chapter 2 intro chemistry

Periodic Table and Orbital Diagrams

• Goals for today: quick chemistry recap to set up chemistry concepts important for biology (periodic table, orbital diagrams, valence electrons), then ionic and covalent bonding, water-related topics, and pH.
• Major theme: in biology we focus on the major elements highlighted in color on the periodic table; these elements make up about 99% of the human body.
• Examples mentioned as major elements: Calcium, Iron, Carbon, Oxygen, Magnesium, Chlorine (as an example of an element used in physiology).
• Two broad categories of elements in body chemistry:
  • Major elements (≈99% of body mass) – e.g., C, H, O, N, Ca, Mg, Cl, etc.
  • Other elements are present but are not the primary focus for physiology in this course.
• Basic element features:
  • Element symbol (usually first letter or first two letters).
  • Atomic number (number of protons; equals number of electrons in a neutral atom).
  • Outer shells are called valence shells; bonding behavior is largely determined by valence electrons.
• Orbital filling rules (the “octet” context here):
  • First shell can hold up to 2 electrons.
  • All subsequent shells can hold up to 8 electrons (octet rule for stability in many atoms).
  • Noble gases have full valence shells and typically do not bond.
• Noble gases and bonding:
  • Helium (He) has 2 electrons total; its first shell is full.
  • Neon (Ne) has 10 electrons (2 in the first shell, 8 in the second) and is also full.
  • Other noble gases include Argon (Ar), Krypton (Kr), Xenon (Xe), etc.; they do not readily form bonds.
• Carbon as a key example:
  • Total electrons for carbon: 6.
  • Shell filling: first shell = 2 electrons; second shell = 4 electrons (2, 4).
  • Carbon can form up to four bonds (e.g., with hydrogen or oxygen), which underpins its central role in organic molecules (e.g., CO₂).
• Quick shell/electron examples to solidify the idea:
  • Carbon: $ext{C: } 2,4 ext{ (or } 1s^2 \, 2s^2 \, 2p^2)$
  • Sodium: $ext{Na: } 2,8,1 ext{ (or } 1s^2 \, 2s^2 \, 2p^6 \, 3s^1)$
  • Chlorine: $ext{Cl: } 2,8,7 ext{ (or } 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^5)$
• Practical takeaway: as you move across the periodic table and up the table, electronegativity generally increases (see below); noble gases are generally inactive because their valence shells are full.

Ionic and Covalent Bonding: Sodium and Chlorine as Examples

• Overview: bonding behavior arises from how atoms achieve stable electron configurations by either sharing or transferring electrons.
• Sodium (Na) example:
  • Neutral Na has 11 electrons: shells arranged as 2, 8, 1.
  • To become stable, it can either gain 7 electrons to fill a shell (energy-inefficient) or lose 1 electron to form a stable configuration with 10 electrons total in the surrounding shells.
  • Lost electron yields a positively charged ion: Na⁺. The atom now has 11 protons and 10 electrons.
  • In short: Na → Na⁺ + e⁻ is the oxidation process; the cation formation is energetically favorable here.
• Chlorine (Cl) example:
  • Neutral Cl has 17 electrons: shells 2, 8, 7.
  • Gaining 1 electron to reach a full third shell (8) is energetically favorable, producing Cl⁻.
  • In short: Cl + e⁻ → Cl⁻; the anion formation is favorable here.
• Ionic bonding concept:
  • Electron transfer creates oppositely charged ions (Na⁺ and Cl⁻) that attract each other via electrostatic forces to form an ionic bond (e.g., in NaCl).
  • In solution (e.g., salt in water), NaCl dissociates into Na⁺ and Cl⁻: ext{NaCl}{(s)} ightarrow ext{Na}^+{(aq)} + ext{Cl}^-_{(aq)}
• Covalent bonding concept (brief): in many biological molecules, atoms share electrons rather than transfer them, forming covalent bonds; the polarity of bonds depends on electronegativity differences (see next section).
• Practical connections:
  • In physiology, Na⁺, K⁺, Cl⁻ are key ions underpinning electrochemical gradients, membrane potentials, and signaling.
  • The idea of electron transfer and sharing helps explain why salts dissociate in water and how ionic vs covalent interactions differ in biological contexts.

Electronegativity, Polarity, and Hydrogen Bonding

• Electronegativity trend:
  • Electronegativity increases as you move up a group (toward the top of the table) and increases as you move from left to right across a period.
  • The left side of the periodic table is less electronegative than the right side.
  • In the visual, elements toward the top-right are more electronegative; toward the bottom-left are less electronegative.
  • Conceptual rule: higher electronegativity means a stronger pull on shared electrons in bonds.
• Polar vs nonpolar bonds and molecules:
  • Polar bonds occur when there is a significant electronegativity difference between bonded atoms, leading to partial positive and partial negative charges.
  • Nonpolar bonds occur when the electronegativity difference is small or zero.
  • A molecule can have polar bonds arranged in a way that yields an overall polar molecule, or can be overall nonpolar despite some polar bonds, depending on symmetry and distribution.
• Water as a polar molecule and hydrogen bonding:
  • Water (H₂O) has polar covalent bonds due to the electronegativity difference between O and H.
  • The molecule has partial positive charges on the hydrogens and a partial negative charge on the oxygen, enabling hydrogen bonding between water molecules.
  • A hydrogen bond is a relatively weak intermolecular interaction where a hydrogen atom covalently bonded to a highly electronegative atom (like O) experiences attraction to another electronegative atom (like O) on a neighboring molecule.
  • Hydrogen bonds are crucial for the structure of DNA (base-pairing) and for water’s properties (liquidity, cohesion, etc.).
• Visualizing polarity and partial charges:
  • In diagrams, a polar bond is often shown with a small delta- (negative) near the electronegative atom and delta+ near the less electronegative atom.
  • For example, in H₂O, O carries a partial negative charge and each H carries a partial positive charge, leading to a bent molecule and strong hydrogen-bonding potential.
• Question prompts from today’s activity:
  • Given Na (2,8,1) and Cl (2,8,7), why does Na tend to lose an electron while Cl tends to gain one?
  • Why does salt dissociate in water but still remain largely in the same place relative to one another inside solution?

Polar and Nonpolar Molecules in the Context of the Cell Membrane

• Amphipathic nature of phospholipids:
  • Phospholipids have polar heads (hydrophilic) and nonpolar tails (hydrophobic).
  • This amphipathic character drives the formation of a bilayer, with polar heads facing water on both sides and nonpolar tails tucked inside.
• The phospholipid bilayer as a membrane structure:
  • The bilayer forms a protective barrier that helps maintain cell shape and integrity.
  • The polar heads interact with cytosol (inside the cell) and interstitial fluid (outside the cell).
  • The nonpolar tails create a hydrophobic interior that discourages many polar and charged molecules from passing freely.
• Membrane components that reinforce structure and transport:
  • Cholesterol (a steroid) is highly hydrophobic and stabilizes the membrane by modulating fluidity and rigidity.
  • Glycolipids and glycoproteins contribute to membrane stability, signaling, and recognition.
  • Together, these components help preserve the membrane architecture and selective permeability.
• Important reminder about polarity in biological context:
  • While water is polar and drives many interactions, biological membranes often involve molecules with mixed polar/nonpolar regions, resulting in complex interactions with water and ions.
• Practical takeaway for summative understanding:
  • Understanding polar vs nonpolar helps explain why certain molecules diffuse through membranes while others require transport mechanisms.
  • It also helps explain the distribution and behavior of lipids, proteins, and carbohydrates in cellular membranes.

Hydrogen Bonds, DNA, and Practical Lab-oriented Notes

• Hydrogen bonds in biology:
  • Hydrogen bonds are a form of weak intermolecular attraction that contribute to DNA base-pair stability and water’s unique properties.
  • They arise from partial positive charges on hydrogen and partial negative charges on electronegative atoms like oxygen.
• DNA context:
  • DNA is a double helix held together by hydrogen bonds between complementary bases; these bonds are strong enough to stabilize the structure yet weak enough to be separated during replication.
• Interactive plan for Monday (as discussed):
  • Instructor will bring water kits to demonstrate hydrogen bonding in real time.
  • Students will have the opportunity to observe hydrogen bonds and related interactions firsthand.

What to Do Before Monday: Exam Prep and Practical Steps

• Two objectives before Monday:
  • Study for the Chapter 1 exam (lab quiz scheduled during the first 20 minutes of class).
  • Bring your device to Connect (online platform) for the exam and reminders.
• Practical tips mentioned by the instructor:
  • If you forget your device, visit the instructor; reminders will be posted on Friday and Monday.
  • The instructor commonly arrives around 07:45–07:50 and is available to help before class.

Quick Reference Highlights (Key Takeaways)

• Electron capacity per shell: $ext{First shell} = 2, ext{ other shells} = 8.$ Oxygen and carbon rely on these rules to form bonds and determine polarity.
• Carbon’s bonding versatility (2,4) enables diverse organic chemistry; it can form up to four covalent bonds.
• Sodium (Na): 2,8,1 configuration; tends to lose one electron to form Na⁺; Chlorine (Cl): 2,8,7 configuration; tends to gain one electron to form Cl⁻; together they form NaCl (ionic bond).
• In water, salt dissociates: ext{NaCl}{(s)} ightarrow ext{Na}^+{(aq)} + ext{Cl}^-_{(aq)}
• Electronegativity trends: increases up and to the right on the periodic table; greater for elements on the top-right; this underpins polarity and hydrogen bonding.
• Polar molecules (like water) have partial charges that enable hydrogen bonding; nonpolar molecules lack such charge separation.
• Phospholipid bilayer is amphipathic: polar heads interact with water; nonpolar tails form a hydrophobic interior, contributing to membrane stability.
• Hydrogen bonds are central to DNA base pairing and to water’s properties; demonstrations (e.g., water kits) can help visualize these interactions.
• Monday plan: exam prep plus a hands-on demonstration of hydrogen bonding; bring your device to connect for updates and reminders.