The Nature of Molecules and the Properties of Water

Introduction

This chapter explores the fundamental chemical principles underlying biological systems, focusing on the structure of atoms, types of chemical bonds, and the unique properties of water that make life possible.

All Matter is Composed of Atoms

Definition and Properties of Matter

• Matter is anything that has mass and occupies space.

• All matter is made from atoms or parts of atoms.

• Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• The three phases of matter are solids, liquids, and gases.

States of Matter

• Solid: Fixed shape and volume.

• Liquid: Fixed volume, no fixed shape.

• Gas: No fixed shape or volume.

Mass vs. Weight

• Mass is the amount of matter in an object; it remains constant regardless of location.

• Weight is the gravitational force acting on an object; it can change depending on location (e.g., on Earth vs. the Moon).

Atomic Structure

Subatomic Particles

• Protons: Positively charged particles located in the nucleus.

• Neutrons: Neutral particles located in the nucleus.

• Electrons: Negatively charged particles found in orbitals surrounding the nucleus.

Atomic Number and Atomic Mass

• Atomic number = number of protons in the nucleus.

• Atomic mass = sum of protons and neutrons.

• Atoms are electrically neutral when the number of protons equals the number of electrons.

Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.

• Some isotopes are radioactive and decay over time, emitting radiation.

• Half-life is the time required for half of the atoms in a sample to decay.

Ions and Electron Arrangement

Ions

• Cation: Atom with more protons than electrons (net positive charge).

• Anion: Atom with more electrons than protons (net negative charge).

Electron Arrangement and Chemical Behavior

• The chemical behavior of an atom depends on the number and arrangement of its electrons.

• Orbitals are areas around the nucleus where electrons are likely to be found.

• Electrons have potential energy related to their location; those further from the nucleus have higher energy.

Redox Reactions

Oxidation and Reduction

• Oxidation: Loss of electrons, resulting in a cation.

• Reduction: Gain of electrons, resulting in an anion.

• Redox reactions involve the transfer of electrons between atoms.

Elements and the Periodic Table

Elements in Biology

• The periodic table organizes elements by atomic number and electron configuration.

• 12 elements are found in living organisms in substantial amounts.

• Four elements—carbon (C), hydrogen (H), oxygen (O), and nitrogen (N)—make up 96.3% of human body weight.

• Trace elements are required in small amounts but are essential for life.

Valence Electrons and the Octet Rule

• Valence electrons are electrons in the outermost energy level.

• Atoms tend to establish completely full outer energy levels (octet rule).

Chemical Bonds

Types of Chemical Bonds

• Ionic bonds: Formed by the attraction of oppositely charged ions (transfer of electrons).

• Covalent bonds: Formed when atoms share one or more pairs of valence electrons.

• Hydrogen bonds: Weak attractions between a partially positive hydrogen atom and a partially negative atom (often oxygen or nitrogen).

• Van der Waals attractions: Weak, non-directional attractions between atoms in close proximity.

Ionic Bonds

• Formed when one atom loses an electron (becomes a cation) and another gains an electron (becomes an anion).

• Example: Sodium (Na) loses an electron to become Na+; Chlorine (Cl) gains an electron to become Cl-; together they form NaCl.

Covalent Bonds

• Atoms share electrons to satisfy the octet rule.

• No net charge is created; electrons are paired.

• The strength of a covalent bond depends on the number of shared electron pairs.

• Examples: , , ,

Electronegativity

• Electronegativity is an atom’s affinity for electrons.

• Differences in electronegativity determine how electrons are distributed in covalent bonds.

• Nonpolar covalent bonds: Equal sharing of electrons.

• Polar covalent bonds: Unequal sharing of electrons.

Hydrogen Bonds

• Formed when a hydrogen atom covalently bonded to a highly electronegative atom (like oxygen) is attracted to another electronegative atom.

• Individually weak and transitory, but collectively significant (e.g., in water).

Van der Waals Attractions

• Weak, non-specific attractions between molecules or parts of molecules that are close together.

• Example: Antibody-antigen binding.

Chemical Reactions

Formation and Breaking of Bonds

• Chemical reactions involve the formation or breaking of chemical bonds.

• Reactants are the original molecules; products are the resulting molecules.

• Many reactions are reversible.

• Influenced by temperature, concentration of reactants/products, and catalysts.

Properties of Water

Structure and Polarity

• Water molecules are polar due to the unequal sharing of electrons between oxygen and hydrogen.

• Oxygen is more electronegative, resulting in partial negative charge on oxygen and partial positive charge on hydrogen.

Hydrogen Bonding in Water

• Hydrogen bonds form between the partially positive hydrogen of one water molecule and the partially negative oxygen of another.

• Responsible for many of water’s unique properties.

Cohesion and Adhesion

• Cohesion: Water molecules stick to each other via hydrogen bonding (surface tension).

• Adhesion: Water molecules stick to other substances.

• Capillary action: Movement of water through small spaces due to cohesion, adhesion, and surface tension.

Thermal Properties of Water

• High specific heat: Large amount of energy required to change temperature.

• High heat of vaporization: Evaporation causes cooling.

• Ice is less dense than liquid water; bodies of water freeze from the top down.

Water as a Solvent

• Water dissolves polar molecules and ions (hydrophilic).

• Nonpolar molecules (hydrophobic) aggregate or assume specific shapes in water.

• Water can form ions:

Acids, Bases, and pH

Definitions

• Acid: Substance that increases when dissolved in water, lowering pH.

• Base: Substance that decreases when dissolved in water, raising pH.

• pH: Negative logarithm of hydrogen ion concentration:

The pH Scale

Buffers

• Buffer: Substance that resists changes in pH by releasing or absorbing hydrogen ions.

• Most biological buffers consist of a pair of molecules, one an acid and one a base.

• Example: The bicarbonate buffer system regulates blood pH.

• Equation:

Blood pH Regulation

• Blood pH is tightly regulated by buffers to maintain homeostasis.

Summary Table: Types of Chemical Bonds

Additional info:

• Some context and definitions were expanded for clarity and completeness.

• Tables were reconstructed and entries inferred where original was incomplete.