Basic Chemistry for Anatomy & Physiology

Basic Chemistry

Introduction

Chemistry is fundamental to understanding Anatomy & Physiology, as it explains the composition and behavior of matter and energy in biological systems. This section covers the essential chemical principles relevant to the human body.

Matter and Energy

Definitions and Properties

• Matter: Anything that has mass and occupies space. It can be seen, smelled, and/or felt.

• Energy: The capacity to do work or put matter into motion. Energy does not have mass or take up space.

• Weight: Mass gravity

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, changeable shape.

• Gas: Changeable shape and volume.

Forms of Energy

• Chemical energy: Stored in bonds of chemical substances.

• Electrical energy: Results from movement of charged particles (ions).

• Mechanical energy: Directly involved in moving matter.

• Radiant (electromagnetic) energy: Travels in waves (e.g., heat, visible light, UV light, X-rays).

Energy Conversions

• Energy can be converted from one form to another, but some energy is lost as heat during conversion.

• Energy conversion is inefficient; some energy becomes unusable.

• Catabolic reactions: Break down molecules for ATP production.

Atoms and Elements

Elements

• All matter is composed of elements, substances that cannot be broken down by ordinary chemical means.

• Four elements make up 96% of body mass: Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N).

• Other elements (e.g., calcium, phosphorus, potassium) are present in smaller amounts.

• The periodic table lists all known elements (currently 118).

Atoms

• Smallest units of elements that retain their properties.

• Composed of protons (positive charge), neutrons (no charge), and electrons (negative charge).

• Protons and neutrons are found in the nucleus; electrons orbit the nucleus in shells.

Atomic Structure

• Atomic number: Number of protons in the nucleus.

• Mass number: Total number of protons and neutrons.

• Isotopes: Variations of an element with the same number of protons but different numbers of neutrons.

• Atomic weight: Average of mass numbers of all isotopes of an element.

Radioisotopes

• Unstable isotopes that decompose to more stable forms, emitting radiation.

• Used in medical imaging (e.g., PET scans) and research.

• All radioactivity can damage living tissue; some types are used to treat cancer.

Combining Matter: Molecules, Compounds, and Mixtures

Molecules and Compounds

• Molecule: Two or more atoms bonded together (e.g., O2).

• Compound: Two or more different kinds of atoms bonded together (e.g., H2O).

Mixtures

• Physical combinations of two or more components.

• Three basic types: solutions, colloids, and suspensions.

Types of Mixtures

• Solutions: Homogeneous mixtures; solute particles are very small and do not settle out.

• Colloids: Heterogeneous mixtures; solute particles are larger and do not settle out (e.g., cytosol).

• Suspensions: Heterogeneous mixtures with large, visible solutes that settle out (e.g., blood).

Concentration of Solutions

• Expressed as percent, milligrams per deciliter (mg/dL), or molarity (M).

• Molarity (M): Number of moles of solute per liter of solvent.

• 1 mole = 6.02 × 1023 molecules (Avogadro's number).

Main Differences Between Mixtures and Compounds

• Mixtures do not involve chemical bonding; compounds do.

• Mixtures can be separated by physical means; compounds require chemical means.

• Mixtures can be homogeneous or heterogeneous; compounds are always homogeneous.

Chemical Bonds

Types of Chemical Bonds

• Ionic bonds: Involve the transfer of electrons from one atom to another, resulting in ions (cations and anions).

• Covalent bonds: Involve the sharing of electrons between atoms.

• Hydrogen bonds: Weak attractions between a hydrogen atom and an electronegative atom (e.g., oxygen or nitrogen).

Role of Electrons in Chemical Bonding

• Electrons occupy energy levels (shells) around the nucleus.

• The outermost shell is called the valence shell and determines chemical reactivity.

• Atoms are most stable when their valence shell is full (usually 8 electrons; "octet rule").

• Atoms will gain, lose, or share electrons to achieve a full valence shell.

Ionic Bonds

• Formed when one atom loses electrons (becomes a cation) and another gains electrons (becomes an anion).

• Example: NaCl (sodium chloride).

Covalent Bonds

• Formed by sharing pairs of electrons between atoms.

• Single, double, or triple bonds depending on the number of shared electron pairs.

• Nonpolar covalent bonds: Electrons are shared equally (e.g., O2).

• Polar covalent bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

Hydrogen Bonds

• Attractions between a hydrogen atom (already covalently bonded to an electronegative atom) and another electronegative atom.

• Important in maintaining the structure of proteins and DNA.

Summary Table: Types of Chemical Bonds

Bond Type	How Formed	Strength	Example
Ionic	Transfer of electrons	Strong (in solid state)	NaCl
Covalent	Sharing of electrons	Strong	H2O, O2
Hydrogen	Attraction between H and electronegative atom	Weak	Between water molecules

Key Equations and Constants

• Avogadro's Number: $6.02 \times 10^{23}$ molecules/mol

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Additional info:

• Understanding chemical principles is essential for grasping physiological processes such as metabolism, nerve conduction, and muscle contraction.

• Radioisotopes are used in diagnostic imaging and cancer treatment.