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Chapter 2: The Chemistry of Life – Study Notes for Anatomy & Physiology

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The Chemistry of Life

Introduction

Chemistry is fundamental to understanding human anatomy and physiology because the chemical level of organization forms the basis for all biological structures and processes. Bones, for example, are strong due to their mineral content and the protein collagen, both of which are chemical substances. Thus, chemistry underpins all concepts in anatomy and physiology.

2.1 Atoms and Elements

Definitions and Importance

  • Matter: Anything that has mass and occupies space.

  • Chemistry: The study of matter and its interactions.

  • Atom: The smallest unit of matter that retains its original properties.

  • Element: A substance composed of one or more identical atoms; cannot be broken down into simpler substances by chemical means.

Atomic Structure

  • Atoms are made of three types of subatomic particles:

    • Protons (p+): Positively charged, located in the atomic nucleus.

    • Neutrons (n0): Uncharged, slightly larger than protons, also in the nucleus.

    • Electrons (e-): Negatively charged, tiny particles that surround the nucleus.

  • Atoms are electrically neutral because the number of protons equals the number of electrons.

  • Most atomic mass comes from protons and neutrons.

  • Electrons occupy electron shells:

    • 1st shell: holds 2 electrons

    • 2nd shell: holds 8 electrons

    • 3rd shell: holds 18 electrons (but is stable with 8)

Elements in the Human Body

  • Elements are defined by their atomic number (number of protons).

  • The Periodic Table organizes elements by atomic number and chemical properties.

  • Chemical symbols are abbreviations for elements.

Type

Examples

Body Mass (%)

Major Elements

Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N)

96%

Mineral Elements

Sodium (Na), Potassium (K), Calcium (Ca), Chlorine (Cl), Magnesium (Mg), Phosphorus (P), Sulfur (S)

<4%

Trace Elements

Iron (Fe), Copper (Cu), Iodine (I), Zinc (Zn), others

small amounts

Isotopes and Radioactivity

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

  • Isotope: Atoms with the same number of protons but different numbers of neutrons.

  • Radioisotopes: Unstable isotopes that release energy as radiation; used in nuclear medicine.

  • Examples: Hydrogen has three isotopes—Hydrogen-1 (no neutrons), Deuterium (1 neutron), Tritium (2 neutrons).

Nuclear Medicine Applications

  • Cancer Radiation Therapy: Uses radioisotopes to damage and kill cancer cells.

  • Radiotracers: Injected into patients to visualize organ function and structure.

  • Treatment of Thyroid Disorders: Iodine-131 is used to treat thyroid gland diseases.

2.2 Matter Combined: Mixtures and Chemical Bonds

Mixtures

  • Mixture: Physical combination of two or more elements; can be separated physically.

  • Molecule: Chemical combination of two or more atoms; can only be separated chemically.

Type

Description

Example

Suspension

Liquid mixed with a solid; large particles settle out

Blood (plasma and red blood cells)

Colloid

Liquid mixed with small solid particles; particles remain dispersed

Milk

Solution

Liquid mixed with a solute; solute dissolves in solvent

Salt water

Chemical Bonds

  • Chemical Bond: Attractive force between atoms.

  • Molecule: Atoms of the same element bonded.

  • Compound: Atoms of different elements bonded.

  • Macromolecule: Large compound made of many atoms.

Valence Electrons and Stability

  • Valence Electrons: Electrons in the outermost shell; involved in bonding.

  • Octet Rule: Atoms are most stable with 8 electrons in the valence shell.

  • Duet Rule: Atoms with 5 or fewer electrons are stable with 2 electrons in the first shell.

Ions and Ionic Bonds

  • Ionic Bond: Electrons are transferred between a metal and a nonmetal.

  • Ion: Charged atom (cation = positive, anion = negative).

  • Oppositely charged ions attract to form salts.

Covalent Bonds

  • Covalent Bond: Electrons are shared between two or more nonmetals; strongest bond type.

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

Bond Type

Electron Pairs Shared

Example

Single

1

CH4 (methane)

Double

2

O2 (oxygen gas)

Triple

3

N2 (nitrogen gas)

Nonpolar vs. Polar Covalent Bonds

  • Nonpolar Covalent Bond: Electrons shared equally (same element, or C-H bonds).

  • Polar Covalent Bond: Electrons shared unequally; creates partial charges (dipoles).

  • Electronegativity: Ability of an atom to attract electrons; higher in atoms with more protons.

Hydrogen Bonds

  • Hydrogen Bond: Weak attraction between partially positive and negative atoms in polar molecules.

  • Responsible for surface tension in water.

Determining Bond Types

  • Ionic: Metal + nonmetal

  • Nonpolar: Identical nonmetals, or only C and H

  • Polar: Two nonmetals with different electronegativities (not C and H)

The Big Picture of Chemical Bonding

  • Ionic bonds result from electron transfer and attraction between ions.

  • Covalent bonds result from electron sharing; can be polar or nonpolar.

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