Notes on Small Molecules and the Chemistry of Life (Ch. 2)

Concept 2.1 An Element’s Atomic Structure Determines Its Properties

  • All matter is composed of atoms.
  • Atom composition:
    • Electrons: negligible mass; negative charge
    • Nucleus:
    • Protons: have mass; positive charge
    • Neutrons: have mass; no charge
  • Element: fundamental substance containing only one kind of atom.
    • Atomic number Z identifies an element (number of protons).
    • The number of protons and electrons helps determine how an element behaves in chemical reactions.
  • Atomic structure and the periodic table (Figure 2.2 reference):
    • Atomic number (number of protons) is shown; chemical symbol; atomic weight.
  • Isotopes and mass number:
    • Mass number A = protons + neutrons.
    • Isotopes: forms of an element with different numbers of neutrons, hence different mass numbers.
    • Examples for carbon:
    • $^{12}$C has 6 neutrons
    • $^{13}$C has 7 neutrons
    • $^{14}$C has 8 neutrons
  • Special names for some isotopes and applications:
    • Some isotopes have specific names; radioactive isotopes (radioisotopes) can be unstable and undergo radioactive decay.
    • Radioactive decay types: alpha, beta, gamma radiation.
  • Isotope analysis in biology:
    • Carbon isotopes ratio $^{13}$C:$^{12}$C can identify origin of biological samples.
  • Investigating LIFE: Determining Beef Source in Big Macs Using Isotope Analysis (1 of 3)
    • Hypothesis: Beef in Big Macs served in different countries comes from the same source.
    • Method: 1) Obtain Big Macs from China and USA. 2) Heat burgers at high temperature to form CO2 gas.
  • Investigating LIFE (2 of 3)
    • Method: 3) Separate and measure isotopes of carbon in a mass spectrometer.
  • Investigating LIFE (3 of 3)
    • Results and conclusion: Big Mac burgers in the two countries have different ratios of $^{13}$C:$^{12}$C; hypothesis is rejected.
  • Mass number and isotopes in practice:
    • Isotopes may be stable or unstable; stable isotopes contribute to molecular identity, while unstable isotopes (radioisotopes) reveal processes and dating information.
  • Electrons and bonding (core idea):
    • The number of electrons determines how an atom will combine with other atoms.
  • Electron shell structure (overview from Figure 2.5 summary):
    • First shell (innermost): 1 orbital; holds 2 electrons
    • Second shell: 4 orbitals; holds 8 electrons
    • Additional shells: 4 orbitals; each shell holds 8 electrons
  • Electron configuration and reactivity:
    • Reactive atoms can share electrons, or lose or gain electrons to achieve stability.
    • Octet rule: tendency to form stable molecules by achieving full valence shells (typically 8 electrons for outer shells, except H/He).
  • Summary of key ideas:
    • Atomic structure sets elemental properties and behavior in reactions via protons/electrons.
    • Isotopes provide mass-number diversity and analytical applications (e.g., tracing origins).
    • Electron arrangements drive bonding and molecular formation using octet stabilization.

Concept 2.2 Atoms Bond to Form Molecules

  • Chemical bond: attractive force linking atoms to form molecules.
  • Covalent bonds: atoms share one or more electron pairs so that outer shells are filled; each atom contributes one partner of each shared pair.
  • Compound: pure substance composed of two or more different elements in a fixed ratio (e.g., $\mathrm{H_2O}$: two H atoms bonded to one O atom).
  • Molecular weight: sum of atomic weights of all atoms in a molecule.
  • Common molecular structures (Figure 2.7 reference):
    • Hydrogen (H2), Oxygen (O2), Water (H2O), Methane (CH4)
  • Covalent bonding capabilities (Table 2.2): usual number of covalent bonds per element
    • Hydrogen (H): 1
    • Oxygen (O): 2
    • Sulfur (S): 2
    • Nitrogen (N): 3
    • Carbon (C): 4
    • Phosphorus (P): 5
  • Covalent bond types:
    • Single bonds: share 1 pair of electrons
    • Double bonds: share 2 pairs of electrons
    • Triple bonds: share 3 pairs of electrons
    • Bond energies are higher for multiple bonds
  • Electronegativity:
    • Attractive force an atomic nucleus exerts on electrons in a bond.
    • Depends on number of protons and distance to electrons.
  • Electronegativity table (Table 2.3): (approximate values from transcript)
    • Oxygen (O): 3.5
    • Chlorine (Cl): 3.1
    • Nitrogen (N): 3.0
    • Carbon (C): 2.5
    • Phosphorus (P): 2.1
    • Hydrogen (H): 2.1
    • Sodium (Na): 0.9
    • Potassium (K): 0.8
  • Bond polarity:
    • Nonpolar covalent bond: electrons shared equally (similar electronegativities)
    • Polar covalent bond: one atom more electronegative; electrons drawn toward that nucleus
  • Polar covalent bonds in water:
    • Figure 2.9 shows water’s covalent bonds are polar with partial charges (δ+ on H, δ− on O).
    • Polar covalent bonds enable strong hydrogen bonding with other polar molecules.
  • Ionic bonding:
    • Occurs when one atom is much more electronegative; electrons are transferred, creating ions with full outer shells.
    • Classic example: formation of NaCl by transfer of electrons.
    • Reaction example (simplified): Na+Cl→Na++Cl−\mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na^+} + \mathrm{Cl^-}
  • Ions:
    • Cations: positively charged
    • Anions: negatively charged
    • Complex ions: groups of covalently bonded atoms carrying a charge (e.g., $\mathrm{NH4^+}$, $\mathrm{SO4^{2-}}$)
    • Ionic bonds arise from electrical attraction between ions
  • Hydration and ions in water:
    • Water molecules surround ions in solution (hydration shells) (Figure 2.11)
  • Hydrogen bonds and Van der Waals forces:
    • Hydrogen bond: attraction between the δ− end of one molecule and the δ+ hydrogen end of another
    • Hydrophilic vs hydrophobic:
    • Polar molecules that form hydrogen bonds with water are hydrophilic
    • Nonpolar molecules (e.g., hydrocarbons) are hydrophobic
    • van der Waals forces: attractions between nonpolar molecules that are in close proximity

Concept 2.3 Chemical Reactions Transform Substances

  • Chemical reactions occur when atoms collide with enough energy to rearrange bonds and form new substances.
  • Example: Combustion of propane
    • Reaction: propane reacts with oxygen to form carbon dioxide and water, releasing energy and light.
    • Balanced representation (conceptual): C<em>3H</em>8+5 O<em>2→3 CO</em>2+4 H2O+energy\mathrm{C<em>3H</em>8 + 5\,O<em>2 \rightarrow 3\,CO</em>2 + 4\,H_2O + \text{energy}}
  • Redox (oxidation–reduction) reactions:
    • Electron transfer between molecules
    • Oxidizing agent gains electrons and is reduced (e.g., oxygen)
    • Reducing agent loses electrons and is oxidized (e.g., propane)
    • Mnemonic: OIL-RIG (Oxidation Is Loss, Reduction Is Gain)

Concept 2.4 The Properties of Water Are Critical to the Chemistry of Life

  • Water’s unusual properties stem from polarity and hydrogen bonding.
  • Ice floats on liquid water:
    • In ice, each molecule hydrogen-bonds to four others in a crystalline, open lattice, creating more space and lower density than liquid water.
  • Water thermodynamics:
    • High specific heat: large amount of energy required to raise the temperature of water due to extensive hydrogen bonding.
    • High heat of fusion: ice requires substantial energy to melt, absorbing heat.
    • High heat of vaporization: substantial energy to convert liquid water to steam.
  • Cohesion and adhesion:
    • Cohesion: hydrogen bonds causing water molecules to stick to each other
    • Adhesion: attraction of water molecules to different substances
  • Aqueous solutions:
    • Life processes occur in aqueous solutions; a solution is a solute dissolved in a solvent.
  • Acids and bases in water:
    • Acids release hydrogen ions: HCl→H++Cl−\mathrm{HCl \rightarrow H^+ + Cl^-}
    • Bases accept H+ ions; strong base example: NaOH→Na++OH−\mathrm{NaOH \rightarrow Na^+ + OH^-}
    • OH⁻ reacts with H⁺ to form water: OH−+H+→H2O\mathrm{OH^- + H^+ \rightarrow H_2O}
    • Weak bases include bicarbonate $\mathrm{HCO3^-}$, ammonia $\mathrm{NH3}$, and amine groups $\mathrm{NH_2}$.
  • pH and the pH scale:
    • pH is the negative logarithm of the hydrogen ion concentration: pH=−log⁡[H+]\mathrm{pH} = -\log [\mathrm{H^+}]
    • In pure water, [H+]=1×10−7 M[\mathrm{H^+}] = 1 \times 10^{-7} \text{ M}, so pH=7\mathrm{pH} = 7 (neutral).
    • Lower pH means higher ${[\mathrm{H^+}]}$ and greater acidity.
  • pH values of familiar substances (from Figure 2.17):
    • The scale ranges from acidic (low pH) to basic (high pH), with neutral at pH 7
    • Examples (approximate from the table):
    • Battery acid ~ pH 1
    • Stomach acid ~ pH around 1–2
    • Lemon juice ~ pH 2
    • Vinegar/cola ~ pH 3
    • Beer ~ pH 4
    • Tomatoes ~ pH ~4
    • Grapes ~ pH ~4–5
    • Black coffee ~ pH ~5
    • Rain ~ pH ~5–6
    • Saliva ~ pH ~6
    • Distilled water ~ pH 7
    • Seawater ~ pH ~8
    • Baking soda ~ pH ~9
    • Milk of magnesia ~ pH ~10
    • Household ammonia ~ pH ~11
    • Household bleach ~ pH ~12
    • Oven cleaner ~ pH ~13–14
  • Homeostasis and buffers:
    • Living organisms maintain relatively constant internal conditions (homeostasis).
    • Buffers help maintain constant pH by resisting changes in H+ concentration.
  • Practical and real-world relevance:
    • Isotope analysis and acid-base chemistry underpin techniques in biology, medicine, environmental science, and forensic applications.
  • Key definitions to remember:
    • Atom, element, isotope, mass number, covalent bond, ionic bond, hydrogen bond, cohesion, adhesion, solution, pH, buffer, homeostasis.
  • Core equations and symbols to recall:
    • Mass number: A=Z+NA = Z + N
    • Atomic number: Z=extnumberofprotonsZ = ext{number of protons}
    • Isotope notation: ZAextX^{A}_{Z} ext{X} where X is the element symbol
    • Covalent bond types refer to shared electron pairs (single, double, triple)
    • Hydrogen ion concentration and pH: pH=−log⁡[H+]\mathrm{pH} = -\log [\mathrm{H^+}]
    • Acid-base reactions (examples):
    • HCl→H++Cl−\mathrm{HCl \rightarrow H^+ + Cl^-}
    • NaOH→Na++OH−\mathrm{NaOH \rightarrow Na^+ + OH^-}
    • OH−+H+→H2O\mathrm{OH^- + H^+ \rightarrow H_2O}
  • Overall takeaways:
    • The arrangement of electrons in atoms governs bonding patterns and molecule formation.
    • Water’s properties enable life-supporting chemistry in aqueous environments and drive important processes like enzyme function, transport, and reaction kinetics.
    • The combination of covalent, ionic, hydrogen bonding, and van der Waals interactions shapes the structure and behavior of biological molecules and systems.