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

Chemistry and Physiological Reactions

Introduction to Chemistry in Physiology

Chemistry is fundamental to understanding all physiological processes in the human body. Every function, from movement to digestion and neural activity, is governed by chemical interactions.

• The human body is composed of many chemicals.

• Chemistry underlies all physiological reactions, including movement, digestion, heart function, and nervous system activity.

Matter and Energy

Definitions and Properties

Understanding matter and energy is essential for grasping how the body functions at the molecular and cellular levels.

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

• Weight: The mass of an object 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.

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

Forms of Energy

Kinetic and Potential Energy

Energy exists in two primary forms, both of which are crucial in biological systems.

• Kinetic energy: Energy in action (e.g., muscle contraction, nerve impulse transmission).

• Potential energy: Stored (inactive) energy (e.g., energy stored in chemical bonds).

• Energy can be transformed from potential to kinetic energy. For example, stored energy in food molecules is released during metabolism to power cellular activities.

Types of Energy in Biological Systems

• Chemical energy: Stored in the bonds of chemical substances. Example: ATP stores energy used for cellular work.

• Electrical energy: Results from the movement of charged particles. Example: Nerve impulses are electrical signals.

• Mechanical energy: Directly involved in moving matter. Example: Muscle contraction moves body parts.

• Radiant or electromagnetic energy: Travels in waves (e.g., heat, visible light, ultraviolet light, X-rays). Example: Light energy is necessary for vision.

States of Matter

Physical States Relevant to the Body

Matter exists in three states, each with distinct properties important for physiological processes.

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

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

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

Atoms and Elements

Basic Chemical Building Blocks

All matter is composed of elements, which are substances that cannot be broken down into simpler substances by ordinary chemical means.

• Elements: Pure substances made of only one type of atom.

• Four elements make up 96% of the human body: carbon (C), oxygen (O), hydrogen (H), and nitrogen (N).

• Other elements, such as calcium, phosphorus, potassium, and sulfur, are present in smaller amounts but are still essential.

Atoms: Structure and Subatomic Particles

• Atoms: The smallest units of elements that retain the properties of the element.

• Composed of three subatomic particles:

  • Protons (p+): Positively charged, located in the nucleus, mass = 1 atomic mass unit (amu).

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

  • Electrons (e-): Negatively charged, orbit the nucleus, negligible mass (0 amu).

• Atoms are electrically neutral when the number of protons equals the number of electrons.

Identifying Elements

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

• Mass number: Total number of protons and neutrons in the nucleus.

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

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

Electron Shells and Chemical Bonding

Electron Arrangement and Energy Levels

Electrons occupy regions around the nucleus called electron shells or energy levels. The arrangement of electrons determines how atoms interact and bond with each other.

• The first shell holds up to 2 electrons, the second up to 8, and the third up to 8 (for the first 18 elements).

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

• Atoms are most stable when their valence shell is full (the "octet rule").

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 one or more electrons (negatively charged).

• Cation: Atom that loses one or more electrons (positively charged).

• The attraction between oppositely charged ions forms an ionic bond. Example: Sodium chloride (NaCl).

Covalent Bonds

Covalent bonds are formed by the sharing of two or more valence shell electrons between atoms.

• Single bond: Sharing of 2 electrons.

• Double bond: Sharing of 4 electrons.

• Triple bond: Sharing of 6 electrons.

• Nonpolar covalent bonds: Electrons are shared equally (e.g., O2, CO2).

• Polar covalent bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

Hydrogen Bonds

Hydrogen bonds are weak attractions between a hydrogen atom (already covalently bonded to an electronegative atom) and another electronegative atom. They are important in maintaining the structure of large molecules like proteins and DNA, and in giving water its unique properties.

• Hydrogen bonds are not true bonds but are important for molecular interactions.

• They are responsible for water's high surface tension and its role as a universal solvent.

Chemical Reactions

Types of Chemical Reactions

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

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

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

• Exchange (Displacement) reactions: Involve both synthesis and decomposition; bonds are both made and broken. General form: $AB + CD \rightarrow AD + CB$

Redox (Oxidation-Reduction) Reactions

• Involve the transfer of electrons between atoms.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Mnemonic: OIL RIG – Oxidation Is Loss, Reduction Is Gain (of electrons).

Energy Flow in Chemical Reactions

• Exergonic reactions: Release energy; products have less potential energy than reactants. Example: Catabolic reactions.

• Endergonic reactions: Absorb energy; products have more potential energy than reactants. Example: Anabolic reactions.

Reversibility of Chemical Reactions

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

• In biological systems, most reactions proceed in one direction due to energy requirements or removal of products.

Summary Table: Types of Chemical Bonds

Bond Type	Description	Example
Ionic	Transfer of electrons from one atom to another, forming ions	NaCl (sodium chloride)
Covalent (Nonpolar)	Equal sharing of electrons between atoms	O2, CO2
Covalent (Polar)	Unequal sharing of electrons, resulting in partial charges	H2O (water)
Hydrogen	Weak attraction between a hydrogen atom and an electronegative atom	Between water molecules

Additional info: This guide expands on the provided slides by including definitions, examples, and academic context to ensure a comprehensive understanding suitable for Anatomy & Physiology students.