Chapter 2: Matter, Atoms, Elements, and the Periodic Table

2.1 Matter, Atoms, Elements, and the Periodic Table

  • Matter has mass and occupies space. Forms of matter include:
    • Solid (e.g., bone)
    • Liquid (e.g., blood)
    • Gas (e.g., oxygen)
  • Atom: the smallest particle exhibiting chemical properties of an element.
  • There are 92 naturally occurring elements that make up matter, organized in the periodic table.
  • Periodic table organizes elements by atomic number and other properties; symbols are unique to each element (e.g., C = carbon).

2.1a Components of an atom

  • Atoms composed of three subatomic particles:
    • Neutrons: mass of one atomic mass unit (amu), no charge.
    • Protons: mass of one amu, positive charge of +1.
    • Electrons: mass ~1/1800 of a proton/neutron, negative charge of −1.
  • Electrons are located in regions around the nucleus called orbitals/shells.

2.1a Elements and the periodic table

  • Chemical symbol: unique to each element; usually the first letter or first letter plus another letter (e.g., C = carbon).
  • Atomic number (Z): number of protons in an atom; located above the symbol on the periodic table.
  • Average atomic mass: mass of protons and neutrons; shown below the element symbol.
  • Elements arranged by atomic number within rows; periodic trends reflect electron configuration.

2.1a Determining subatomic particles

  • Proton number = atomic number = Z
  • Neutron number = atomic mass − atomic number
    • Example: for Na, neutron number = 23 − 11 = 12
  • Electron number = proton number (in neutral atoms) → e⁻ = Z

2.1a Diagramming atomic structures

  • Atoms have shells of electrons surrounding the nucleus.
  • Each shell has a given energy level and can hold a limited number of electrons.
  • Innermost shell holds up to 2 electrons; second shell up to 8; shells closest to the nucleus must be filled first.

2.1b Isotopes

  • Isotopes are different atoms of the same element:
    • Same number of protons and electrons; different number of neutrons.
    • Identical chemical characteristics; different atomic masses.
  • Example: Carbon isotopes – Carbon-12 (6 neutrons, majority), Carbon-13 (7 neutrons), Carbon-14 (8 neutrons).
  • The weighted average atomic mass takes all isotopes into account.
  • Radioisotopes: contain excess neutrons, unstable, emit high-energy radiation (α, β, γ).
  • Physical half-life: time for 50% of a radioisotope to decay.
  • Biological half-life: time required for half of the radioactive material to be eliminated from the body.

2.1b Isotopes (continued)

  • Medical imaging example: use of radioisotopes to trace metabolic reactions; iodine uptake in thyroid imaging shows nodule activity.

2.1c Chemical Stability and the Octet Rule

  • Periodic table organized into columns by number of electrons in the outer (valence) shell.

  • Elements tend to lose, gain, or share electrons to obtain a complete outer shell of eight electrons (the octet).

  • Example trends:

    • Column IA (group 1): elements with one electron in their valence shell (e.g., H, Li, Na, K).
    • Column VIIA (group 17): elements with a full valence shell, tending toward stability.

  • Noble gases (e.g., He, Ne) are chemically inert due to complete valence shells.

  • Octet Rule: Elements tend to achieve a full outer shell of eight electrons to reach chemical stability.

2.2 Ions and Ionic Compounds

  • Chemical compounds: stable associations between two or more elements in fixed ratios; can be ionic or molecular.
  • Ionic compounds are lattices held together by ionic bonds.

2.2a Ions

  • Ions: atoms with positive charge (cations) or negative charge (anions) produced by loss or gain of one or more electrons.
  • Ions have significant physiological functions (e.g., K⁺ in sweat and in physiological processes).
  • Example: Sodium (Na) loses an electron to form Na⁺ (+1).

2.2a Ions (continued)

  • Formation details:
    • Losing electrons: Na → Na⁺ + e⁻ (cation formation; +1 charge)
    • Gaining electrons: Cl + e⁻ → Cl⁻ (anion formation; −1 charge)
  • Polyatomic ions: ions composed of more than one atom (e.g., bicarbonate HCO₃⁻, phosphate PO₄³⁻).

2.2b Ionic Bonds

  • Ionic bonds occur between cations and anions and form salts.
  • Examples:
    • Sodium chloride: NaCl (Na⁺ and Cl⁻ in a lattice).
    • Magnesium chloride: MgCl₂ (Mg²⁺ with two Cl⁻).

2.3 Covalent Bonding, Molecules, and Molecular Compounds

  • Covalent bond: atoms share electrons; typically occurs when both atoms require electrons to achieve stability.
  • Common in the human body: H, O, N, C form covalent bonds.
  • Carbon has a valence of four and can form up to four covalent bonds, creating carbon skeletons (chains or rings).
  • Electronegativity determines how electrons are shared; highly electronegative atoms attract electrons more strongly.
  • Polar covalent bonds: electrons shared unequally, leading to partial charges (δ⁺ and δ⁻).
  • Nonpolar covalent bonds: electrons shared equally (e.g., O=O, C–H with similar electronegativities).
  • Exception to the polar-covalent rule: C–H bonds are often treated as nonpolar because electronegativity difference is small.

2.3c Nonpolar, Polar, and Amphipathic Molecules

  • Amphipathic molecules contain both polar (hydrophilic) and nonpolar (hydrophobic) regions, e.g., phospholipids.
  • Nonpolar molecules: e.g., O–O, C–H bonds are nonpolar.
  • Polar molecules: e.g., O–H bonds in water.

2.3d Intermolecular Attractions

  • Intermolecular attractions are weak bonds between molecules but collectively significant for structure and function.
  • Hydrogen bonds: between a polar molecule's partially positive hydrogen and a partially negative atom on another molecule; water’s properties rely on this.
  • Other attractions: dipole interactions; hydrophobic interactions (nonpolar molecules in polar environments); intramolecular attractions when large molecules fold in space.

2.4 Molecular Structure and Properties of Water

  • Water is either organic or inorganic; here focus on inorganic-water relationships.
  • Water is a polar molecule: one O atom bonded to two H atoms; O has partial negative charge; H atoms have partial positive charge.
  • Each water molecule can form up to four hydrogen bonds with neighbors.
  • Water’s properties underpin body temperature regulation, transport, lubrication, cushioning, and solvation.

2.4a Molecular Structure of Water

  • Water: H₂O; O is partially negative, H is partially positive; maximum hydrogen bonding gives water unique properties.

2.4b Properties of Water

  • Phases: gas (vapor), liquid (water), solid (ice) depending on temperature.

  • Functions of liquid water:

    • Transports dissolved substances throughout the body
    • Lubricates and cushions joints and tissues
    • Excretes wastes; dissolves substances for elimination

  • Cohesion and adhesion; surface tension; surfactants prevent alveolar collapse in lungs.

  • High specific heat and high heat of vaporization due to hydrogen bonding; important for temperature regulation (sweating).

  • Heat-related terms:

    • Specific heat: energy required to raise 1 g of substance by 1°C; for water, very high due to hydrogen bonding.
    • Heat of vaporization: energy required to convert 1 g of liquid to gas; high for water.

2.4c Water as the Universal Solvent

  • Water dissolves many substances; solvents are substances that dissolve in water; water is called the universal solvent because most substances dissolve in it.
  • Hydrophilic substances dissolve or form hydration shells; electrolytes dissociate and conduct current; nonelectrolytes dissolve but do not conduct.
  • Substances that do not dissolve in water are hydrophobic; lipids are examples.
  • Amphipathic molecules (e.g., phospholipids) have polar heads that interact with water and nonpolar tails that avoid water, leading to membrane assembly (bilayers) or micelles.

2.5 pH, Neutralization, and the Action of Buffers

  • pH: a measure of H⁺ concentration in solution; ranges from 0 to 14.
  • Pure water at 25°C has pH ~ 7 and [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M.
  • pH and [H⁺] are inversely related: pH = −log₁₀[H⁺].
  • Moving one pH unit represents a 10-fold change in H⁺ concentration.
  • Neutralization: acids neutralized by adding base; bases neutralized by adding acid.
  • Buffers: resist pH changes by accepting H⁺ from excess acid or donating H⁺ to neutralize base; e.g., carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻) buffer blood pH.
  • Blood pH target range: ~7.35 to 7.45.

Section 2.5 What did you learn?

  • Possible exam-style prompts include: (1) how water contributes to pH balance, (2) buffer function, (3) neutralization, (4) examples of buffers in blood, etc.

2.6 Water Mixtures

  • Mixtures: two or more substances combined without chemical change; can be separated by physical means (evaporation, filtering).
  • Categories of water mixtures:
    • Suspension: larger particles (>1 mm); appears cloudy; particles settle unless in motion (e.g., blood cells in plasma).
    • Colloid: smaller particles than a suspension but larger than in a solution; remains mixed without motion; scatters light.
    • Solution: homogeneous, particles <1 nm; dissolves in water and does not scatter light; examples include sugar water, saline.
  • Emulsion: special colloid of water and a nonpolar liquid (e.g., oil and vinegar; breast milk).

2.6b Expressions of Solution Concentration

  • Concentration is expressed as:
    • Mass/volume: mass of solute per volume of solution.
    • Mass/volume percent: grams of solute per 100 mL of solution.
    • Molarity: M = rac{n}{V} where n is moles of solute and V is volume of solution.
    • Molality: m = rac{n}{kg ext{ solvent}} (moles solute per kilogram solvent; temperature independent).
    • Osmolarity: number of particles per liter of solution.
    • Osmolality: number of particles per kilogram of water.
  • Avogadro’s number: N_A = 6.022 imes 10^{23} particles per mole.
  • One mole concept: mass in grams equals the molecular mass; for carbon, 1 mole of C = 12.01 ext{ g}; molecular mass is the sum of atomic masses of constituent atoms.

2.7 Biological Macromolecules: General Characteristics

  • Biological macromolecules are large organic molecules synthesized by the body; contain C, H, O; often N, P, S.
  • Carbon skeletons can take various forms; functional groups confer properties; many are polar and form hydrogen bonds; some act as acids (e.g., carboxyl) or bases (e.g., amine).
  • Polymers: molecules made of monomers; carbohydrate monomers (sugars), nucleic acid monomers (nucleotides), protein monomers (amino acids).
  • Dehydration synthesis (condensation): subunits join by removing —H and —OH to form a covalent bond and release water.
  • Hydrolysis: water is used to cleave covalent bonds, yielding subunits.

2.7a Lipids

  • Lipids: diverse, fatty, water-insoluble molecules; functions include energy storage, cellular membranes, hormones.
  • Four primary classes:
    1) Triglycerides
    2) Phospholipids
    3) Steroids
    4) Eicosanoids

2.7b Lipids: Details

  • Triglycerides: long-term energy storage; formed from glycerol and three fatty acids; fatty acids vary in length and degree of unsaturation.
    • Saturated: no double bonds.
    • Unsaturated: one double bond.
    • Polyunsaturated: two or more double bonds.
    • Stored in adipose tissue; lipogenesis (formation) and lipolysis (breakdown).
  • Phospholipids: amphipathic; form cellular membranes; glycerol backbone with phosphate-containing polar head and nonpolar fatty acid tails; form bilayers and micelles.
  • Steroids: multi-ring hydrocarbon structures; include cholesterol (membrane component and precursor to other steroids) and steroid hormones (e.g., testosterone, estrogen) and bile salts.
  • Eicosanoids: 20-carbon fatty acids derived from arachidonic acid; local signaling molecules involved in inflammation and nervous system signaling; classes include prostaglandins, prostacyclins, thromboxanes, leukotrienes.

2.7c Carbohydrates

  • General formula: a CH₂O unit for each carbon; formula often written as (CH₂O)ₙ.
  • Monosaccharides: simple sugar monomers (e.g., glucose).
  • Disaccharides: two monosaccharides linked (e.g., sucrose, lactose, maltose).
  • Polysaccharides: many monosaccharides (e.g., glycogen, starch, cellulose).
  • Glycogen: storage form of glucose in liver and skeletal muscle; glycogenesis (formation) and glycogenolysis (breakdown).
  • Glucose: primary energy source; concentration must be tightly maintained.
  • Glucose isomers include galactose and fructose; hexose monosaccharides.
  • Pentose sugars include ribose and deoxyribose; components of nucleic acids.
  • Other features: glycosaminoglycans (GAGs) and proteoglycans in connective tissue.

Section 2.7c What did you learn?

  • Possible prompts: identify monosaccharide vs disaccharide vs polysaccharide; glucose isomer differences; glycogen structure and storage.

2.7d Nucleic Acids

  • Nucleic acids store and transfer genetic information; two classes:
    • DNA (deoxyribonucleic acid)
    • RNA (ribonucleic acid)
  • Both are polymers of nucleotide monomers linked by phosphodiester bonds.

Nucleotide Monomer

  • Components of a nucleotide:
    • Sugar: a five-carbon pentose (deoxyribose in DNA; ribose in RNA)
    • Phosphate group
    • Nitrogenous base: purines (A, G) or pyrimidines (C, U in RNA; C, T in DNA)
  • Bases:
    • Pyrimidines: cytosine (C), uracil (U), thymine (T)
    • Purines: adenine (A), guanine (G)

DNA vs RNA

  • DNA: double-stranded, located in nucleus and mitochondria; uses deoxyribose sugar and bases A, C, G, T; thymine; strands held together by hydrogen bonds (A–T, G–C).
  • RNA: single-stranded; uses ribose sugar and bases A, C, G, U; thymine is replaced by uracil.
  • ATP (adenosine triphosphate): a nucleotide-based energy currency with three phosphate groups; high-energy bonds between the last two phosphates release energy when broken.
  • Other important nucleotide-containing molecules: NAD⁺ and FAD; coenzymes involved in ATP production.

2.7d Nucleic Acids (continued)

  • RNA and DNA structures are illustrated to show how nucleotides polymerize and how base-pairing stabilizes the DNA double helix.

2.7e Proteins

  • Protein functions include:
    • Enzymatic catalysis (synthesis, digestion)
    • Structural support (cytoskeleton)
    • Movement (actin, myosin)
    • Transport (hemoglobin carries O₂)
    • Membrane transport (carrier proteins)
    • Protection (antibodies)

2.7e General protein structure

  • Proteins are polymers of amino acids linked by peptide bonds.
  • 20 amino acids with common backbone: amine group, carboxyl group, hydrogen, and distinctive R group.
  • N-terminal end has a free amine group; C-terminal end has a free carboxyl group.
  • Oligopeptides: 3–20 amino acids; polypeptides: longer chains; proteins are one or more polypeptide chains.
  • Glycoproteins: proteins with carbohydrate attached; e.g., ABO blood group determinants on erythrocytes.

Section 2.7e What did you learn?

  • Possible prompts: identify monomers and bonds in proteins; define oligomeric vs polymeric forms; explain glycoproteins.

2.8a Categories of Amino Acids

  • Grouped by their R group:
    • Nonpolar amino acids: R groups with hydrophobic characteristics (interaction by hydrophobic forces).
    • Polar amino acids: R groups that are polar and interact with water.
    • Charged amino acids: R groups carry negative or positive charges; form ionic bonds.
  • Some amino acids have special functions:
    • Proline can introduce bends in the protein chain.
    • Cysteine can form disulfide bonds.
    • Methionine is one of the first amino acids in protein synthesis.

2.8a Amino Acids (Figure 2.25)

  • Visual representations show the diversity of R groups and how they influence protein folding and function.

2.8b Amino Acid Sequence and Protein Conformation

  • Primary structure: linear sequence of amino acids.
  • Conformation (three-dimensional shape): essential for function; achieved through folding aided by chaperone proteins.
  • Intramolecular interactions contributing to conformation:
    • Hydrophobic exclusion: nonpolar residues fold away from water.
    • Hydrogen bonds: between polar R groups and between amine and carboxyl groups.
    • Ionic bonds: between negatively and positively charged R groups.
    • Disulfide bonds: covalent links between cysteine residues.

Secondary structures

  • Alpha helix: spiral coil, provides elasticity in fibrous proteins (e.g., skin, hair).
  • Beta sheet: planar pleated arrangement, provides flexibility in globular proteins (e.g., enzymes).

Tertiary structure

  • Three-dimensional folding of a single polypeptide chain; two categories:
    • Globular proteins: compact shapes.
    • Fibrous proteins: elongated, extended chains.

Quaternary structure

  • Present in proteins with two or more polypeptide chains (e.g., hemoglobin has four polypeptide chains).

Prosthetic groups and denaturation

  • Prosthetic groups: nonprotein groups covalently bonded to proteins (e.g., heme in hemoglobin).
  • Denaturation: conformational change that disturbs protein activity; usually irreversible and can be caused by high temperature or pH changes.

Section 2.8 What did you learn?

  • Possible prompts include differences among primary, secondary, tertiary, and quaternary structures; what happens during denaturation; how pH affects protein structure.

Quick reference: Key equations and concepts (LaTeX)

  • Electron shell capacity (first two shells): 2, \, 8 electrons respectively.
  • Neutron number: N = A - Z where A is atomic mass number and Z is atomic number.
  • Ionic bond example: Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻.
  • Covalent bonds: electrons shared; single, double, triple bonds (e.g., H–H, O=O, N≡N).
  • Octet rule: aim for outer shell to have 8 electrons.
  • Water structure: ext{H}_2 ext{O} with hydrogen bonding; water can form up to four hydrogen bonds per molecule.
  • pH and hydrogen ion concentration: ext{pH} = -\, ext{log}_{10} [ ext{H}^+] ; neutral pH ~7; blood pH ~7.35–7.45.
  • Molarity: M = rac{n}{V}; molality: m = rac{n}{kg ext{ solvent}}.
  • Avogadro's number: N_A = 6.022 imes 10^{23} particles per mole.
  • Carbohydrate general formula:
      • -
    • Monosaccharides, disaccharides, polysaccharides follow the pattern: $(CH2O)n$.
  • Nucleotides consist of: sugar (pentose), phosphate, and a nitrogenous base; nucleic acids use phosphodiester linkages.
  • Proteins: amino acids connected by peptide bonds; secondary structures include the α-helix and β-pleated sheet; tertiary and quaternary structures involve hydrophobic interactions, hydrogen bonds, ionic bonds, and disulfide bridges.