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Chapter 3: States of Matter and Atomic Structure (Review)

Exam logistics and study approach

  • There are three questions you can answer; you should answer one. If you answer more than one, the instructor will take your lowest score among those questions.
  • Do not hedge your bets by answering all of them; follow the steps and focus on quality for one question.
  • Calculator use: one simple dosage question will appear; not hard math (example: 4 ÷ 2). A calculator may be allowed; be prepared to use it.
  • Tools: calculators and metric conversions may be limited; expect only straightforward metric-to-metric conversions (e.g., cm to mm).
  • Exam strategy: write as much as possible on the exam, cross things out, circle, underline, draw pictures, use notes on the exam. This is allowed; do not feel you’re restricted from annotating.
  • There will be a shift to Chapter 3: States of Matter.
  • A quick break note: breaks or time checks may be included; stay focused on the material during lecture time.

States of matter: overview and key ideas

  • Matter is anything that takes up space; at a microscopic level, it is made up largely of atoms.
  • When molecules are cold, movement is reduced, allowing closer packing into a solid.
  • When heat energy is added, molecules gain energy and move more; this can lead to a liquid as molecules slide past each other.
  • With more heat, molecules move farther apart, density decreases, and the substance becomes a gas.
  • Heat energy acts on the bonds holding matter together; heating can break bonds, changing the structure (e.g., cooking an egg).
  • Denaturation: breaking of bonds in a material (e.g., proteins) due to heat, altering the material’s structure.
  • Phase changes: solid → liquid (melting), liquid → gas (vaporization/boiling), gas → liquid (condensation), solid → gas (sublimation). The reverse is sublimation for gas → solid.
  • Humidity vs. condensation: humidity is water vapor in air; condensation is water vapor turning into liquid on surfaces.
  • Dry ice (CO₂) sublimates at room temperature (solid directly to gas); this is a classic example of sublimation.
  • Practical relevance: understanding states of matter helps explain bodily regulation and how temperature affects biological processes and materials.

Dalton’s atomic theory and basic chemistry concepts

  • Four main ideas (Dalton):
    • All matter is made of atoms.
    • Atoms are indivisible and indestructible (in Dalton’s view) — in modern chemistry, atoms can be split in nuclear reactions, but for chemical bonding this idea is a good starting point.
    • Atoms of a given element have the same mass and properties.
    • Compounds are composed of two or more different atoms (i.e., different elements).
  • If a chemical reaction occurs, bonds are formed, broken, or rearranged; RXN is shorthand for a chemical reaction.
  • Atoms are the smallest units of matter; molecules are bonded atoms; compounds are molecules containing at least two different elements.
  • Distinguishing terms:
    • Molecule: one or more bonded atoms.
    • Compound: a molecule with at least two different elements.
  • Examples:
    • H₂O is a compound (two different elements: H and O).
    • C₆H₁₂O₆ (glucose) is a compound with three elements (C, H, O).
    • O₂ is a molecule but not a compound (same element, O).
  • Atom basics connect to molecular structure and bonding, which drive chemical behavior in biology.

Biologically important elements and bonding concepts

  • Key elements to know (mnemonics and significance):
    • Major elements: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N) — these make up most organic molecules and water.
    • Other commonly referenced elements include Phosphorus (P), Potassium (K), Iodine (I), Sulfur (S), Calcium (Ca), Iron (Fe), Chlorine (Cl), Magnesium (Mg).
    • A common mnemonic the instructor used is C H O P K I N S Ca Fe Cl Mg (with occasional shorthand variants).
  • Potassium note: symbol K comes from the Latin Kalium; confusion between P (phosphorus) and K (potassium) is common; keep their symbols straight.
  • The major group CHON (and often P, S) form the backbone of organic and biological molecules; C, H, O, N are the bulk of life’s mass (especially water: H and O).
  • Goiter and iodine: iodine is essential for thyroid function; iodine deficiency can cause goiter; iodine is concentrated by the thyroid; iodized salt introduction was to prevent deficiency.
  • Calcium, iron, chlorine, magnesium, and other elements have specific biological roles (bone structure, oxygen transport, electrolyte balance, etc.).
  • The emphasized takeaway: CHON are the core building blocks of most organic life; many biological molecules are composed primarily of these elements.
  • Reading the periodic table:
    • The interactive resource ptable.com is mentioned as a helpful tool for studying element properties.
  • A note on mass and isotopes (lead into the next topic): atoms of the same element can have different numbers of neutrons (isotopes), which affects atomic mass and sometimes stability.

Reading formulas, atomic structure, and isotopes

  • Two major kinds of formulas:
    • Molecular formula: tells you how many atoms of each element are in a molecule (the recipe).
    • Structural formula: shows how atoms are connected and how bonds are arranged (the actual structure).
    • Example discussion: acetic acid (vinegar) and ethanol (drinking alcohol).
    • Acetic acid (molecular formula): typically written as $C2H4O2$; structural formula shows the actual arrangement CH₃-COOH. Note: In the transcript, a specific count given for acetic acid was inconsistent; the correct molecular formula is $C2H4O2$.
    • Ethanol (EtOH): molecular formula $C2H6O$; common structural representation is CH₃-CH₂-OH.
    • Important conceptual point: molecular formulas do not convey the connectivity or arrangement of atoms; structural formulas do.
  • Atomic numbers and masses:
    • Atomic number Z = number of protons (and, for neutral atoms, equals the number of electrons).
    • Mass number A = Z + N, where N is the number of neutrons.
    • Electrons weigh negligibly; protons and neutrons each weigh ~1 Dalton (amu).
    • Example: carbon has Z = 6; a neutral carbon atom has 6 protons and 6 electrons.
  • Isotopes and isotopic mass:
    • Isotope: same element (same Z) with a different number of neutrons (N). Example: carbon-12 (Z=6, N=6, A=12) and carbon-14 (Z=6, N=8, A=14).
    • The atomic mass listed for elements on the periodic table is a weighted average of isotope masses based on natural abundance:
    • Weighted average mass formula: $$ar{A} = ext{(fractional abundance of isotope 1)} imes A1 + ext{(fractional abundance of isotope 2)} imes A2 + \