Chapter Two Part One Vocabulary

Matter and Elements

• Matter is anything that takes up space and has mass (e.g., car, desk, textbook, computer).
• States of matter: solid, liquid, or gas.
• All matter (living or non-living) is made up of elements.
• Elements are substances that cannot be broken down into a simpler substance; examples include carbon, oxygen.
• Common elements making up about 95% of living organisms:
  • Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N), Phosphorus (P), and Sulfur (S) (the list can vary slightly, calcium (Ca) is sometimes listed as a newer addition for bones).
• One-letter abbreviations are used for element symbols (e.g., O, C, H, N, P, S).

Atomic Structure and the Periodic Table

• Elements are organized in the periodic table with:
  • Periods (horizontal rows) labeled 1 through 7.
  • Groups (vertical columns) that reflect similar chemical properties; groups are often labeled using Roman numerals in some contexts.
• Each tile in the periodic table displays:
  • Atomic number (Z): number of protons in the nucleus.
  • Atomic symbol (e.g., H, C, O).
  • Atomic mass (approximate; often the average mass of naturally occurring isotopes).
• Atomic number concept:
  • Determines the identity of the element (e.g., 1 proton = hydrogen, 6 protons = carbon, 8 protons = oxygen).
• The nucleus contains protons (positive charge) and neutrons (neutral charge); electrons orbit in electron shells surrounding the nucleus (negative charge).

Atomic Numbers, Mass, and Neutrons

• Atomic number Z = number of protons.
• In a neutral atom, number of electrons ≈ number of protons (to balance charge).
• Mass number A (also called atomic mass in some contexts) = number of protons + number of neutrons: A = Z + N.
• Neutrons N can be found from N = A − Z.
• For example:
  • Hydrogen: Z = 1; in a common isotope, A ≈ 1, so N = A − Z = 0; electrons ≈ 1.
  • Carbon: Z = 6; if A = 12, then N = 12 − 6 = 6 neutrons; electrons = 6.
  • Lithium: Z = 3; if A = 7, then N = 7 − 3 = 4 neutrons; electrons = 3.
• Isotopes are atoms of the same element (same Z) with different N (and thus different A).
• Example: Carbon-12 vs Carbon-14 differ in neutrons; both have 6 protons, 6 electrons, but different neutron counts and masses.
• Notation: atomic structure can be represented as if writing the isotope as $^{A}_{Z} ext{X}$, where X is the element symbol.
• Important concept: proton number (Z) is invariant for a given element; neutron number (N) can vary, leading to isotopes; electrons can vary in ions but protons do not change.

Electron Arrangements and Valence

• Electron shells (orbitals) hold electrons with capacity limits:
  • First shell (innermost) holds up to 2 electrons.
  • Second shell can hold up to 8 electrons.
  • Third shell can also hold up to 8 (often more in heavier elements, but commonly explained with the 2-8-8 rule for light biologically relevant elements).
• The outer shell (valence shell) determines bonding behavior: the number of electrons in the outer shell predicts how many other atoms an atom tends to bond with.
• Example: Carbon (Z = 6) has 2 electrons in the first shell and 4 in the second shell; outer shell holds 4 electrons, so carbon tends to bond with four other atoms.
• Nitrogen (Z = 7) has 2 in the first shell and 5 in the second; outer shell has 5 electrons, leaving 3 spaces for bonding.
• Oxygen (Z = 8) has 2 in the first shell and 6 in the second; outer shell has 6 electrons, leaving 2 spaces for bonding.
• These valence patterns explain why carbon forms up to four bonds, nitrogen up to three, and oxygen up to two (in many common molecules).
• Comparison with silicon (Z = 14): outer shell configurations show similar bonding patterns in a column of the periodic table (group trends).
• The number of outer-shell electrons helps predict bonding type and molecular structure in living matter (C, N, O, P, S, H are especially important here).

Isotopes and Radiocarbon Dating

• Isotopes: atoms of the same element with different numbers of neutrons (different mass numbers A).
• Carbon isotopes example: $^{12} ext{C}$, $^{13} ext{C}$, $^{14} ext{C}$.
• Carbon-14 is produced in the upper atmosphere via cosmic radiation and decays over time to carbon-12, emitting energy in the process.
• Radiocarbon dating uses the decay of $^{14} ext{C}$ to estimate the age of organic material by comparing remaining $^{14} ext{C}$ to its initial amount.
• Key takeaway: The proton number Z stays constant for a given element; neutron number N can change, leading to different isotopes with different masses.

Ionic and Covalent Bonding

• Ionic bonding (transfer of electrons): one atom donates one or more electrons to another atom, creating ions with opposite charges that attract each other.
  • Example: Sodium (Na, Z = 11) donates one electron to chlorine (Cl, Z = 17).
  • Results: Na becomes Na⁺; Cl becomes Cl⁻ (chloride).
  • In water, ionic bonds can be broken as water wedges between ions, forming hydrated ions and driving dissolution (electrolytes).
• Covalent bonding (sharing of electrons): atoms share electrons so that each atom attains a stable outer shell.
  • Non-polar covalent bonds: equal sharing of electrons; very stable (example: H−H in H₂ or C−H bonds in methane are largely non-polar in certain contexts).
  • Polar covalent bonds: unequal sharing of electrons; one end of the bond becomes partially positive while the other becomes partially negative.
  • Water (H₂O) is a classic example of polar covalent bonding because the shared electrons are drawn more toward the oxygen atom, creating partial charges.
• Hydrogen bonds: a special interaction between a slightly positive hydrogen atom (attached to an electronegative atom like O) and a slightly negative atom on another molecule or part of the same molecule.
  • They are weaker than covalent bonds but collectively can be strong, especially in large networks (e.g., water molecules bonding to each other).
  • In water, hydrogen bonds help explain water's high cohesion and high boiling point relative to other small molecules.

Water: Properties, Bonding, and Roles in Biology

• Water as a solvent: water’s polar covalent bonds allow it to dissolve many ionic and other polar substances; water can break ionic bonds and form ions in solution.
• Ions and electrolytes: dissociated species in solution (e.g., Na⁺ and Cl⁻) that conduct electricity in solution; many bodily processes rely on ions.
• Water’s cohesion and adhesion:
  • Cohesion: attraction between water molecules; helps water molecules stick together.
  • Adhesion: attraction between water molecules and other surfaces (e.g., tubes, plant xylem).
  • Together they enable capillary action, allowing water movement through narrow spaces (e.g., from soil to plant leaves).
• Surface tension: arises from cohesive forces at the surface; enables phenomena like water striders and the formation of a barrier between air and water.
• Temperature effects on surface tension: higher temperatures can increase surface tension in some contexts; temperature affects cohesion and adhesion dynamics.
• Density and phase change of water:
  • Liquid water is denser than ice; ice is less dense due to crystalline structure forming when water freezes, causing ice to float on liquid water.
  • This density difference is crucial for life: ice floats, insulating liquid water below and allowing life to persist under the ice.
• High heat capacity and vaporization:
  • Water heats up and cools down slowly (high heat capacity), buffering environmental temperatures and stabilizing ecosystems.
  • Vaporization (liquid to gas) requires substantial energy; sweating cools the body as water on the skin absorbs heat and vaporizes.
• Importance for life: the combination of solvent properties, cohesion/adhesion, surface tension, density behavior, and heat capacity make water essential for biochemical processes, climate stability, and ecological systems.

Water Dissolution, Ions, and pH

• When salts like NaCl dissolve in water, water molecules intervene between ions, enabling dissociation into Na⁺ and Cl⁻ ions; the process is influenced by hydrating (solvation).
• Water can also promote the breakdown of other ionic compounds, generating various ions (e.g., OH⁻, H⁺) depending on the solute.
• pH concept: a measure of how acidic or basic (alkaline) a solution is, determined by the concentration of hydrogen ions (H⁺) or hydroxide ions (OH⁻) in solution.
• Acids and bases (Arrhenius viewpoint):
  • Acids increase H⁺ (or H₃O⁺) in solution when dissolved (e.g., lemon juice contains acids that release hydrogen ions).
  • Bases increase OH⁻ in solution (e.g., oven cleaners and many cleaners are basic).
  • Extreme acidity or basicity can be dangerous; the pH scale ranges roughly from 0 to 14 in aqueous solutions.
• Neutral pH: pure water is around pH 7 (neutral);
  • Milk is slightly acidic (around pH ~6.5).
  • pH 6 is acidic (slightly); pH 13 is basic/alkaline.
  • A very acidic example: pH 1.2 (highly acidic).
• It is important to note that acids and bases are not inherently “bad” or “good”; extreme values are harmful in biological contexts.
• pH and ions in biological systems influence many processes (enzyme activity, physiology, etc.).

Practice Implications and Connections

• Protons (Z) determine the identity of an element; changing Z would change the element (e.g., 7 protons = nitrogen, not carbon or oxygen).
• Neutron number can vary (isotopes) without changing the element’s identity, affecting atomic mass and nuclear properties.
• The outer electron configuration (valence) drives chemical bonding type and molecular structure; this explains why carbon forms diverse and complex biomolecules, why water is polar, and why many biological macromolecules rely on hydrogen bonding networks.
• Water’s properties underpin essential life processes: solvent capabilities enable biochemical reactions; cohesive/adhesive properties support transport in organisms; surface tension and capillary action support circulation and nutrient transport in plants; high heat capacity stabilizes climate and bodily temperatures.
• Understanding acids, bases, salts, and pH is fundamental for predicting reactions, buffer systems, and the behavior of biological fluids in health and disease.

Quick recap of key formulas and concepts

• Atomic notation: $^{A}_{Z} ext{X}$ where A = mass number, Z = atomic number, X = element symbol.
• Protons = Electrons (in neutral atoms) = Z.
• Neutrons: N = A - Z
• Mass number relation: A = Z + N
• Shell capacities (approximate): 1st shell max 2 electrons; 2nd shell max 8 electrons; 3rd shell often 8 in simple bio contexts.
• Isotopes differ in neutron number but share Z and hence the same element.
• Ionic bond example: Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻; ionic bond via electrostatic attraction.
• Covalent bond example: H–H sharing electrons; polar covalent bond in H₂O due to unequal sharing.
• Hydrogen bond: interaction between partially positive H and a neighboring electronegative atom (e.g., O in water).
• pH scale reference points: Neutral ~7; acidic
• Water’s essential properties: solvent, cohesion, adhesion, surface tension, capillary action, density differences between ice and liquid, high heat capacity, and high heat of vaporization.