Chapter 2: The Chemical Context of Life

Introduction

Understanding biology at the molecular level requires knowledge of chemistry and physics. This chapter explores the structure of atoms, properties of matter, types of chemical bonds, and their real-world applications in biological systems.

Matter and Its Forms

Definition and Properties of Matter

• Matter is anything that occupies space and has mass.

• Matter exists in different phases: solid, liquid, and gas. For example, water can exist as ice, liquid water, or water vapor.

• The transfer of matter and energy is fundamental to biological processes, from cellular structures to ecosystems.

Atoms and Molecules

• An atom is the smallest unit of matter that retains the properties of an element.

• A molecule consists of two or more atoms connected by chemical bonds.

Elements and Compounds

• Elements are pure substances consisting of only one type of atom and cannot be broken down by chemical reactions.

• Compounds are substances composed of two or more elements in a fixed ratio, exhibiting emergent properties distinct from their constituent elements.

Pure Substances vs. Mixtures

• A compound has elements in fixed ratios (e.g., NaCl, H2O, C6H12O6).

• A mixture contains elements or compounds not in fixed ratios (e.g., salt water is uniform; water and oil is nonuniform).

Emergent Properties of Compounds

• Compounds often have properties very different from the elements that compose them. For example, sodium (a reactive metal) and chlorine (a poisonous gas) form sodium chloride (table salt), which is safe to eat.

Salts and Their Biological Importance

• Salts are compounds formed from the ionic assembly of a cation (+) and an anion (−).

• Example: Magnesium sulfate (Epsom salt, MgSO4) is used as a plant supplement to improve chlorophyll production.

Toxic Elements

• The toxicity of an element is relative; for example, oxygen is toxic to anaerobic organisms.

• Some plants adapt to toxic soils and can be used for environmental cleanup (phytoremediation).

Atomic Structure and Properties

Subatomic Particles

• Proton: Positive charge, mass ≈ 1 Dalton.

• Neutron: No charge, mass ≈ 1 Dalton.

• Electron: Negative charge, negligible mass compared to protons and neutrons.

Atomic Number and Mass

• Atomic number: Number of protons in an atom; defines the element.

• Atomic mass: Total number of protons and neutrons in the nucleus.

Electron Orbitals and Valence Electrons

• Electrons occupy orbitals around the nucleus; the number of orbitals corresponds to the row in the periodic table.

• Valence electrons (in the outermost orbital) determine an atom’s chemical behavior.

Isotopes and Ions

• Isotopes: Atoms of the same element with different numbers of neutrons (thus different masses).

• Ions: Atoms that have gained or lost electrons, resulting in a net charge (cation: positive, anion: negative).

Types of Isotopes

• Stable isotopes: Nuclei do not lose subatomic particles.

• Radioactive isotopes: Nuclei decay spontaneously, releasing energy and particles.

Types of Radioactive Decay

• Alpha decay: Loss of 2 protons and 2 neutrons; stopped by paper.

• Beta decay: Conversion of a neutron to a proton and emission of an electron; stopped by wood.

• Gamma decay: Release of energy only; stopped by lead.

Applications of Radioactive Isotopes

• Radioactive tracers are used in medicine to trace biochemical pathways and visualize tissues/organs (e.g., CT and PET scans).

Radiometric Dating

• Radiometric dating measures the ratio of isotopes to determine the age of fossils or rocks.

• Half-life: The time required for 50% of a radioactive isotope to decay.

• Carbon-14 dating is used for dating biological materials; uranium isotopes are used for geological dating.

Chemical Bonds and Molecular Structure

Chemical Bonds

• Chemical bonds hold atoms together in molecules.

• Main types: Covalent bonds (sharing electrons), Ionic bonds (transfer of electrons), Hydrogen bonds, and Van der Waals interactions.

Ionic Bonds

• Formed when one atom donates an electron to another, creating oppositely charged ions that attract each other.

• Example: Sodium chloride (NaCl) forms from Na+ and Cl−.

Covalent Bonds

• Formed when two atoms share one or more pairs of valence electrons.

• Can be single, double, or triple bonds depending on the number of shared electron pairs.

• Can be nonpolar (equal sharing) or polar (unequal sharing).

Hydrogen Bonds

• Weak bonds formed between a hydrogen atom covalently bonded to an electronegative atom (like oxygen or nitrogen) and another electronegative atom.

• Important for stabilizing the structures of DNA and proteins, and for water’s unique properties.

Van der Waals Interactions

• Weak, transient attractions between molecules due to temporary shifts in electron density.

• Individually weak but collectively significant in stabilizing molecular structures.

Energy in Chemical Bonds

• Chemical bonds store energy; breaking bonds releases energy used in biological reactions.

• The stronger the bond, the more energy is released when it is broken.

Structure and Function of Molecules

• The three-dimensional structure of a molecule determines its function in biological systems.

• Example: Adenosine causes drowsiness; caffeine blocks adenosine receptors, promoting alertness.

Chemical Reactions and Equilibrium

Chemical Reactions

• Reactants: Starting substances in a reaction.

• Products: Substances formed as a result of the reaction.

• Activation energy: The minimum energy required to initiate a reaction.

Chemical Equilibrium

• At equilibrium, the forward and reverse reactions occur at the same rate, so the concentrations of reactants and products remain constant.

Types of Chemical Reactions

• Endothermic reactions: Absorb energy (e.g., sweating, cold packs).

• Exothermic reactions: Release energy (e.g., burning wood, hot packs).

Summary Table: Key Terms and Concepts