Chapter 2: The Chemical Context of Life

Introduction

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

Concept 2.1: Matter and Chemical Elements

Definition of Matter and Elements

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

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

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

Example: Water (H2O) is a compound made of hydrogen and oxygen in a 2:1 ratio.

Essential and Trace Elements in Life

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

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

• Trace Elements: Required in minute quantities (<0.01%), e.g., Iron (Fe), Iodine (I)

Concept 2.2: Atomic Structure and Properties

Structure of Atoms

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

• Subatomic Particles:

  • Neutrons: No electrical charge; contribute to isotopes.

  • Protons: Positive charge; determine the element's identity.

  • Electrons: Negative charge; involved in chemical bonding.

• Atoms are electrically neutral overall (number of protons = number of electrons).

Atomic Number and Atomic Mass

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

• Atomic Mass: Sum of protons and neutrons in the nucleus (measured in daltons).

• Electrons have negligible mass compared to protons and neutrons.

Example: Carbon-12 has 6 protons and 6 neutrons; atomic mass = 12.

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.

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

Equation for Remaining Mass after n Half-lives:

Applications: Radiometric dating, medical diagnostics.

Concept 2.3: Chemical Bonds and Molecular Interactions

Chemical Bonds

• Covalent Bonds: Sharing of electron pairs between atoms.

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

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

• Ionic Bonds: Transfer of electrons from one atom to another, resulting in charged ions (cations and anions) that attract each other.

• Hydrogen Bonds: Weak 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 charges when electrons are distributed unevenly.

Intramolecular vs. Intermolecular Forces

• Intramolecular Forces: Forces within a molecule (e.g., covalent and ionic bonds).

• Intermolecular Forces: Forces between molecules (e.g., hydrogen bonds, van der Waals forces).

Polarity of Bonds

• Nonpolar Covalent Bond: Electrons are shared equally between atoms.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges.

Example: Water (H2O) has polar covalent bonds, leading to hydrogen bonding.

Properties of Ionic Compounds (Salts)

• Often form crystals in nature.

• Not considered molecules; formula indicates element ratio in the crystal.

• Stable when dry, but dissociate easily in water.

Concept 2.4: Electron Configuration and Chemical Reactivity

Electron Shells and Orbitals

• Electrons are arranged in shells around the nucleus; each shell has a specific energy level.

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

• First shell holds up to 2 electrons; subsequent shells hold up to 8 electrons.

• Orbitals: Three-dimensional spaces where electrons are found 90% of the time.

• Electron configuration is written as: (example for neon).

Hybridization

• Occurs when atomic orbitals mix to form new, hybrid orbitals during bonding.

• Determines the shape and function of molecules.

Example: Carbon forms sp3 hybrid orbitals in methane (CH4).

Concept 2.5: Chemical Reactions

Making and Breaking Chemical Bonds

• Chemical Reaction: The process of making and breaking chemical bonds.

• Reactants: Starting molecules in a reaction.

• Products: Resulting molecules after the reaction.

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

Example Equation:

Chemical Equilibrium: The relative concentrations of reactants and products remain constant.

Summary Table: Types of Chemical Bonds

Practice Problems

• Calculate the number of neutrons in an isotope given atomic number and mass number.

• Determine the remaining mass of a radioactive isotope after a given number of half-lives.

• Write electron configurations for elements.

Additional info: Some explanations and examples have been expanded for clarity and completeness based on standard biology curriculum.