Atoms, Molecules, and Water: Foundations of Life

Atoms, Molecules, and Water

Introduction

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

Atoms: The Building Blocks of Matter

Definition and Structure of Atoms

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

• Element: A pure substance made of only one kind of atom; cannot be broken down by ordinary chemical reactions.

• Atoms are composed of subatomic particles:

  • Protons: Positively charged particles found in the nucleus.

  • Neutrons: Uncharged particles found in the nucleus.

  • Electrons: Negatively charged particles orbiting the nucleus in electron shells.

• Atomic Number: Number of protons in the nucleus; defines the element.

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

Isotopes and Radioactivity

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Some isotopes are radioactive, meaning they spontaneously break apart, releasing energy or subatomic particles.

• Example: Carbon has three isotopes: C-12, C-13, and C-14 (C-14 is radioactive).

• Radioactive isotopes are used in medical technology, such as detecting brain tumors.

Periodic Table Organization

• Elements are organized by atomic number and chemical properties.

• Horizontal rows are called periods; vertical columns are groups.

• Noble gases (e.g., helium, argon) are chemically inert due to full outer electron shells.

Subatomic Particles: Mass and Charge

Subatomic Particle	Mass (atomic mass units)	Charge
Proton (p)	1	+1
Neutron (n)	1	0
Electron (e)	~0	-1

Electron Shells and Chemical Behavior

Electron Shells

• Electrons occupy shells (energy levels) around the nucleus.

• First shell: up to 2 electrons; second: up to 8; third: up to 16.

• Outer shell electrons (valence electrons) determine chemical reactivity.

• Atoms with full outer shells are stable and unreactive.

Energy Capture and Release

• Electrons can absorb energy and move to higher-energy shells.

• When electrons fall back to lower shells, they release energy as heat or light.

Formation of Molecules

Chemical Bonds

• Chemical bonds hold atoms together in molecules.

• Bonds form when atoms gain, lose, or share electrons.

• Three major types of chemical bonds:

  1. Ionic bonds

  2. Covalent bonds (polar and nonpolar)

  3. Hydrogen bonds

Ionic Bonds

• Formed between ions (atoms that have gained or lost electrons).

• Oppositely charged ions attract each other.

• Example: Sodium (Na+) and Chloride (Cl-) form NaCl (table salt).

Covalent Bonds

• Formed by sharing electrons between atoms.

• Can be nonpolar (equal sharing) or polar (unequal sharing).

• Nonpolar covalent bonds: Occur between atoms of the same element (e.g., H2, O2, N2).

• Polar covalent bonds: Occur when one atom attracts electrons more strongly (e.g., H2O).

Hydrogen Bonds

• Attractive forces between slightly positive and slightly negative poles of adjacent polar molecules.

• Common in water, giving it unique properties.

• Hydrogen bonds also occur in molecules containing nitrogen or fluorine.

Summary Table: Types of Bonds

Type of Bond	How It Forms	Example
Ionic	Transfer of electrons; attraction between oppositely charged ions	NaCl (table salt)
Nonpolar Covalent	Equal sharing of electrons	H2, O2, N2
Polar Covalent	Unequal sharing of electrons	H2O (water)
Hydrogen	Attraction between polar molecules	Between water molecules

Water: Essential for Life

Cohesion and Adhesion

• Cohesion: Tendency of water molecules to stick together due to hydrogen bonding.

• Surface tension: Resistance of water surface to being broken.

• Adhesion: Tendency of water molecules to stick to other substances, especially those with charged regions.

• Capillary action: Movement of water into narrow spaces, important for water transport in plants.

Water as a Solvent

• Solvent: Substance that dissolves other substances (solutes).

• Water dissolves ionic and polar molecules (hydrophilic), but not nonpolar molecules (hydrophobic).

• Examples: Table sugar (sucrose), NaCl dissolve in water; fats and oils do not.

Temperature Moderation

• Water has a high specific heat: requires a lot of energy to change temperature.

• High heat of vaporization: requires much energy to evaporate.

• Helps organisms maintain stable internal temperatures.

Density of Ice

• Ice is less dense than liquid water; it floats.

• Allows aquatic life to survive under ice in cold climates.

Acids, Bases, and pH

• Water can dissociate into hydrogen ions (H+) and hydroxide ions (OH-):

$\mathrm{H_2O \rightarrow OH^- + H^+}$

• Acid: Releases H+ in solution; pH < 7.

• Base: Releases OH- or accepts H+; pH > 7.

• Neutral: Equal concentrations of H+ and OH-; pH = 7.

pH Scale

pH Value	Type of Solution
0-6	Acidic (H+ > OH-)
7	Neutral (H+ = OH-)
8-14	Basic (OH- > H+)

Buffers

• Buffer: Substance that helps maintain a stable pH by accepting or releasing H+.

• Bicarbonate (HCO3-): Important buffer in human blood; can accept or release H+ as needed.

$\mathrm{HCO_3^- + H^+ \leftrightarrow H_2CO_3}$

Summary

• Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Elements are defined by atomic number; isotopes differ in neutron number.

• Chemical bonds (ionic, covalent, hydrogen) form molecules essential for life.

• Water's unique properties—cohesion, adhesion, solvent ability, temperature moderation, and density—are vital for biological systems.

• Acids, bases, and buffers regulate pH, crucial for cellular function.