Chapter 2: The Chemistry of Microbiology – Study Notes

Chapter 2: The Chemistry of Microbiology

Objectives

• State the number of protons, neutrons, and electrons in an atom using the periodic table.

• List and describe the three types of chemical bonds.

• Read an equation for a chemical reaction and identify the type of reaction.

• Sketch and give examples of a carbohydrate, a protein, a lipid, and a nucleic acid.

Atoms and Elements

Basic Definitions

• Matter: Anything that takes up space and has mass.

• Atoms: The smallest chemical units of matter.

Atomic Structure

• Electrons: Negatively charged subatomic particles circling a nucleus.

• Nucleus: Structure containing neutrons and protons.

• Neutrons: Uncharged (neutral) particles.

• Protons: Positively charged particles.

• The number of protons and electrons in an atom is always equal, balancing the overall charge.

Elements

• Composed of a single type of atom (e.g., hydrogen, oxygen, sodium).

• Atomic number: Number of protons in the nucleus.

• Atomic mass (atomic weight): Sum of masses of protons and neutrons (electrons have negligible mass).

Examples

• Boron: Atomic number 5, atomic mass 11 → 6 neutrons ($11-5=6$).

• Iron: Atomic number 26, atomic mass 56 → 30 neutrons ($56-26=30$).

Common Elements of Life

Element	Symbol	Atomic Number	Atomic Mass*	Biological Significance
Hydrogen	H	1	1	Component of organic molecules and water; H+ released by acids
Carbon	C	6	12	Backbone of organic molecules
Nitrogen	N	7	14	Component of amino acids, proteins, and nucleic acids
Oxygen	O	8	16	Component of many organic molecules and water; O2- released by bases; necessary for aerobic metabolism
Sodium	Na	11	23	Principal cation outside cells
Phosphorus	P	15	31	Component of nucleic acids and ATP
Sulfur	S	16	32	Component of proteins

Additional info: Table includes other elements such as Mg, Cl, K, Ca, Fe, Zn, etc., each with specific biological roles.

Isotopes

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

• Stable and unstable (radioactive) isotopes exist.

• Radioactive isotopes release energy during radioactive decay.

Electron Configuration

• Only electrons interact during chemical reactions.

• Electrons occupy shells; valence electrons are in the outermost shell and determine chemical behavior.

Chemical Bonds

Valence

• Combining capacity of an atom, determined by the number of electrons in the outer shell.

• Atoms are stable when their outer shell contains eight electrons (octet rule).

Types of Chemical Bonds

• Covalent Bonds: Sharing of electron pairs between atoms.

• Nonpolar Covalent Bonds: Electrons shared equally; no charge separation (e.g., H2, O2).

• Polar Covalent Bonds: Unequal sharing of electrons; partial charges develop (e.g., H2O).

• Ionic Bonds: Transfer of electrons from one atom to another, forming charged ions (cations and anions) that attract each other (e.g., NaCl).

• Hydrogen Bonds: Weak electrical attraction between partially charged H+ and a full or partial negative charge on another atom; important for stabilizing large molecules (e.g., DNA structure).

Relative Strength of Bonds

Type of Bond	Description	Relative Strength
Nonpolar covalent	Electrons shared equally	Strong
Polar covalent	Electrons shared unequally	Strong
Ionic	Electrons transferred, ions formed	Weaker in aqueous environments
Hydrogen	Partial charges attract	Weaker than ionic

Chemical Reactions

Types of Reactions

• Synthesis Reactions: Build larger molecules from smaller ones; require energy (endothermic). Example: Dehydration synthesis (removal of water).

• Decomposition Reactions: Break down molecules into smaller units; release energy (exothermic). Example: Hydrolysis (addition of water).

• Exchange Reactions: Atoms are exchanged between molecules; involve both breaking and forming bonds. Example: AB + CD → AD + CB.

• The sum of all chemical reactions in an organism is called metabolism.

Water, Acids, Bases, and Salts

Water

• Most abundant substance in organisms.

• Properties due to polar covalent bonds:

  • Cohesive (surface tension)

  • Excellent solvent

  • Remains liquid over a wide temperature range

  • High heat capacity

  • Participates in many chemical reactions

Acids and Bases

• Acids: Release H+ ions in water (low pH).

• Bases: Bind H+ or release OH- ions (high pH).

• pH Scale: Measures H+ concentration; ranges from 0 (acidic) to 14 (basic), with 7 as neutral.

• Buffers help maintain stable pH in organisms.

Salts

• Compounds that dissociate in water into cations and anions other than H+ and OH-.

• Salts are important as electrolytes for electrical signaling and metabolic processes.

Organic Macromolecules

Functional Groups

• Specific groupings of atoms within molecules that have characteristic properties (e.g., hydroxyl, carboxyl, amino, phosphate).

Types of Macromolecules

• Lipids: Hydrophobic molecules, not composed of regular subunits. Types include fats (triglycerides), phospholipids, waxes, and steroids.

• Carbohydrates: Composed of C, H, and O. Types include monosaccharides (e.g., glucose), disaccharides (e.g., sucrose), and polysaccharides (e.g., glycogen). Functions: energy source, cell wall structure.

• Proteins: Made of amino acids (monomers). Functions: structure, enzymes, regulation, transport. Peptide bonds link amino acids.

• Nucleic Acids: DNA and RNA are polymers of nucleotides (phosphate, pentose sugar, nitrogenous base). Store and transfer genetic information. ATP is a nucleotide used for energy.

Examples

• Carbohydrate: Glucose (C6H12O6), cellulose.

• Lipid: Triglyceride, phospholipid (major component of cell membranes).

• Protein: Enzyme (e.g., DNA polymerase), structural protein (collagen).

• Nucleic Acid: DNA, RNA, ATP.

Additional info: The structure and function of these macromolecules are foundational for understanding microbial physiology and genetics.