Basic Chemistry for Anatomy & Physiology: Atoms, Subatomic Particles, and Chemical Bonds

Basic Chemistry (Part 1 of 2)

Introduction

This module introduces foundational chemistry concepts essential for understanding biological processes in Anatomy & Physiology. Topics include the structure of atoms, subatomic particles, isotopes, and the nature of chemical bonds.

Atoms

Definition and Structure

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

• Atoms are the building blocks of all substances, including those found in the human body.

• Each element is composed of only one kind of atom, which gives the element its unique properties.

• Examples:

  • Carbon (C): Found in all living things; forms the backbone of organic molecules.

  • Sodium (Na): Essential for nerve function and fluid balance.

Additional info: If atoms are divided further into subatomic particles, they no longer retain the properties of the element.

Subatomic Particles of an Atom

Types and Properties

• Protons: Positively charged particles located in the nucleus; mass = 1 atomic mass unit (amu).

• Neutrons: Neutral particles (no charge) also found in the nucleus; mass = 1 amu.

• Electrons: Negatively charged particles orbiting the nucleus; mass is negligible compared to protons and neutrons.

Summary Table:

Particle	Charge	Location	Relative Mass
Proton	+1	Nucleus	1 amu
Neutron	0	Nucleus	1 amu
Electron	-1	Orbitals/Shells	~0 amu

Atomic Neutrality and Ions

Electrical Neutrality

• Atoms are electrically neutral when they have equal numbers of protons and electrons.

• If an atom loses or gains electrons, it becomes an ion (charged particle).

• Cation: An atom that has lost one or more electrons (positively charged).

• Anion: An atom that has gained one or more electrons (negatively charged).

Example: If a sodium atom (Na) loses one electron, it becomes a Na+ cation.

Radioactive Isotopes

Definition and Applications

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

• Radioactive isotopes: Unstable isotopes that emit radiation as they decay to become more stable forms.

• Radiation: The energy or particles released during radioactive decay.

• Medical Applications:

  • Radioactive isotopes are used in diagnostic imaging (e.g., PET scans) and cancer treatment.

  • Example: Radioactive iodine is used to treat thyroid disorders.

Electron Shells and Valence Electrons

Organization and Chemical Reactivity

• Electrons orbit the nucleus in energy levels called shells.

• The valence shell is the outermost electron shell and determines an atom's chemical reactivity.

• Atoms are inert (non-reactive) if their valence shell is full; otherwise, they are reactive and tend to form bonds to achieve stability.

• Example: Hydrogen, carbon, oxygen, and nitrogen all have vacancies in their valence shells and readily form bonds.

Molecules and Compounds

Definitions and Examples

• Molecule: Two or more atoms bonded together (can be the same or different elements).

• Compound: A molecule that contains atoms of more than one element in fixed proportions.

• Examples:

  • Water (H2O): Compound of hydrogen and oxygen.

  • Carbon dioxide (CO2): Compound of carbon and oxygen.

  • Methane (CH4): Compound of carbon and hydrogen.

Chemical Bonds

Types of Chemical Bonds

• Ionic Bonds: Formed when one atom donates electrons to another, resulting in attraction between oppositely charged ions (cations and anions).

• Covalent Bonds: Formed when two atoms share one or more pairs of electrons.

• Hydrogen Bonds: Weak attractions between a hydrogen atom (already covalently bonded to an electronegative atom) and another electronegative atom (often oxygen or nitrogen).

Ionic Bonds

• Formed between metals and nonmetals (e.g., NaCl).

• Result in the formation of crystalline solids (salts).

• Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl-, which are held together by electrostatic attraction.

Covalent Bonds

• Formed between nonmetals.

• Atoms share electrons to fill their valence shells.

• Single covalent bond: One pair of shared electrons (e.g., H2).

• Double covalent bond: Two pairs of shared electrons (e.g., O2).

• Polar covalent bond: Unequal sharing of electrons, leading to partial charges (e.g., H2O).

• Nonpolar covalent bond: Equal sharing of electrons (e.g., O2).

Hydrogen Bonds

• Very weak compared to ionic and covalent bonds.

• Important in stabilizing the structures of proteins and nucleic acids, and in the properties of water.

• Example: Hydrogen bonds between water molecules give water its high surface tension and boiling point.

ATP: The Energy Molecule

Structure and Function

• Adenosine Triphosphate (ATP): The primary energy carrier in cells.

• Composed of adenosine (adenine + ribose) and three phosphate groups.

• Energy is stored in the high-energy phosphate bonds and released when ATP is converted to ADP (adenosine diphosphate) and inorganic phosphate (Pi):

$\text{ATP} \rightarrow \text{ADP} + \text{P}_i + \text{energy}$

• This energy is used for cellular processes such as muscle contraction, nerve impulse transmission, and biosynthesis.

Additional info: ATP is regenerated from ADP and Pi during cellular respiration, primarily in the mitochondria.