Chapter 2: Chemistry Comes Alive

Introduction

Chemistry is fundamental to understanding physiological processes in the human body. This chapter explores the basic principles of chemistry and biochemistry as they relate to anatomy and physiology, including the structure of matter, types of chemical bonds, and the properties of water, salts, acids, bases, and organic molecules essential for life.

2.1 Matter and Energy

States of 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), and 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. It exists as kinetic energy (energy in action) and potential energy (stored energy).

• Forms of energy include chemical, electrical, mechanical, and radiant (electromagnetic) energy.

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

2.2 Atoms and Elements

Elements and Atoms

• Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods. Four elements (carbon, oxygen, hydrogen, nitrogen) make up 96% of the human body.

• Atoms are the smallest units of elements, consisting of protons, neutrons, and electrons.

Atomic Structure

• Protons: Positive charge, 1 amu, found in nucleus.

• Neutrons: No charge, 1 amu, found in nucleus.

• Electrons: Negative charge, virtually no mass, orbit the nucleus.

Atomic Number, Mass Number, Isotopes

• Atomic number: Number of protons in the nucleus.

• Mass number: Total number of protons and neutrons.

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

2.3 Combining Matter

Molecules, Compounds, and Mixtures

• Molecule: Two or more atoms bonded together.

• Compound: Molecule with two or more different kinds of atoms.

• Mixtures: Physical combinations of two or more substances. Types include solutions, colloids, and suspensions.

Types of Mixtures

• Solutions: Homogeneous mixtures; solute particles are very tiny and do not settle out or scatter light. Example: mineral water.

• Colloids: Heterogeneous mixtures; solute particles are larger and scatter light but do not settle out. Example: Jell-O.

• Suspensions: Heterogeneous mixtures with large, visible solutes that settle out. Example: blood.

2.4 Chemical Bonds

Role of Electrons in Bonding

• Electrons occupy energy levels called electron shells. The outermost shell is the valence shell and determines chemical reactivity.

• The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell (except H and He, which require two).

Types of Chemical Bonds

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

• Covalent bonds: Formed by sharing 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).

Polar vs. Nonpolar Covalent Bonds

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

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

Hydrogen Bonds

• Hydrogen bonds are important in maintaining the structure of water, proteins, and DNA.

2.5 Chemical Reactions

Types of Chemical Reactions

• Synthesis (combination) reactions: Atoms or molecules combine to form larger molecules (anabolic).

• Decomposition reactions: Molecules are broken down into smaller molecules (catabolic).

• Exchange (displacement) reactions: Bonds are both made and broken.

Energy Flow in Chemical Reactions

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

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

Factors Affecting Reaction Rates

• Temperature, concentration, particle size, and catalysts (enzymes) affect the speed of chemical reactions.

2.6 Inorganic Compounds

Water

• Most abundant and important inorganic compound in the body (60–80% of cell volume).

• Properties: high heat capacity, high heat of vaporization, polar solvent, reactivity, and cushioning.

Salts

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

• 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

• Buffers resist abrupt changes in pH by releasing or binding H+ ions. The bicarbonate buffer system is important in blood.

2.7 Organic Compounds

General Characteristics

• Contain carbon (except CO2 and CO).

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

• Many are polymers formed by dehydration synthesis and broken down by hydrolysis.

2.8 Carbohydrates

Classification

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

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

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

2.9 Lipids

Types of Lipids

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

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

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

• Eicosanoids: Derived from arachidonic acid; include prostaglandins involved in inflammation and other functions.

Additional info: Unsaturated fatty acids have one or more double bonds, causing kinks in the chain and making them liquid at room temperature (e.g., plant oils).