Chemistry and Physiological Reactions

Introduction to Chemistry in Anatomy & Physiology

Chemistry is fundamental to understanding physiological processes in the human body. All bodily functions, from movement to digestion and neural activity, are driven by chemical reactions.

• Chemistry underlies all physiological reactions: Movement, digestion, heart function, and nervous system activity depend on chemical changes.

• The body is composed of chemicals: Every structure and function is based on chemical composition and interaction.

Matter and Energy

Definitions and Properties

Matter and energy are the basic building blocks of the universe and are essential concepts in anatomy and physiology.

• Matter: Anything that has mass and occupies space. Matter can be seen, smelled, and/or felt.

• Weight: Mass plus the effects of gravity.

• Energy: The capacity to do work or put matter into motion. Energy does not have mass nor does it occupy space.

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

Forms of Energy

Kinetic and Potential Energy

Energy exists in two main forms, each playing a role in physiological processes.

• Kinetic energy: Energy in action, such as muscle contraction or nerve impulse transmission.

• Potential energy: Stored (inactive) energy, such as energy stored in chemical bonds or in gradients across membranes.

• Transformation: Energy can be transformed from potential to kinetic, enabling action (e.g., stored energy in ATP is released for cellular work).

Types of Energy in the Human Body

Several forms of energy are important for physiological functions.

• Chemical energy: Stored in bonds of chemical substances; released during chemical reactions (e.g., metabolism).

• Electrical energy: Results from movement of charged particles; essential for nerve impulses and muscle contractions.

• Mechanical energy: Directly involved in moving matter; seen in muscle movement and blood flow.

• Radiant or electromagnetic energy: Travels in waves; includes heat, visible light, ultraviolet light, and X-rays. Important for processes like vision and synthesis of vitamin D.

States of Matter

Physical States Relevant to Physiology

Matter exists in three primary states, each with distinct properties affecting physiological processes.

• Solid: Has a definite shape and volume (e.g., bones).

• Liquid: Has a changeable shape but definite volume (e.g., blood, plasma).

• Gas: Has changeable shape and volume (e.g., oxygen, carbon dioxide in lungs).

Atoms and Elements

Basic Chemical Units

Atoms are the smallest units of matter that retain the properties of an element. Elements are pure substances composed of only one type of atom.

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

• Major elements in the body: Carbon, oxygen, hydrogen, and nitrogen make up 96% of body mass.

• Minor elements: Eleven elements make up less than 0.01% of body mass but are essential for life.

Periodic Table and Element Classification

The periodic table organizes elements by their properties and atomic structure. In human physiology, only a subset of elements is biologically relevant.

• Common elements in the human body: See table below.

Atomic Structure

Subatomic Particles

Atoms are composed of three types of subatomic particles, each with distinct properties.

• Protons (p+): Carry a positive charge; located in the nucleus; mass = 1 amu.

• Neutrons (n0): No charge; located in the nucleus; mass = 1 amu.

• Electrons (e-): Carry a negative charge; orbit the nucleus; mass is negligible (0 amu).

Atomic Number, Mass Number, and Isotopes

Atoms of different elements are distinguished by their number of protons, neutrons, and electrons.

• Atomic number: Number of protons in the nucleus; unique to each element.

• Mass number: Total number of protons and neutrons.

• Isotopes: Variants of the same element with different numbers of neutrons; same atomic number, different mass number.

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

Electron Shells and Chemical Bonding

Electron Arrangement and Energy Levels

Electrons occupy specific energy levels or shells around the nucleus. The arrangement of electrons determines chemical reactivity.

• Electron shells: Regions around the nucleus where electrons are likely to be found.

• Shell capacity: First shell holds 2 electrons; second and third shells hold up to 8 electrons each.

• Valence shell: Outermost electron shell; electrons here have the most potential energy and are involved in chemical reactions.

Octet Rule and Chemical Reactivity

Atoms tend to interact in ways that fill their valence shell, usually with 8 electrons (octet rule).

• Octet rule: Atoms desire 8 electrons in their valence shell (except H and He, which want 2).

• Noble gases: Have full valence shells and are chemically inert.

• Reactive elements: Do not have full valence shells and will gain, lose, or share electrons to achieve stability.

Types of Chemical Bonds

Ionic Bonds

Ionic bonds form when electrons are transferred from one atom to another, resulting in charged particles called ions.

• Anion: Atom that gains electrons; carries a negative charge.

• Cation: Atom that loses electrons; carries a positive charge.

• Ionic compounds: Form crystalline structures (e.g., NaCl).

Covalent Bonds

Covalent bonds are formed by the sharing of electrons between atoms. They can be single, double, or triple bonds depending on the number of shared electrons.

• Nonpolar covalent bonds: Equal sharing of electrons; results in electrically balanced molecules (e.g., CO2).

• Polar covalent bonds: Unequal sharing of electrons; results in molecules with partial charges (e.g., H2O).

• Electronegativity: Atom's ability to attract electrons; higher electronegativity leads to partial negative charge.

Hydrogen Bonds

Hydrogen bonds are weak attractions between a hydrogen atom (attached to an electronegative atom) and another electronegative atom. They are important for the structure of water and biological macromolecules.

• Not true bonds: More like weak magnetic attractions.

• Common in water: Responsible for water's surface tension and its liquid state.

• Intramolecular bonds: Help maintain three-dimensional shapes of large molecules (e.g., proteins, DNA).

Chemical Reactions

Types and Representation

Chemical reactions involve the formation, rearrangement, or breaking of chemical bonds. They are represented by chemical equations.

• Reactants: Substances entering into a reaction.

• Products: Substances resulting from the reaction.

• Balanced equations: Show equal numbers of atoms for each element on both sides.

Example:

Types of Chemical Reactions

• Synthesis (Combination): Atoms or molecules combine to form a larger, more complex molecule.

• Decomposition: Molecule is broken down into smaller molecules or atoms.

• Exchange (Displacement): Bonds are both made and broken; atoms are exchanged between molecules.

Redox (Oxidation-Reduction) Reactions

Redox reactions involve the transfer of electrons between atoms.

• Oxidation: Loss of electrons (OIL: Oxidation Is Loss).

• Reduction: Gain of electrons (RIG: Reduction Is Gain).

Energy Flow in Chemical Reactions

Chemical reactions can either release or absorb energy.

• Exergonic reactions: Net release of energy; products have less potential energy than reactants (e.g., catabolic reactions).

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

Reversibility of Chemical Reactions

Most chemical reactions are theoretically reversible, but in biological systems, many are not due to energy requirements or removal of products.

• Equilibrium: Occurs when neither forward nor reverse reaction is dominant.

• Irreversible reactions: Common in biology due to high energy barriers or product removal.

Additional info:

• Understanding basic chemistry is essential for grasping more advanced topics in anatomy and physiology, such as metabolism, cellular respiration, and molecular biology.