Overview: A Chemical Connection to Biology

Biology is deeply connected to the principles of chemistry and physics. Living organisms are composed of matter, which is organized into a hierarchy from atoms to molecules to cells. At each level, new properties emerge that are not present at the previous level. Understanding the chemical context of life is essential for studying biological processes.

Concept 2.1: Matter, Elements, and Compounds

Definition of Matter, Elements, and Compounds

• Matter is anything that takes up space and has mass.

• Elements are substances that cannot be broken down into other substances by chemical reactions.

• Compounds are substances consisting of two or more elements in a fixed ratio. Compounds have properties different from their constituent elements (emergent properties).

Example: Table salt (NaCl) is a compound formed from sodium (a metal) and chlorine (a gas), resulting in an edible substance with new properties.

The Elements of Life

• About 20–25% of the 92 natural elements are essential for life.

• Four elements—carbon, hydrogen, oxygen, and nitrogen—make up about 96% of living matter.

• Other important elements include calcium, phosphorus, potassium, sulfur, sodium, chlorine, and magnesium.

• Trace elements are required in minute quantities (e.g., iron, iodine).

Concept 2.2: Atomic Structure and Properties

Atoms and Subatomic Particles

• An atom is the smallest unit of matter that retains the properties of an element.

• Atoms are composed of protons (positive charge), neutrons (no charge), and electrons (negative charge).

• Protons and neutrons form the atomic nucleus; electrons form a cloud around the nucleus.

• Proton and neutron masses are nearly identical and measured in daltons.

Atomic Number, Mass Number, and Isotopes

• Atomic number: Number of protons in the nucleus (also equals number of electrons in a neutral atom).

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

• Atomic mass: Approximate total mass of an atom (in daltons).

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

Formula:

$\text{Mass number} = \text{Number of protons} + \text{Number of neutrons}$

$\text{Number of neutrons} = \text{Mass number} - \text{Atomic number}$

Energy Levels and Electron Shells

• Energy is the capacity to cause change; potential energy is energy due to position or structure.

• Electrons have different potential energies depending on their distance from the nucleus.

• Electrons occupy electron shells with discrete energy levels.

• Electrons can move between shells by absorbing or releasing energy equal to the difference between shells.

Electron Distribution and the Periodic Table

• The chemical behavior of an atom is determined by the distribution of electrons in its shells, especially the outermost shell (valence shell).

• The periodic table arranges elements by increasing atomic number and electron configuration.

Electron Orbitals

• An orbital is a three-dimensional space where an electron is likely to be found.

• Each shell contains a specific number of orbitals, each holding up to two electrons.

• Atoms interact to complete their valence shells, often by sharing or transferring electrons.

Concept 2.3: Chemical Bonds and Molecular Structure

Covalent Bonds

• Covalent bonds involve the sharing of pairs of valence electrons between atoms.

• Each atom contributes one electron to the shared pair, helping both achieve a full valence shell.

• Covalent bonds can be single (one pair shared) or double (two pairs shared).

• Valence is the number of covalent bonds an atom can form (e.g., H: 1, O: 2, N: 3, C: 4).

Electronegativity and Types of Covalent Bonds

• Electronegativity is an atom’s attraction for electrons in a covalent bond.

• Nonpolar covalent bonds: Electrons are shared equally (e.g., H2, O2).

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

Ionic Bonds and Ionic Compounds

• Ionic bonds form when one atom transfers electrons to another, creating oppositely charged ions (cations and anions).

• Ionic compounds (salts) are formed by the attraction between cations and anions (e.g., NaCl).

• In NaCl, sodium donates an electron to chlorine, resulting in Na+ and Cl– ions.

Weak Chemical Interactions: Hydrogen Bonds and Van der Waals Forces

• Hydrogen bonds are weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom.

• Hydrogen bonds are crucial for stabilizing DNA, protein structure, and water’s properties.

• Van der Waals interactions are weak attractions due to temporary partial charges from uneven electron distribution. Many together can be significant (e.g., gecko adhesion).

Molecular Shape and Function

• A molecule’s shape is determined by the positions of its atoms’ orbitals, often hybridized in covalent bonds.

• Molecular shape is critical for biological recognition and function (e.g., enzyme-substrate, hormone-receptor interactions).

• Water has a bent (V-shaped) structure; methane is tetrahedral.

Summary Table: Types of Chemical Bonds

Additional info: The concepts of atomic structure, chemical bonding, and molecular shape are foundational for understanding all biological molecules and processes, including metabolism, genetics, and cellular structure.