AP BIOLOGY II | Chapter 2 Notes: The Chemical Context of Life

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

• Matter is anything that takes up space and has mass; organisms are composed of matter.
• An element is a substance that cannot be broken down to other substances by chemical reactions.
• A compound is a substance consisting of two or more elements in a fixed ratio; a compound has emergent properties, characteristics different from those of its elements.
• Elements and Compounds:
  • Matter is made up of elements.
  • A compound is a substance with two or more elements in a fixed ratio.
  • Emergent properties arise when elements form compounds (e.g., NaCl).
• The Elements of Life:
  • Of 90+ natural elements, about 20–25% are essential elements needed by an organism to live and reproduce.
  • Trace elements are required in minute quantities.
  • Example: iodine (I) is required for normal activity of the thyroid gland; iodine deficiency can cause goiter.
• Emergent Properties of a Compound (Figure 2.2): Na + Cl → NaCl (sodium chloride) illustrates how compound properties differ from constituent elements.
• Elements in the Human Body (Table 2.1): by mass (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)
• Evolution of Tolerance to Toxic Elements:
  • Some naturally occurring elements are toxic to organisms.
  • Arsenic can be lethal in humans; some species adapt to toxic environments.
  • Example: sunflowers can take up lead, zinc, and other heavy metals in concentrations lethal to many organisms; used to detoxify contaminated soils after Hurricane Katrina.
• Key concepts to connect: matter, elements, compounds, emergent properties, essential and trace elements, and environmental adaptations.

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 other elements.
• An atom is the smallest unit of matter that retains the properties of an element.
• Subatomic particles (relevant): neutrons (no electrical charge), protons (positive charge), electrons (negative charge).
• Atom structure:
  • Neutrons and protons form the atomic nucleus.
  • Electrons form a cloud of negative charge around the nucleus.
  • Neutron mass and proton mass are almost identical and are measured in daltons.
• Isotopes (1 of 2):
  • All atoms of an element have the same number of protons but may differ in the number of neutrons.
  • Isotopes are atomic forms that differ in neutron number.
  • Radioactive isotopes decay spontaneously, giving off particles and energy.
• Isotopes (2 of 2):
  • Applications: dating fossils, tracing atoms through metabolic processes, diagnosing medical disorders.
  • Hazards: radiation can damage cellular molecules.
• Atomic Number and Atomic Mass (1 of 2):
  • Atomic number = number of protons in the nucleus.
  • Mass number = protons + neutrons in the nucleus.
  • Atomic mass approximated by the mass number.
• Atomic Number and Atomic Mass (2 of 2):
  • Example: Sodium (Na) has mass about 23 daltons; neutrons and protons each have mass ≈ 1 dalton, so total atomic mass ≈ mass number.
• Subatomic particles and the nucleus:
  • Protons (p+) and neutrons (n) reside in the nucleus.
  • Electrons (e−) orbit the nucleus in electron shells.
• Summary: the identity and properties of an element are determined by its atomic number, mass number, isotopes, and the arrangement/ratio of subatomic particles.

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 other atoms to achieve stability, leading to chemical bonds.
• Covalent Bonds (1 of 7):
  • 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 covalent bonds.
• Covalent Bonds (2 of 7) – Molecular representations:
  • Molecular formula shows the types and numbers of atoms in a molecule.
  • Electron distribution diagrams and structural formulas show shared electrons.
• Covalent Bonds (3 of 7):
  • A single bond = sharing of one pair of electrons (represented by a single line).
  • A double bond = sharing of two pairs of electrons (represented by a double line).
• Covalent Bonds (4 of 7):
  • Bonding capacity (valence) = the number of bonds an atom can form, typically corresponding to the electrons needed to complete the outermost shell.
• Covalent Bonds (5 of 7):
  • Pure elements are composed of molecules of one type of atom (e.g., O2).
  • Compounds are molecules with two or more different types of atoms (e.g., H2O, CO2).
• Covalent Bonds (6 of 7):
  • Electronegativity = an atom’s attraction for the electrons in a covalent bond.
  • Higher electronegativity pulls shared electrons more strongly.
• Covalent Bonds (7 of 7):
  • Nonpolar covalent bonds: electrons shared equally.
  • Polar covalent bonds: unequal sharing, leading to partial charges (dipoles).
• Polar Covalent Bonds in Water (Figure 2.9): water is polar due to unequal sharing of electrons between O and H.
• Ionic Bonds (1 of 3):
  • Atoms can transfer electrons, forming ions (cations and anions).
  • An ionic bond is the attraction between oppositely charged ions.
• Ionic Bonds (2 of 3):
  • Example: Na transferring its valence electron to Cl → Na+ and Cl−; both achieve full valence shells.
• Ionic Bonds (3 of 3):
  • Compounds formed by ionic bonds are ionic compounds or salts (e.g., NaCl); salts often form crystals in nature.
• Weak Chemical Interactions:
  • Most strong bonds in organisms are covalent, but many functional forms are held by weak bonds.
  • Ionic bonds in water, hydrogen bonds, and van der Waals interactions are examples of weak interactions.
• Hydrogen Bonds: (Definition)
  • A hydrogen bond forms when a hydrogen atom covalently bonded to an electronegative atom is also attracted to another electronegative atom nearby.
  • In living cells, the electronegative partners are usually oxygen or nitrogen.
• van der Waals Interactions:
  • Caused by momentary uneven distribution of electrons, creating transient positive/negative regions.
  • These weak interactions occur when molecules are very close; can be cumulatively strong (e.g., gecko adhesion).
• Molecular Shape and Function (1 of 2):
  • Molecule size and shape are key to function.
  • Shape determines how biological molecules recognize/respond to one another.
• Molecular Shape and Function (2 of 2):
  • Biological molecules may bind temporarily through weak interactions if shapes are complementary.
  • Molecules with similar shapes can have similar biological effects.
• Molecular Mimic Example (Figure 2.14): endorphin vs morphine; morphine can bind to endorphin receptors due to mimicked shape.

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

• Chemical reactions involve making and breaking chemical bonds.
• Reactants are starting molecules; products are final molecules.
• Photosynthesis (important chemical reaction): sunlight powers conversion of carbon dioxide and water to glucose and oxygen.
• Reversibility: all chemical reactions are theoretically reversible; forward and reverse reactions occur.
• Chemical Equilibrium: reached when forward and reverse reaction rates are equal.

Concept 2.5: Hydrogen Bonding Gives Water Properties That Help Make Life Possible on Earth

• Water is a polar molecule: oxygen region is partially negative; hydrogens are partially positive.
• Hydrogen bonds form between water molecules; at any instant, most water molecules are hydrogen-bonded to neighbors.
• Four emergent properties of water support life on Earth:
  • Cohesive behavior
  • Ability to moderate temperature
  • Expansion upon freezing
  • Versatility as a solvent
• Cohesion and Adhesion:
  • Water molecules stick to each other (cohesion) and to other surfaces (adhesion).
  • Cohesion contributes to water transport in plants; adhesion helps water cling to cell walls, counteracting gravity.
• Surface Tension: a measure of how hard it is to break the surface of a liquid; related to cohesion.
• Water Transport in Plants (Figure 2.17): evaporation pulls water upward; two types of water-conducting cells; cohesion and adhesion support transport.
• 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.
• Temperature and Heat (1 of 2):
  • Kinetic energy = energy of motion; thermal energy = total kinetic energy of molecular motion.
  • Temperature represents average kinetic energy; heat is thermal energy transferred between bodies.
• Temperature and Heat (2 of 2):
  • Calorie (cal) = amount of heat required to raise the temperature of 1 g of water by 1°C.
  • Food Calories (Calories with capital C) are kilocalories (kcal): 1 kcal = 1000 cal; 1 kcal raises the temperature of 1 kg of water by 1°C.
• Water’s High Specific Heat (1 of 2):
  • Specific heat of water is about 4.18 J g⁻¹ °C⁻¹ (often rounded as 4.18 J g⁻¹ K⁻¹ or 1 cal g⁻¹ °C⁻¹).
  • Water resists temperature change due to hydrogen bonding.
• Water’s High Specific Heat (2 of 2):
  • Heat is absorbed when hydrogen bonds break; heat is released when hydrogen bonds form.
  • This property helps stabilize temperatures in bodies of water and organisms.
• Evaporative Cooling:
  • Evaporation is the transformation of a substance from liquid to gas; heat of vaporization is the energy required to convert 1 g to gas.
  • Evaporation cools surfaces and helps stabilize temperatures.
• Floating of Ice on Liquid Water (1 of 2):
  • Ice floats because its hydrogen bonds are more ordered, making ice less dense than liquid water.
  • Water reaches its greatest density at 4°C.
• Floating of Ice on Liquid Water (2 of 2):
  • Floating ice insulates the water below and supports life beneath the frozen surface.
  • Ice also provides a habitat for polar animals; concern over disappearance due to climate change.
• Water: The Solvent of Life (1 of 3):
  • A solution is a homogeneous mixture; solvent is the dissolving agent; solute is dissolved substance; an aqueous solution has water as solvent.
• Water: The Solvent of Life (2 of 3):
  • Water’s polarity allows it to form hydrogen bonds and act as a solvent for many solutes.
  • When ionic compounds dissolve, each ion is surrounded by a hydration shell of water molecules.
• Water: The Solvent of Life (3 of 3):
  • Nonionic polar molecules can dissolve in water if they have ionic and polar regions on their surface.
• Hydrophilic and Hydrophobic Substances:
  • Hydrophilic substances have an affinity for water.
  • Hydrophobic substances lack affinity for water (e.g., oils) due to nonpolar bonds.
• Solute Concentration in Aqueous Solutions (1 of 2):
  • Most biological reactions involve solutes dissolved in water.
  • Reactions depend on solute concentration in a solution.
• Solute Concentration in Aqueous Solutions (2 of 2):
  • Molecular mass is the sum of all atomic masses in a molecule.
  • Numbers of molecules are measured in moles (mol).
  • Avogadro’s number: 1 mol contains N_A = 6.022 × 10²³ entities.
  • The unit dalton (Da) is used for atomic/molecular masses; 1 Da ≈ 1 g/mol.
  • Molarity (M) = moles of solute per liter of solution: M = rac{n{ ext{solute}}}{V{ ext{solution}}}
• Acids and Bases (1 of 6):
  • Sometimes a hydrogen ion is transferred between water molecules, producing a hydronium ion (H₃O⁺).
  • By convention, the hydronium ion is denoted as H₃O⁺.
  • Hydration yields H₂O ⇌ H⁺ + OH⁻ with hydronium involved in acid-base chemistry.
• Acids and Bases (2 of 6):
  • Water dissociation is rare but essential to life; strong acids and bases are highly reactive.
  • Acids increase [H⁺] in water; bases decrease [H⁺].
• Acids and Bases (3 of 6):
  • A strong acid such as hydrochloric acid, HCl, dissociates completely in water to H⁺ and Cl⁻.
  • Ammonia, NH₃, acts as a weak base by accepting a hydrogen ion to form NH₄⁺; this reaction is reversible.
• Acids and Bases (4 of 6):
  • Sodium hydroxide, NaOH, dissociates completely to form OH⁻, which combines with H⁺ to form water.
  • A solution with equal concentrations of H⁺ and OH⁻ is neutral.
• Acids and Bases (5 of 6):
  • Weak acids act reversibly and can donate back hydrogen ions; carbonic acid, H₂CO₃, acts as a weak acid.
• Acids and Bases (6 of 6):
  • Weak bases are also reversible in accepting ions; examples include the role of carbonic acid as a buffer precursor.
• The pH Scale (1 of 2):
  • In any aqueous solution, the product [H⁺][OH⁻] is constant, related to water’s autoionization.
  • The pH is defined as ext{pH} = -\log [ ext{H}^+].
  • For a neutral solution, [H⁺] = [OH⁻] so pH = 7; neutral solutions have [H⁺] = 10^{-7} M at 25°C.
• The pH Scale (2 of 2):
  • Acidic solutions have pH < 7; basic (alkaline) solutions have pH > 7.
  • Most biological fluids have pH values in the range 6–8.
• Buffers (1 of 2):
  • Buffers help maintain internal pH by resisting changes in [H⁺] and [OH⁻].
  • Buffers typically contain a weak acid and its conjugate base, which convert reversibly to maintain pH.
• Buffers (2 of 2):
  • Carbonic acid (H₂CO₃) acts as a buffer that contributes to pH stability in human blood.
• Acidification: A Threat to Our Oceans (1 of 2):
  • Human activities (fossil fuel combustion) increase atmospheric CO₂, about 25% of which is absorbed by oceans.
  • Dissolved CO₂ forms carbonic acid, leading to ocean acidification.
• Acidification (2 of 2):
  • With increasing acidity, H⁺ reacts with carbonate (CO₃²⁻) to form bicarbonate (HCO₃⁻); carbonate ion concentrations are projected to decline by about 40% by 2100.
  • Organisms that build coral reefs or shells rely on carbonate ions; their availability decreases under acidification.
• Ocean Acidification (Figure 2.25):
  • CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻; H⁺ + CO₃²⁻ ⇌ HCO₃⁻; Ca²⁺ + CO₃²⁻ ⇌ CaCO₃ (illustrating the carbonate equilibrium affecting CaCO₃ formation).

Connections to Foundational Principles and Real-World Relevance

• Structure–function relationship: atomic structure and bonding explain molecular geometry and biological activity.
• Emergent properties: properties of compounds cannot be deduced from elements alone (e.g., NaCl’s properties differ from Na or Cl).
• Water as a solvent: polarity and hydrogen bonding enable diverse solubility, essential for biochemistry and physiology.
• Energy and matter: chemical bonds store energy; energy changes drive metabolism, photosynthesis, and thermal regulation.
• Environmental chemistry: ocean acidification links atmospheric chemistry to marine biology and ecosystem stability.
• Ethical/philosophical: understanding molecular interactions informs medicine, environmental stewardship, and technological advances.

Key Equations and Quantities (LaTeX)

• pH definition: ext{pH} = -\log [ ext{H}^+]
• Water autoionization and neutrality (at 25°C): K_w = [ ext{H}^+][ ext{OH}^-] = 10^{-14}, \ [ ext{H}^+] = [ ext{OH}^-] = 10^{-7} ext{ M}, \ ext{pH} = 7
• Covalent bonding (symbolic): A–A' sharing electrons to form a molecule; for example, H–H, O=O, H–O–H, C–H with varying bond orders.
• Covalent bond representations:
  • Structural formula uses lines to indicate shared electron pairs.
• Common ion charges: Na⁺, Cl⁻, Ca²⁺, CO₃²⁻, H⁺, OH⁻.
• Hydration shell: when a salt dissolves in water, each ion is surrounded by a shell of water molecules.
• Specific heat of water: c( ext{water}) \,\approx\, 4.18\ \text{J}\ \text{g}^{-1}\ \text{°C}^{-1}
• Calorie relationships: 1\ \text{cal} = 1\ \text{g} \cdot 1^{\circ}\text{C} \quad\text{and}\quad 1\ \text{kcal} = 1000\ \text{cal}
• Molarity: M = \frac{n{ ext{solute}}}{V{ ext{solution}}}
• Avogadro’s number and the mole concept: NA = 6.022\times 10^{23} \quad\text{and}\quad 1\text{ mol} = NA\
• Reaction reversibility and equilibrium: forward rate = reverse rate at equilibrium.
• Ocean carbonate system (simplified):
  ext{CO}2 + ext{H}2 ext{O} \rightleftharpoons ext{H}2 ext{CO}3 \rightleftharpoons ext{H}^+ + ext{HCO}3^- \rightleftharpoons ext{H}^+ + ext{CO}3^{2-}
  ext{Ca}^{2+} + ext{CO}3^{2-} \rightarrow ext{CaCO}3
• Density of ice vs water (4°C): ice is less dense due to an expanded hydrogen-bond network compared to liquid water.
• Neutral, acidic, basic pH ranges: ext{pH} < 7 ext{ (acidic)}, \ 7 ext{ (neutral)}, \ ext{pH} > 7 ext{ (basic)}$$