Life's Chemistry and Water – Key Concepts Summary
2.1 An Element's Atomic Structure Determines Its Properties
All matter is made of atoms; element identity is defined by the number of protons in the nucleus (atomic number, Z).
Atoms have a nucleus (protons, neutrons) and electrons in orbit around the nucleus in electron shells/orbitals.
Isotopes differ in neutron number but not in identity (same Z); isotope mass affects atomic mass but not element identity.
Atomic mass unit: dalton, Da; electron mass \approx 0.0005 Da.
Bohr model vs orbitals: shells have capacity (innermost \le 2 e⁻, next \le 8, then 18, etc.); orbitals (s and p) within shells determine energy and shape.
Outer (valence) electrons determine chemical properties; atoms tend to achieve full outer shells (octet rule): 2 in s + 6 in p = 8 in outer shell; noble gases have full shells and are largely inert.
Electronegativity increases from bottom-left to top-right of the periodic table; more electronegative atoms attract electrons more strongly.
Common abundant elements in life: H, C, N, O; these are highly electronegative and form diverse bonds.
Element identity is set by the number of protons; atomic mass varies with neutrons (isotopes).
Ions and ionic bonds: loss/gain of electrons forms cations and anions; oppositely charged ions attract to form ionic bonds.
In water-rich biology, ions are typically found in solution, where hydration shells reduce lattice strength and enable ionic interactions.
Examples: NaCl forms a lattice in the solid; in solution, ions are solvated by water.
2.2 Atoms Bond to Form Molecules
Covalent bonds: sharing of electrons; strong bonds.- Nonpolar covalent: electronegativity difference _\lesssim 0.5; electrons shared roughly equally.
Polar covalent: electronegativity difference _\approx 0.5\text{–}2; electrons pulled toward more electronegative atom; uneven distribution creates partial charges.
Directionality and bond length depend on the bonded atoms; carbon often forms four covalent bonds in a tetrahedral geometry.
Ionic bonds: transfer of electrons creating cations and anions; strong electrostatic attraction forms ionic compounds (often crystalline lattices).
Hydrogen bonds: weak electrostatic interactions between a slightly positive H (from a polar bond) and a electronegative atom (N, O, or F) on another molecule or within a molecule; collectively strong in biology.
Van der Waals interactions: weak, transient dipole-induced dipole attractions; significant when many such contacts occur.
Polar and nonpolar concepts:- Polar covalent bonds create dipoles (e.g., H–O in water).
Nonpolar molecules interact mainly via van der Waals; hydrophobic substances resist water.
Water as solvent: water dissolves many substances via hydration shells; ions dissociate and are stabilized by water’s dipoles.
Amphipathic molecules: have both hydrophilic and hydrophobic regions (e.g., soaps).
Electronegativity scale (examples): O 3.4, Cl 3.2, N 3.0, C 2.6, H 2.2, Na 0.9, K 0.8. If akeupside electronegativity _\le 0.5 \rightarrow nonpolar covalent; _> 2 \rightarrow ionic.
Summary: bonds determine molecule structure, polarity, and interactions; covalent bonds link atoms, while weaker interactions (H-bonds, van der Waals) influence structure and function.
2.3 Chemical Transformations Involve Energy and Energy Transfers
Energy basics:- Kinetic energy: energy of motion (thermal, sound, electromagnetic).
Potential energy: stored energy (chemical-bond energy, gravity, elasticity).
Energy is conserved: \text{Total energy before} = \text{Total energy after}} (First Law of Thermodynamics).
Energy transformations drive biological change: e.g., light energy _\rightarrow chemical-bond energy (photosynthesis); chemical-bond energy _\rightarrow motion (muscle contraction).
Second Law of Thermodynamics: entropy tends to increase; usable energy decreases after each transformation.
Chemical bonds and free energy: - Exergonic (spontaneous) if _\triangle G < 0; energy is released.
Endergonic if _\triangle G > 0; energy input is required.
Activation energy (E_a): energy required to start a reaction; even exergonic reactions need an input to proceed.
Reaction rate factors: temperature (more collisions and higher energy), concentration (more collisions), and activation energy (lowered by catalysts such as enzymes).
Hydrolysis vs condensation:- Hydrolysis: water-added; large molecules broken into smaller ones; generally exergonic.
Condensation (dehydration): water produced; builds larger molecules; typically endergonic.
Pathways: many biological processes occur in steps, with overall energy changes and entropy increases at each step.
2.4 Chemical Reactions Transform Substances
Reactions involve breaking and forming bonds; energy differences and entropy changes determine - Exergonic vs endergonic nature via _\triangle G.
Activation energy (E_a) governs whether a reaction proceeds spontaneously and at what rate.
Rates depend on activation energy, temperature, and reactant concentrations.
Enzymes reduce E_a, increasing reaction rates in living systems.
Sucrose hydrolysis example: - Reactants: sucrose + H_2O; products: glucose + fructose; energy is released and entropy increases.
Condensation to form sucrose is endergonic and has higher activation energy; it can occur in reverse under appropriate conditions.
Equilibrium: forward and reverse reaction rates balance at equilibrium; concentrations determine net direction.
2.5 The Properties of Water Are Critical to the Chemistry of Life
Water properties arise from polarity of the O–H bonds and hydrogen bonding.
Water as a solvent:- Hydrophilic substances dissolve in water via interactions with partial charges.
Hydrophobic substances do not dissolve well and tend to aggregate.
Hydration shells stabilize ions in solution.
Water’s heat properties:- High specific heat: large amount of heat required to change temperature.
High heat of vaporization: large energy required to vaporize.
These properties help moderate temperature in organisms and environments.
Cohesion and adhesion:- Cohesion: water molecules attract to each other (surface tension).
Adhesion: water interacts with polar surfaces; enables capillary movement in plants.
Self-ionization and pH:- Pure water _\rightleftharpoons H\text{⁺} + OH\text{⁻}; in water, H\text{⁺} often exists as H_3O\text{⁺} (hydronium).
pH is the negative log of [H\text{⁺}] (pH scale).
Acids increase [H\text{⁺}] (lower pH); bases increase OH\text{⁻} (higher pH).
Buffers: resist pH changes by absorbing excess H\text{⁺} or OH\text{⁻}; important in biology (e.g., blood buffering around pH 7.35–7.45).
Acid rain and seawater chemistry: dissolution of CaCO_3 affects ocean pH; implications for marine life.
Water’s role in chemistry: dissolves ions and polar molecules; supports aqueous reactions essential for life.
2.6 Functional Groups Give Molecules Specific Properties
Carbon-based biomolecules: carbohydrates, lipids, proteins, nucleic acids; function and structure shaped by functional groups.
Functional groups (examples and general properties):
Methyl (-CH3): nonpolar; hydrophobic.
Hydroxyl (-OH) (alcohols): polar; can form hydrogen bonds; often increases solubility in water.
Sulfhydryl (-SH) (thiols): polar-ish; forms disulfide bridges; can stabilize protein structure.
Aldehyde (-CHO): polar; highly reactive; important in energy-releasing reactions.
Ketone (C=O within carbon skeleton): polar; reactive; central in carbohydrate metabolism.
Carboxyl (-COOH): acidic; readily ionizes to -COO⁻; participates in condensation (forms peptide bonds).
Amine (-NH₂): basic; accepts H\text{⁺}; participates in condensation (peptide bonds).
Phosphate (-PO₄³⁻/⁻²): acidic; participates in condensation and stores energy (e.g., ATP-related chemistry).
Isomers:
Structural isomers: same formula, different connectivity.
Stereoisomers: same connectivity, different 3D arrangement.
Functional groups confer characteristic reactivity and influence molecular shape and interactions with water and other molecules.
Example note: phosphate groups are highly polar and interact with water; adding a phosphate can turn a hydrophobic molecule into a hydrophilic, polar one.
Visual Summary:
Functional groups help explain molecular behavior in biological contexts.
Isomerism adds molecular diversity beyond a single formula.