Chapter 2 Notes — The Chemical Context of Life
- Matter definition: anything that takes up space and has mass; exists in many forms (rocks, metals, oils, gases, living organisms).
- Element: a substance that cannot be broken down into other substances by chemical reactions.
- 92 naturally occurring elements (examples: gold, copper, carbon, oxygen).
- Element symbol: usually the first letter or two of the name; some symbols derive from Latin or German (e.g., Na from natrium).
- Compound: a substance consisting of two or more different elements in a fixed ratio (e.g., NaCl, water H2O).
- Examples: table salt NaCl (Na:Cl = 1:1), water H2O (H:O = 2:1).
- Emergent properties: compounds have chemical and physical characteristics different from their constituent elements.
- The Elements of Life: about 20–25% of natural elements are essential for life; humans ~25 essential elements, plants ~17.
- Major elements in the human body (approximate mass percentages including water):
- Oxygen (O): ~65.0%
- Carbon (C): ~18.5%
- Hydrogen (H): ~9.5%
- Nitrogen (N): ~3.3%
- Calcium (Ca): ~1.5%
- Phosphorus (P): ~1.0%
- Potassium (K): ~0.4%
- Sulfur (S): ~0.3%
- Sodium (Na): ~0.2%
- Chlorine (Cl): ~0.2%
- Magnesium (Mg): ~0.1%
- Trace elements (less than 0.01%): Boron (B), chromium (Cr), cobalt (Co), copper (Cu), fluorine (F), iodine (I), iron (Fe), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vanadium (V), zinc (Zn)
- Question prompts and concepts:
- Interpret data: given the body’s makeup, which compound accounts for the high oxygen percentage?
- Evolution of tolerance to toxic elements: sunflowers detoxifying soils by taking up heavy metals; natural selection in polluted environments.
- Concept check highlights:
- 2.1.1: Is a trace element essential? Explain.
- 2.1.2: Iron in humans as a trace element—deficiency effects?
- 2.1.3: Everyday implications of essential vs trace elements.
Concept 2.2: An element's properties depend on the structure of its atoms
- Atom basics: the smallest unit of matter that retains the properties of an element.
- Atom components:
- Nucleus: protons (+ charge) + neutrons (no charge)
- Electron cloud: electrons (− charge) surrounding the nucleus
- Atomic number (Z): number of protons in the nucleus; defines the element and equals the number of electrons in a neutral atom.
- Mass number (A): total number of protons and neutrons in the nucleus.
- Isotopes: atoms of the same element with different neutron numbers; same number of protons, different mass numbers.
- Examples for carbon: ${}^{12}\mathrm{C}$ (6 protons, 6 neutrons), ${}^{13}\mathrm{C}$ (6 protons, 7 neutrons), ${}^{14}\mathrm{C}$ (6 protons, 8 neutrons).
- Isotopes may be stable or radioactive (radioactive isotopes decay and transform into other elements).
- Atomic mass vs mass number: mass number is the total count of protons + neutrons; atomic mass is a weighted average of isotopes, measured in daltons (1 Da = 1 g/mol).
- Daltons and atomic mass unit (amu): masses of protons and neutrons ~1 Da; electrons contribute negligibly to atomic mass.
- Isotope applications in biology:
- Stable isotopes used as tracers in metabolism.
- Radioactive isotopes used in fossil dating (e.g., radiometric dating) and medical imaging (e.g., PET scanners).
- Subatomic particles:
- Protons: +1 charge; Neutrons: 0; Electrons: -1.
- Nucleus contains protons and neutrons; electron cloud held by electrostatic attraction to the nucleus.
- Atomic structure models:
- Early models show nucleus and electron shells; electrons occupy fixed energy levels (electron shells).
- Energy levels and shell structure determine chemical behavior; electrons exist only at discrete energy levels.
- Electron energy and distance from nucleus:
- More distant electrons have higher potential energy.
- The first shell holds up to 2 electrons; second shell up to 8; third shell capacity increases with atomic number.
- Electron distribution diagrams (periodic table context):
- Periods (rows) reflect filling of electron shells.
- Valence electrons: electrons in the outermost shell; determine chemical reactivity.
- Atoms with complete valence shells (e.g., He, Ne, Ar) are inert (chemically unreactive).
- Concept checks:
- 2.2.1: Build atomic number and mass number for a given isotope example.
- 2.2.2: Determine electrons for an element (e.g., fluorine) in a neutral atom and count electron shells.
- Bonding overview: atoms with incomplete valence shells bond to complete their valence shells; bonds can be covalent or ionic.
- Covalent bonds:
- Definition: sharing of a pair of valence electrons by two atoms.
- Example: H2 forms when two hydrogen atoms share one electron each; both atoms effectively complete their valence shells.
- Bonding capacity (valence): the number of covalent bonds an atom can form based on the number of electrons needed to complete its valence shell (e.g., O has valence 2; H has valence 1; C has valence 4).
- Bond representations: electron distribution diagrams, structural formulas (e.g., H-H, O=O, H-O-H), space-filling models.
- Nonpolar covalent bonds: equal sharing of electrons (e.g., H2, O2).
- Polar covalent bonds: unequal sharing due to electronegativity differences; partial charges develop (e.g., in H2O, O is partially negative, H is partially positive).
- Electronegativity:
- The attractiveness of an atom for shared electrons in a covalent bond.
- Higher electronegativity pulls electrons closer, creating partial charges in polar covalent bonds.
- Ionic bonds:
- Occur when electron transfer results in oppositely charged ions (cations and anions) that attract each other.
- Example: NaCl formation via transfer of one electron from Na to Cl, yielding Na+ and Cl− that bond ionically.
- Ionic compounds (salts) form crystal lattices; their formula indicates the ratio of ions, not discrete molecules.
- Common example: MgCl2 with Mg2+ and Cl− in a 1:2 ratio.
- In solution, ionic bonds can weaken due to hydration (water surrounds ions), facilitating dissolution.
- Ionic bonds in biological systems:
- NaCl dissolution in water involves hydration shells around Na+ and Cl−, separating ions.
- Weak bonds and their roles in biology:
- Hydrogen bonds: forms between a hydrogen atom covalently bonded to an electronegative atom (often O or N) and another electronegative atom; crucial in water structure and many biomolecular interactions.
- Van der Waals interactions: transient charges in molecules create weak attractions; cumulatively strong for large molecules (e.g., gecko adhesion).
- Hydrogen bonds and van der Waals interactions contribute to molecular recognition and shape-dependent binding (e.g., endorphins binding to brain receptors).
- Molecular shape and function:
- Molecules have characteristic shapes (e.g., H2O ~ V shape with ~104.5° angle; CH4 tetrahedral).
- Shape determines complementarity and binding with other molecules; binding often involves weak, reversible interactions.
- Concept checks:
- 2.3.1: Compare nitrogen bonding context with isotopes and valence.
- 2.3.2: Determine how many covalent bonds can be formed by an atom from its valence.
Concept 2.4: Chemical reactions make and break chemical bonds
- Chemical reactions rearrange matter by forming and breaking chemical bonds; matter is conserved (atoms are not created or destroyed).
- Reaction notation:
- Reactants → Products; coefficients indicate the number of molecules involved.
- Example: 2 H2 + O2 → 2 H2O (stoichiometry shows amounts of reactants and products).
- Photosynthesis: a key biological example of chemical rearrangement powered by light energy:
- Overall equation: 6\mathrm{CO2} + 6\mathrm{H2O} \rightarrow \mathrm{C6H{12}O6} + 6\mathrm{O2}
- Reactants: carbon dioxide and water; Products: glucose and oxygen.
- Reversibility and chemical equilibrium:
- All reactions are theoretically reversible; forward and reverse reactions can occur.
- At chemical equilibrium, forward and reverse reaction rates are equal; concentrations of reactants and products stabilize at particular ratios.
- Example: 3\mathrm{H2} + \mathrm{N2} \rightleftharpoons 2\mathrm{NH_3} (ammonia synthesis is reversible).
- Reaction rates and concentration:
- Increasing the concentration of reactants increases the rate of the forward reaction (more collisions).
- As products accumulate, the reverse reaction becomes more frequent.
- Real-world relevance:
- Balance of forward and reverse reactions underpins metabolic pathways and environmental chemistry.
Concept 2.5: Hydrogen bonding gives water properties that help make life possible on Earth
- Water as a chemical and physical marvel:
- Water is a polar molecule due to its covalent bonds and V-shaped geometry; overall charge is unevenly distributed: O is δ−; H is δ+.
- Hydrogen bonds form between water molecules (O δ− of one water to H δ+ of a neighboring water); these bonds are weaker than covalent bonds yet numerous and dynamic.
- Emergent properties of water due to hydrogen bonding:
- Cohesion: water molecules stay attached via hydrogen bonds; important for transport in plants (transpiration pull) and surface tension.
- Adhesion: water adheres to surfaces (e.g., cell walls), aiding capillary action and movement against gravity.
- Surface tension: high due to cumulative hydrogen bonding; allows droplets and organisms (e.g., water-walking spiders).
- Temperature regulation: water has a high specific heat (specific heat capacity) and a high heat of vaporization; water resists temperature change and moderates climate.
- Evaporative cooling: loss of the most energetic molecules as water evaporates cools surfaces.
- Ice density and floating: ice is less dense than liquid water due to a crystalline lattice formed by hydrogen bonding, causing ice to float; this insulates aquatic life in winter.
- Water as a versatile solvent:
- Polar molecules and ions dissolve readily in water due to hydration shells forming around solutes.
- Hydrophilic substances: have affinity for water and may dissolve or remain dispersed due to hydrogen bonding.
- Hydrophobic substances: nonionic and nonpolar, do not dissolve well in water (e.g., oils); important in cell membrane structure.
- Solutes, solvents, and solutions:
- A solution is a homogeneous mixture of solute in solvent (e.g., sugar in water).
- Aqueous solution: solvent is water.
- Hydration shells: water molecules surround dissolved ions (e.g., Na+ and Cl−) and ions are shielded from each other.
- Molarity and moles:
- Substances are often measured in moles; one mole contains Avogadro’s number of molecules: N_A = 6.02 \times 10^{23}
- The mass of one mole of a substance (in grams) equals its molecular mass in Daltons (Da).
- Example: sucrose, C12H22O11, has molecular mass 342 Da; to make a 1 M solution, dissolve 342 g of sucrose in water to reach 1 L total solution.
- Acids, bases, and pH:
- Water autoionization: \mathrm{H_2O \rightleftharpoons H^+ + OH^-} (often represented as H+ and OH− in solution; hydronium H3O+ is commonly used).
- Equilibrium constant for water at 25°C: K_w = [H^+][OH^-] = 10^{-14}
- pH scale: \mathrm{pH} = -\log [H^+]; neutral water at 25°C has pH ~7 with [H+] = [OH−] = 10^{-7} M.
- Acids increase [H+]; bases decrease [H+] or increase [OH−].
- Strong acids/bases dissociate completely in water; weak acids/bases do so reversibly (e.g., carbonic acid H2CO3 ⇌ H+ + HCO3−).
- Buffer systems minimize pH changes by absorbing or releasing H+; common biological buffer: carbonic acid-bicarbonate system in blood:
- Carbonic acid dissociation: \mathrm{H2CO3 \rightleftharpoons H^+ + HCO_3^-}
- In bloodstream, CO2 + H2O forms H2CO3, which dissociates to regulate pH (buffer reaction shifts left or right with changes in H+).
- Ocean acidification as a chemical consequence of elevated atmospheric CO2:
- Atmospheric CO2 dissolves in seawater and forms carbonic acid: \mathrm{CO2 + H2O \rightarrow H2CO3}
- Carbonic acid dissociates: \mathrm{H2CO3 \rightarrow H^+ + HCO_3^-}
- Increased H+ reduces carbonate ions (CO3^{2-}) availability, affecting CaCO3 formation and calcification (e.g., corals).
- Balance with calcium carbonate: CO3^{2-} + H^+ \rightarrow HCO3^- and related precipitation-dissolution dynamics.
- Practical implications and examples discussed in the chapter:
- Water’s properties underpin life on Earth: support for plant water transport, climate stability, and biochemistry in cells.
- Hydrophilic vs hydrophobic interactions influence solubility, protein folding, and membrane structure.
- The concept of hydration shells helps explain ion dissolution and transport in biological systems.
- The role of buffers in maintaining intracellular pH and blood pH homeostasis.
- The interplay between chemistry and biology, including real-world examples like ant defense chemicals, insect communication, and the multidisciplinary nature of biology, chemistry, and physics.
Connections, examples, and broader implications
- Multidisciplinary nature of biology: chemistry and physics underpin biological phenomena (e.g., ant chemical defenses, biomolecule binding, water’s solvent properties).
- Real-world relevance:
- Ant defense: formic acid as a defense mechanism and disinfectant; highlights chemistry’s role in survival and ecology.
- Water’s properties enable life-supporting processes: cohesion/adhesion for nutrient transport in plants; high heat capacity for climate stability; solvent properties essential for biochemical reactions.
- Ocean acidification links atmospheric CO2 to marine biology; affects calcifying organisms and ecosystems.
- Photosynthesis (overall):6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2
- Atomic/molecular mass and Avogadro’s number:
- Avogadro’s number:N_A = 6.02 \times 10^{23} molecules per mole
- Relationship: mass of one mole (in grams) equals molecular mass in Daltons (Da).
- Water autoionization and Kw:
- \mathrm{H_2O \rightleftharpoons H^+ + OH^-}
- K_w = [H^+][OH^-] = 10^{-14}\quad (25^{\circ}\mathrm{C})
- pH definition:
- \mathrm{pH} = -\log [H^+]
- Carbonic acid–bicarbonate buffer:
- \mathrm{H2CO3 \rightleftharpoons H^+ + HCO_3^-}
- Ocean acidification chemistry:
- \mathrm{CO2 + H2O \rightarrow H2CO3}
- \mathrm{H2CO3 \rightarrow H^+ + HCO_3^-}
- \mathrm{CO3^{2-} + H^+ \rightarrow HCO3^-}
- Ionic bond formation (NaCl example):
- \mathrm{Na \rightarrow Na^+ + e^-}
- \mathrm{Cl + e^- \rightarrow Cl^-}
- \mathrm{Na^+ + Cl^- \rightarrow NaCl}
- Hydration shell concept (descriptive): water molecules surround dissolved ions, separating them and shielding them from each other.
- Subatomic particle masses (approximate): protons and neutrons ≈ 1.7 × 10^{-24} g each; electrons much lighter (mass negligible for atomic mass).
Visual and conceptual reminders
- Emergent properties arise when atoms combine into compounds (e.g., NaCl properties differ from Na and Cl individually).
- The valence shell determines reactivity; full valence shells imply low reactivity (inert/neutral elements at far right of the periodic table).
- The strength and directionality of bonds (covalent vs ionic; polar vs nonpolar) shape the three-dimensional structure and function of biomolecules.
- Water’s unique properties—cohesion, adhesion, high specific heat, high heat of vaporization, ice density, and solvent capability—collectively make Earth suitable for life and shape biological processes from cellular to ecological scales.
Quick study prompts (concept checks summarized)
- Concept Check 2.1: Are trace elements always essential? Explain. What might iron deficiency do to hemoglobin function?
- Concept Check 2.2: Determine atomic number, number of protons/electrons, and valence electrons for a given element; relate to electron shells.
- Concept Check 2.3: Why does a nitrogen atom with 7 protons and a radioactive isotope with 8 neutrons differ in mass number but not in bonding behavior? How many bonds can fluorine form?
- Concept Check 2.4: Which type of chemical reaction is faster at equilibrium: forward or reverse? How would adding products affect the equilibrium?
- Concept Check 2.5: Explain how hydrogen bonding leads to water’s properties that support life.
Connections to broader topics
- This chapter establishes the foundation for understanding biochemistry, physiology, environmental science, and medicine by linking atomic structure, bonding, reactions, and water chemistry to living systems and the environment.
