Chapter 2: The Chemical Context of Life - Practice Flashcards

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

• Organisms are composed of matter, which is anything that takes up space and has mass; matter exists in many forms (rocks, metals, oils, gases, living organisms).

• Elements and compounds:

  • An element is a substance that cannot be broken down into other substances by chemical reactions.
  • Today, there are 92 naturally occurring elements. Examples: gold (Au), copper (Cu), carbon (C), oxygen (O).
  • Each element has a symbol (usually the first letter or two); some symbols come from Latin or German (e.g., sodium = Na from natrium).
  • A compound is a substance consisting of two or more different elements fixed in a ratio (e.g.,
  • table salt: sodium chloride, NaCl (1:1 Na:Cl)
  • water: H2O (2 H : 1 O)

• Emergent properties: a compound has characteristics different from its constituent elements (e.g., NaCl is edible, whereas Na is a reactive metal and Cl2 is a poisonous gas).

• The Formic acid example (Formic acid CH2O2) illustrates that a compound’s properties depend on its atoms and their bonding; in formic acid, an O atom attracts H’s electron, releasing H+ and making the compound an acid capable of stinging.

• The 4 most abundant elements in living matter (approx. 96% total):

  • Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N)
  • Relative percentages in the human body: O
    ightarrow 65.0 ext{\%},\, C
    ightarrow 18.5 ext{\%},\, H
    ightarrow 9.5 ext{\%},\, N
    ightarrow 3.3 ext{\%}

• Other significant elements that make up most of the remaining ~4%: Ca,\, P,\, K,\, S, ext{ and a few others}

• Trace elements are required in minute quantities. Examples and notes:

  • Some trace elements (e.g., iron, Fe) are needed by all forms of life.
  • In vertebrates, iodine (I) is essential for thyroid hormone production. A daily intake of 0.15 ext{ mg} is adequate for normal thyroid activity; iodine deficiency can cause goiter.
  • Consuming seafood or iodized salt reduces goiter risk.

• Table 2.1 (Elements in the Human Body) lists relative abundances and the concept of trace elements vs essential elements.

• Some naturally occurring elements are toxic (e.g., arsenic, As). Arsenic exposure via groundwater occurs in parts of the world (e.g., southern Asia) from well water; efforts exist to reduce arsenic levels.

• Mastering Biology prompts: interpret data about element composition.

• Practical implications:

  • Understanding elemental composition informs about organismal biology, nutrition, and environmental health.
  • Emergent properties explain why compounds behave differently from their constituent atoms.

• Key terms to remember:

  • Element, compound, atom, nucleon (protons, neutrons), electron, electron cloud, valence electrons, atomic symbol, atomic number, mass number, isotope, radioactive isotope, dalton/amu.

• The Formic acid illustration and the concept of emergent properties emphasize how chemical bonding determines a compound’s properties and reactivity.

• Quick reference equations and numbers:

  • Formula for table salt: ext{NaCl}
  • Water: ext{H}_2 ext{O}
  • Formic acid (example): ext{CH}2 ext{O}2
  • Emergent-property example: Na (metal) + Cl2 (gas) → NaCl (edible compound).
  • Percentage data for major body elements: ext{O} = 65.0 ext{ ext{%}}, ext{ C} = 18.5 ext{ ext{%}}, ext{ H} = 9.5 ext{ ext{%}}, ext{ N} = 3.3 ext{ ext{%}}

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

• An element is defined by the type of atom it contains; atoms are the smallest unit that retains an element’s properties.

• Subatomic particles relevant here:

  • Protons: positively charged (+1)
  • Neutrons: neutral (no charge)
  • Electrons: negatively charged (−1); electrons form a cloud around the nucleus.

• Nucleus: contains protons and neutrons; positively charged due to protons.

• Electron cloud: rapidly moving electrons; attraction between opposite charges keeps electrons near the nucleus.

• Atomic number (Z): number of protons in an atom; determines the element’s identity and, for neutral atoms, equals the number of electrons.

• Mass number (A): total number of protons and neutrons in the nucleus; used to deduce neutron count.

• Isotopes: atoms of the same element with different numbers of neutrons; same proton number but different mass numbers.

  • Examples for carbon (Z = 6): ^{12}{6}C (6 neutrons), ^{13}{6}C (7 neutrons), ^{14}_{6}C (8 neutrons).
  • Isotopes have nearly identical chemical behavior but different masses.
  • Atomic mass is an average of isotope masses weighted by abundance; carbon’s average mass is approx 12.01 ext{ Da}.

• Radioactive isotopes: isotopes that are unstable and decay, emitting particles and energy; useful in medicine and research but hazardous due to radiation.

• Neutrons and protons have masses ~1.7 imes 10^{-24} ext{ g} each; electrons have negligible mass in mass calculations.

• Dalton (Da) or atomic mass unit (amu): standard unit for atomic/molecular masses; protons and neutrons ≈ 1 Da; electron mass is ~1/2000 of a Da.

• Atomic structure concepts:

  • Atom neutrality implies equal numbers of protons and electrons.
  • Atomic number identifies the element and counts protons/electrons in a neutral atom.
  • The mass number identifies a specific isotope (A = protons + neutrons).

• Isotopes and biological relevance:

  • Radioactive carbon-14 used in radiometric dating and tracing metabolic processes.
  • Radioactive isotopes in medicine (e.g., PET scanning) help diagnose and monitor diseases.
  • Radiation poses hazards; dose and type matter for cellular damage.

• Mass and isotope examples to remember:

  • Sodium-23 symbol: ^{23}_{11}Na; neutrons = 23-11=12.
  • The simplest atom: hydrogen-1, symbol ^{1}_{1}H, no neutrons.

• The concept of isotopes and atomic mass ties into how elements behave chemically (same Z, different A can yield the same chemistry but different mass).

• Key definitions and numbers:

  • Atomic number: number of protons (and electrons in neutral atoms).
  • Mass number: total protons + neutrons; neutron count = A − Z.
  • Isotopes: same Z, different A (and neutron count).
  • Carbon isotopes: ^{12}{6}C, ^{13}{6}C, ^{14}_{6}C with mass numbers 12, 13, 14 respectively.
  • Carbon-14 half-life: t_{1/2} = 5{,}730 ext{ years}.

• Applications mentioned:

  • Radiometric dating uses fixed half-lives to estimate fossil/rock ages.
  • PET scans visualize metabolic activity using radioactive tracers.
  • Radioactive decay curves can be calibrated and used to interpret dating data (example: Neanderthal dating exercise).

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

• Atoms with incomplete valence shells interact to complete valence shells via chemical bonds, forming molecules or ionic compounds.

• Types of bonds:

  • Covalent bonds: sharing of valence electrons between atoms; results in molecules.
  • Ionic bonds: transfer of electrons leading to oppositely charged ions that attract each other.

• Covalent bonding details:

  • The shared electrons produce a bond; the bonding capacity (valence) equals the number of electrons needed to complete the outer shell.
  • Example: Oxygen has valence 2; it typically forms two covalent bonds to complete its valence shell.
  • Bond types:
  • Nonpolar covalent bonds: equal sharing of electrons (examples: H–H in H2, O=O in O2).
  • Polar covalent bonds: unequal sharing due to different electronegativities (example: O–H in H2O).

• Electronegativity: the tendency of an atom to attract shared electrons in a covalent bond; higher electronegativity means a stronger pull on shared electrons.

  • In H2O, O is more electronegative than H, giving partial negative on O and partial positives on H.
  • CH4 has less polar covalent bonds because C and H have similar electronegativities.

• Ionic bonds:

  • Occur when one atom is much more electronegative and effectively transfers an electron to another atom, creating ions.
  • Cations (positive) and anions (negative) attract to form ionic bonds.
  • Example: NaCl forms as Na+ and Cl− ions; MgCl2 forms with Mg2+ and two Cl−.
  • Ionic compounds are typically crystalline salts; NaCl is often used as an example.
  • In solid form, ionic bonds are strong; in water, dissociation occurs and bonds are weaker due to solvation.

• The distinction between molecules and salts:

  • Molecules consist of two or more covalently bonded atoms.
  • Ionic compounds do not consist of discrete molecules; they form a lattice of ions.

• Weak interactions (reinforcement of structure):

  • Hydrogen bonds: attraction between a hydrogen atom with partial positive charge and a highly electronegative atom with partial negative charge (commonly O or N).
  • Van der Waals interactions: transient dipole attractions between nonpolar molecules; can be collectively strong (e.g., Gecko adhesion via numerous weak contacts).

• Molecular geometry and function:

  • The shape of a molecule influences function, recognition, and interactions.
  • Orbital hybridization creates specific geometries (e.g., tetrahedral for carbon in CH4; bent shape for H2O).
  • Opiates mimic endorphins by resembling their shape, binding to specific receptors, illustrating structure–function relationships in biology.

• Notable example compounds and concepts:

  • Water (H2O): polar covalent bonds due to electronegativity difference between O and H; hydrogen bonds play a critical role in water’s properties and biology.
  • Methane (CH4): tetrahedral geometry due to sp3-like hybridization around carbon; four equivalent C–H bonds.
  • Salt crystallography: NaCl lattice with alternating Na+ and Cl− ions.

• Important definitions and figures:

  • Electronegativity: the pull an atom exerts on shared electrons in a bond.
  • Valence: bonding capacity of an atom; related to the number of electrons needed to complete the valence shell (e.g., H = 1, O = 2, N = 3, C = 4).
  • The concept of inert elements: Helium (He), Neon (Ne), and Argon (Ar) have full valence shells and are generally unreactive.

• Concept checks and applications:

  • Why does a crystal of MgCl2 involve ionic bonds but dissolve in water? (Dissolution involves ion exposure to solvent and hydration overshadowing lattice energy.)
  • How does molecular shape influence binding to receptors or enzymes? (Shape complementarity governs specificity.)

• Key formulas and numbers:

  • Water polarity: H2O has partial charges: O partial negative, H partial positive (illustrated in diagrams).
  • Covalent bond examples:
  • H–H single bond
  • O=O double bond
  • H–O–H angles in H2O depend on hybridization (approximately 104.5°).
  • Ionic lattice concept: NaCl lattice with Na+ and Cl− arranged in a 3D crystal.

Concept 2.4: Chemical reactions make and break chemical bonds

• Chemical reactions rearrange matter by breaking and forming chemical bonds; matter is conserved (no atoms created or destroyed).
• Reversibility and equilibrium:
  • Reactions are reversible in principle; forward and reverse reaction rates can equalize, reaching chemical equilibrium.
  • At equilibrium, concentrations of reactants and products stabilize in a ratio, not necessarily equal amounts.
  • Some reactions may be driven essentially to completion if products are favored (concentration differences push the reaction forward).
• Photosynthesis as a key biological example:
  • Overall simplified equation (summarized):
  • Reactants: 6 ext{CO}2 + 6 ext{H}2 ext{O} + ext{light energy}
  • Products: ext{C}6 ext{H}{12} ext{O}6 + 6 ext{O}2
  • Real processes are a sequence of steps; energy from sunlight drives the rearrangement of atoms from CO2 and H2O into glucose and O2.
• The concept of “chemical equilibrium” underpins how biological systems balance synthesis and degradation pathways.
• The role of energy in reactions:
  • Energy is required to move electrons to higher energy levels during bond formation or breaking.
  • In biological systems, energy flows drive endergonic and exergonic reactions, enabling complex metabolism.
• The photosynthesis summary above emphasizes the two-way dependence: life relies on photosynthesis for organic molecules and oxygen, while respiration uses these molecules to release energy.
• Example of a basic equilibrium representation:
  • For ammonia synthesis: ext{N}2 + 3 ext{H}2
    ightleftharpoons 2 ext{NH}_3
• Important phrase: “Matter is conserved in chemical reactions.”

The Elements of Life (and the Body’s Elemental Composition)

• Four elements make up about 96% of living matter: ext{O}, ext{C}, ext{H}, ext{N}.
• The remaining ~4% is composed of Ca, P, K, S, and a few other elements.
• Trace elements are required in minute quantities; some are essential only for certain species.
• Iodine example:
  • Essential for thyroid hormone; deficiency leads to goiter.
  • Daily iodine intake: 0.15 ext{ mg} sufficient for normal thyroid activity.
  • Sea food or iodized salt reduces goiter incidence.
• Table 2.1: Elements in the Human Body (selected data)
  • Oxygen (O): 65.0 ext{ ext{%}} of body mass
  • Carbon (C): 18.5 ext{ ext{%}}
  • Hydrogen (H): 9.5 ext{ ext{%}}
  • Nitrogen (N): 3.3 ext{ ext{%}}
  • Calcium (Ca): 1.5 ext{ ext{%}}
  • Phosphorus (P): 1.0 ext{ ext{%}}
  • Potassium (K): 0.4 ext{ ext{%}}
  • Sulfur (S): 0.3 ext{ ext{%}}
  • Sodium (Na): 0.2 ext{ ext{%}}
  • Chlorine (Cl): 0.2 ext{ ext{%}}
  • Magnesium (Mg): 0.1 ext{ ext{%}}
  • Trace elements: less than 0.01 ext{ ext{%}} each (examples: B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn)
• Essential elements: ~20–25% of natural elements are essential; in humans, about 25 elements are needed; in plants, about 17.
• Toxic elements and environmental health:
  • Arsenic (As) toxicity and groundwater exposure (e.g., southern Asia). Efforts are underway to reduce arsenic exposure.
• Concept connections:
  • Elemental composition explains bulk properties of organisms, nutrition, and environmental interactions.
  • Emergent properties in compounds arise from how atoms bond and spatial arrangement.

Isotopes, Radioactivity, and Radiometric Dating

• Isotopes are atoms of the same element with different numbers of neutrons and thus different mass numbers, but the same number of protons.
• Radioactive isotopes decay at fixed rates (half-lives) and can transform into different elements:
  • Example: Carbon-14 decays to Nitrogen-14 via neutron-to-proton conversion when decaying, changing the element.
• Radioactive tracers in biology: usable as diagnostic tools by incorporating them into biologically active molecules and tracking their distribution/metabolism (e.g., PET imaging).
• Radiometric dating relies on radioactive decay to estimate the age of fossils or rocks:
  • For a radioactive isotope with a fixed half-life, after t half-lives, fraction remaining is ext{fraction} = igl( frac{1}{2}igr)^t.
  • The fraction remaining for carbon-14 in a fossil can be converted to years using the known half-life: t_{1/2} ext{(C-14)} = 5{,}730 ext{ years}.
  • A fraction around 0.0078 of the original 14C remaining corresponds to a specific age (Neanderthal case study example).
• Radioactivity hazards: Radiation can damage cellular molecules; dose and type of radiation determine risk; medical uses (diagnostic imaging) are typically at safe levels.
• Long half-lives: Uranium-238 half-life is t_{1/2} = 4.5 imes 10^{9} ext{ years}; useful in dating very old materials but less relevant for short-term biology.
• Practical applications:
  • PET scans (visualize high metabolic activity) and radiometric dating.
  • In clinical settings, radiopharmaceuticals help diagnose organ function, cancer, and other diseases.

Energy, Electron Distribution, and Atomic Structure

• Energy and electron shells:

  • Electrons exist in discrete energy levels (shells); they cannot occupy energies between levels.
  • Energy can be absorbed to move electrons to higher shells or released when electrons drop to lower shells.
  • The first shell (closest to the nucleus) holds up to 2 electrons; the second shell can hold up to 8 electrons; electrons fill the lowest available shell first.
  • The distribution of electrons across shells determines chemical behavior and reactivity (valence electrons).

• The concept of an orbital:

  • An orbital is a three-dimensional region where an electron is likely to be found; orbitals form the basic components of electron shells (e.g., s and p orbitals).
  • The first shell has one s orbital (1s). The second shell has one 2s orbital and three 2p orbitals (2px, 2py, 2p_z).
  • Each orbital can hold a maximum of two electrons.

• Valence and reactivity:

  • Atoms with incomplete valence shells are reactive; those with complete valence shells are inert (e.g., noble gases like He, Ne, Ar).
  • Valence electrons determine the number of bonds an atom can form (e.g., H has valence 1, O has valence 2, N has valence 3, C has valence 4).

• The periodic table organization:

  • Elements are arranged in rows (periods) corresponding to the filling of electron shells.
  • The left-to-right progression reflects increasing atomic number and electron addition.
  • Visuals show electron distributions (1s, 2s, 2p) and the corresponding valence electrons.

• Molecular geometry and orbital hybridization:

  • Covalent bonding involves rearrangement of valence-shell orbitals into hybrid orbitals (e.g., tetrahedral for carbon in CH4).
  • Water (H2O) uses two occupied hybrid orbitals from oxygen to form two O–H bonds; two lone pairs occupy the remaining hybrid orbitals, giving a bent shape with an angle of about 104.5°.
  • Methane (CH4) forms a tetrahedral geometry because all four hybrid orbitals on carbon are used to bond with hydrogens.

• Shape and function:

  • The 3D shape of molecules affects recognition and binding to other molecules; examples include interactions with receptors and endorphin signaling.
  • Molecular mimicry: morphine resembles endorphins and can bind to endorphin receptors, illustrating how shape relates to function and physiological effects.

• Visual and conceptual tools mentioned:

  • Electron distribution diagrams for elements (neon, argon, etc.) illustrate electron shells and valence behavior.
  • Lewis dot structures and line-notations (e.g., H:H for H2) to represent covalent bonds.
  • Space-filling models and ball-and-stick models illustrate three-dimensional molecular geometry.

Covalent Bonds, Ionic Bonds, and Molecular Interactions

• Covalent bonds: sharing of a pair of valence electrons between atoms; bond strength and polarity depend on electronegativity differences.

  • Nonpolar covalent bonds: equal sharing (e.g., H–H in H2, O═O in O2).
  • Polar covalent bonds: unequal sharing due to electronegativity differences (e.g., O in H2O pulls electrons more strongly, creating partial negative O and partial positive H).

• Ionic bonds: transfer of electrons leading to oppositely charged ions (cations and anions) that attract to form an ionic bond.

  • Examples: NaCl (Na+ and Cl−) and MgCl2 (Mg2+ with two Cl−).
  • Ionic compounds form crystal lattices; NaCl is commonly used as a canonical example.
  • In aqueous environments, ionic bonds can be weakened by solvation and hydration.

• Weak interactions and their roles:

  • Hydrogen bonds: attraction between a hydrogen attached to an electronegative atom (e.g., O or N) and another electronegative atom with a lone pair; central to the properties of water and biomolecules.
  • Van der Waals interactions: transient, nonpolar attractions between molecules or regions of large molecules; collectively can be strong (e.g., gecko adhesion).

• Molecular recognition and shape: the shape and charge distribution of molecules determine how they recognize and interact with each other; complementary shapes enable specific binding (e.g., drug-receptor interactions).

• Practical examples and figures from the chapter:

  • Water’s polarity and hydrogen bonding contribute to its solvent properties and its role in biology.
  • O in H2O is more electronegative than H, leading to a dipole and partial charges (d− on O, d+ on H).
  • The covalent bonds in H2, O2, H2O, and CH4 illustrate different bonding and shapes.
  • Ionic compounds such as NaCl consist of ions held together by ionic bonds in a lattice; in water, these bonds dissociate due to hydration.

Chemical Reactions, Equilibrium, and Photosynthesis

• Chemical reactions involve breaking old bonds and forming new bonds to rearrange atoms; matter is conserved.
• Equilibrium and reversibility:
  • Reactions can proceed forward or reverse; at equilibrium, forward and reverse reaction rates are equal, resulting in stable concentrations of reactants and products.
  • The relative amounts of reactants and products reflect their ratio at equilibrium, not necessarily equal concentrations.
• Photosynthesis (summary):
  • Raw materials: ext{CO}2 and ext{H}2 ext{O}
  • Energy input: sunlight
  • Products: ext{C}6 ext{H}{12} ext{O}6 (glucose) and ext{O}2
  • Overall reaction, in simplified form: 6 ext{CO}2 + 6 ext{H}2 ext{O}
    ightarrow ext{C}6 ext{H}{12} ext{O}6 + 6 ext{O}2
  • Photosynthesis is a multi-step process; energy from sunlight drives the rearrangement of atoms. All chemical reactions are technically reversible, but biological systems exploit energy inputs to push reactions in favorable directions.
• The concept of equilibrium is central to metabolism and biochemistry.
• Visual concepts to recall:
  • Reactants on the left, products on the right, with an arrow to denote the direction of reaction and to represent reversibility.
  • In the context of biology, many reactions are catalyzed by enzymes that influence reaction rates toward equilibrium in living systems.

The Energy and Structure of Atoms: Orbitals, Shells, and Shapes

• Energy levels and atomic stability:
  • Electrons exist in discrete energy levels; a lower energy level is more stable.
  • The distance from the nucleus correlates with energy level; electrons occupy the lowest available energy levels.
  • The first shell is closest to the nucleus and holds up to 2 electrons; the second shell holds up to 8 electrons.
• Atomic orbitals and hybridization:
  • Orbitals are the regions where electrons are most likely to be found.
  • For the second shell, there are four orbitals: one 2s and three 2p orbitals.
  • When atoms bond, their valence orbitals can hybridize to form new orbital shapes (e.g., sp3 hybrids leading to tetrahedral geometry in CH4).
• Concept of valence and reactivity:
  • The chemical behavior of an atom is largely determined by the number of electrons in its outer (valence) shell.
  • Atoms with incomplete valence shells are reactive; those with full valence shells are generally inert (e.g., He, Ne, Ar).
• Examples of common valence and orbitals:
  • Hydrogen: 1 electron in 1s orbital; valence = 1.
  • Oxygen: valence = 2; forms two bonds (e.g., H2O).
  • Carbon: valence = 4; can form four bonds (e.g., CH4).
• Molecular geometry and function:
  • Shapes such as bent water (H2O) vs tetrahedral methane (CH4) arise from hybridization and lone-pair effects.
  • Shape determines molecular recognition, binding affinity, and biological function (e.g., receptor-ligand interactions).
• Visual notes:
  • Electron distribution diagrams illustrate how electrons occupy shells and orbitals.
  • The classic Lewis dot notation and line drawings (H:H, H–H; H2O; CH4) capture essential bonding concepts.

Molecular Shape, Endorphins, and Chemical Recognition

• Molecular shape governs recognition: biological molecules bind to specific partners only if their shapes are complementary.
• Endorphins, morphine, and receptors:
  • Endorphins are endogenous signaling molecules that bind to receptors to relieve pain and produce euphoria.
  • Morphine mimics endorphins in shape and can bind to the same receptors, explaining why opiates have similar effects.
• The concept of structure–function relationship is a unifying theme in biology: molecular shape determines function and interactions.

Weak Interactions Reinforcing 3D Structure and Function

• Van der Waals interactions:
  • Arise from transient dipoles in molecules; individually weak but collectively can be significant (e.g., gecko adhesion).
• Hydrogen bonds:
  • A hydrogen atom covalently bonded to a highly electronegative atom (like O or N) is attracted to another electronegative atom nearby.
  • Central to the properties of water and to the 3D structure of biomolecules (e.g., DNA base pairing, protein folding).
• Combined, weak interactions help stabilize biological macromolecules and enable transient, reversible interactions crucial for biology.

Practical Implications and Connections

• The Chemical Context of Life establishes the foundational ideas for biology:
  • Matter is composed of elements and compounds; the properties of compounds depend on their constituent atoms and bonds.
  • The arrangement of electrons determines atom identity, reactivity, and bonding.
  • Bonds (covalent and ionic) and weak interactions (hydrogen bonds, van der Waals) shape molecules, their stability, and their interactions in cells.
  • The geometry and 3D structure of molecules drive recognition, signaling, and function in biological systems.
• Real-world relevance:
  • Understanding element abundance helps explain nutrition and health risks (e.g., iodine deficiency and goiter).
  • Isotopes and radiometric dating provide tools for archaeology, geology, and medicine.
  • Covalent vs ionic bonding and molecular geometry explain why certain drugs interact with receptors or why salts dissolve in water.
• Foundational numerical and formula references to remember:
  • Photosynthesis overall equation (simplified): 6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2
  • Half-life of Carbon-14: t{1/2} = 5{,}730\text{ years}; general decay relationship: N = N0\left(\tfrac{1}{2}\right)^{t/t_{1/2}}
  • Uranium-238 half-life: t_{1/2} = 4.5 \times 10^{9} \text{ years}
  • Atomic number and mass number concepts:
  • Example: ^{4}_{2}\mathrm{He} has 2 protons and 2 neutrons; mass approx 4 Da.
  • Isotopes all have the same number of protons but different numbers of neutrons, e.g., ^{12}{6}C, ^{13}{6}C, ^{14}_{6}C.
  • The mass of electrons is negligible in atomic mass calculations; protons and neutrons are ~1 Da each; electron mass is about 1/2000 Da.
• Case study prompts mentioned:
  • Serpentine plant communities show local adaptation to soils rich in toxic elements (Cr, Ni, Co); natural selection leads to species adapted to such soils, illustrating evolution and tolerance to toxins.
  • The Evolution case highlights natural selection and local adaptation as drivers of biodiversity.

Quick reference glossary and numbers

• Elements, compounds, and emergent properties: compounds have properties distinct from their constituent elements.

• Major biological elements: ext{O}, ext{C}, ext{H}, ext{N} (approx. 96%); other essential elements: ext{Ca}, ext{P}, ext{K}, ext{S}, ext{others}; trace elements: present in very small amounts.

• Essential elements in organisms: humans ~25; plants ~17.

• Iodine and thyroid function: daily iodine needs ~0.15\text{ mg}; deficiency leads to goiter; iodized salt and seafood help.

• Key bonds: covalent bonds (shared electrons), ionic bonds (electrons transferred to form ions), hydrogen bonds (between H and electronegative atoms), and van der Waals interactions (transient dipoles).

• Bond polarity and electronegativity:

  • Polar covalent bonds: unequal sharing (e.g., O–H in H2O).
  • Nonpolar covalent bonds: equal sharing (e.g., H2, O2).

• Molecular shapes and functions are tightly integrated with orbital hybridization and the spatial arrangement of atoms.

• Photosynthesis and energy flow: chemical energy from sunlight drives the formation of glucose from CO2 and H2O; oxygen is a by-product.

• Radiometric dating and isotopes:

  • Carbon-14 dating is effective up to about 75,000 years.
  • Half-lives underpin dating curves and age estimates for fossils and archaeological sites.

• Connections to previous and future topics:

  • Builds the foundation for understanding biochemistry, cellular metabolism, and physiology covered in later chapters.
  • Sets the stage for discussing water chemistry, macromolecules, enzyme function, and energy transformations in cells.