Matter, Energy & Bonds

Introduction

Understanding the basic principles of matter, energy, and chemical bonds is essential for studying Anatomy & Physiology. These concepts explain the structure of atoms, the formation of molecules, and the interactions that underlie biological processes.

Subatomic Particles and Atomic Structure

Electrons, Protons, and Neutrons

Atoms are the fundamental units of matter, composed of three main subatomic particles:

• Protons: Positively charged particles located in the nucleus.

• Neutrons: Neutral particles also found in the nucleus.

• Electrons: Negatively charged particles that orbit the nucleus in electron shells.

Atomic Number vs. Mass Number

• Atomic Number: The number of protons in an atom; defines the element.

• Mass Number: The sum of protons and neutrons in the nucleus.

Atoms are usually electrically neutral, meaning the number of protons equals the number of electrons.

Summary Table: Subatomic Particles

Electron Shells

• Electron shells are regions around the nucleus where electrons are found.

• Each shell can hold a specific number of electrons:

  • First shell: 2 electrons

  • Second shell: 8 electrons

  • Third shell: 8 electrons (satisfied with 8, known as the octet rule)

• Atoms may have more than three shells, especially in larger elements.

Atoms, Elements, and Major Elements in the Body

Atoms vs. Elements

• Atom: The smallest unit of matter with unique chemical properties.

• Element: A substance made entirely of one type of atom.

• Examples:

  • Hydrogen and Oxygen are elements; a water molecule (H2O) is a compound made of these elements.

Major Elements in the Human Body

• The body is primarily composed of a few major elements: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N).

• These elements are essential for forming biomolecules such as proteins, carbohydrates, lipids, and nucleic acids.

Isotopes and Radioisotopes

Definitions and Properties

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

• Isotopes have identical chemical properties but may differ in stability.

• Radioisotopes: Unstable isotopes that decay over time, emitting radiation.

Radioisotope Applications

• Used in medical imaging and treatment (e.g., Iodine-131 for thyroid disorders).

• Examples:

  • Gallium-67: Medical diagnosis

  • Iodine-123: Diagnosing thyroid and metabolic disorders

  • Iodine-125: Clinical tests and biomedical research

Mixtures: Solutions, Solvents, Colloids, and Suspensions

Definitions

• Solution: Homogeneous mixture where solute is completely dissolved in solvent (e.g., salt water).

• Solvent: Substance that dissolves the solute (e.g., water).

• Colloid: Mixture with small, evenly distributed particles that do not settle out (e.g., milk).

• Suspension: Heterogeneous mixture with large particles that settle out over time (e.g., muddy water).

Comparison Table: Mixture Types

Valence Shells and Chemical Reactivity

Valence Electrons and the Octet Rule

• Valence shell: The outermost electron shell of an atom.

• Valence electrons determine chemical reactivity and bonding behavior.

• Octet rule: Atoms are most stable when they have eight electrons in their valence shell (except for small atoms like hydrogen and helium, which follow the duet rule).

Ions and Electrolytes

Definitions

• Ion: An atom or molecule with a net electric charge due to loss or gain of electrons.

• Electrolyte: A substance that dissociates into ions in water and can conduct electricity (e.g., sodium chloride).

Chemical Bonds

Types of Chemical Bonds

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

• Covalent Bonds: Formed by the sharing of electrons between atoms.

  • Non-polar covalent bonds: Equal sharing of electrons (e.g., O2 molecule).

  • Polar covalent bonds: Unequal sharing of electrons, creating partial charges (e.g., water, H2O).

• Hydrogen Bonds: Weak attractions between a hydrogen atom in one molecule and an electronegative atom (such as oxygen or nitrogen) in another. Example: Hydrogen bonds in water and DNA structure.

• Van der Waals Forces: Weak attractions due to temporary polarization in nonpolar molecules. Example: Interactions between noble gases.

Summary Table: Types of Chemical Bonds

Additional info:

• Double and triple covalent bonds are stronger than single covalent bonds.

• Hydrogen bonds are essential for the structure of proteins and nucleic acids.

• Van der Waals forces include London dispersion forces, dipole-dipole interactions, and dipole-induced dipole forces.

Key Equations

• Atomic mass: