Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the chemical foundations essential for understanding biological processes. It covers the nature of matter, elements, atoms, chemical bonds, and reactions, providing the groundwork for all biological phenomena.

Concept 2.1: Matter, Elements, and Compounds

Definitions and Relationships

• Matter: Anything that takes up space and has mass.

• Element: A substance that cannot be broken down into other substances by chemical reactions.

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

• Compound: A substance consisting of two or more different elements combined in a fixed ratio.

• Molecule: Two or more atoms held together by covalent bonds.

• Mixture: A combination of substances not held together by chemical bonds.

• Avogadro’s Number: The number of particles (atoms, molecules) in one mole, approximately .

Compounds have emergent properties that differ from those of their constituent elements. For example, sodium (a reactive metal) and chlorine (a poisonous gas) combine to form sodium chloride (table salt), which is edible and essential for life.

The Elements of Life

• About 20–25% of the 92 natural elements are essential for life.

• Major elements: Carbon, hydrogen, oxygen, and nitrogen make up about 96% of living matter.

• Other important elements: Calcium, phosphorus, potassium, and sulfur account for most of the remaining 4%.

• Trace elements: Required in minute quantities (e.g., iron, iodine).

Concept 2.2: Atomic Structure and Properties

Subatomic Particles

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles orbiting the nucleus in electron shells.

The number of protons determines the element’s identity, while the number of neutrons can vary, resulting in isotopes. Electrons determine chemical behavior.

Atomic Number, Mass Number, and Isotopes

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

• Mass number (A): Sum of protons and neutrons.

• Atomic mass: Approximate total mass of an atom (in daltons).

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Radioactive isotopes: Unstable isotopes that decay, emitting particles and energy.

Application: Radioactive isotopes are used in medical imaging (e.g., PET scans) and radiometric dating of fossils.

Electron Energy Levels and Shells

• Electrons occupy energy levels (shells) around the nucleus.

• Energy is absorbed or lost as electrons move between shells.

• Valence electrons (in the outermost shell) determine chemical reactivity.

Electron Orbitals

• An orbital is a three-dimensional space where an electron is found 90% of the time.

• Each shell contains a specific number of orbitals (e.g., s, p).

• No more than two electrons can occupy a single orbital.

Concept 2.3: Chemical Bonds and Molecular Structure

Covalent Bonds

• Covalent bond: Sharing of a pair of valence electrons between two atoms.

• Single bond: Sharing of one pair of electrons (e.g., H–H).

• Double bond: Sharing of two pairs of electrons (e.g., O=O).

• Electronegativity: Atom’s attraction for shared electrons in a bond.

• Nonpolar covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally, creating partial charges.

Ionic Bonds

• Formed when one atom transfers electrons to another, creating ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic bond: Attraction between cations and anions.

• Ionic compounds (salts): Compounds formed by ionic bonds, often crystalline in structure.

Weak Chemical Interactions

• Hydrogen bonds: Attraction between a hydrogen atom covalently bonded to an electronegative atom (usually O or N) and another electronegative atom.

• Van der Waals interactions: Weak attractions due to transient local partial charges.

• Weak bonds are crucial for the structure and function of large biological molecules (e.g., DNA, proteins).

Molecular Shape and Function

• Molecular shape is determined by the positions of atoms’ orbitals and is critical for biological function.

• Shape determines how molecules interact and recognize each other (e.g., hormone-receptor binding).

• Hybridization of orbitals can create specific molecular geometries (e.g., tetrahedral in methane).

Concept 2.4: Chemical Reactions

Making and Breaking Bonds

• Chemical reaction: The making and breaking of chemical bonds, transforming reactants into products.

• All chemical reactions are reversible; equilibrium is reached when forward and reverse reactions occur at the same rate.

• Example: Formation of water:

Photosynthesis: A Key Biological Reaction

• Photosynthesis converts carbon dioxide and water into glucose and oxygen using sunlight.

• Equation:

Chemical Equilibrium

• At equilibrium, the concentrations of reactants and products remain constant.

• Indicated by double arrows () in chemical equations.

Summary Table: Atom Components and Their Roles

Key Vocabulary

• Matter: Anything that takes up space and has mass.

• Element: A substance that cannot be broken down by chemical reactions.

• Atom: The smallest unit of matter retaining element properties.

• Compound: Substance of two or more different elements in a fixed ratio.

• Molecule: Group of two or more atoms held together by attractive forces.

• Avogadro’s number: Number of particles in one mole ().