Chemistry Comes Alive: Basic Chemistry for Anatomy & Physiology 2A

Introduction to Chemistry in Anatomy & Physiology

Chemistry is fundamental to understanding physiological processes in the human body. All biological functions, from movement to digestion, are governed by chemical reactions. This section introduces the essential chemical principles that underpin anatomy and physiology.

Matter and Energy

Definition of 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.

Definition of Energy

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

• Energy 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.

• Potential energy: Stored (inactive) energy.

• Energy can be transformed from potential to kinetic energy (e.g., stored energy released to cause movement).

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 Conversions

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

• Energy conversion is inefficient; some energy is "lost" as heat, which is partly unusable.

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: carbon, oxygen, hydrogen, and nitrogen.

• Other elements are present in smaller amounts.

• The periodic table lists all known elements.

Atoms

• Atoms are the unique building blocks of each element.

• They are the smallest particles of an element that retain the properties of that element.

• Atoms give each element its particular physical and chemical properties.

Chemical Symbols

• Each element is represented by a one- or two-letter symbol (e.g., O for oxygen, C for carbon).

• Some symbols are derived from Latin names (e.g., Na for sodium, K for potassium).

Structure of Atoms

Subatomic Particles

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

• Neutrons: No charge (0), 1 amu.

• Electrons: Negative charge (-), virtually no weight (0 amu).

Atomic Models

• Planetary model: Electrons move in fixed orbits around the nucleus (simplified and outdated, but useful for illustrations).

• Orbital model: Electrons are found in regions of probability called orbitals, forming an electron cloud around the nucleus.

Atomic Number and Mass Number

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

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

Isotopes and Atomic Weight

• 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

• Radioisotopes are unstable isotopes that decompose to more stable forms, emitting radiation (radioactivity).

• Used in medical diagnosis and treatment, but can also cause tissue damage.

Combining Matter: Molecules, Compounds, and Mixtures

Molecules and Compounds

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

• Compound: A molecule containing two or more different kinds of atoms (e.g., H2O).

Mixtures

• Mixtures are physical combinations of two or more components.

• Three basic types: solutions, colloids, and suspensions.

Types of Mixtures

• Solutions: Homogeneous mixtures; solute particles are evenly distributed and do not settle out (e.g., saline solution).

• Colloids (emulsions): Heterogeneous mixtures; larger particles that do not settle out, often cloudy (e.g., cytosol).

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

Concentration of Solutions

• Percent (%): Parts of solute per 100 parts of solution.

• Milligrams per deciliter (mg/dL): Mass of solute per 100 mL of solution.

• Molarity (M): Moles of solute per liter of solution. 1 mole = molecular weight in grams; contains $6.02 \times 10^{23}$ molecules (Avogadro's number).

Differences Between Mixtures and Compounds

• Mixtures do not involve chemical bonding; compounds do.

• Mixtures can be separated by physical means; compounds require chemical reactions to separate.

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

Chemical Bonds

Role of Electrons in Chemical Bonding

• Electrons occupy energy levels called electron shells around the nucleus.

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

• Atoms are most stable when their valence shell is full (usually 8 electrons; octet rule).

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

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

Types of Chemical Bonds

• Ionic bonds: Formed by transfer of electrons from one atom to another, creating ions (cations and anions). Example: NaCl.

• Covalent bonds: Formed by sharing electrons between atoms. Can be single, double, or triple bonds.

• Nonpolar covalent bonds: Equal sharing of electrons (e.g., O2, CO2).

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

• Hydrogen bonds: Weak attractions between a hydrogen atom (already covalently bonded to an electronegative atom) and another electronegative atom. Important in water and in stabilizing large molecules like proteins and DNA.

Chemical Reactions

Chemical Equations

• Chemical reactions involve the formation, rearrangement, or breaking of chemical bonds.

• Reactants: Substances entering into the reaction.

• Products: Substances produced by the reaction.

• Equations are balanced to show the conservation of mass.

Types of Chemical Reactions

• Synthesis (combination) reactions: Atoms or molecules combine to form a larger, more complex molecule. $A + B \rightarrow AB$

• Decomposition reactions: A molecule is broken down into smaller molecules or atoms. $AB \rightarrow A + B$

• Exchange (displacement) reactions: Involve both synthesis and decomposition. $AB + C \rightarrow AC + B$

• Redox (oxidation-reduction) reactions: Involve the transfer of electrons between atoms. Atoms are reduced when they gain electrons and oxidized when they lose electrons.

Energy Flow in Chemical Reactions

• Exergonic reactions: Release energy; products have less potential energy than reactants (e.g., catabolic and oxidative reactions).

• Endergonic reactions: Absorb energy; products have more potential energy than reactants (e.g., anabolic reactions).

Reversibility of Chemical Reactions

• Many chemical reactions are theoretically reversible: $A + B \rightleftharpoons AB$

• In biological systems, most reactions are not easily reversible due to energy requirements or removal of products.

Factors Affecting the Rate of Chemical Reactions

• Temperature: Higher temperature increases reaction rate.

• Concentration: Higher concentration of reactants increases reaction rate.

• Particle size: Smaller particles increase reaction rate.

• Catalysts: Substances that increase the rate of reaction without being consumed. Enzymes are biological catalysts.

Additional info: This guide provides foundational chemistry concepts essential for understanding physiological processes in the human body, as required for Anatomy & Physiology students.