Chapter 2: Chemistry Comes Alive – Foundations for Anatomy & Physiology

Chemistry and Physiological Reactions

Introduction to Chemistry in Physiology

• Body Composition: The human body is composed of many chemicals, and all physiological reactions are underpinned by chemistry.

• Examples of Physiological Reactions: Movement, digestion, heart function, and nervous system activity all depend on chemical processes.

• Branches of Chemistry:

  • Basic Chemistry: Study of matter and energy at the most fundamental level.

  • Biochemistry: Focuses on the chemical processes within and related to living organisms.

2.1 Matter and Energy

Matter

• Definition: Matter is anything that has mass and occupies space.

• Properties:

  • 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

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

• Properties:

  • Does not have mass or occupy space.

  • The greater the work done, the more energy is used up.

Kinetic vs. Potential Energy

• Kinetic Energy: Energy in action (e.g., movement).

• Potential Energy: Stored (inactive) energy that can be released to do work.

• Transformation: Energy can be transformed from potential to kinetic and vice versa.

Forms of Energy

• Chemical Energy: Stored in bonds of chemical substances.

• Electrical Energy: Results from movement of charged particles.

• Mechanical Energy: Directly involved in moving matter.

• Radiant/Electromagnetic Energy: Travels in waves (e.g., heat, visible light, ultraviolet light, X-rays).

2.2 Elements and Atoms

Elements

• Definition: Substances that cannot be broken down into simpler substances by ordinary chemical methods.

• Body Composition:

  • Four elements (carbon, oxygen, hydrogen, nitrogen) make up 96% of the body.

  • 9 elements make up 3.9% of the body.

  • 11 elements make up less than 0.01% of the body.

Atoms

• Definition: Smallest particles of an element with properties of that element.

• Structure:

  • Unique building blocks for each element.

  • Give each element its physical and chemical properties.

Table: Common Elements Composing the Human Body

Element	Atomic Symbol	Approx. % Body Mass	Function
Oxygen	O	65.0	Major component of organic and inorganic molecules; needed for ATP production
Carbon	C	18.5	Primary component of all organic molecules
Hydrogen	H	9.5	Influences pH of body fluids
Nitrogen	N	3.2	Component of proteins and nucleic acids

Atomic Structure

• Subatomic Particles:

  • Protons: Positive charge, 1 atomic mass unit (amu).

  • Neutrons: No charge, 1 amu.

  • Electrons: Negative charge, virtually no weight.

Models of Atomic Structure

• Orbital Model: Electrons as a cloud of negative charge.

• Planetary Model: Electrons as small spheres orbiting the nucleus.

Atomic Number vs. Mass Number

• Atomic Number: Number of protons in the nucleus (written as subscript to the left of atomic symbol).

• Mass Number: Total number of protons and neutrons in the nucleus (written as superscript to the left of atomic symbol).

Isotopes and Radioisotopes

• Isotopes: Structural variations of the same element; same number of protons, different number of neutrons.

• Atomic Weight: Average of mass numbers of all isotope forms of an atom.

• Radioisotopes: Unstable isotopes that decompose to more stable forms, releasing energy (used in imaging and cancer treatments).

2.3 Combining Matter

Molecules and Compounds

• Molecule: General term for 2 or more atoms bonded together.

• Compound: Specific molecule with 2 or more different kinds of atoms bonded together (e.g., C6H12O6).

Types of Mixtures

Type	Description	Example
Solution	Homogeneous mixture; solute particles very tiny, do not settle or scatter light	Mineral water
Colloid	Heterogeneous mixture; solute particles larger, scatter light, do not settle	Jell-O
Suspension	Heterogeneous mixture; solute particles very large, settle out	Blood

Solutions

• Solvent: Substance present in greatest amount (usually water).

• Solute: Substance dissolved in solvent (present in smaller amounts).

• Example: Blood sugar – glucose is solute, blood plasma is solvent.

Solution Concentrations

• Percent of solute in total solution (e.g., 10% salt solution).

• Milligrams per deciliter (mg/dl): e.g., blood glucose ~80 mg/dl.

• Molarity (M): Number of moles of solute per liter of solvent.

  • 1 mole = molecular weight in grams.

  • Avogadro's number: $6.02 \times 10^{23}$ molecules per mole.

Other Mixtures

• Colloids (Emulsions): Heterogeneous, particles do not settle out, can undergo sol-gel transformations (e.g., cytosol).

• Suspensions: Heterogeneous, large particles settle out (e.g., blood cells in plasma).

Difference Between Mixtures and Compounds

• Mixtures do not involve chemical bonding between components; compounds do.

• Mixtures can be separated by physical means; compounds only by breaking chemical bonds.

• Mixtures can be heterogeneous or homogeneous; compounds are only homogeneous.

2.4 Chemical Bonds

Introduction to Chemical Bonds

• Chemical Bonds: "Energy relationships" between electrons of reacting atoms; not actual physical structures.

• Electrons determine if and what type of chemical bond is formed.

Role of Electrons in Chemical Bonding

• Electron Shells: Electrons occupy areas around the nucleus called shells or energy levels.

• Shells have specific capacities:

  • Shell 1: 2 electrons

  • Shell 2: 8 electrons

  • Shell 3: 18 electrons

• Valence Shell: Outermost shell, involved in chemical reactions.

Octet Rule

• Atoms desire 8 electrons in their valence shell (exceptions: H and He want 2).

• Noble gases have full valence shells and are chemically inert.

• Other atoms gain, lose, or share electrons to achieve stability.

Chemically Inert vs. Reactive Elements

• Inert Elements: Full valence shell (e.g., Helium, Neon), not reactive.

• Reactive Elements: Incomplete valence shell (e.g., Hydrogen, Carbon, Oxygen, Sodium), tend to form bonds.

Types of Chemical Bonds

• Ionic Bonds: Transfer of electrons from one atom to another, resulting in ions.

  • Anion: Atom that gained electrons (negative charge).

  • Cation: Atom that lost electrons (positive charge).

  • Attraction of opposite charges forms the bond (e.g., NaCl).

• Covalent Bonds: Sharing of two or more valence electrons.

  • Single bond: 2 electrons shared.

  • Double bond: 4 electrons shared.

  • Triple bond: 6 electrons shared.

  • Types: Polar (unequal sharing, e.g., H2O) and Nonpolar (equal sharing, e.g., CO2).

• Hydrogen Bonds: Weak attraction between electropositive hydrogen and an electronegative atom (e.g., between water molecules).

Table: Comparison of Bond Types

Bond Type	Electron Sharing/Transfer	Example
Ionic	Complete transfer	NaCl
Polar Covalent	Unequal sharing	H2O
Nonpolar Covalent	Equal sharing	CO2

Importance of Hydrogen Bonds

• Responsible for surface tension in water (e.g., water strider walking on water).

• Stabilize structures of proteins and DNA.

Additional info:

• Hydrogen bonds, while weak individually, are crucial for the three-dimensional structure of biological macromolecules.

• Understanding chemical bonds is foundational for later topics such as cell structure, metabolism, and physiology.