Intro to Chemistry for Biology: Atoms, Bonds, Isotopes, and Basic Biochemistry

Overview: Why chemistry matters in biology

  • Blood pH and oxygen measurements relate to chemistry: a finger clip (oximeter) measures oxygen levels; oxygen is a chemical; most medical treatments use chemicals. A basic understanding of chemistry helps explain how the body works and how different treatments interact with it.
  • Matter and essential elements in physiology:
    • All matter is composed of elements. In this lecture, key body-related elements discussed include: Carbon (C), Hydrogen (H), Oxygen (O), Phosphorus (P), Potassium (K), Iodine (I), Sulfur (S), Calcium (Ca), Iron (Fe), Magnesium (Mg).
    • Note on symbols and origins: Potassium is symbolized by K (from kalium); Iodine is I; Sulfur is S; Calcium is Ca; Iron is Fe (from ferrous); Magnesium is Mg. A few remarks about spelling/origins were given in the lecture, including a claim that calcium starts with a K due to Greek etymology, which is a point discussed but not essential for core understanding.
  • Subatomic particles and atomic structure (very basics):
    • Three subatomic particles: Protons, Neutrons, Electrons.
    • Proton = positive charge; Electron = negative charge; Neutron = neutral.
    • The nucleus contains protons and neutrons; electrons exist in shells surrounding the nucleus.
  • Atomic number, mass, and isotopes:
    • Atomic number (Z) equals the number of protons in an atom.
    • In a neutral atom, number of electrons equals the number of protons.
    • Isotopes: same element (same number of protons) but different numbers of neutrons. This changes mass but not the element identity.
    • Mass number (A) = protons + neutrons. Common carbon isotopes:
    • Carbon-12: Z=6,\, N=6,\, A=12
      ight.
    • Carbon-13: Z=6,\, N=7,\, A=13
      ight.
    • Carbon-14: Z=6,\, N=8,\, A=14
      ight.
    • The mass of most of the atom comes from protons and neutrons; electrons contribute negligibly to atomic mass.
  • Electron arrangement and shells (simplified):
    • Electrons reside in electron shells with capacity limits (in this lecture’s model):
    • First shell holds 22 electrons.
    • Second shell holds 88 electrons.
    • Third shell holds 88 electrons.
    • For carbon (Z = 6), electrons fill as 2 in the first shell and 4 in the second shell (2, 4).
  • Neutral atoms and bonding basics:
    • Atoms in nature are typically electrically neutral; number of protons equals number of electrons.
    • When atoms gain or lose electrons, they become ions (charged atoms/atoms with unequal numbers of protons and electrons).
    • Ionic bonds result from transfer of electrons between atoms, creating oppositely charged ions that attract.
  • Noble gases and valence shells:
    • The noble gases (group 18: He, Ne, Ar, Kr, Xe, Rn) are inert/non-reactive because their valence shell is full.
    • Example: Helium has a full first shell (2 electrons). Neon has a full second shell (2 in the first, 8 in the second = 10 electrons total), making it stable.
  • Covalent bonding: sharing electrons
    • Covalent bonds form when two atoms share electrons. A single covalent bond is represented by one shared pair of electrons (one line in diagrams).
    • Carbon’s typical bonding capacity: a carbon atom seeks access to 4 electrons to complete its outer shell, so it forms up to 4 covalent bonds.
    • Hydrogen and the octet rule: hydrogen needs 1 bond (to have access to 2 electrons in its first shell), and then it’s “done.” Hydrogen is treated as a dead end beyond one bond.
  • Examples of covalent bonding in simple molecules:
    • Methane (CH₄): one carbon atom covalently bonded to four hydrogen atoms. Each H forms one bond; carbon forms four bonds total, filling its valence shell.
    • Visualization: single bonds shown as lines; each line represents a covalent bond.
    • Carbon dioxide (CO₂): one carbon atom double-bonded to each of two oxygens. Oxygen has valence of 6 (two in the first shell, six in the outer shell in this model) and needs two bonds to satisfy its electron needs.
    • Double bonds: two shared electron pairs between atoms (shown as two lines between atoms). Example: O=C=O in CO₂.
    • Nitrogen gas (N₂): diatomic molecule with a triple covalent bond (N≡N). Each nitrogen has five valence electrons and needs three more to reach eight.
  • Important nuance about bonds and how to decide when to use double or triple bonds:
    • If carbon has insufficient bonds to fill its octet, a double bond may be formed (as in C₂H₄, ethylene, where a C=C bond exists and each carbon is bonded to two hydrogens).
    • For some configurations, triple bonds (as in N₂) or even multiple bonds are required to satisfy octets.
    • The tendency to form double bonds depends on the number of available bonds after connecting to other atoms; hydrogen cannot form multiple bonds because it can only form one bond.
  • Organic versus inorganic molecules:
    • Organic molecules are typically defined as carbon-hydrogen compounds.
    • Carbon dioxide (CO₂) is not considered organic in this lecture because it lacks hydrogen.
  • Ionic compounds and example NaCl:
    • In ionic compounds, atoms gain or lose electrons to form ions with charges.
    • Cations (positive ions) form by losing electrons (e.g., Na⁺, Ca²⁺).
    • Anions (negative ions) form by gaining electrons (e.g., Cl⁻).
    • A salt like sodium chloride (NaCl) consists of Na⁺ and Cl⁻ held together by electrostatic attraction.
  • Protons, electrons, and element identity:
    • The number of protons (the atomic number) defines the element.
    • Changing the number of protons changes the element itself (e.g., changing carbon (Z=6) to nitrogen (Z=7) yields a different element).
    • The number of electrons in a neutral atom equals the number of protons, but this balance can be altered to form ions.
  • Practical implications and exploratory notes:
    • The simplest way to think about bonding is whether atoms share electrons (covalent) or transfer electrons to form ions (ionic).
    • The outermost shell (valence shell) determines reactivity and bonding patterns; full valence shells lead to noble gases that are largely inert.
    • In biology and medicine, understanding these basic principles helps explain how drugs interact with targets, how metabolic pathways operate, and how dating of biological samples can be inferred from isotope decay.
  • Carbon dating (brief):
    • Living organisms contain a small amount of ${}^{14}$C (a radioactive isotope of carbon).
    • ${}^{14}$C is unstable and decays over time; measuring the amount remaining in a sample allows estimation of age, given a relatively constant decay rate.
    • In practice, this is used to estimate the age of fossils and archaeological materials by assessing how much ${}^{14}$C has decayed since the organism died.
  • Quick recap of key numeric and symbolic facts (for study):
    • First shell capacity: 22 electrons; second shell capacity: 88 electrons; third shell capacity: 88 electrons (in this model).
    • Carbon: Z=6,<br/>≠N=6oA=12Z=6,<br />\neq N=6 o A=12 for the common isotope; other isotopes: $N=7$ for ${}^{13}$C, $N=8$ for ${}^{14}$C.
    • Oxygen valence: outer shell requires 22 electrons; forms up to 22 covalent bonds in many common molecules.
    • Methane: CH₄ uses four covalent bonds from carbon to four hydrogens.
    • Carbon dioxide: CO₂ uses two double bonds (O=C=O).
    • Nitrogen gas: N₂ uses a triple covalent bond (N≡N).
    • Noble gases (He, Ne, Ar, Kr, Xe, Rn) have full valence shells and are largely nonreactive.
    • Ionic examples: Na⁺, Cl⁻ in NaCl; cations form by losing electrons, anions by gaining electrons.
    • Hydrogen behavior: typically forms a single covalent bond; acts as a terminal “dead end” in many organic structures.
  • Connections to broader topics:
    • The discussion of electron shells and valence underpins organic chemistry (chapter on organic molecules) and biochemistry (how macromolecules form).
    • Isotopes and carbon dating connect chemistry to geology, paleontology, and archaeology.
    • Ionic versus covalent bonding influences the properties of materials, electrolytes in biology, and drug interactions.
  • Ethical and practical implications (brief):
    • When discussing radiation and isotopes, safety and proper handling are important in real-world labs; the lecture referenced working with radioactivity and memory of exposure, underscoring the importance of safety protocols in labs.
  • Quick study prompts (to test understanding):
    • What is the atomic number of carbon, and how many protons, neutrons, and electrons does a neutral ${}^{12}$C atom have?
    • Answer: Z=6;
      ightarrow p=6,
      e=6,
      n=6.
    • What kind of bond would you expect between two nitrogen atoms in N₂? Why?
    • Answer: A triple covalent bond to satisfy octets (N has 5 valence electrons and needs 3 more).
    • Why is O₂ not a “compound” in the strict sense used here?
    • Answer: It is a diatomic molecule consisting of a single element (oxygen), not a compound formed from two different elements.
    • How does carbon dating work in principle?
    • Answer: It relies on the decay of the radioactive isotope ${}^{14}$C at a relatively constant rate to estimate elapsed time since death of an organism.
  • Note for study: This summary consolidates the lecture’s key points. If you want, I can generate a condensed cheat-sheet with only the essential equations and the most commonly tested facts (e.g., valence shell capacities, common bonding patterns, isotope notation).