Chapter 2: The Chemical Level of Organization (Principles of Anatomy and Physiology)
Basic Principles of Chemistry
Chemistry is the science of the structure and interactions of matter.
Matter is anything that has mass and takes up space.
Mass is the amount of matter contained in a substance; weight is the force of gravity acting on that mass.
Matter and Forms of Matter
Matter exists in three fundamental forms:
Solid: definite shape and volume
Liquid: definite volume, takes the shape of its container
Gas: no definite shape or volume
All forms of matter are composed of chemical elements.
Chemical Elements
Elements are given chemical symbols (examples):
O = oxygen, C = carbon, H = hydrogen, N = nitrogen
These four elements make up the majority of the human body.
Atoms
Atoms are the smallest units of matter that retain the properties of an element.
Subatomic particles:
Protons (p+), located in the nucleus
Neutrons (n0), located in the nucleus
Electrons (e−), negatively charged, orbit the nucleus in the electron cloud
Atoms are composed of a nucleus (protons and neutrons) and surrounding electron cloud.
Atomic Structures (key isotopes and values)
Hydrogen (H):
Atomic number Z = 1; Mass number A = 1 or 2; Atomic mass ≈ 1.01
Carbon (C):
Z = 6; A = 12 or 13; Atomic mass ≈ 12.01
Nitrogen (N):
Z = 7; A = 14 or 15; Atomic mass ≈ 14.01
Oxygen (O):
Z = 8; A = 16, 17, or 18; Atomic mass ≈ 16.00
Sodium (Na):
Z = 11; A = 23; Atomic mass ≈ 22.99
Chlorine (Cl):
Z = 17; A = 35 or 37; Atomic mass ≈ 35.45
Potassium (K):
Z = 19; A = 39, 40, or 41; Atomic mass ≈ 39.10
Iodine (I):
Z = 53; A = 127; Atomic mass ≈ 126.90
Isotope note: Mass number A = number of protons + neutrons; boldface in slides indicates the most common isotope.
Ions, Molecules, and Compounds
Ions: atoms that have lost or gained electrons, resulting in a positive (cation) or negative (anion) charge.
Ion formation depends on valence electrons to achieve stability.
Molecule: two or more atoms sharing electrons.
Compound: a substance composed of two or more different elements.
Free Radicals
Free radicals have an unpaired electron in the outer shell, making them unstable, highly reactive, and potentially destructive to nearby molecules.
Sources include UV light, ozone, x-rays, pollution, cigarette smoke.
Antioxidants help neutralize reactive oxygen species.
Chemical Bonds
A chemical bond occurs when atoms are held together by forces of attraction.
The likelihood of bond formation is largely determined by the number of electrons in an atom’s valence (outer) shell.
The octet rule (rule of eight): a stable valence shell has 8 electrons.
Bonding outcomes: gain, loss, or sharing of electrons depending on valence electron count and energetics.
The periodic table helps predict valence electron counts and bonding behavior.
Ionic Bonds
Ionic bonds form when a positively charged ion (cation) is attracted to a negatively charged ion (anion).
No electrons are shared; instead, electrons are transferred to achieve stable electron configurations.
Example: sodium chloride, NaCl.
Na tends to lose one electron to form Na+ with a complete valence shell.
Cl tends to gain an electron to form Cl− with a complete valence shell.
Cations vs. Anions: note that the electron transferred during bond formation may come from a different atom than the one that accepts it.
Ionic bonds create compounds with overall neutral charge but composed of oppositely charged ions.
Covalent Bonds
Covalent bonds form when two or more atoms share electrons.
Can share one (single), two (double), or three (triple) pairs of electrons.
Strength generally increases with the number of shared electron pairs.
Covalent bonds can occur between atoms of the same element or different elements.
Polar vs. Nonpolar Covalent Bonds:
Nonpolar covalent: electrons shared equally; no partial charges.
Polar covalent: electrons pulled more toward one nucleus, creating partial charges (dipole).
Hydrogen Bonds
Hydrogen bonds arise from the attraction between oppositely charged parts of polar molecules.
They are much weaker than ionic or covalent bonds but are crucial for the properties of water and biomolecules.
Structural Organization in Water and Hydrogen Bonding
Water’s polarity makes it an excellent solvent for ionic and polar substances (hydrophilic) and poor solvent for nonpolar substances (hydrophobic).
Hydrogen bonding between water molecules provides cohesion and surface tension.
Water plays a key role in biochemical reactions as a solvent and participant (hydrolysis and dehydration synthesis).
Chemical Reactions
Reactions occur when bonds are formed or broken; reactants become products.
Metabolism refers to all chemical processes in organisms, including energy transformations.
Forms of Energy & Chemical Reactions
Energy types:
Potential energy: stored energy
Kinetic energy: energy of motion
Chemical energy: energy stored in chemical bonds
Law of conservation of energy: energy cannot be created or destroyed, only converted.
Reactions may release energy (exergonic) or require energy input (endergonic).
Cellular coupling: exergonic reactions power endergonic ones (e.g., glucose catabolism releases energy; ATP bonds store energy released by breakdown).
Activation Energy and Catalysts
Activation energy (E_a): energy required to start a reaction; reactants must collide with sufficient energy/orientation to react.
Catalysts speed up reactions; most are enzymes; catalysts are not consumed in the reaction and can catalyze many cycles.
Body conditions (temperature, concentration) influence reaction rates; enzymes help overcome these limitations.
Types of Chemical Reactions
Synthesis (anabolism): two or more substances combine to form a larger, more complex product.
Decomposition (catabolism): larger molecules split into smaller components.
Exchange (displacement): parts of molecules exchange places.
Reversible: reactions can proceed in both directions; products can revert to reactants.
Oxidation-reduction (redox): electron transfer between reactants; oxidation involves loss of electrons and energy release; reduction involves gain of electrons and energy gain.
Inorganic vs Organic Compounds and Solutions
Inorganic compounds: typically lack carbon and are simpler (water is the most important inorganic compound in living things).
Organic compounds: contain carbon, usually with hydrogen, and form covalent bonds; typically large molecules with carbon chains.
Water properties and role in solutions are central to many biological processes.
Properties of Water and Solvent Capabilities
Water is a versatile solvent for ionic and polar substances (hydrophilic) due to its polarity.
Nonpolar substances (lipids) are hydrophobic.
Water participates in reactions as a reactant or product (hydrolysis and dehydration synthesis).
Water exhibits high heat capacity and high heat of vaporization, enabling thermal regulation in organisms.
Water acts as a lubricant in body cavities and joints.
Percentages and Molarity in Solutions
Percentage (mass per volume): grams of solute per 100 milliliters of solution.
Example: to make a 10% NaCl solution, dissolve 10 g NaCl in enough water to make 100 mL solution.
Molarity (M, moles per liter): 1 M = 1 mole of solute per liter of solution.
Example: to make 1 M NaCl, dissolve 1 mole of NaCl (58.44 g) in water to make 1 L of solution.
Note: A mole is defined as the amount of substance containing as many elementary entities as there are atoms in 12 g of carbon-12; the slide humorously adds, “A mole is the amount of any substance that has a mass in grams equal to the sum of the atomic masses of all its atoms,” and includes a joking aside about the mole’s origin.
A key aside from the slides: the exact numerical value 58.44 g/mol is the molar mass of NaCl.
Acids, Bases, & Salts
Dissociation in water:
Acids dissociate to yield H+ and anions (proton donors).
Bases dissociate to yield OH− (and cations) (proton acceptors).
Salts dissociate into cations and anions, neither being H+ or OH−.
Examples (illustrated in slides): HCl (acid), KOH (base), KCl (salt).
pH and Buffers
pH measures the hydrogen ion concentration in solution.
pH scale ranges from 0 to 14; lower values mean more acidic; 7 is neutral; higher values are basic.
The scale is logarithmic: a change of one pH unit represents a tenfold change in hydrogen ion concentration.
Expressions:
ext{pH} = -
\log_{10} [H^+][H^+] = 10^{- ext{pH}}
A neutral pH at 25°C corresponds to [H^+] = [OH^-] = 1.0 imes 10^{-7} ext{ M}.
Buffers maintain blood and body fluid pH within a narrow range by converting strong acids or bases into weak acids or bases, often involving H+ donation or removal. Example: bicarbonate ion (
ext{HCO}3^-) neutralizes an acid by forming ext{H}2 ext{CO}3: ext{HCO}3^- + ext{H}^+
ightarrow ext{H}2 ext{CO}3.
Overview of Organic Compounds
Carbon is the backbone of organic chemistry; carbon forms four bonds (valence = 4).
Carbon can form chains, rings, and various shapes; carbon compounds are often hydrophobic in water.
Carbon compounds are key energy sources; hydrocarbon skeletons may be bonded to functional groups that determine reactivity.
Major Functional Groups of Organic Molecules (selected)
Hydroxyl group: –OH (polar, hydrophilic; many –OH groups increase water solubility)
Sulfhydryl group: –SH (polar, can stabilize protein structure via disulfide bonds; e.g., cysteine)
Carbonyl group: C=O (polar; present in ketones and aldehydes)
Aldehydes: carbonyl at the end of the carbon skeleton
Ketones: carbonyl within the carbon skeleton
Carboxyl group: –COOH (acidic; amino acids have –COOH; deprotonates at body pH to become –COO−)
Ester: –COO– (common in fats/oils; triglycerides)
Phosphate: –PO4^{2−} (very hydrophilic; energy carrier in ATP)
Amines: –NH2 (acts as a base in physiology; amino acids contain –NH2)
Carbohydrates
Composed of C, H, O; H:O ratio ≈ 2:1 (author calls it “watered carbon”).
Main energy source for cells; include sugars, glycogen, starch, cellulose.
Classes based on size: monosaccharides, disaccharides, polysaccharides.
Monosaccharides examples: Glucose, Fructose, Galactose; also ribose and deoxyribose (pentoses).
Disaccharides: Sucrose (glucose + fructose), Lactose (glucose + galactose), Maltose (glucose + glucose).
Polysaccharides: Glycogen, Starch, Cellulose.
Lipids
Lipids are hydrophobic and insoluble in water.
Classes include fatty acids, triglycerides, phospholipids, steroids, and more.
Fatty acids can be saturated or unsaturated.
Triglycerides: energy storage, insulation, protection; made from glycerol plus three fatty acids via ester linkages.
Phospholipids: major components of cell membranes; amphipathic with polar heads and nonpolar tails.
Steroids: cholesterol (membrane component; precursor to bile salts, vitamin D, steroid hormones); other steroids include bile salts, cortisol, and sex hormones.
Eicosanoids: prostaglandins and leukotrienes; regulate inflammation, immunity, etc.
Other lipids: carotene (vitamin A precursor/antioxidant), vitamin E (antioxidant and wound healing), vitamin K (blood clotting), lipoproteins (lipid transport in blood).
Fatty Acids and Triglycerides (structural notes)
Palmitic acid (saturated, C16:0) vs. oleic acid (unsaturated, C18:1).
Dehydration synthesis links fatty acids to glycerol to form triglycerides; hydrolysis can release fatty acids.
Phospholipids and Membranes
Phospholipids have polar heads (hydrophilic) and nonpolar tails (hydrophobic).
Arranged in a bilayer to form cell membranes; polar heads face aqueous environments, nonpolar tails face inward.
Steroids
Four-ring hydrocarbon skeleton.
Examples: Cholesterol; Estradiol; Testosterone; Cortisol.
Proteins
Proteins are large nitrogen-containing molecules with many functions:
Structural: collagen, keratin
Regulatory: hormones (e.g., insulin), signaling molecules
Contractile: actin, myosin
Immunological: antibodies
Transport: hemoglobin
Catalytic: enzymes (e.g., amylase, sucrase, ATPase)
Amino Acids: building blocks of proteins; each amino acid has:
Central carbon (C)
Hydrogen atom (H)
Amino group (–NH2)
Carboxyl group (–COOH)
Side chain (R group) that defines identity
Peptide bonds: link amino acids; a protein is a chain of amino acids connected by peptide bonds.
Protein Structure (four levels)
Primary structure: unique amino acid sequence (linear chain).
Secondary structure: local folding/alignment (alpha helices and beta pleated sheets) stabilized by hydrogen bonds.
Tertiary structure: three-dimensional folding of a single polypeptide chain.
Quaternary structure: assembly of two or more polypeptide chains (many proteins stop at tertiary structure).
Enzymes
Enzymes are biological catalysts that speed up reactions.
Names typically end in -ase.
Properties:
Highly specific: each enzyme fits one substrate.
Extremely efficient: can increase reaction rates dramatically (up to billions of times).
Regulated by cellular controls: enzymes can be activated or inhibited.
Active site: region where substrate binds; formation of the enzyme–substrate complex enables the chemical reaction.
Nucleic Acids
Nucleic acids are chains of nucleotides.
Each nucleotide consists of:
Nitrogenous base (adenine A, thymine T, cytosine C, guanine G, or uracil U in RNA)
Pentose sugar (deoxyribose in DNA, ribose in RNA)
Phosphate group
DNA forms the genetic code in cell nuclei; RNA relays instructions from genes to guide protein synthesis.
Components of a Nucleotide (illustrative)
Nitrogenous bases: purines (A, G) and pyrimidines (C, T, U).
Sugar: deoxyribose (DNA) or ribose (RNA).
Phosphate group connects sugars to form the backbone.
DNA and RNA Structures and Pairs
DNA is a double helix with two strands of nucleotides.
Bases pair via hydrogen bonds: A pairs with T (2 hydrogen bonds), G pairs with C (3 hydrogen bonds).
RNA is typically single-stranded; in DNA vs RNA distinctions:
DNA: deoxyribose sugar; bases A, T, C, G; two strands; A–T and G–C pairings.
RNA: ribose sugar; bases A, U, C, G; usually single strand; A–U and G–C pairings.
Adenosine Triphosphate (ATP)
ATP is the principal energy-storing molecule in the body.
Structure: adenosine attached to three phosphate groups (adenosine triphosphate).
High-energy phosphate bonds: cleavage releases usable energy for cellular processes.
General note: ATP hydrolysis provides energy for muscle contraction, active transport, and many biosynthetic reactions.
Cellular ATP Production (overview)
Cellular respiration catabolizes glucose to release energy that is stored in ATP bonds.
Two main phases:
Anaerobic respiration: without oxygen; glucose partially broken down to pyruvic acid; yields 2 ATP per glucose.
Aerobic respiration: with oxygen; glucose is broken down to CO2 and H2O; yields approximately 30–32 ATP per glucose.
Detailed chapters on respiration are covered later in the course.
Connections to Foundational Principles
Matter and energy transformation: chemical reactions involve conversion between potential, kinetic, and chemical energy.
The covalent/ionic/hydrogen bonding framework explains molecular stability, structure, and function in biomolecules.
The property of water as a solvent underpins hydration, ionic interactions, and biochemical reaction environments.
Organic chemistry underpins biomolecular structure and function (carbohydrates, lipids, proteins, nucleic acids).
Buffers and pH homeostasis are essential for maintaining protein structure and enzyme activity.
Practical and Ethical/Contextual Notes
The study of chemical principles is essential for understanding physiology, medicine, and biochemistry.
Some slides include lighthearted asides (e.g., about molarity humor) — educational context should prioritize scientific accuracy in exams.
Real-world relevance: acid-base balance, energy metabolism, and macromolecule function underpin health, disease, and pharmacology.
[H^+] = 10^{- ext{pH}}
[OH^-] = 10^{- ext{pOH}}
ext{pH} = -
\log_{10}[H^+]
1 ext{ M} = rac{1 ext{ mole}}{1 ext{ liter}}
58.44 ext{ g/mol} for NaCl
Example buffer reaction: ext{HCO}3^- + ext{H}^+ ightarrow ext{H}2 ext{CO}_3
For a typical neutral body pH: [H^+] \approx 1.0 \times 10^{-7} \text{ M} and [OH^-] \approx 1.0 \times 10^{-7} \text{ M} at 25°C.
Key practical values from the slides:
Gastric juice pH: 1.2 - 3.0
Lemon juice: pH \approx 2.3
Vinegar: pH \approx 3.0
Orange juice: pH \approx 3.5
Blood: pH \approx 7.35 - 7.45
Cerebrospinal fluid: pH \approx 7.4
Urine: pH \approx 4.6 - 8.0
Saliva: pH \approx 6.35 - 6.85
Milk: pH \approx 6.8
Distilled water: pH = 7.0
Lye (strong base): pH = 14.0
Major groups and example molecules provide a framework to analyze biomolecules in exams (e.g., identify functional groups in a given molecule, predict solubility, or predict bonding patterns).
Summary of Key Formulas and Concepts
Atomic structure and isotopes: Z = number of protons; A = Z + number of neutrons; common isotopes are noted by mass number.
Ionic vs covalent bonding; polarity and partial charges in covalent bonds influence solubility and biological interactions.
Hydrogen bonding, cohesion, and surface tension are essential for water properties and biomolecule interactions.
Energy changes in reactions: exergonic vs endergonic; activation energy; catalysis by enzymes.
Carbohydrates, lipids, proteins, and nucleic acids form the core of biological macromolecules; their monomers, polymers, and functional roles define cellular structure and function.
ATP and cellular respiration as central energy pathways; understanding glucose catabolism and ATP synthesis is foundational for physiology and bioenergetics.