Chapter 2: The Chemical Context of Life

Introduction

This chapter introduces the fundamental chemical principles that underlie biological processes. Understanding the structure of atoms, elements, and compounds is essential for studying how life functions at the molecular level.

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

Definitions and Key Concepts

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

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

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

Organisms are composed of matter, which is made up of elements. Compounds are combinations of elements bonded together in specific ratios, giving rise to unique properties.

Essential and Trace Elements in Life

• Essential Elements (make up ~96% of living matter): Carbon (C), Oxygen (O), Hydrogen (H), Nitrogen (N)

• Other Essential Elements (~4%): Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg)

• Trace Elements: Required in minute quantities (less than 0.01%)

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

Atomic Structure

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

• Subatomic Particles:

  • Neutrons: No electrical charge; contribute to atomic mass and isotopes.

  • Protons: Positive charge; determine the element’s identity (atomic number).

  • Electrons: Negative charge; involved in chemical bonding.

• Atoms are electrically neutral overall because they have equal numbers of protons and electrons.

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in the nucleus; unique to each element.

• Mass Number: Sum of protons and neutrons in the nucleus.

• Atomic Mass: Approximate total mass of an atom, measured in daltons.

• Electrons are so small they do not significantly contribute to atomic mass.

Isotopes and Radioactivity

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

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

• Radioactive isotopes are used in medicine (diagnostic tools) and radiometric dating.

Half-life: The time required for half the atoms of a radioactive isotope to decay.

Example: Calculating Half-Lives

• If 60 grams of Np-240 (half-life = 1 hour) are present, after 4 hours (4 half-lives), the remaining mass is:

  • Number of half-lives:

  • Remaining mass: grams

Electron Behavior and Energy Levels

Electron Shells and Orbitals

• Energy: The capacity to cause change; electrons have potential energy based on their position relative to the nucleus.

• Electron Shells: Electrons are arranged in shells around the nucleus; the further from the nucleus, the higher the energy.

• Electrons can move between shells by absorbing or releasing energy in fixed amounts.

• Valence Electrons: Electrons in the outermost shell; determine chemical behavior.

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

Electron Orbitals

• Orbital: A three-dimensional space where an electron is found 90% of the time.

• Each shell contains a specific number and type of orbitals:

  • s orbitals: hold up to 2 electrons

  • p orbitals: hold up to 6 electrons

  • d orbitals: hold up to 10 electrons

  • f orbitals: hold up to 14 electrons

• Electron configuration is written with the number of electrons as a superscript (e.g., for carbon).

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

Chemical Bonds

• Chemical Bond: An attraction between two atoms resulting from sharing or transferring valence electrons.

• Types of chemical bonds:

  • Covalent Bonds (strong)

  • Ionic Bonds (strong in dry conditions)

  • Hydrogen Bonds (weak)

  • Van der Waals Interactions (weak)

Covalent Bonds

• Involve sharing pairs of valence electrons between atoms.

• Single Covalent Bond: Sharing of one pair of electrons (e.g., H—H).

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

• Electronegativity: The attraction of an atom for the electrons in a covalent bond.

• Nonpolar Covalent Bond: Electrons are shared equally.

• Polar Covalent Bond: Electrons are shared unequally, leading to partial charges.

Ionic Bonds

• Formed when one atom strips electrons from another, creating ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Oppositely charged ions attract to form ionic bonds (e.g., Na+ and Cl- form NaCl).

Properties of Ionic Compounds (Salts)

• Often form crystals in nature.

• Not considered molecules; formula indicates ratio of elements.

• Stable when dry, but dissociate easily in water.

Hydrogen Bonds

• Form when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (usually oxygen or nitrogen).

• Important in the structure of water and biological molecules like DNA.

Van der Waals Interactions

• Weak attractions due to transient local partial charges as electrons move around atoms.

• Collectively, these interactions can be significant (e.g., gecko’s toe hairs adhering to surfaces).

Molecular Shape and Hybridization

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

• Hybridization: The mixing of atomic orbitals to form new hybrid orbitals, influencing molecular geometry.

• Molecular shape is crucial for function (e.g., enzyme-substrate interactions).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

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

• Reactions are reversible; products of the forward reaction can become reactants in the reverse reaction.

• Chemical Equilibrium: The point at which forward and reverse reactions occur at the same rate; concentrations of reactants and products remain constant.

Example: Water Formation

• Hydrogen and oxygen gases react to form water.

Summary Table: Types of Chemical Bonds

Key Terms

• Atom, Element, Compound, Isotope, Ion, Cation, Anion, Covalent Bond, Ionic Bond, Hydrogen Bond, Van der Waals Interaction, Electronegativity, Valence Electron, Hybridization, Chemical Reaction, Chemical Equilibrium

Example Application: The properties of water, such as cohesion, adhesion, and high specific heat, arise from hydrogen bonding between water molecules.

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