Chemistry and Biochemistry Foundations for Anatomy & Physiology

Introduction to Chemistry in Anatomy & Physiology

Chemistry and biochemistry are foundational to understanding physiological processes in the human body. Knowledge of these subjects helps explain how the body maintains homeostasis, supports life, and responds to changes in the environment.

Why This Matters

• Understanding chemistry and biochemistry aids in determining effective solutions for dehydration and fluid loss.

• All physiological reactions—such as movement, digestion, and nerve impulses—depend on chemical processes.

• Chemistry is divided into basic chemistry and biochemistry.

Part 1 – Basic Chemistry

2.1 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.

• Energy does not have mass or take up space.

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

• Kinetic energy: energy in action

• Potential energy: stored (inactive) energy

• Energy can be transformed from potential to kinetic energy.

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)

• Energy can be converted from one form to another (e.g., electrical energy to light energy in a lamp)

2.2 Atoms and Elements

• All matter is composed of elements, which are substances that cannot be broken down by ordinary chemical methods.

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

• Periodic table lists all known elements.

• Smallest particles of an element with properties of that element are atoms.

• Atomic symbol: one- or two-letter shorthand for each element (e.g., C for carbon, O for oxygen).

Structure of Atoms

• Atoms are composed of three subatomic particles:

  • Protons: positive charge (+), mass = 1 amu

  • Neutrons: no charge, mass = 1 amu

  • Electrons: negative charge (-), virtually no mass

• Protons and neutrons are found in the nucleus; electrons orbit around the nucleus.

• Two models for atomic structure:

  • Planetary model: simplified, depicts electrons in orbits (not accurate for real electron behavior)

  • Orbital model: depicts regions (orbitals) where electrons are likely to be found

Identifying Elements

• Elements differ by number of subatomic particles.

• Atomic number: number of protons in nucleus

• Mass number: total number of protons and neutrons

• Isotopes: structural variations of same element (different number of neutrons)

• Atomic weight: average of mass numbers of all isotopes of an atom

Radioisotopes

• Isotopes that decompose to more stable forms, releasing energy (radioactivity).

• Used in biological research and medicine (e.g., cancer treatment, diagnosis).

2.3 Combining Matter

Molecules and Compounds

• Molecule: two or more atoms bonded together

• Compound: specific molecule with two or more different kinds of atoms bonded together

Mixtures

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

• Types of mixtures:

  • Solutions: homogeneous mixtures, particles evenly distributed

  • Colloids: heterogeneous mixtures, particles not evenly distributed

  • Suspensions: heterogeneous mixtures with large, visible solutes that settle out

Solutions

• Solvent: substance present in greatest amount (e.g., water)

• Solute: substance dissolved in solvent (e.g., sugar in water)

• Solutions are usually transparent.

• Concentration of solutions can be expressed as:

  • Percent of solute in total solution

  • Milligrams per deciliter (mg/dl)

  • Molarity (moles of solute per liter of solvent):

    • 1 mole = molecular weight (sum of atomic weights) in grams

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

Colloids

• Also called emulsions; heterogeneous mixtures with large particles that do not settle out.

• Can undergo sol-gel transformations (e.g., cytosol in cells).

Suspensions

• Heterogeneous mixtures with large, visible solutes that settle out (e.g., blood).

Difference Between Mixtures and Compounds

• Mixtures do not involve chemical bonding; compounds do.

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

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

2.4 Chemical Bonds

• Chemical bonds are "energy relationships" between electrons of reacting atoms.

• Electrons occupy regions called electron shells around the nucleus.

• Shells have different energy levels; the outermost shell is the valence shell.

• Octet rule: atoms tend to have 8 electrons in their valence shell (except for the first shell, which is full with 2 electrons).

Types of Chemical Bonds

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

• Covalent bonds: sharing of two or more valence electrons between atoms

  • Nonpolar covalent bonds: equal sharing

  • Polar covalent bonds: unequal sharing, resulting in partial charges (e.g., water)

• Hydrogen bonds: attractive force between electropositive hydrogen and electronegative atom (e.g., water molecules)

2.5 Chemical Reactions

• Chemical reactions occur when chemical bonds are formed, rearranged, or broken.

• Represented as chemical equations:

  • Reactants: substances entering reaction

  • Products: resulting substances

  • Example: $H_2 + O_2 \rightarrow H_2O$

Types of Chemical Reactions

1. Synthesis (combination): atoms or molecules combine to form larger molecules

   • Example: $A + B \rightarrow AB$

2. Decomposition: molecule broken down into smaller molecules or atoms

   • Example: $AB \rightarrow A + B$

3. Exchange: bonds are both made and broken

   • Example: $AB + C \rightarrow AC + B$

Redox (Oxidation-Reduction) Reactions

• Atoms are reduced when they gain electrons and oxidized when they lose electrons.

• Example: $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O + ATP$

Energy Flow in Chemical Reactions

• Exergonic reactions: release energy

• Endergonic reactions: absorb energy

Reversibility and Rate of Chemical Reactions

• All chemical reactions are theoretically reversible.

• Rate can be affected by temperature, concentration, particle size, and catalysts.

• Catalysts: increase reaction rate without being changed (biological catalysts are enzymes).

Part 2 – Biochemistry

2.6 Inorganic Compounds

Water

• Most abundant inorganic compound in living cells (60–80% of cell volume).

• Properties:

  • High heat capacity

  • High heat of vaporization

  • Polar solvent properties

  • Reactivity

  • Cushioning

Classification of Chemicals

• Inorganic compounds: water, salts, acids, and bases (do not contain carbon)

• Organic compounds: carbohydrates, fats, proteins, nucleic acids (contain carbon, usually large, covalently bonded)

Summary Table: Types of Chemical Bonds

Bond Type	Description	Example
Ionic	Transfer of electrons from one atom to another, forming ions	NaCl (table salt)
Covalent	Sharing of electrons between atoms	O2, H2O
Hydrogen	Attraction between hydrogen and electronegative atom	Between water molecules

Additional info:

• These notes provide a foundation for understanding the chemical basis of life, which is essential for further study in anatomy and physiology.