Chemistry Comes Alive: Foundations for Anatomy & Physiology

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Chapter 2: Chemistry Comes Alive

Introduction

This chapter provides foundational chemical concepts essential for understanding human anatomy and physiology. It covers the nature of matter and energy, atomic structure, chemical bonding, types of chemical reactions, and the properties of water and other compounds critical to life.

Matter and Energy

Matter

Matter is anything that has mass and occupies space.
Matter can be seen, smelled, and/or felt.
Weight is mass plus the effects of gravity.

States of Matter

Solid: Definite shape and volume.
Liquid: Changeable shape; definite volume.
Gas: Changeable shape and volume.

Energy

Energy is the capacity to do work or put matter into motion.
The greater the work done, the more energy is used up.
Energy form conversions: Energy may be converted from one form to another, but conversion is inefficient (some energy is lost as heat).

Atoms and Elements

All elements are made up of atoms, which are the unique building blocks for each element.
Atoms are the smallest particles of an element that retain its properties.
Atoms give each element its particular physical and chemical properties.
Atoms are composed of three subatomic particles:
- Protons (positively charged)
- Neutrons (neutral)
- Electrons (negatively charged)

Common Elements in the Human Body

The human body is composed primarily of a few major elements, with several lesser and trace elements playing important roles.

Element	Atomic Symbol	Approx. % Body Mass	Functions
Oxygen	O	65.0	Component of organic and inorganic molecules; needed for ATP production.
Carbon	C	18.5	Component of all organic molecules (carbohydrates, lipids, proteins, nucleic acids).
Hydrogen	H	9.5	Component of all organic molecules; as an ion, influences pH.
Nitrogen	N	3.3	Component of proteins and nucleic acids.

Additional info: Lesser elements include calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium; trace elements include iodine, iron, and others required in minute amounts for enzyme function and other roles.

Identifying Elements

Atomic symbol: One- or two-letter chemical shorthand for each element (e.g., H for hydrogen).
Atomic number: Number of protons in the nucleus; written as a subscript to the left of the atomic symbol (e.g., 3Li).
Mass number: Total number of protons and neutrons; written as a superscript to the left of the atomic symbol (e.g., 7Li).
Isotopes: Structural variations of the same element with the same number of protons but different numbers of neutrons.
Atomic weight: Average of mass numbers of all isotope forms of an atom.

Radioisotopes

Radioisotopes are unstable isotopes that emit radiation as they decay to more stable forms.
Used in biological research and medicine (e.g., imaging, cancer treatment).
All radioactivity can damage living tissue.

Combining Matter: Molecules, Compounds, and Mixtures

Molecules and Compounds

Molecule: Two or more atoms bonded together (e.g., O2).
Compound: Molecule with two or more different kinds of atoms bonded together (e.g., H2O).
Most matter exists as mixtures: two or more components physically intermixed.

The Three Basic Types of Mixtures

Type	Description	Example
Solution	Solute particles are very tiny, do not settle or scatter light	Mineral water
Colloid	Solute particles are larger than in a solution and scatter light; do not settle	Jell-O
Suspension	Very large solute particles, settle out, may scatter light	Blood (plasma and cells)

Difference Between Mixtures and Compounds

Mixtures do not involve chemical bonding; compounds do.
Mixtures can be separated by physical means; compounds require chemical means to separate.
Mixtures can be heterogeneous or homogeneous; compounds are always homogeneous.

Chemical Bonds

Overview

Chemical bonds are energy relationships between electrons of reacting atoms.
Electrons in the outermost shell (valence shell) are involved in bonding.
Three main types: ionic, covalent, and hydrogen bonds.

Role of Electrons in Chemical Bonding

Octet rule: Atoms tend to have 8 electrons in their valence shell (except H and He, which are stable with 2).

Types of Chemical Bonds

Ionic Bonds

Ions are atoms that have gained or lost electrons and become charged.
Ionic bonds involve the transfer of valence electrons from one atom to another, resulting in oppositely charged ions (cations and anions).
Example: NaCl (sodium chloride)

Covalent Bonds

Covalent bonds are formed by sharing two or more valence electrons between atoms.
Single bond: sharing 2 electrons; double bond: sharing 4; triple bond: sharing 6.
Types:
- Nonpolar covalent bonds: Electrons shared equally (e.g., O2, CO2).
- Polar covalent bonds: Electrons shared unequally, creating partial charges (e.g., H2O).

Hydrogen Bonds

Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule (often oxygen or nitrogen).
Important in maintaining the structure of proteins and nucleic acids.

Summary Table: Major Chemical Bond Types

Type	Description	Strength
Covalent bonds	Sharing of pairs of electrons (polar or nonpolar)	Strongest
Ionic bonds	Attraction between two oppositely charged ions	Intermediate
Hydrogen bonds	Attraction between a hydrogen atom and an electronegative atom	Weakest

Chemical Reactions

Chemical Equations

Chemical reactions occur when chemical bonds are formed, rearranged, or broken.
Written as chemical equations showing reactants and products.
Amounts of reactants and products are shown in balanced equations.

Types of Chemical Reactions

Synthesis (Combination) reactions: Atoms or molecules combine to form a larger, more complex molecule. Example:
Decomposition reactions: Molecule is broken down into smaller molecules or atoms. Example:
Exchange (Displacement) reactions: Bonds are both made and broken. Example:
In living systems, many reactions are reduction-oxidation (redox) reactions where electrons are transferred. Example:

Energy Flow in Chemical Reactions

Exergonic reactions: Net release of energy (catabolic and oxidative reactions).
Endergonic reactions: Net absorption of energy (anabolic reactions).

Reversibility of Chemical Reactions

All chemical reactions are theoretically reversible:
Many biological reactions are not easily reversible due to high energy requirements or removal of products.

Rate of Chemical Reactions

Temperature: Increased temperature usually increases reaction rate.
Concentration of reactants: Higher concentration increases rate.
Particle size: Smaller particles increase rate.
Catalysts: Increase reaction rate without being chemically changed or becoming part of the product.

Biochemistry: Inorganic and Organic Compounds

Biochemistry

Study of chemical composition and reactions of living matter.
All chemicals are either organic (contain carbon, usually large, covalently bonded) or inorganic (do not contain carbon; includes water, salts, acids, bases).
Both types are essential for life.

Inorganic Compounds: Water

Water is the most abundant inorganic compound, accounting for 60–80% of the volume of living cells.
Its unique properties make it vital for life.

Properties of Water

High heat capacity: Absorbs and releases heat with little temperature change; prevents sudden temperature changes.
High heat of vaporization: Evaporation requires large amounts of heat; important for cooling (e.g., sweating).
Polar solvent properties: Dissolves and dissociates ionic substances; forms hydration layers around large charged molecules (e.g., proteins); major transport medium in the body.
Reactivity: Necessary for hydrolysis and dehydration synthesis reactions.
Cushioning: Protects organs from physical trauma (e.g., cerebrospinal fluid cushions the nervous system).