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 (C) has 6 protons, 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 and similar chemical properties.

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

Molecules vs. Compounds

Definitions and Differences

• Molecule: Two or more atoms covalently bonded together (can be same or different elements). Example: O2 (oxygen gas).

• Compound: A substance formed from two or more different elements chemically bonded. Example: H2O (water).

All compounds are molecules, but not all molecules are compounds.

Ionic and Covalent Bonds in Biological Molecules

Comparison and Biological Significance

• Ionic Bonds: Formed by the transfer of electrons from one atom to another, resulting in oppositely charged ions that attract each other. Example: NaCl (sodium chloride).

• Covalent Bonds: Formed by the sharing of electron pairs between atoms. Example: The bonds between hydrogen and oxygen in H2O.

• Ionic bonds are generally weaker in aqueous environments, while covalent bonds provide stability to biological molecules.

Unique Properties of Water

Key Properties and Their Biological Importance

• Cohesion: Water molecules stick to each other due to hydrogen bonding, enabling surface tension.

• High Heat Capacity: Water can absorb or release large amounts of heat with little temperature change, stabilizing environments.

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

• High Heat of Vaporization: Large amounts of energy are required to convert water from liquid to gas, aiding in temperature regulation.

Each property supports life by maintaining stable conditions and enabling essential processes.

Hydrogen Bonding and Water's Polarity

Origin and Biological Importance

• Polarity: Water is a polar molecule with partial positive (hydrogen) and partial negative (oxygen) ends.

• Hydrogen Bonds: Weak attractions between the hydrogen atom of one water molecule and the oxygen atom of another.

• Hydrogen bonds are crucial for the structure of proteins, nucleic acids, and for water's unique properties.

Molecular Polarity and Solubility

Effect of Polarity on Solubility

• Polar molecules dissolve well in polar solvents (like water) due to favorable interactions.

• Nonpolar molecules are insoluble in water but dissolve in nonpolar solvents.

• "Like dissolves like" principle governs solubility.

pH, Acids, Bases, and Buffers in Biological Systems

Definitions and Importance

• pH: A measure of hydrogen ion concentration;

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

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

• Buffers: Solutions that resist changes in pH by absorbing or releasing H+ ions. Example: The bicarbonate buffer system in blood.

Buffers are essential for maintaining homeostasis in biological systems.

Capillary Action

Definition, Mechanism, and Examples

• Capillary Action: The movement of liquid within narrow spaces due to adhesive and cohesive forces.

• Occurs when adhesion to the walls is stronger than cohesion between liquid molecules.

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