Fundamental Chemistry Concepts for Microbiology

Basic Atomic Structure and Subatomic Particles

Subatomic Particles: Protons, Neutrons, and Electrons

Atoms are the fundamental units of matter, composed of three main subatomic particles: protons, neutrons, and electrons. Understanding their properties is essential for grasping chemical behavior relevant to microbiology.

• Protons: Positively charged particles located in the nucleus. They determine the atomic number and identity of an element. Charge: +1; Mass: ~1 atomic mass unit (amu).

• Neutrons: Neutral particles also found in the nucleus. They contribute to atomic mass but do not affect charge. Charge: 0; Mass: ~1 amu.

• Electrons: Negatively charged particles orbiting the nucleus in electron shells. Charge: -1; Mass: ~1/1836 amu (much lighter than protons/neutrons).

Example: A carbon atom (atomic number 6) has 6 protons, typically 6 neutrons, and 6 electrons.

Isotopes and Their Biological Applications

Definition and Use of Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different atomic masses. Some isotopes are stable, while others are radioactive.

• Radioisotopes are used in biological research for tracing biochemical pathways, dating fossils, and medical diagnostics.

• Example: Carbon-14 is a radioactive isotope used in radiocarbon dating and as a tracer in metabolic studies.

Valence Electrons and Chemical Bonding

Role of Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. They determine an atom's chemical reactivity and bonding behavior.

• Elements in the same group of the periodic table have the same number of valence electrons, leading to similar chemical properties.

• Example: All Group 1 elements (alkali metals) have one valence electron and tend to form +1 ions.

Molecules vs. Compounds

Definitions and Differences

A molecule is two or more atoms covalently bonded together. A compound is a substance composed of two or more different elements chemically combined in fixed proportions.

• Molecule Example: O2 (oxygen gas) – two oxygen atoms bonded together.

• Compound Example: H2O (water) – hydrogen and oxygen atoms bonded together.

Ionic and Covalent Bonds in Biological Molecules

Comparison and Biological Significance

Ionic bonds form when electrons are transferred from one atom to another, creating charged ions that attract each other. Covalent bonds form when atoms share electron pairs.

• Ionic bonds: Often found in salts (e.g., NaCl). Important for nerve impulse transmission.

• Covalent bonds: Found in most biological molecules (e.g., proteins, DNA, carbohydrates).

• Effect on Structure/Function: Covalent bonds provide stability; ionic bonds can dissociate in water, affecting molecule interactions.

Example: The peptide bond in proteins is a covalent bond.

Unique Properties of Water

Key Properties and Their Biological Roles

• Cohesion: Water molecules stick together via hydrogen bonds, aiding in transport in plants.

• High Heat Capacity: Water absorbs heat with minimal temperature change, stabilizing environments.

• Solvent Abilities: Water dissolves many substances, facilitating biochemical reactions.

• High Heat of Vaporization: Evaporation of water cools organisms (e.g., sweating).

Example: Water's solvent property allows nutrients and waste to be transported in blood.

Hydrogen Bonds and Water's Polarity

Formation and Biological Importance

Hydrogen bonds form between the slightly positive hydrogen atom of one water molecule and the slightly negative oxygen atom of another, due to water's polarity.

• Polarity arises from the unequal sharing of electrons in the O-H bond.

• Hydrogen bonds are crucial for the structure of DNA, proteins, and the unique properties of water.

Biological Importance: Hydrogen bonds stabilize macromolecular structures and enable water's role as a universal solvent.

Molecular Polarity and Solubility

Effect of Polarity on Solubility

Polar molecules dissolve well in water (a polar solvent) due to favorable interactions, while nonpolar molecules do not.

• "Like dissolves like": Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.

• Example: Glucose (polar) dissolves in water; oils (nonpolar) do not.

pH, Acids, Bases, and Buffers in Biology

Definitions and Biological Importance

• pH: A measure of hydrogen ion concentration; $pH = -\log_{10}[H^+]$

• Acids: Substances that increase [H+] in solution (pH < 7).

• Bases: Substances that decrease [H+] (pH > 7).

• Buffers: Solutions that resist changes in pH by absorbing or releasing H+ ions.

Biological Importance: Buffers (e.g., bicarbonate buffer system) maintain homeostasis by stabilizing pH in blood and cells.

Example: The bicarbonate buffer system in blood: $CO_2 + H_2O \leftrightarrow H_2CO_3 \leftrightarrow HCO_3^- + H^+$

Capillary Action in Nature

Definition, Mechanism, and Examples

Capillary action is the movement of liquid within narrow spaces due to the combination of cohesive and adhesive forces.

• Why it happens: Water molecules adhere to the walls of a tube (adhesion) and pull other water molecules along (cohesion).

• Examples from nature: Water transport in plant xylem; movement of water in soil.