Atoms, Molecules, and Life: Foundations of Biological Chemistry

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Atoms, Molecules, and Life

Introduction

This chapter introduces the fundamental chemical principles that underlie all biological processes. Understanding atoms, molecules, and their interactions is essential for studying the structure and function of living organisms.

Atoms: The Basic Units of Matter

Definition and Structure

Atom: The smallest unit of an element that retains the chemical properties of that element.
Element: A substance that cannot be broken down into simpler substances by ordinary chemical reactions.
Each element is defined by its atomic number, which is the number of protons in its nucleus.
Atoms are composed of subatomic particles:
- Protons: Positively charged particles found in the nucleus.
- Neutrons: Neutral particles found in the nucleus.
- Electrons: Negatively charged particles that orbit the nucleus in electron shells.
Most atoms are electrically neutral, with equal numbers of protons and electrons.

Atomic Number and Mass Number

Atomic number (Z): Number of protons in the nucleus.
Mass number (A): Total number of protons and neutrons in the nucleus.
Electrons have negligible mass compared to protons and neutrons.

Isotopes and Radioactivity

Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different mass numbers.
Some isotopes are radioactive, meaning their nuclei are unstable and break apart, releasing energy (radioactive decay).
Radioactive isotopes have important applications in medicine and research (e.g., cancer treatment, medical imaging).

Electron Shells and Chemical Properties

Electron Arrangement

Electrons occupy electron shells (energy levels) around the nucleus.
The shell closest to the nucleus can hold up to 2 electrons; subsequent shells can hold up to 8 electrons (octet rule).
Atoms are most stable when their outermost electron shell is full.
Atoms with incomplete outer shells tend to react with other atoms to achieve stability.

Valence Electrons and Reactivity

Valence electrons: Electrons in the outermost shell, responsible for chemical bonding and reactivity.
Atoms with full outer shells are chemically inert (e.g., noble gases).
Atoms with partially filled outer shells are reactive and seek to gain, lose, or share electrons.

Chemical Bonds and Molecules

Types of Chemical Bonds

Ionic Bonds: Formed when electrons are transferred from one atom to another, creating oppositely charged ions that attract each other.
- Example: Sodium chloride (NaCl)
Covalent Bonds: Formed when two atoms share one or more pairs of electrons.
- Can be nonpolar (equal sharing, e.g., O2, N2) or polar (unequal sharing, e.g., H2O).
Hydrogen Bonds: Weak attractions between the slightly positive hydrogen atom of one polar molecule and the slightly negative atom (often oxygen or nitrogen) of another polar molecule.
- Important in water, DNA, and proteins.

Table: Types of Chemical Bonds

Type of Bond	Description	Example
Ionic	Electrons transferred; attraction between oppositely charged ions	NaCl (table salt)
Nonpolar Covalent	Electrons shared equally	O2, N2
Polar Covalent	Electrons shared unequally	H2O
Hydrogen Bond	Attraction between polar molecules (H with O or N)	Between water molecules

Water: The Molecule of Life

Cohesion and Adhesion

Cohesion: Tendency of water molecules to stick together due to hydrogen bonding; responsible for surface tension.
Adhesion: Tendency of water molecules to stick to other polar or charged substances; enables capillary action.
These properties are essential for water transport in plants.

Water as a Solvent

Solvent: A substance that dissolves other substances (solutes) to form a solution.
Water is an excellent solvent for polar molecules and ions (hydrophilic substances).
Nonpolar molecules (hydrophobic), such as fats and oils, do not dissolve in water.

Thermal Properties of Water

Water has a high specific heat, meaning it can absorb or release large amounts of heat with little temperature change.
High heat of vaporization: Requires a lot of energy to convert water from liquid to gas, which helps regulate temperature in organisms and environments.
Evaporation of sweat cools the body by absorbing heat.

Density and Ice Formation

Unlike most substances, water is less dense as a solid (ice) than as a liquid.
Hydrogen bonds cause water molecules to spread apart when frozen, allowing ice to float.
This insulates aquatic environments, allowing life to persist under ice in winter.

Acids, Bases, and pH

Ionization of Water

Water molecules can dissociate into hydrogen ions (H+) and hydroxide ions (OH-):
Acid: Substance that increases the concentration of H+ in solution.
Base: Substance that increases the concentration of OH- or decreases H+ in solution.

pH Scale

pH: A measure of the hydrogen ion concentration in a solution.
pH 7 is neutral (H+ = OH-).
pH < 7 is acidic (H+ > OH-); pH > 7 is basic (OH- > H+).
Common acids: lemon juice, vinegar. Common bases: sodium hydroxide, antacids.

Buffers

Buffer: A molecule that helps maintain a stable pH in a solution by accepting or releasing H+ as needed.
Buffers are essential for maintaining homeostasis in biological systems.

Case Study: Unstable Atoms and Radioactivity

Radioactive Isotopes in the Environment

Natural disasters, such as earthquakes, can damage nuclear power plants, releasing radioactive isotopes into the environment.
Radioactive elements differ from stable elements by having unstable nuclei that emit radiation.
These isotopes can contaminate air, water, and living organisms, posing health risks.

Summary Table: Properties of Water

Property	Biological Importance
Cohesion	Enables water transport in plants; surface tension
Adhesion	Capillary action; water movement in soil and tissues
High Specific Heat	Stabilizes temperature in organisms and environments
High Heat of Vaporization	Evaporative cooling (sweating)
Lower Density of Ice	Ice floats, insulating aquatic life in winter
Solvent Properties	Dissolves ions and polar molecules; medium for biochemical reactions