The Chemical Context of Life - Comprehensive Study Notes (Concepts 2.1–2.5)

Concept 2.1: Matter consists of chemical elements in pure form and in combinations called compounds

• Matter definition: anything that takes up space and has mass; organisms are composed of matter.
• Element: a substance that cannot be broken down to other substances by chemical reactions.
• Compound: a substance consisting of two or more elements in a fixed ratio.
• Emergent properties: a compound has characteristics different from those of its constituent elements.
• Organisms are composed of matter.
• Visual references: Figure 2.1 (illustrative in slides).

The Elements of Life

• Of 92 natural elements, about 20–25% are essential elements needed by organisms to live and reproduce.
• Trace elements: required in minute quantities.
• Example: iodine (I) is required for normal activity of the thyroid gland in vertebrates.
• In humans, iodine deficiency can cause goiter.

Evolution of Tolerance to Toxic Elements

• Some naturally occurring elements are toxic to organisms.
• Arsenic is linked to many diseases and can be lethal in humans.
• Some species adapt to environments with elements that are usually toxic.
• Example: sunflowers can uptake lead, zinc, and other heavy metals in concentrations lethal to most organisms; used to detoxify contaminated soils after disasters (e.g., Hurricane Katrina).

Concept 2.2: An element’s properties depend on the structure of its atoms

• Each element consists of a specific type of atom, different from atoms of other elements.

• Atom: the smallest unit of matter that retains the properties of an element.

• Subatomic particles to know:

  • Neutrons (no electrical charge)
  • Protons (positive charge)
  • Electrons (negative charge)

• Atomic nucleus contains neutrons and protons; electrons form a cloud around the nucleus.

• Neutron mass and proton mass are almost identical; masses are measured in daltons (Da).

Subatomic Particles and Atomic Structure

• Atoms are composed of subatomic particles: neutrons, protons, electrons.

• Neutrons and protons form the atomic nucleus; electrons form a cloud around the nucleus.

• Masses: protons and neutrons have masses close to 1 Da each; electrons have negligible mass relative to nucleons.

• Visual: Figure 2.3 shows an electron cloud around a nucleus (simplified representations).

Atomic Number and Atomic Mass

• Atomic number (Z): the number of protons in the nucleus.

• Mass number (A): the sum of protons and neutrons in the nucleus, i.e., A = Z + N.

• Atomic mass: the atom’s total mass; can be approximated by the mass number A.

• Example: Sodium-23 (
  ^{23}_{11} ext{Na})

• For sodium: mass number A = 23, atomic number Z = 11; neutrons N = A − Z = 12.

• Since protons and neutrons each have mass ~1 Da, the atomic mass of sodium is ~23 Da.

Isotopes

• All atoms of an element have the same number of protons (Z) but may differ in the number of neutrons (N).

• Isotopes: two atoms of an element that differ in neutron number.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

• Applications in biology: dating fossils, tracing atoms through metabolic processes, diagnosing medical disorders.

Concept 2.3: The formation and function of molecules depend on chemical bonding between atoms

• Atoms with incomplete valence shells can share or transfer valence electrons with certain other atoms.
• This sharing or transfer leads to atoms staying close together and being held by chemical bonds.

Covalent Bonds

• Covalent bond: sharing of a pair of valence electrons by two atoms.

• Shared electrons count as part of each atom’s valence shell.

• A molecule consists of two or more atoms held together by valence bonds.

• Notation: structural formula (e.g., H—H);

• Molecular formula: e.g., H2.

• Bonding capacity (valence): the number of bonds an atom can form; usually corresponds to the number of electrons needed to complete the shell.

• Pure elements are composed of molecules of one type of atom (e.g., H2, O2).

• Compounds: molecules with two or more types of atoms (e.g., H2O, CH4).

• Electronegativity: an atom’s attraction for electrons in a covalent bond.

  • More electronegative atoms pull shared electrons more strongly.
  • Nonpolar covalent bond: equal sharing of electrons.
  • Polar covalent bond: unequal sharing, leading to partial charges on atoms.

Ionic Bonds

• Some atoms strip electrons from bonding partners, forming ions.
• Example: Na transfers an electron to Cl, forming Na+ and Cl−.
• Cation: positively charged ion; Anion: negatively charged ion.
• Ionic bond: attraction between a cation and an anion.
• Ionic compounds (salts) are formed by ionic bonds; often occur as crystals (e.g., NaCl).

Weak Chemical Bonds

• The strongest bonds in organisms are covalent bonds forming molecules, but weaker bonds (ionic, hydrogen bonds) are also important.
• Many large biological molecules are stabilized by weak bonds.

Hydrogen Bonds

• Hydrogen bond forms when a hydrogen covalently bonded to an electronegative atom is also attracted to another electronegative atom (usually O or N in biological systems).

• Visual: Figure 2.12 shows examples like H2O and NH3 with δ+ and δ− regions.

Molecular Shape and Function

• Molecular shape is crucial to function; shapes determine recognition and response between biological molecules.

• Similar shapes can have similar biological effects (structure–function relationship).

• Example: Endorphin and morphine interactions with receptors (Figure 2.14): natural endorphin vs morphine binding to brain cell receptors.

Concept 2.4: Chemical reactions make and break chemical bonds

• Chemical reactions involve making and breaking bonds.

• Reactants: starting molecules.

• Products: final molecules.

• Example: Photosynthesis (an important chemical reaction):
  6 \, \mathrm{CO}2 + 6 \, \mathrm{H}2\mathrm{O} \rightarrow \mathrm{C}6\mathrm{H}{12}\mathrm{O}6 + 6 \, \mathrm{O}2.

• All chemical reactions are reversible: forward and reverse reaction rates can be equal.

• Chemical equilibrium is reached when forward and reverse reaction rates are equal.

Concept 2.5: Hydrogen bonding gives water properties that help make life possible on Earth

• All organisms are made mostly of water and live in a water-dominated environment.

• Water is polar: oxygen region is δ−, hydrogen region is δ+.

• Two water molecules are held together by a hydrogen bond.

• Four emergent properties of water contributing to life:

  • Cohesive behavior
  • Ability to moderate temperature
  • Expansion upon freezing
  • Versatility as a solvent

Cohesion and Adhesion

• Water molecules are linked by multiple hydrogen bonds (cohesion).
• Cohesion helps transport water and nutrients against gravity in plants.
• Adhesion: clinging of water to surfaces, also relevant for movement of water in plants.
• Visual: Figure 2.17 shows cohesion and adhesion and their roles in water transport.
• Surface tension is a measure of how hard it is to break the surface of a liquid and is related to cohesion.

Moderation of Temperature by Water

• Water absorbs heat from warmer air and releases stored heat to cooler air.

• Water can absorb or release a large amount of heat with only a small change in its own temperature.

• Definitions:

  • Kinetic energy: energy of motion.
  • Thermal energy: total kinetic energy due to molecular motion.
  • Temperature: average kinetic energy of molecules.
  • Heat: thermal energy transfer between bodies.

• Units:

  • Celsius (°C) for temperature;
  • Calorie (cal): amount of heat to raise 1 g of water by 1°C.
  • Kilocalorie (kcal): 1 kcal = 1000 cal; on food packages, typically kcal.
  • Joule (J): 1 J = 0.239 cal; 1 cal = 4.184 J.

• Specific heat of water: 1 cal/(g·°C).

• High specific heat is due to hydrogen bonding: heat absorbed when bonds break; heat released when bonds form; this helps stabilize temperatures in bodies of water and organisms.

Evaporative Cooling

• Evaporation: liquid to gas transformation.
• Heat of vaporization: heat required to convert 1 g of liquid to gas.
• As water evaporates, the remaining surface cools (evaporative cooling), helping stabilize temperatures in organisms and in bodies of water.

Floating of Ice on Liquid Water

• Ice floats because hydrogen bonds in ice are more ordered, making ice less dense than liquid water.
• Water reaches maximum density at 4°C.
• If ice sank, many bodies of water would freeze solid, making life impossible on Earth.

Water as the Solvent of Life

• A solution is a homogeneous mixture of substances; solvent is the dissolving agent; solute is the substance dissolved.

• An aqueous solution has water as the solvent.

• Water’s polarity enables it to form hydrogen bonds and dissolve many substances.

• When an ionic compound dissolves, each ion is surrounded by a hydration shell of water molecules.

• Visual: Figure 2.21 demonstrates hydration shells around ions.

• Water can dissolve nonionic polar molecules as well, provided they have ionic and polar regions.

• Hydrophilic vs Hydrophobic:

  • Hydrophilic: affinity for water.
  • Hydrophobic: does not have an affinity for water (e.g., oils, nonpolar molecules).

• A colloid is a stable suspension of fine particles in a liquid.

Solute Concentration in Aqueous Solutions

• Most biochemical reactions occur in water; reaction rates depend on solute concentrations.
• Molecular mass is the sum of masses of all atoms in a molecule.
• moles: a quantity used to count molecules; 1 mole = $6.02 \times 10^{23}$ molecules (Avogadro’s number).
• Avogadro’s number and the unit dalton are such that $6.02 \times 10^{23}$ daltons = 1 g.
• Molarity (M): number of moles of solute per liter of solution.

Acids and Bases

• In water, sometimes a hydrogen ion (H+) is transferred to another water molecule, producing a hydroxide ion (OH−) and a hydronium ion (H3O+). By convention, H+ represents H3O+ in solution.

• Water dissociation is rare but important in life chemistry.

• H+ and OH− are highly reactive.

• Acids increase H+ concentration in solution; bases decrease H+ concentration.

• Examples:

  • Strong acid: HCl dissociates completely into H+ and Cl− in water: ext{HCl} \rightarrow \text{H}^+ + \text{Cl}^-.
  • Strong base: NaOH dissociates completely to form Na+ and OH−; OH− neutralizes H+ to form water: \text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-.
  • Ammonia: NH3 acts as a relatively weak base by accepting an H+ to form NH4+: \text{NH}3 + \text{H}^+ \rightleftharpoons \text{NH}4^+.
  • Carbonic acid: H2CO3 can reversibly release or accept H+: \text{H}2\text{CO}3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+.n

• pH Scale: In aqueous solution at 25°C, the product of [H+] and [OH−] is constant: [\text{H}^+] [\text{OH}^-] = 10^{-14}. The pH is defined as \text{pH} = -\log [\text{H}^+]. For neutral water, [H+] = 10^−7, so pH = 7.

• Acidic solutions have pH < 7; basic solutions have pH > 7; most biological fluids have pH roughly between 6 and 8.

Buffers

• Buffers help maintain stable internal pH by minimizing changes in H+ and OH− concentrations.
• Most buffers consist of an acid-base pair that reversibly binds H+.
• Example: carbonic acid (H2CO3) acts as a buffer contributing to pH stability in human blood.

Acids donate H+ in aqueous solutions; bases donate OH− or accept H+

• Basic diagrammatic summary: [H+] < [OH−] neutral; [H+] > [OH−] acidic; [H+] < [OH−] basic etc. (visual reference: Figure 2.UN03 and pH diagrams).

Acidification: A Threat to Our Oceans

• Human activities, especially burning fossil fuels, increase CO2 in the atmosphere.
• About 25% of human-generated CO2 is absorbed by the oceans.
• Dissolved CO2 forms carbonic acid, leading to ocean acidification.
• As seawater acidifies, H+ ions combine with carbonate ions (CO3^{2−}) to form bicarbonate (HCO3−), reducing carbonate ion availability needed by organisms that build coral reefs or shells.
• Predicted carbonate ion decline by ~40% by 2100, raising concern for calcifying organisms.
• Chemical pathway illustrated: \mathrm{CO}2 + \mathrm{H}2\mathrm{O} \rightarrow \mathrm{H}2\mathrm{CO}3 \rightarrow \mathrm{H}^+ + \mathrm{HCO}3^- \ \mathrm{H}^+ + \mathrm{CO}3^{2-} \rightarrow \mathrm{HCO}3^- \ \mathrm{CO}3^{2-} + \mathrm{Ca}^{2+} \rightarrow \mathrm{CaCO}_3.

Recap of Key Equations and Concepts

• Chemical reaction: 6\, \mathrm{CO}2 + 6\, \mathrm{H}2\mathrm{O} \rightarrow \mathrm{C}6\mathrm{H}{12}\mathrm{O}6 + 6\, \mathrm{O}2. (Photosynthesis)
• Mass and composition:
  • A = Z + N (mass number = protons + neutrons).
  • Z = number of protons (atomic number).
  • Atomic mass ≈ mass number A.
• Isotopes: same Z, different N; radioactive isotopes decay emitting particles/energy.
• Bond types:
  • Covalent bond: sharing of a pair of electrons; can be single (H–H), double (O=O), etc.
  • Ionic bond: transfer of electrons creating cations and anions; attraction between ions.
  • Hydrogen bond: attractive interaction involving H with δ+ and electronegative atoms like O or N.
• Solubility concepts:
  • Hydrophilic: water-loving; hydrophobic: water-fearing.
  • Hydration shell: water molecules surround dissolved ions.
• pH and buffers are central to biological systems; oceans are vulnerable to acidification due to CO2 uptake.

Connections and Implications

• Structure governs function: electron distribution and molecular geometry dictate reactivity and interaction.
• Weak bonds, though individually weak, collectively shape large biomolecules and cellular structures.
• Water’s properties underpin all of life: temperature regulation, solvent capabilities, and the ability to sustain stable environments.
• Balancing pH and buffering capacity is essential for enzyme activity, metabolism, and overall cellular homeostasis.
• Human activities impacting ocean chemistry have broad ecological and biogeochemical consequences, potentially altering reef formation and marine food webs.

Quick Reference for Exam Prep

• Key terms: matter, element, compound, emergent properties, essential elements, trace elements.
• Atomic concepts: atomic number Z, mass number A, isotopes, radioactive isotopes.
• Bonding: covalent, ionic, hydrogen bonds; electronegativity; polarity; valence.
• Reactions: reactants, products, reversibility, chemical equilibrium.
• Water physics: cohesion, adhesion, surface tension, specific heat, heat of vaporization, evaporative cooling, density of ice, solvent properties.
• Acids and bases: H+, OH−, hydronium H3O+, pH, buffers, carbonic acid as a buffer.
• Environmental chemistry: ocean acidification via CO2 dissolution and carbonate chemistry.