Chapter 2 Notes: The Chemical Context of Life
CONCEPT 2.1: Matter and elements
Matter is anything that takes up space and has mass; organisms are composed of matter.
A chemical 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 that are different from its constituent elements.
A compound’s properties depend on its atoms and how they are bonded together.
Bond formation and the arrangement of atoms determine the behavior of compounds (e.g., formic acid).
The Elements of Life
About 20–25% of the 92 natural elements are essential (essential elements).
Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) make up 96% of living matter.
Most of the remaining 4% consists of calcium (Ca), phosphorus (P), potassium (K), and sulfur (S).
Trace elements are required in only minute quantities.
Table 2.1: Elements in the Human Body
Oxygen (O): 65.0% of body mass
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% of mass): B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn.
Case Study: Evolution of Tolerance to Toxic Elements
Some elements are toxic, yet some species adapt to environments containing toxic elements.
Example: certain plant communities are adapted to serpentine soils.
CONCEPT 2.2: Atomic structure determines element properties
An element consists of unique atoms; an atom is the smallest unit of matter that retains the properties of an element.
Subatomic Particles (Figure 2.4)
Neutrons: no electrical charge (neutral)
Protons: positive charge
Electrons: negative charge
Protons and neutrons form the atomic nucleus; electrons form a cloud around the nucleus.
Neutron mass and proton mass are nearly identical and measured in daltons; electrons are very light and are ignored when calculating total atomic mass.
Atomic Number and Atomic Mass
Atomic number: number of protons in the nucleus.
Mass number: sum of protons and neutrons in the nucleus.
Atomic mass (approximately) can be estimated by the mass number.
Isotopes and Radioactivity
All atoms of an element have the same number of protons but may differ in neutrons.
Isotopes are atoms of an element that differ in neutron number.
Some isotopes are radioactive and decay spontaneously, emitting particles and energy.
Radioactive isotopes are used as diagnostic tools in medicine (radiotracers) and in research.
They can track atoms through metabolism and are used with imaging instruments (e.g., PET scanners monitor growth and metabolism of cancers).
Radiometric Dating
A parent isotope decays into a daughter isotope at a fixed rate; expressed as a half-life.
In radiometric dating, the ratio of isotopes is measured to calculate how many half-lives have passed since the fossil or rock formed.
Half-life values vary from seconds or days to billions of years.
The Energy Levels of Electrons
Energy is the capacity to cause change; potential energy arises from location or structure.
Matter tends to move toward the lowest possible state of potential energy.
Electrons differ in potential energy based on distance from the nucleus.
Changes in electron potential energy occur in fixed steps; electrons reside in electron shells, each with a characteristic distance and energy.
An energy-level diagram or shell diagram depicts these levels.
Electron Distribution and the Periodic Table
The chemical behavior of an atom is determined by the distribution of electrons in its electron shells.
The periodic table reflects electron distributions; left-to-right across a row corresponds to sequential addition of electrons and protons.
Valence electrons are the electrons in the outermost shell (valence shell); they largely determine an atom’s chemical behavior.
Atoms with a full valence shell are chemically inert.
Electron Orbitals
An orbital is the three-dimensional space where an electron is found 90% of the time.
Each electron shell contains a specific number of orbitals.
No more than 2 electrons can occupy a single orbital.
Atoms interact to complete their valence shells.
Examples include visible diagrams of s and p orbitals (1s, 2s, 2p; 3s, 3p, etc.).
CONCEPT 2.3: Bonding and the formation of molecules/ionic compounds
Atoms with incomplete valence shells can share or transfer valence electrons with certain other atoms.
These interactions usually keep atoms close together, held by chemical bonds.
Covalent Bonds
Covalent bond: sharing 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.
A single covalent bond = sharing one pair of valence electrons.
A double covalent bond = sharing two pairs of valence electrons.
Structural formulas show bonds (e.g., H–H for H2, O═O for O2); molecular formulas are shorthand (e.g., H2).
Bonding capacity is called valence. Covalent bonds can form between atoms of the same element or different elements.
A compound is a combination of two or more different elements.
Electronegativity and Bond Polarity
Electronegativity: an atom’s attraction for electrons in a covalent bond.
The more electronegative an atom is, the more strongly it pulls shared electrons toward itself.
Nonpolar covalent bond: equal sharing of electrons.
Polar covalent bond: unequal sharing; results in partial positive/negative charges on atoms.
Ionic Bonds and Ionic Compounds
Some atoms strip electrons from bonding partners, forming ions.
Cation: positively charged ion; Anion: negatively charged ion.
Oppositely charged ions attract, forming an ionic bond.
Compounds formed by ionic bonds are salts (e.g., NaCl).
Salts are often found as crystals; NaCl’s formula indicates the ratio of elements in a crystal.
Most salts are stable when dry but dissociate easily in water.
Weak Chemical Interactions
The strongest bonds in organisms are covalent, but many biological molecules are stabilized by weak bonds.
The reversibility of weak bonds can be advantageous for dynamic biological processes.
Types of weak interactions: hydrogen bonds and van der Waals interactions.
Hydrogen Bonds
Hydrogen bond forms when a hydrogen covalently bonded to an electronegative atom is attracted to another electronegative atom.
In cells, electronegative partners are usually oxygen or nitrogen.
Van der Waals Interactions
Uneven electron distribution can create transient regions of positive/negative charge.
Van der Waals interactions are attractions between molecules close to each other due to these charges.
Can be collectively strong (e.g., gecko toe hairs and wall surfaces).
Molecular Shape and Function
A molecule’s size and shape are critical to its function.
Geometry is determined by the positions of atoms’ orbitals.
Hybridization of s and p orbitals can create specific molecular shapes (e.g., tetrahedral geometry in CH4).
Space-filling and ball-and-stick models illustrate shapes.
Molecular shape influences biological recognition and response; e.g., morphine and natural endorphins have similar shapes and bind the same brain receptors.
Figure references illustrate shapes and binding interactions.
CONCEPT 2.4: Chemical reactions involve making and breaking bonds
Chemical reactions create or break chemical bonds to transform reactants into products.
Reactants are the starting molecules; products are the resulting molecules.
Reactions are often reversible: forward and reverse reactions occur.
Examples of Chemical Reactions
Photosynthesis: sunlight powers the conversion of carbon dioxide and water to glucose and oxygen.
Balanced equation: 6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2
Reversibility and Equilibrium
All chemical reactions are reversible: products of the forward reaction can become reactants of the reverse reaction.
Example of a reversible reaction: 3\,H2 + N2 \rightleftharpoons 2\,NH_3
Chemical equilibrium is reached when forward and reverse reactions occur at the same rate.
At equilibrium, the relative concentrations of reactants and products do not change.
Practical and Real-World Relevance
Understanding chemical bonds and reactions explains how biological molecules form, function, and interact.
Electronegativity and polarity explain solubility and molecular interactions in biology.
Hydrogen bonds and water’s properties underpin many physiological processes (protein folding, DNA structure, etc.).
Weak interactions enable dynamic processes (enzyme-substrate binding, DNA base pairing, etc.).
Isotopes and radiometric dating connect chemistry to Earth history and archaeology.
Radiotracers and PET imaging connect chemistry to medicine and diagnostics.
Photosynthesis equation highlights the fundamental chemical basis of life’s energy capture.
Connections to Foundational Principles
Atomic theory: elements, isotopes, atomic number, mass, electrons, and nucleon structure.
Chemical bonding: covalent, ionic, and weak interactions shape molecular behavior.
Electron structure and the periodic table explain reactivity and valence.
Energy and thermodynamics: electrons occupy shells; systems favor low potential energy.
Chemical reactions and equilibrium underpin metabolism and biosynthesis.
Ethical, Philosophical, and Practical Implications
Use of radioactive isotopes in medicine raises safety, ethical, and regulatory considerations.
Radiometric dating informs our understanding of Earth’s history and poses interpretive limitations.
The reversibility of reactions and dynamic bonding are central to biology and can influence drug design and toxicology.
The study of molecular recognition and binding has implications for pharmacology, neuroscience, and personalized medicine.
Key Formulas and Notation
Photosynthesis (balanced): 6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2
Ammonia synthesis (example of a reversible reaction): 3\,H2 + N2 \rightleftharpoons 2\,NH_3
Atomic concepts: atomic number = number of protons; mass number = protons + neutrons; atomic mass ≈ mass number.
Electron configuration and valence: valence electrons reside in the outermost shell and determine chemical behavior; full valence shell → inert.
Electronegativity: attraction of an atom for electrons in a covalent bond; more electronegative atoms pull electrons more strongly.
Ionic bond: transfer of electrons leading to cations and anions; attraction between ions forms the bond.
Covalent bond: sharing of electron pairs between atoms; single bond = one pair, double bond = two pairs.
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