Notes on Matter, Atoms, and Molecules

2.1a Matter, Atoms, Elements, and the Periodic Table

• Matter has mass and occupies space.
  • 3 forms: solid (e.g., bone), liquid (e.g., blood), gas (e.g., oxygen).
• Atom = smallest particle exhibiting chemical properties of an element.
• 92 naturally occurring elements make up matter.
• Elements are organized in the periodic table of elements.
• Periodic table features and trends:
  • Arranged by atomic number (number of protons).
  • Increasing electronegativity across a period (left to right) and generally decreasing down a group.
  • Each element has a chemical symbol (e.g., H for hydrogen) and an average atomic mass.
• Key concepts to know:
  • Atomic number Z = number of protons.
  • Mass number A ≈ number of protons + neutrons.
  • Average atomic mass is shown below the symbol on the table.
• Components of an atom (Page 5–8):
  • Protons: mass ~ 1 amu, +1 charge; located in the nucleus.
  • Neutrons: mass ~ 1 amu, 0 charge; located in the nucleus.
  • Electrons: mass ~ 1/1800 amu, −1 charge; located in electron orbitals surrounding the nucleus.
• Determining subatomic particle numbers (Page 7):
  • Proton number = atomic number Z.
  • Neutron number = atomic mass A − atomic number Z.
  • Example: Na has A = 23, Z = 11 → Neutrons = 23 − 11 = 12.
  • Electron number = proton number for a neutral atom; equals Z.
• Diagramming atomic structure (Page 8–9):
  • Atoms have electron shells/orbitals with specific energy levels.
  • Inner shells fill first: the innermost shell holds 2 electrons; the second shell can hold up to 8; outer shells tend to be filled to satisfy stability rules before outer shells.
• The shell model vs. the cloud model (Figure 2.2): nucleus contains protons and neutrons; electron shells or electron clouds surround the nucleus.
• The Octet Rule (Page 10):
  • Elements tend to lose, gain, or share electrons to obtain a complete outer shell of 8 electrons (an octet).
  • Noble gases (column VIIIA) have complete outer shells and are generally stable and unreactive.
• Organization of the periodic table by valence electrons (Page 11):
  • Elements are organized in a way that reflects the number of electrons in their outer (valence) shell.
• Summary of key equations and concepts:
  • Neutron number: $N = A - Z$
  • Electron number (neutral atom): $e = Z$
  • Octet rule: outer shell stable when it contains 8 electrons (except for very light elements with complete outer shells).

2.2 Ions and Ionic Compounds

• Chemical compounds: stable associations between two or more elements in fixed ratios.
  • Classified as ionic or molecular.
  • Ionic compounds are structures of ions held together by ionic bonds in a lattice (crystal) structure.
• Ions:
  • Atoms with positive charge (cations) or negative charge (anions).
  • Formed by gain or loss of one or more electrons.
  • Important physiological functions (e.g., electrolytes in body fluids).
  • In some contexts, large doses of certain ions are used for specific purposes (e.g., potassium chloride in lethal injections; electrolytes in sports drinks).
• Formation of ions and ionic compounds (Page 14–15):
  • Example: Sodium (Na) donates an electron to reach stability (octet).
  • Na becomes Na⁺; atomic number Z = 11, but now 11 protons and 10 electrons.
  • Chlorine (Cl) gains an electron to reach stability; becomes Cl⁻.
  • Na⁺ and Cl⁻ form NaCl, an ionic compound held together by ionic bonds.
  • Conceptual diagram: Na atom → Na⁺ + e⁻; Cl atom + e⁻ → Cl⁻; Na⁺/Cl⁻ attract to form NaCl.
• Polyatomic ions (Page 16):
  • Ions composed of more than one atom (e.g., bicarbonate HCO₃⁻, phosphate PO₄³⁻).
• Ionic bonds (Page 17):
  • Cations and anions are bound by electrostatic attraction.
  • Examples: table salt NaCl; magnesium chloride MgCl₂ (Mg²⁺ donates electrons to Cl⁻).
• Real-world relevance:
  • Electrolyte balance is essential for nerve conduction, muscle function, and fluid balance.

2.3 Covalent Bonding, Molecules, and Molecular Compounds

• Covalent bonding: electrons are shared between atoms.
  • Forms molecules; molecules with two or more different elements are molecular compounds (e.g., CO₂).
  • O₂ is a molecule but not a molecular compound because it consists of the same element (diatomic oxygen).
• Chemical formulas: molecular vs structural (Page 19):
  • Molecular formula: indicates the number and type of atoms (e.g., H₂CO₃).
  • Structural formula: shows arrangement of atoms and bonds (e.g., O=C=O for CO₂).
• Isomers (Page 20):
  • Glucose, galactose, and fructose have the same molecular formula (C₆H₁₂O₆) but different arrangements, leading to different properties.
• Covalent bonds (Page 21–23):
  • Atoms form covalent bonds to fulfill outer-shell (octet) needs.
  • Commonly form bonds in the human body with H, O, N, and C.
  • Bond types:
  • Single covalent bond: one shared electron pair (e.g., H–H).
  • Double covalent bond: two shared electron pairs (e.g., O=O).
  • Triple covalent bond: three shared electron pairs (e.g., N≡N).
  • Carbon has a valence of four, allowing it to form up to four covalent bonds.
• Polar and nonpolar covalent bonds (Pages 25–27):
  • Electronegativity determines how electrons are shared.
  • Nonpolar covalent bonds: equal sharing, typically between atoms of the same element (e.g., H–H, O–O).
  • Polar covalent bonds: unequal sharing, creating partial charges (δ+ and δ−).
  • Most common living-element pairings in biology: H, C, N, O (in order of increasing electronegativity: H < C < N < O).
  • Carbon–hydrogen bonds are exception: sharing is relatively nonpolar even though atoms differ.
• Partial charges and polarity (Page 27):
  • More electronegative atom gains a partial negative charge (δ−); less electronegative atom gains a partial positive charge (δ+).
  • Represented with δ− and δ+.
• Nonpolar, polar, and amphipathic molecules (Pages 28–31):
  • Nonpolar molecules contain nonpolar covalent bonds (e.g., O–O; C–H).
  • Polar molecules contain polar covalent bonds (e.g., O–H in H₂O).
  • Some molecules contain polar bonds that cancel out, resulting in a nonpolar molecule (e.g., CO₂).
  • Amphipathic molecules have both hydrophilic (polar) and hydrophobic (nonpolar) regions (e.g., phospholipids, glycerol backbone with phosphate head and fatty acid tails).
  • Glycerol and phospholipid examples illustrate amphipathic properties and membrane architecture (bilayers and micelles).
• Intermolecular attractions (Pages 32–33):
  • Hydrogen bonds: a specific, relatively weak attraction between polarized molecules, particularly when H is bonded to O, N, or F.
  • Hydrogen bonds form between water molecules and contribute to water’s unique properties; they are also crucial in DNA and proteins structure.

2.4 Molecular Structure and Properties of Water

• Water in biology:
  • Water is a polar molecule: one O atom with partial negative charge and two H atoms with partial positive charge.
  • The molecule has a bent shape and forms four hydrogen bonds with neighboring molecules, contributing to water’s unique properties.
  • Structural illustration: H–O–H with partial charges on H and O.
• Water’s classification of substances (Page 34):
  • Organic molecules contain carbon and are carbon-based; inorganic molecules include water, salts, acids, and bases.
• Water structure and polarity (Page 35–37):
  • Central to water’s properties; polar covalent bonds create dipole moments.
  • Water can form four hydrogen bonds per molecule, contributing to properties such as high cohesion and high surface tension.
• Phases and functions of water (Page 37–41):
  • Three phases: gas (water vapor), liquid (water), solid (ice).
  • Functions of liquid water: transports dissolved substances, lubricates, cushions, excretes wastes, and dissolves many substances.
  • Cohesion: attraction between water molecules due to hydrogen bonding.
  • Surface tension: inward pulling of cohesive forces at the water surface; surfactants prevent alveolar collapse in lungs.
  • Adhesion: attraction between water and other substances.
  • High specific heat and high heat of vaporization:
  • Specific heat = energy required to raise the temperature of 1 g of a substance by 1°C.
  • Water has a very high specific heat due to strong hydrogen bonding, helping stabilize body temperature.
  • Heat of vaporization: energy required to vaporize 1 g of a liquid; water’s value is high, enabling sweating to cool the body.
• Water as the universal solvent (Page 42–45):
  • Water dissolves many substances; called the universal solvent given its solvent properties.
  • Hydrophilic substances dissolve in water and form a hydration shell; some dissolve but do not dissociate (e.g., glucose, ethanol).
  • Electrolytes dissolve and dissociate into ions (NaCl → Na⁺ + Cl⁻); electrolytes conduct current in solution.
  • Hydrophobic (water-fearing) substances do not dissolve in water (e.g., fats, cholesterol); hydrophobic exclusion occurs.
  • Amphipathic molecules partially dissolve in water and may form bilayers or emulsions (e.g., phospholipids in membranes; bile salts).
  • Hydration shells surround dissolved polar molecules or ions, aiding transport in body fluids.
• Substances and their interactions with water (Fig. 2.13):
  • Hydrophilic substances (e.g., glucose) dissolve and may dissociate (electrolytes) or remain intact (nonelectrolytes).
  • Hydrophobic substances do not dissolve; nonpolar molecules interact via hydrophobic interactions.
  • Amphipathic molecules can form bilayers (phospholipids) or micelles in aqueous environments.

2.5a Water: A Neutral Solvent

• Water self-ionizes (autoprotolysis):
  • Water dissociates to form ions: ext{2 H}2 ext{O} ightleftharpoons ext{H}3 ext{O}^+ + ext{OH}^-
  • Very low ion concentration in pure water (roughly 1 in 10,000,000 per liter).
  • Hydronium ion: $ext{H}_3 ext{O}^+$; hydroxide ion: $ext{OH}^-$.
  • Net charge remains zero (neutral solution).

2.5b Acids and Bases

• Arrhenius definitions (Page 47–48):
  • Acid: dissociates in water to produce H⁺ (a proton) and an anion; proton donor.
  • Base: accepts a proton (H⁺) in solution; proton acceptor.
  • Examples:
  • Strong acid: HCl → ext{HCl}
    ightarrow ext{H}^+ + ext{Cl}^- (in stomach).
  • Weak acid: carbonic acid (H₂CO₃) in blood; less dissociation.
  • Strong base: ammonia or bleach (NH₃ or NaOH as bases in solution).
• Acid-base reactions and buffer bases (noted in content):
  • Buffers help resist pH changes by accepting H⁺ or donating H⁺ as needed.

2.5c pH, Neutralization, and the Action of Buffers

• pH scale basics (Page 49–51):
  • pH = measure of relative [H⁺] in solution; scale ranges 0–14.
  • Neutral solution has pH 7; acidic solutions have pH < 7; basic/alkaline solutions have pH > 7.
  • The pH and hydrogen ion concentration are inversely related: as [H⁺] increases, pH decreases.
  • A 10-fold change in hydrogen ion concentration corresponds to a one-unit change in pH.
• Neutralization and buffers (Page 52):
  • Neutralization: turning an acidic or basic solution back toward neutral (pH 7) by adding the opposite.
  • Buffers: substances that resist pH changes by accepting H⁺ from excess acid or donating H⁺ to neutralize base.
  • Blood pH is maintained in a narrow range (7.35–7.45) by buffers such as carbonic acid (weak acid) and bicarbonate (weak base).

2.6 Water Mixtures

• Mixtures defined: two or more substances physically combined without chemical change; can be separated by physical means.
• Categories of water mixtures (Page 54–56):
  • Solution: homogeneous, particles < 1 nm; substances dissolve in water; light does not scatter; e.g., sugar water, salt water, blood plasma.
  • Colloid: particles larger than those in a solution but smaller than suspensions; remains mixed; scatters light; e.g., cytosol fluid, plasma components.
  • Suspension: larger particles (>1 mm) mixed with water; not mixed unless in motion; appears cloudy.
  • Emulsion: a special colloid formed when a polar substance (water) interacts with a nonpolar liquid (oil); requires agitation to mix (e.g., oil and vinegar, breast milk).
• Visuals and examples (Figure descriptions):
  • Blood, gelatin, soda as suspensions or colloids depending on context.
  • Emulsion example: oil and water when agitated.

2.7 Biological Macromolecules: General Characteristics

• Macromolecules are large organic molecules synthesized by the body.
  • Always contain carbon (C), hydrogen (H), and oxygen (O); some contain nitrogen (N), phosphorus (P), or sulfur (S).
  • Carbon skeletons can take many forms; polymers are built from monomers.
  • Macromolecules are often polar and can hydrogen-bond; many act as acids (e.g., carboxyl group) or bases (e.g., amine group).
• Polymers and monomers (Page 58):
  • Monomers: repeating subunits (e.g., sugar for carbohydrates, nucleotides for nucleic acids, amino acids for proteins).
  • Dimers form when two monomers bond.
• Dehydration synthesis (condensation) vs hydrolysis (Page 59):
  • Dehydration synthesis: a subunit loses an —H and another subunit loses an —OH; new covalent bond forms; a molecule of water is produced. E.g., forming a polymer from monomers.
  • Hydrolysis: water is used to split polymers; an —H is added to one subunit and an —OH to another.

2.7b Lipids

• Lipids: diverse, nonpolar, water-insoluble molecules with various roles (Page 60):
  • Stored energy, components of cellular membranes, and hormones.
  • Four primary classes: triglycerides, phospholipids, steroids, eicosanoids.
• Triglycerides (Page 61–62):
  • Used for long-term energy storage.
  • Formed from glycerol and three fatty acids; fatty acids vary in length and number of double bonds.
  • Saturated fatty acids: no double bonds; Unsaturated fatty acids: one double bond; Polyunsaturated: two or more double bonds.
  • Adipose tissue stores triglycerides.
  • Lipogenesis: synthesis of triglycerides when nutrients are in excess.
  • Lipolysis: breakdown of triglycerides when energy is needed;
  • Visual: triglyceride structure shows glycerol backbone with three fatty acid chains.
• Phospholipids (Page 63):
  • Amphipathic molecules that form chemical barriers of cell membranes.
  • Structure: glycerol, phosphate group, and two fatty acid tails.
  • Forms a hydrophilic (polar) head and hydrophobic (nonpolar) tails, enabling bilayer membranes.
• Steroids (Page 64):
  • Four fused rings forming a multiringed structure; differ by side chains.
  • Examples: cholesterol (membrane component and precursor to other steroids), steroid hormones (testosterone, estrogen), bile salts.

2.7c Carbohydrates

• General features (Page 65):
  • Carbohydrates typically have a general formula with hydrogen and hydroxyl groups attached to carbon.
  • Monosaccharides: simple sugar monomers.
  • Disaccharides: two monosaccharides bound together.
  • Polysaccharides: three or more sugars; examples include glycogen, starch, cellulose.
• Glucose and energy (Page 66–67):
  • Glucose: a six-carbon (hexose) carbohydrate; primary energy source for cells.
  • Glycogen: stored in liver and skeletal muscle; glycogenesis forms glycogen; glycogenolysis breaks it down; gluconeogenesis forms glucose from noncarbohydrate sources in the liver.
• Isomers (Page 68–69):
  • Glucose, galactose, and fructose share the same molecular formula (C₆H₁₂O₆) but have different structures and properties.
• Other carbohydrates (Page 69–70):
  • Pentoses (e.g., ribose, deoxyribose).
  • Disaccharides: sucrose (glucose + fructose), lactose (glucose + galactose), maltose (glucose + glucose).
  • Polysaccharides: glycogen (animals), starch (plants), cellulose (plants; dietary fiber).
  • Glycosaminoglycans (GAGs) and proteoglycans in connective tissue.

2.7d Nucleic Acids

• DNA and RNA (Page 71–76):
  • Two classes: DNA (deoxyribonucleic acid) and RNA (ribonucleic acid).
  • Both are polymers of nucleotide monomers.
  • Each nucleotide consists of three components:
  • Sugar (pentose)
  • Phosphate group
  • Nitrogenous base (purines: A, G; pyrimidines: C, T, U in RNA)
  • Nitrogenous bases:
  • Purines: Adenine (A), Guanine (G).
  • Pyrimidines: Cytosine (C), Uracil (U) in RNA, Thymine (T) in DNA.
  • DNA structure:
  • Double-stranded; sugar–phosphate backbone; hydrogen bonds between complementary bases (A–T and G–C).
  • located in nucleus and mitochondria.
  • RNA structure:
  • Single-stranded; contains ribose; U replaces T.
  • ATP (Page 77):
  • Adenosine triphosphate: a nucleotide with adenine, ribose, and three phosphates; central molecule for energy transfer.
  • The last two phosphate groups form energy-rich bonds; hydrolysis releases energy.
• Note on bases and pairing:
  • A pairs with T (DNA) via two hydrogen bonds; A pairs with U (RNA) via two hydrogen bonds; G pairs with C via three hydrogen bonds.

2.7e Proteins

• Functions of proteins (Page 78):
  • Catalyze chemical reactions (enzymes).
  • Provide structural support (cytoskeleton).
  • Enable body movement (actin, myosin in muscle).
  • Transport substances in blood (e.g., hemoglobin for O₂).
  • Mediate membrane transport via carrier proteins.
  • Protect the body (antibodies).
• Protein structure (Page 79):
  • Proteins are polymers of 20 amino acids.
  • Each amino acid has:
  • An amine group (–NH₂)
  • A carboxyl group (–COOH)
  • A hydrogen atom
  • A variable side chain (R group) that distinguishes amino acids.
  • Amino acids link via peptide bonds to form polymers.
• Visuals (Figure references):
  • Dipeptide and peptide bond formation show condensation reaction releasing water (H₂O).
  • Primary structure is the linear sequence of amino acids in a protein.

2.8 Amino Acid Sequence and Protein Conformation (2.8b)

• Primary structure (Page 81):

  • The linear sequence of amino acids in a polypeptide chain.

• Prosthetic groups (Page 82):

  • Nonprotein molecules covalently bonded to a protein required for normal function (e.g., the heme group in hemoglobin).

• Denaturation (Page 82–83):

  • Conformational change that disrupts protein activity; often irreversible.

• Denaturation triggers (Page 83):

  • Changes in temperature or pH can disrupt electrostatic interactions and intramolecular bonds, potentially lethal in physiological contexts (e.g., blood pH changes).

• Overall, these notes summarize the essential concepts of matter, atomic structure, chemical bonding, water chemistry, mixtures, and the major biological macromolecules (carbohydrates, lipids, proteins, nucleic acids) along with their properties, examples, and relevance to physiology. Be sure you understand how ionic and covalent bonds differ, how water’s properties support life, and how macromolecules’ structure informs their function.