomers join by removing OH from one monomer and H from the other, forming a covalent bond and releasing water (H2O).
    • Monomer 1 + Monomer 2 → Dimer/Polymer + H2O
  • Hydrolysis: large molecules are broken down by the addition of water; water is added to split covalent bonds.
    • Polymer + H2O → Monomer 1 + Monomer 2

Chemical Equations and Reaction Types

  • Metabolism encompasses all biochemical reactions in living systems.
  • Chemical reactions occur when existing bonds are broken and new bonds are formed.
  • General form: Reactants → Products with an arrow indicating direction.
    • Example: A + B → C
  • Classification of chemical reactions (Table 3.1 concepts):
    • Change in chemical structure:
    • Decomposition: complex structures broken into simpler structures (e.g., Sucrose → glucose + fructose)
    • Synthesis: simple structures bonded to form more complex structures (e.g., amino acids → dipeptide)
    • Exchange: atoms, molecules, ions, or electrons exchanged between structures (e.g., creatine phosphate + ADP → creatine + ATP)
    • Changes in chemical energy:
    • Exergonic: energy released (e.g., glucose and oxygen → carbon dioxide and water)
    • Endergonic: energy required (e.g., amino acids → dipeptide)
    • Net direction:
    • Irreversible or Reversible (equilibrium).
  • Net biochemical reactions may be considered in the context of overall cellular metabolism and energy balance.

Reaction Rates and Activation Energy

  • Reaction rate: the speed at which a chemical reaction proceeds.
  • Activation energy (Ea): the minimum energy required to rearrange existing bonds for a reaction to occur; a primary determinant of reaction rate.
  • Factors influencing Ea include temperature, concentration, catalysts (enzymes), and particle orientation during collisions.

Enzymes and Their Action

  • Enzymes are biological catalysts that accelerate physiological reactions by lowering Ea.
  • They do not alter the thermodynamics of the reaction; they simply provide an alternative pathway with a lower activation energy.
  • Enzymes are selective: each enzyme catalyzes a specific reaction or type of reaction.
  • Structure: most enzymes are globular proteins with an active site that binds a substrate.
  • Induced-fit model: enzyme undergoes a conformational change upon substrate binding to achieve a tight fit.
  • Enzyme–substrate complex formation accelerates product formation and releases products, enabling the enzyme to catalyze additional reactions.
  • Examples discussed: lactose breakdown by lactase; synthesis of glycogen by glycogen synthase.

Enzyme Structure and Function: Key Concepts

  • Active site specificity ensures only a particular substrate binds.
  • Size and shape of the active site determine catalysis.
  • Enzymes may require cofactors (non-protein helpers) and/or coenzymes for activity.
  • Protein structure is organized into four levels relevant to enzymes: primary, secondary, tertiary, and quaternary.

Energy Overview: ATP, Metabolism, and Cellular Energy Exchange

  • Energy types in biology: kinetic, potential, and chemical energy.
  • ATP is the primary energy currency; energy stored in high-energy phosphate bonds.
  • ATP cycle exemplified by two reactions:
    • Endergonic ATP formation: ADP + Pi → ATP (requires energy input)
    • Exergonic ATP hydrolysis: ATP → ADP + Pi (energy release)
  • Cellular respiration converts organic molecules to ATP via glycolysis, the intermediate stage, the citric acid cycle, and the electron transport chain.
  • Overall cellular respiration yields energy in the form of ATP along with CO2 and H2O as waste products:
    extC<em>6extH</em>12extO<em>6+6extO</em>2<br/>ightarrow6extCO<em>2+6extH</em>2extO+extATP(energy).ext{C}<em>6 ext{H}</em>{12} ext{O}<em>6 + 6 ext{O}</em>2 <br /> ightarrow 6 ext{CO}<em>2 + 6 ext{H}</em>2 ext{O} + ext{ATP (energy)}.

Major Biomolecule Classes (Four Classes of Organic Molecules)

  • Organic biomolecules contain carbon and hydrogen (and often oxygen), and sometimes nitrogen, phosphorus, or sulfur:
    • Lipids
    • Carbohydrates
    • Nucleic acids
    • Proteins
  • Biomolecules form polymers or list-like structures created by monomers; dehydration synthesis builds polymers, hydrolysis breaks them down.
  • Carbon skeletons can take various forms; functional groups influence properties like polarity and hydrogen-bonding capacity.
  • Functional groups are typically polar and hydrophilic; some act as acids (e.g., carboxyl) or bases (e.g., amino groups).

Lipids (Detailed)

  • Lipids are diverse and not all polymers; they are hydrophobic or amphipathic.
  • Four primary lipid classes:
    1) Triglycerides
    2) Phospholipids
    3) Steroids
    4) Eicosanoids
  • Triglycerides:
    • Structure: glycerol + three fatty acids; formed via dehydration synthesis; energy-dense storage form in adipose tissue; provide cushioning and insulation.
    • Fatty acid types:
    • Saturated: no double bonds
    • Unsaturated: one or more double bonds (monounsaturated/polyunsaturated)
  • Phospholipids:
    • Amphipathic: polar hydrophilic head and nonpolar hydrophobic tails.
    • Key component of cell membranes; form lipid bilayers as barrier structures.
  • Steroids:
    • Four-ring hydrocarbon skeleton; cholesterol as membrane component and precursor to steroids (hormones, bile acids).
  • Prostaglandins (a type of eicosanoid):
    • Derived from polyunsaturated fatty acids; regulate inflammation, blood flow, clotting, and labor.
  • Other lipids:
    • Glycolipids, fat-soluble vitamins (A, E, K).

Carbohydrates (Detailed)

  • Carbohydrates: general formula extCH<em>2extOext{CH}<em>2 ext{O} (empirical formula) or $( ext{CH}2 ext{O})_n$; carbohydrate class includes:
    • Monosaccharides (simple sugars)
    • Disaccharides (two monosaccharides linked)
    • Polysaccharides (many monosaccharides linked)
  • Glucose: a six-carbon (hexose) monosaccharide; primary energy source for cells.
  • Glycogen: storage form of glucose in liver and skeletal muscle; glycogenesis (formation) and glycogenolysis (breakdown).
  • Gluconeogenesis: synthesis of glucose from noncarbohydrate sources in the liver.
  • Isomers: molecules with the same molecular formula but different structural arrangements; e.g., glucose, galactose, and fructose (all C6H12O6) with different properties.
  • Plant starch and cellulose:
    • Starch: energy storage in plants (glucose polymers)
    • Cellulose: structural carbohydrate in plant cell walls; dietary fiber; indigestible by humans.

Nucleic Acids

  • Two main classes: DNA (deoxyribonucleic acid) and RNA (ribonucleic acid).
  • Both are polymers composed of nucleotide monomers linked by phosphodiester bonds.
  • Nucleotide components:
    • Five-carbon sugar: pentose (deoxyribose in DNA; ribose in RNA)
    • Phosphate group
    • Nitrogenous base (purines A, G; pyrimidines C, T, U)
  • Bases by category:
    • Purines: Adenine (A), Guanine (G)
    • Pyrimidines: Cytosine (C), Thymine (T) in DNA, Uracil (U) in RNA
  • DNA is double-stranded with hydrogen-bonded base pairs; RNA is typically single-stranded.
  • Adenosine triphosphate (ATP) is a nucleotide and a key energy carrier in cells: structure includes adenine, ribose, and three phosphate groups; energy stored in the bond between the last two phosphates and released upon hydrolysis.

Proteins (Detailed)

  • Functions:
    • Catalysts (enzymes) in metabolism
    • Defense, transport, structural support, movement, regulation, and storage
  • Protein structure and organization:
    • Amino acids: 20 standard amino acids with amino group (-NH2), carboxyl group (-COOH), hydrogen, and distinctive side chain (R).
    • Peptide bonds link amino acids into polypeptides.
  • Levels of structure:
    • Primary: amino acid sequence
    • Secondary: alpha helices and beta sheets formed by hydrogen bonding
    • Tertiary: three-dimensional folding of a single polypeptide
    • Quaternary: arrangement of multiple polypeptide chains (subunits) in a protein such as hemoglobin
  • Denaturation: disruption of protein structure due to heat or pH changes; often irreversible.
  • Prosthetic groups: nonprotein components covalently bonded to proteins (e.g., heme in hemoglobin).
  • The amino acid side chains determine protein folding, stability, and interactions.
  • The concept of chaperone proteins helps proteins fold into correct conformations.

Metabolism and Energy: Consolidated View

  • Energy forms: kinetic, potential, and chemical energy.
  • Chemical energy is stored in the bonds of molecules and is released during chemical reactions.
  • ATP cycling couples exergonic and endergonic reactions to power cellular activities.
  • Cellular respiration is the major pathway to convert glucose into ATP, with oxygen required for efficient energy production.

Summary of Key Equations and Notations (LaTeX)

  • Water molecule: extH2extOext{H}_2 ext{O}
  • Glucose + Oxygen reaction (cellular respiration):
    extC<em>6extH</em>12extO<em>6+6extO</em>2<br/>ightarrow6extCO<em>2+6extH</em>2extO+extATPext{C}<em>6 ext{H}</em>{12} ext{O}<em>6 + 6 ext{O}</em>2 <br /> ightarrow 6 ext{CO}<em>2 + 6 ext{H}</em>2 ext{O} + ext{ATP}
  • pH relationship: extpH=−<br/>olinebreaklog⁡[H+]ext{pH} = -<br /> olinebreak \log[H^+]
  • Ionic bond general form: extA++extB−<br/>ightarrowextABext{A}^+ + ext{B}^- <br /> ightarrow ext{AB} (ionic compound)
  • Ionic notation examples: $^{12}{6} ext{C}$, $^{13}{6} ext{C}$, $^{14}_{6} ext{C}$
  • Dehydration synthesis (generic): extMonomer<em>1+extMonomer</em>2<br/>ightarrowextPolymer+extH2extOext{Monomer}<em>1 + ext{Monomer}</em>2 <br /> ightarrow ext{Polymer} + ext{H}_2 ext{O}
  • Hydrolysis (generic): extPolymer+extH<em>2extOightarrowextMonomer</em>1+extMonomer2ext{Polymer} + ext{H}<em>2 ext{O} ightarrow ext{Monomer}</em>1 + ext{Monomer}_2
  • ATP cycle basics:
    • Formation: extADP+extPi<br/>ightarrowextATPext{ADP} + ext{P}_i <br /> ightarrow ext{ATP} (endergonic)
    • Hydrolysis: extATP<br/>ightarrowextADP+extPiext{ATP} <br /> ightarrow ext{ADP} + ext{P}_i (exergonic)
  • Nucleotide structure (general): sugar + phosphate + nitrogenous base (e.g., ATP, AMP, etc.)

Connections to Foundational Principles and Real-World Relevance

  • The chemical level provides the foundation for understanding how biomolecules interact to build cells, tissues, and organs.
  • Bond types (ionic vs covalent) underlie the stability, reactivity, and function of molecules in physiology.
  • Water’s properties (polarity, hydrogen bonding, solvent capabilities) explain transport, hydration, temperature regulation, and biochemical reactions in the body.
  • The concept of energy flow (ATP) connects chemistry to cellular physiology, growth, muscle contraction, nerve signaling, and metabolism.
  • Enzymes illustrate the principle that biological systems optimize reaction rates to sustain life, emphasizing specificity, active sites, and regulation.
  • Macromolecules (lipids, carbohydrates, nucleic acids, proteins) reveal how structure dictates function, enabling energy storage, information transfer, and catalysis.

Practical Implications and Ethical/Philosophical Points (Brief)

  • Understanding isotopes and radioactive tracers informs medical imaging, cancer treatment, and historical dating methods.
  • pH balance and buffers illustrate the body's homeostatic regulation and the fragile limits of physiological conditions.
  • Denaturation highlights the importance of environmental stability in maintaining protein function, relevant to food science, medicine, and disease.
  • The balance of energy production and consumption is central to health, athletic performance, and metabolic disorders.
  • Knowledge of biomolecules underpins nutrition science, pharmacology, and biotechnology.

Quick Reference: Key Concepts by Topic

  • Matter and Atoms: mass, space, three forms; atoms, subatomic particles; atomic number and mass.
  • Bonding: ionic vs covalent; polar vs nonpolar; amphipathic molecules; hydrogen bonding.
  • Water: solvent, high heat capacity, hydrogen bonding, hydration shells, pH, buffers.
  • Acids/Bases: proton donors/acceptors; pH scale; neutralization.
  • Macromolecules: lipids, carbohydrates, nucleic acids, proteins; monomers, polymers, dehydration/hydrolysis.
  • Energy: ATP generation and use; cellular respiration steps; energy carriers (NADH, FADH2).
  • Genetics: DNA/RNA structure and function; nucleotides; transcription/translation foreshadowed by these biomolecules.

End of Chapter 2 Notes