Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes. Understanding the chemical context of life is essential for grasping how living organisms function at the molecular level.

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

Definition of Matter

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

• All organisms are composed of matter.

Elements and Compounds

• 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.

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

• Example: Sodium (Na) and chlorine (Cl) are both dangerous in pure form, but together they form sodium chloride (NaCl), or table salt, which is safe to eat.

Elements of Life

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

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

• Other important elements: Calcium (Ca), phosphorus (P), potassium (K), and sulfur (S) make up most of the remaining 4%.

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

Table: Elements in the Human Body

Adaptation to Toxic Elements

• Some elements can be toxic to organisms.

• Certain species can adapt to environments containing toxic elements (e.g., plants in serpentine soils).

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.

• Atoms are composed of subatomic particles:

  • Protons: Positive charge

  • Neutrons: No charge

  • Electrons: Negative charge

• Protons and neutrons form the atomic nucleus; electrons form a cloud around the nucleus.

• Proton and neutron mass are nearly identical and measured in daltons.

• Electrons are much lighter and usually ignored in atomic mass calculations.

Atomic Number and Atomic Mass

• Atomic number: Number of protons in the nucleus.

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

• Atomic mass: Total mass of the atom, approximately equal to the mass number.

Isotopes

• Atoms of the same element may have different numbers of neutrons, called isotopes.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

Applications of Isotopes

• Radioactive tracers: Used in medicine to track atoms through metabolism (e.g., PET scans).

• Radiometric dating: Measures the ratio of different isotopes to determine the age of fossils and rocks.

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

Concept 2.2 (continued): Energy Levels of Electrons

Potential Energy and Electron Shells

• Energy: The capacity to cause change.

• Potential energy: Energy that matter possesses due to its location or structure.

• Electrons have different amounts of potential energy depending on their distance from the nucleus.

• Electrons are arranged in electron shells, each with a characteristic energy level.

• Changes in electron energy occur in fixed amounts (quantized).

Electron Distribution and Chemical Properties

• The chemical behavior of an atom is determined by the distribution of electrons in its shells.

• The periodic table arranges elements by electron configuration.

• Valence electrons: Electrons in the outermost shell (valence shell) are most important for chemical behavior.

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

Electron Orbitals

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

• Each electron shell consists of a specific number of orbitals.

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

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 valence electrons.

• These interactions result in chemical bonds that hold atoms together.

Covalent Bonds

• Covalent bond: The sharing of a pair of valence electrons by two 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).

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

• Valence: The bonding capacity of an atom.

• Covalent bonds can form between atoms of the same or different elements.

Electronegativity and Bond Polarity

• Electronegativity: An atom’s attraction for electrons in a covalent bond.

• Nonpolar covalent bond: Electrons are shared equally.

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges.

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

Ionic Bonds

• Atoms may strip electrons from their bonding partners, forming ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic bond: Attraction between oppositely charged ions.

• Ionic compounds (salts): Compounds formed by ionic bonds, often found as crystals (e.g., NaCl).

Weak Chemical Interactions

• Many biological molecules are held together by weak bonds, which allow for flexibility and reversibility.

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

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

Molecular Shape and Function

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

• Molecular shape is crucial for biological recognition and response (e.g., morphine and endorphins have similar shapes and bind the same brain receptors).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Chemical reaction: The making and breaking of chemical bonds.

• Reactants: Starting molecules in a reaction.

• Products: Resulting molecules from a reaction.

• Example: Formation of water:

Photosynthesis

• Photosynthesis is a key chemical reaction in biology, powered by sunlight.

• Equation:

• Reactants: Carbon dioxide and water; Products: Glucose and oxygen.

Reversibility and Equilibrium

• Chemical reactions are often reversible.

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

• Example:

Summary of Key Concepts

• The properties of elements depend on atomic structure.

• Chemical bonding (covalent, ionic, and weak interactions) determines the formation and function of molecules and compounds.

• Chemical reactions rearrange matter and are fundamental to biological processes.

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