Chapter 2: Chemistry Comes Alive

Introduction

Chemistry is fundamental to understanding physiological processes in the human body. All living and nonliving matter is composed of chemicals, and the interactions between these chemicals underlie all bodily functions, from muscle contraction to nerve signaling. This chapter provides an essential foundation in basic chemistry and biochemistry for students of anatomy and physiology.

2.1 Matter and Energy

Matter

• Matter is anything that has mass and occupies space.

• It can exist in three states:

  • Solid: Definite shape and volume.

  • Liquid: Changeable shape, definite volume.

  • Gas: Changeable shape and volume.

• Weight is mass plus the effects of gravity.

Energy

• Energy is the capacity to do work or put matter into motion.

• Two main forms:

  • Kinetic energy: Energy in action.

  • Potential energy: Stored (inactive) energy.

• Energy can be transformed from one form to another, but some is always lost as heat.

• Major forms of energy in the body:

  • Chemical energy: Stored in bonds of chemical substances.

  • Electrical energy: Movement of charged particles.

  • Mechanical energy: Directly involved in moving matter.

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

2.2 Atoms and Elements

Elements

• Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods.

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

• Other important elements include calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium, and trace elements like iron and iodine.

Table: Common Elements in the Human Body

Atoms

• Atoms are the smallest units of an element that retain the properties of that element.

• Composed of three subatomic particles:

  • Protons (p+): Positive charge, 1 amu, in nucleus

  • Neutrons (n0): No charge, 1 amu, in nucleus

  • Electrons (e-): Negative charge, ~0 amu, orbit nucleus

Atomic Structure Examples

• Hydrogen: 1 proton, 0 neutrons, 1 electron

• Helium: 2 protons, 2 neutrons, 2 electrons

• Lithium: 3 protons, 4 neutrons, 3 electrons

Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Hydrogen, Deuterium, and Tritium are isotopes of hydrogen.

2.3 Combining Matter

Molecules and Compounds

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

• Compound: Molecule with two or more different kinds of atoms (e.g., H2O).

Mixtures

• Most matter exists as mixtures: two or more components physically intermixed.

• Three types:

  • Solutions: Homogeneous, solute particles do not settle out (e.g., mineral water).

  • Colloids: Heterogeneous, larger particles that do not settle out (e.g., Jell-O).

  • Suspensions: Heterogeneous, large particles that settle out (e.g., blood).

Solution Example

Colloid Example

Suspension Example

2.4 Chemical Bonds

Role of Electrons in Chemical Bonding

• Electrons occupy energy levels called electron shells.

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

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve 8 electrons in their valence shell (except H and He, which are stable with 2).

Chemically Inert vs. Reactive Elements

• Inert elements have complete valence shells and are stable (e.g., Helium, Neon).

• Reactive elements have incomplete valence shells and tend to form bonds (e.g., Hydrogen, Carbon, Oxygen, Sodium).

Types of Chemical Bonds

• Ionic bonds: Transfer of electrons from one atom to another, forming ions (cations and anions).

• Covalent bonds: Sharing of electrons between atoms. Can be single, double, or triple bonds.

• Hydrogen bonds: Weak attractions between a hydrogen atom and an electronegative atom (e.g., between water molecules).

Ionic Bond Example: Sodium Chloride Formation

Covalent Bond Examples

Nonpolar vs. Polar Covalent Bonds

• Nonpolar: Equal sharing of electrons (e.g., CO2).

• Polar: Unequal sharing, resulting in partial charges (e.g., H2O).

Bond Comparison Table

Hydrogen Bonding in Water

2.5 Chemical Reactions

Chemical Equations

• Represent the process of reactants forming products.

• Example:

Types of Chemical Reactions

• Synthesis (Combination): Atoms/molecules combine to form larger molecules. (Anabolic)

• Decomposition: Molecules broken down into smaller molecules/atoms. (Catabolic)

• Exchange (Displacement): Bonds are both made and broken; atoms are exchanged between molecules.

Redox Reactions

• Involve the transfer of electrons between atoms.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

Energy Flow in Chemical Reactions

• Exergonic: Release energy (products have less energy than reactants).

• Endergonic: Absorb energy (products have more energy than reactants).

Factors Affecting Reaction Rate

• Temperature (higher = faster)

• Concentration (higher = faster)

• Particle size (smaller = faster)

• Catalysts (e.g., enzymes) increase reaction rate without being consumed.

2.6 Inorganic Compounds

Water

• Makes up 60–80% of cell volume; most abundant inorganic compound in the body.

• Key properties:

  • High heat capacity: Absorbs/releases heat with little temperature change.

  • High heat of vaporization: Evaporation requires much energy (cooling effect).

  • Polar solvent properties: Dissolves and dissociates ionic substances; forms hydration layers.

  • Reactivity: Involved in hydrolysis and dehydration synthesis reactions.

  • Cushioning: Protects organs (e.g., cerebrospinal fluid).

Salts

• Ionic compounds that dissociate in water to form electrolytes (conduct electricity).

• Vital for nerve impulse transmission, muscle contraction, and water balance.

• Examples: NaCl, CaCO3, KCl.

Acids and Bases

• Acids: Proton donors; release H+ ions (e.g., HCl).

• Bases: Proton acceptors; release OH- ions (e.g., NaOH).

• pH scale: Measures H+ concentration; ranges from 0 (acidic) to 14 (basic), with 7 as neutral.

• Buffers: Resist abrupt changes in pH by releasing or binding H+ ions. Example: Bicarbonate buffer system in blood.

2.7 Organic Compounds: Synthesis and Hydrolysis

• Organic molecules contain carbon (except CO2 and CO).

• Major classes: Carbohydrates, lipids, proteins, nucleic acids.

• Many are polymers (chains of monomers).

• Dehydration synthesis: Joins monomers by removing water.

• Hydrolysis: Breaks polymers into monomers by adding water.

2.8 Carbohydrates

• Include sugars and starches; contain C, H, O (2:1 ratio of H:O).

• Three classes:

  • Monosaccharides: Simple sugars (e.g., glucose, fructose, ribose).

  • Disaccharides: Two monosaccharides joined (e.g., sucrose, maltose, lactose).

  • Polysaccharides: Long chains of monosaccharides (e.g., starch, glycogen).

2.9 Lipids

• Contain C, H, O (less O than carbohydrates); insoluble in water.

• Main types:

  • Triglycerides: Three fatty acids + glycerol; energy storage, insulation, protection.

  • Phospholipids: Glycerol + two fatty acids + phosphate group; major component of cell membranes.

  • Steroids: Four interlocking rings; cholesterol is the most important.

  • Eicosanoids: Derived from arachidonic acid; roles in inflammation, blood clotting.

Summary Table: Types of Mixtures

Additional info: This summary covers the foundational chemical principles necessary for understanding anatomy and physiology, including atomic structure, chemical bonding, types of mixtures, and the major classes of biological molecules. These concepts are essential for further study of cellular structure, metabolism, and physiological processes.