Atoms, Molecules, and Life

Introduction to Chemistry in Biology

Understanding chemistry is essential in biology because the interactions between atoms form the basis of all biological structures and processes. Atoms are the fundamental units of matter, and their properties and interactions underpin the organization of life.

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

• Element: A substance that cannot be broken down into other substances by chemical reactions; composed of one type of atom.

Subatomic Particles

Atoms are composed of three main subatomic particles, each with distinct properties and locations within the atom.

• Protons determine the atomic number and identity of an element.

• Neutrons contribute to atomic mass and can vary in number, forming isotopes.

• Electrons occupy electron shells and are involved in chemical bonding.

The Atomic Nucleus and Isotopes

• Atomic Number: Number of protons in an atom; unique for each element.

• Atomic Weight (Mass Number): Sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different atomic weights.

Electron Shells and Chemical Reactivity

Electrons are arranged in shells around the nucleus, with each shell corresponding to a specific energy level. The chemical behavior of an atom is largely determined by the number of electrons in its outermost shell (valence shell).

• Atoms with incomplete valence shells are reactive.

• Atoms with full valence shells are inert (nonreactive).

Chemical Bonds: Joining Atoms to Make Molecules

Types of Chemical Bonds

Atoms interact to fill their valence shells, forming chemical bonds that create molecules and compounds.

• Chemical Bond: An attraction that holds two atoms together.

• Molecule: Two or more atoms held together by chemical bonds; smallest unit of a compound.

• Compound: Substance composed of two or more elements in a fixed ratio (e.g., NaCl).

• Mixture: Combination of two or more elements and/or compounds in variable ratios (e.g., air).

Ions and Ionic Bonds

• Ion: A charged atom or molecule.

• Anion: Negatively charged ion.

• Cation: Positively charged ion.

• Ionic Bond: Attraction between oppositely charged ions after electron transfer.

Ionic compounds (salts) form crystals and are strong in solid form but dissociate easily in water.

Covalent Bonds

• Covalent Bond: Formed when two atoms share electrons.

• Single Covalent Bond: Shares one pair of electrons.

• Double Covalent Bond: Shares two pairs of electrons.

• Triple Covalent Bond: Shares three pairs of electrons.

Nonpolar vs. Polar Covalent Bonds

• Nonpolar Covalent Bond: Electrons are shared equally between atoms.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges on the molecule.

Polar molecules have regions of positive and negative charge.

Hydrogen Bonds

• Hydrogen Bond: Weak attraction between a hydrogen atom (covalently bonded to one atom) and another electronegative atom.

• Hydrogen bonds are much weaker than covalent bonds but are crucial for stabilizing biological molecules (e.g., DNA, proteins).

• They are reversible and can form between or within molecules.

Water and Life

Properties of Water

Water is essential for life, making up 70-95% of cells and covering most of Earth's surface. Its unique properties arise from its polar covalent bonds and ability to form hydrogen bonds.

Cohesion and Surface Tension

• Cohesion: Attraction between molecules of the same kind (e.g., water molecules stick together).

• Surface Tension: The force that allows the surface of a liquid to resist external force; water has high surface tension due to hydrogen bonding.

Solvent Properties

• Hydrophilic: Substances with an affinity for water (e.g., polar and ionic compounds) dissolve easily in water.

• Hydrophobic: Substances that do not interact with water (e.g., nonpolar molecules) tend to clump together in aqueous environments.

Temperature Moderation

• Specific Heat: Amount of energy required to change the temperature of 1 gram of a substance by 1°C.

• Water has a high specific heat (1 cal/g°C), allowing it to buffer temperature changes.

• Heat of Vaporization: Energy required to convert 1 gram of liquid to gas; water's high heat of vaporization enables evaporative cooling.

• Vaporization (Evaporation): Transformation from liquid to gas, cooling surfaces as water evaporates.

Expansion Upon Freezing

• Water expands when it freezes, making ice less dense than liquid water. This property insulates aquatic environments in cold climates.

Solutions, Concentration, and Acidity

Solute Concentration

• Solute Concentration: Amount of a substance in a given volume of solution.

• Molarity (M): Number of moles of solute per liter of solution.

• Mole: 6.022 × 1023 particles (Avogadro's number); a counting unit like "dozen" but for atoms/molecules.

• Molecular Weight: Mass of one mole of a molecule, calculated by summing the atomic masses of its atoms.

Acids, Bases, and the pH Scale

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

• In pure water:

• Acid: Substance that increases H+ concentration in solution.

• Base: Substance that decreases H+ concentration (often by increasing OH-).

The pH Scale

• pH Scale: Measures acidity from 0 (most acidic) to 14 (most basic); pH 7 is neutral.

• Relationship:

• Each pH unit represents a tenfold change in H+ concentration.

• Most biological fluids have pH between 6 and 8.

Buffers and pH Stability

• Buffer: Substance that minimizes changes in pH by accepting or donating H+ ions.

• Buffers are crucial in biological systems to maintain homeostasis (e.g., bicarbonate buffer in blood).

Example: Bicarbonate Buffer System

• Helps maintain blood pH by balancing H+ and HCO3- concentrations.