Chapter 2A: Chemistry Comes Alive

Introduction to Chemistry in Physiology

Chemistry forms the foundation of all physiological processes in the human body, including movement, digestion, heart function, and nervous system activity. Understanding the chemical basis of life is essential for comprehending how the body operates at the molecular level.

Atoms and Elements

Elements: Basic Building Blocks

All matter is composed of elements, which are substances that cannot be broken down into simpler substances by ordinary chemical methods. Four elements—Carbon (C), Oxygen (O), Hydrogen (H), and Nitrogen (N)—make up 96% of the human body. Elements are represented by atomic symbols, often derived from Latin names (e.g., Na for sodium, K for potassium).

• Atomic Symbol: One or two letters representing an element.

• Major Elements: Oxygen, Carbon, Hydrogen, Nitrogen.

• Lesser Elements: Calcium, Phosphorus, Potassium, Sulfur, Sodium, Chlorine, Magnesium, Iodine, Iron.

• Trace Elements: Required in tiny amounts; often part of enzymes or necessary for enzyme activation.

Structure of Atoms

Subatomic Particles

Atoms are composed of three subatomic particles:

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

• Neutrons: No charge (0), 1 amu.

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

Protons and neutrons are located in the nucleus, while electrons orbit around the nucleus. The number of protons equals the number of electrons in a neutral atom.

Atomic Structure of Smallest Atoms

Different elements contain different numbers of subatomic particles. For example:

• Hydrogen: 1 proton, 0 neutrons, 1 electron

• Helium: 2 protons, 2 neutrons, 2 electrons

• Lithium: 3 protons, 4 neutrons, 3 electrons

Isotopes

Isotopes are structural variations of the same element, differing in the number of neutrons. For example, hydrogen has three isotopes: protium (1H), deuterium (2H), and tritium (3H). Tritium is radioactive and used in research.

Identifying Elements

Atomic Number and Mass Number

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

• Mass Number: Total number of protons and neutrons in the nucleus; represents the atom's mass.

Radioisotopes

Definition and Applications

Radioisotopes are isotopes that decompose to more stable forms, emitting particles and energy (radioactivity). They are used in biological research, medical diagnosis, and treatment, but can also damage living tissue.

Combining Matter: Molecules and Compounds

Molecules and Compounds

Most atoms combine to form molecules (two or more atoms bonded together) or compounds (molecules with two or more different kinds of atoms). For example, glucose (C6H12O6) is a compound, while O2 is a molecule.

Solutions

Components of Solutions

• Solute: Substance dissolved in a solvent (can be gas, liquid, or solid).

• Solvent: Liquid in which the solute is dissolved; water is the main solvent in the body.

• Solution: Homogeneous mixture of solute and solvent.

Example: In a sodium chloride solution, sodium chloride is the solute and water is the solvent.

Chemical Bonds

Role of Electrons in Chemical Bonding

Electrons occupy energy levels (shells) around the nucleus. The shell closest to the nucleus is filled first. The outermost shell is called the valence shell, and its electrons are involved in chemical reactions.

• Shell 1: 2 electrons

• Shell 2: up to 8 electrons

• Shell 3: up to 18 electrons (8 is stable)

Octet Rule

Atoms are stable when their outermost electron shell is full (2 for small atoms, 8 for others). Inert atoms (e.g., helium, neon) have full valence shells and do not form bonds. Most atoms form bonds to achieve stability.

Types of Chemical Bonds

Ionic Bonds

Ionic bonds are formed when electrons are transferred from one atom to another, resulting in charged ions. Anions have a negative charge (gain electrons), while cations have a positive charge (lose electrons). The attraction between oppositely charged ions forms an ionic bond.

Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms. There are two main types:

• Nonpolar Covalent Bonds: Equal sharing of electrons; molecule is electrically neutral (e.g., methane, carbon dioxide).

• Polar Covalent Bonds: Unequal sharing of electrons; results in electrically polar molecules (e.g., water).

Hydrogen Bonds

Hydrogen bonds are weak bonds formed between an electropositive hydrogen atom and an electronegative atom of another molecule. They are important for:

• Properties of water

• Maintaining protein structure

• Holding DNA strands together

Chemical Equations

Representation of Chemical Reactions

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

• Reactants: Substances entering the reaction

• Products: Resulting chemical end products

Example equations:

• (Hydrogen gas)

• (Methane)

Subscripts indicate the number of atoms joined by bonds; prefixes indicate the number of unjoined atoms.

Patterns of Chemical Reactions

Synthesis and Decomposition

• Synthesis Reaction: Atoms or molecules combine to form a larger molecule; energy is absorbed (anabolic).

• Decomposition Reaction: Molecule is broken down into smaller components; energy is released (catabolic).

General equations:

• Synthesis:

• Decomposition:

Biochemistry: Organic and Inorganic Compounds

Organic Compounds

Organic compounds contain carbon (and hydrogen) and are mostly covalently bonded. Examples include carbohydrates, proteins, and fats/lipids.

Inorganic Compounds

Inorganic compounds lack carbon and include water, salts, acids, and bases. Water is the main solvent in the body, making up 60–80% of cell volume.

• Examples of organic molecules: carbohydrates, proteins, fats/lipids

• Examples of inorganic molecules: sodium chloride, water, oxygen