Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes. Understanding the nature of matter, elements, atoms, and chemical bonds is essential for studying life at the molecular level.

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

Definition and Properties of Matter

• Matter is anything that takes up space and has mass.

• All organisms are composed of matter.

Elements and Compounds

• An element is a substance that cannot be broken down into other substances by chemical reactions.

• A compound is a substance consisting of two or more elements in a fixed ratio.

• Compounds have emergent properties that are different from those of their constituent elements.

Example: Sodium (Na, a metal) and chlorine (Cl, a poisonous gas) combine to form sodium chloride (NaCl, table salt), which is edible and safe.

Elements Essential to Life

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

• Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) make up approximately 96% of living matter.

• The remaining 4% is mostly calcium (Ca), phosphorus (P), potassium (K), and sulfur (S).

• Trace elements are required in minute quantities (e.g., iron, iodine).

Table: Major Elements in the Human Body

Concept 2.2: An Element's Properties Depend on the Structure of Its Atoms

Atomic Structure

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

• Atoms are composed of subatomic particles:

  • Protons (positive charge, found in the nucleus)

  • Neutrons (no charge, found in the nucleus)

  • Electrons (negative charge, orbit the nucleus in electron shells)

• The number of protons determines the element's identity (atomic number).

• The sum of protons and neutrons is the mass number.

• Atomic mass is approximately equal to the mass number (measured in daltons).

Isotopes and Radioactivity

• Isotopes are atoms of the same element with different numbers of neutrons.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

• Radioactive isotopes are used as tracers in medicine and for radiometric dating of fossils.

Example: PET scans use radioactive tracers to monitor cancer growth and metabolism.

Energy Levels of Electrons

• Energy is the capacity to cause change; potential energy is energy due to position or structure.

• Electrons have potential energy based on their distance from the nucleus.

• Electrons occupy electron shells with characteristic energy levels.

• Only certain energy levels are allowed; electrons can move between shells by absorbing or releasing energy.

Electron Distribution and Chemical Properties

• The chemical behavior of an atom is determined by the distribution of electrons, especially those in the valence shell (outermost shell).

• Elements with a full valence shell are chemically inert (e.g., noble gases).

Electron Orbitals

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

• Each electron shell contains a specific number of orbitals; each orbital holds up to 2 electrons.

Concept 2.3: The Formation and Function of Molecules and Ionic Compounds Depend on Chemical Bonding Between Atoms

Chemical Bonds

• Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds.

• The main types of chemical bonds are covalent bonds and ionic bonds.

Covalent Bonds

• A covalent bond is the sharing of a pair of valence electrons between two atoms.

• A single bond involves one pair of shared electrons; a double bond involves two pairs.

• Electronegativity is an atom's attraction for shared electrons.

• Nonpolar covalent bonds: electrons are shared equally.

• Polar covalent bonds: electrons are shared unequally, leading to partial charges.

Example: In water (H2O), oxygen is more electronegative than hydrogen, resulting in a polar molecule.

Ionic Bonds

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

• Cation: positively charged ion; anion: negatively charged ion.

• The attraction between cations and anions forms an ionic bond.

• Ionic compounds (salts) are often found as crystals (e.g., NaCl).

Weak Chemical Interactions

• Weak bonds, such as hydrogen bonds and van der Waals interactions, are important in biological molecules.

• Hydrogen bonds form when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (often O or N).

• Van der Waals interactions occur due to transient local partial charges.

Example: Hydrogen bonds stabilize the structure of DNA and proteins.

Molecular Shape and Function

• The shape of a molecule is determined by the positions of its atoms' orbitals.

• Molecular shape is crucial for biological function, such as enzyme-substrate interactions.

Example: Morphine mimics endorphins by fitting into the same brain receptors due to similar shapes.

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Chemical reactions involve the making and breaking of chemical bonds.

• Reactants are the starting materials; products are the resulting substances.

• All chemical reactions are reversible; products can become reactants in the reverse reaction.

• Chemical equilibrium is reached when the forward and reverse reactions occur at the same rate.

Examples of Chemical Reactions

• Formation of water:

• Photosynthesis:

• Reversible reaction (ammonia synthesis):

Summary Table: Types of Chemical Bonds

Additional info: Some explanations and examples have been expanded for clarity and completeness, following standard introductory biology textbooks.