Introduction to Chemistry in Anatomy & Physiology

Overview

Chemistry is fundamental to understanding the structure and function of the human body. This chapter introduces the language and core concepts of chemistry, explains how matter is organized, describes the formation of chemical bonds and reactions, and compares organic and inorganic compounds as they relate to physiological processes.

• Key Focus: The relationship between chemistry and physiological reactions in living organisms.

• Applications: Medicine, pharmacology, biotechnology, and healthcare advancements.

Basic Principles of Chemistry

Definition of Matter

• Matter: Anything that has mass and occupies space.

• Mass: The amount of matter in a substance.

• Weight: The force of gravity acting on mass.

Elements and Compounds

• Element: A pure substance consisting of only one type of atom; cannot be broken down by chemical means.

• Compound: A substance formed from two or more elements combined in a fixed ratio; has different properties from its constituent elements.

• Example: Water (H2O) is a compound of hydrogen and oxygen.

Atoms and Molecules

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

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

Organization of Matter

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, shape of container.

• Gas: No definite shape or volume.

Chemical Elements in the Human Body

• There are 92 naturally occurring elements; about 25 are found in living organisms.

• Major elements: Oxygen, carbon, hydrogen, and nitrogen make up 96% of body mass.

Atomic Structure

Subatomic Particles

• Proton: Positively charged, found in nucleus.

• Neutron: No charge, found in nucleus.

• Electron: Negatively charged, orbits nucleus.

Atomic Number, Mass Number, and Isotopes

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Sum of protons and neutrons.

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

• Radioisotopes: Unstable isotopes that decay, emitting radiation; used in medical imaging and cancer treatment.

Electron Shells and Valence Electrons

• Electrons are arranged in shells around the nucleus.

• The outermost shell is the valence shell; electrons here determine chemical reactivity.

• Atoms are most stable when their valence shell is full (octet rule).

The Periodic Table

Organization and Use

• Elements are arranged by increasing atomic number.

• Groups (columns) contain elements with similar chemical properties due to the same number of valence electrons.

• Periods (rows) show trends in properties across the table.

Ions, Molecules, and Compounds

Ions

• Cation: Positively charged ion (lost electrons).

• Anion: Negatively charged ion (gained electrons).

• Example: Na+ (sodium ion), Cl- (chloride ion).

Molecules and Compounds

• Molecule: Two or more atoms held together by chemical bonds.

• Compound: Molecule containing atoms of different elements.

Mixtures and Solutions

• Solution: Homogeneous mixture of solute dissolved in solvent.

• Solvent: The dissolving agent (e.g., water).

• Solute: The substance dissolved (e.g., salt).

• Example: Blood plasma is a solution with water as the solvent.

Mole and Avogadro's Number

• Mole: The amount of substance containing particles (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

• Example: 1 mole of H2O = 18.02 g.

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.

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

• Nonpolar Covalent: Electrons shared equally.

• Polar Covalent: Electrons shared unequally, creating partial charges.

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

Octet Rule

• Atoms tend to form bonds to achieve eight electrons in their valence shell, leading to stability.

Chemical Reactions

Types of Chemical Reactions

• Synthesis (Anabolism): Two or more substances combine to form a more complex product.

• Decomposition (Catabolism): A complex molecule breaks down into simpler substances.

• Exchange: Parts of molecules are exchanged.

• Reversible: Products can revert to reactants.

• Oxidation-Reduction (Redox): Involves transfer of electrons; oxidation is loss, reduction is gain.

Energy in Chemical Reactions

• Energy: The capacity to do work; exists as kinetic (in motion) or potential (stored).

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

• Activation Energy: The minimum energy required to start a reaction.

• Catalyst: Substance that speeds up a reaction by lowering activation energy; in biology, these are enzymes.

Factors Affecting Reaction Rates

• Temperature (higher = faster)

• Concentration (higher = faster)

• Particle size (smaller = faster)

• Catalysts (increase rate)

Inorganic vs. Organic Compounds

Definitions and Examples

• Inorganic Compounds: Usually do not contain carbon; examples include water, salts, acids, and bases.

• Organic Compounds: Contain carbon, often bonded to hydrogen, oxygen, nitrogen, or other elements; examples include carbohydrates, lipids, proteins, and nucleic acids.

Water: Importance and Properties

Properties of Water

• Polarity: Water is a polar molecule, allowing it to dissolve many substances (universal solvent).

• Cohesion and Adhesion: Water molecules stick to each other (cohesion) and to other substances (adhesion).

• High Specific Heat: Water can absorb or release large amounts of heat with little temperature change.

• High Heat of Vaporization: Large amount of energy required to convert water from liquid to gas.

• Role in Chemical Reactions: Medium for most body reactions; participates in hydrolysis and dehydration synthesis.

Acids, Bases, and Salts

Definitions

• Acid: Substance that releases hydrogen ions (H+); proton donor.

• Base: Substance that accepts hydrogen ions or releases hydroxide ions (OH-); proton acceptor.

• Salt: Ionic compound formed from acid-base reactions; dissociates into ions in water.

pH Scale

• Measures hydrogen ion concentration; ranges from 0 (acidic) to 14 (basic), with 7 as neutral.

• Each pH unit represents a tenfold change in H+ concentration.

• Buffers: Substances that resist changes in pH by absorbing or releasing H+ or OH-.

Overview of Organic Compounds

Major Classes

• Carbohydrates: Sugars and starches; provide energy and structural support.

• Lipids: Fats, oils, phospholipids, steroids; energy storage, cell membranes, hormones.

• Proteins: Polymers of amino acids; structure, enzymes, transport, defense.

• Nucleic Acids: DNA and RNA; store and transmit genetic information.

Functional Groups

• Specific groups of atoms within molecules that determine chemical reactivity (e.g., hydroxyl, carboxyl, amino, phosphate).

Polymers and Monomers

• Many organic molecules are polymers, made by joining monomers via dehydration synthesis; broken down by hydrolysis.

Summary Table: Comparison of Inorganic and Organic Compounds

Key Equations and Concepts

• Avogadro's Number: particles/mol

• Molar Mass Calculation:

• pH Calculation:

Additional info:

• Free radicals are highly reactive molecules with unpaired electrons; antioxidants neutralize them to prevent cellular damage.

• Enzymes are biological catalysts, usually proteins, that speed up chemical reactions without being consumed.

• ATP (adenosine triphosphate) is the primary energy carrier in cells, releasing energy when hydrolyzed to ADP and phosphate.