CHAPTER 2: Chemical Context of Life

Introduction

The study of biology requires an understanding of the chemical principles that govern living systems. This chapter introduces the basic chemical concepts essential for understanding life, including the nature of matter, elements, atoms, and the types of chemical bonds that form the molecules of life.

Elements and Compounds

Definition of Matter, Elements, and Compounds

• Matter: Anything that takes up space and has mass. All organisms are composed of matter.

• Element: A substance that cannot be broken down into other substances by chemical reactions. Each element is made of unique atoms.

• Compound: A substance consisting of two or more elements combined in a fixed ratio. Compounds have characteristics different from those of their constituent elements.

Example: Sodium (Na) and chlorine (Cl) are elements; when combined, they form sodium chloride (NaCl), a compound with properties distinct from either element.

Elements Essential to Life

Major, Minor, and Trace Elements

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

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

• Minor elements: Most of the remaining 4% consists of calcium (Ca), phosphorus (P), potassium (K), and sulfur (S).

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

Additional info: Trace elements include iron (Fe), iodine (I), and zinc (Zn), which are vital for specific biological functions.

Atomic Structure

Subatomic Particles

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

• Atoms are composed of three main subatomic particles:

  • Neutrons: No electrical charge.

  • Protons: Positive charge; determines the atom’s identity (atomic number).

  • Electrons: Negative charge; found in a cloud around the nucleus.

• Neutrons and protons form the atomic nucleus.

• Electrons occupy regions of space called electron shells around the nucleus.

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: The atom’s total mass, approximately equal to the mass number (measured in daltons).

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

Isotopes

• Atoms of the same element with different numbers of neutrons are called isotopes.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

• Applications:

  • Dating fossils (e.g., carbon-14)

  • Tracing atoms through metabolic processes

  • Diagnosing medical disorders

Example: Hydrogen has three isotopes: protium (1 proton), deuterium (1 proton, 1 neutron), and tritium (1 proton, 2 neutrons).

Electron Energy Levels and Chemical Behavior

Electron Shells and Energy

• Potential energy: Energy that matter possesses because of its location or structure.

• Electrons have potential energy due to their distance from the nucleus; electrons in outer shells have more energy.

• Electrons occupy specific energy levels or electron shells.

Valence Electrons and Reactivity

• Valence electrons: Electrons in the outermost shell; determine the chemical behavior of an atom.

• Atoms with full valence shells are chemically inert (e.g., noble gases: helium, neon, argon).

• Atoms with incomplete valence shells are reactive and tend to form chemical bonds.

Chemical Bonds

Covalent Bonds

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

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

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

• Structural formula: Shows the arrangement of atoms (e.g., H—H, O=O).

• Molecular formula: Indicates the number and type of atoms (e.g., H2, O2).

Example: Water (H2O) is a molecule formed by covalent bonds between hydrogen and oxygen atoms.

Electronegativity and Polarity

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

• Nonpolar covalent bond: Electrons are shared equally (e.g., O2, CH4).

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

Example: In water, oxygen is more electronegative than hydrogen, creating a polar molecule with partial positive and negative charges.

Ionic Bonds

• Formed when one atom transfers an electron to another, resulting in oppositely charged ions.

• Cation: Positively charged ion; Anion: Negatively charged ion.

• Ionic bond: Attraction between a cation and an anion.

• Ionic compounds (salts) often form crystalline structures (e.g., NaCl).

Weak Chemical Bonds

• Weak bonds, such as hydrogen bonds and van der Waals interactions, are crucial for the structure and function of large biological molecules.

• Hydrogen bond: Attraction between a hydrogen atom covalently bonded to an electronegative atom (usually oxygen or nitrogen) and another electronegative atom.

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

Example: Hydrogen bonds hold water molecules together and contribute to the properties of DNA and proteins.

Chemical Reactions

Making and Breaking Bonds

• Chemical reaction: The making and breaking of chemical bonds, leading to changes in the composition of matter.

• Reactants: Starting materials; Products: Resulting materials.

Example: Photosynthesis:

Summary Table: Types of Chemical Bonds

Additional info: The chemical context of life is foundational for understanding biological molecules and processes, as all living things are governed by the principles of chemistry.