Chapter 2: The Chemical Context of Life

Concept 2.1: Matter consists of chemical elements in pure form and in combinations called compounds

All living organisms are composed of matter, which is anything that occupies space and has mass. Matter is made up of elements, and combinations of these elements form compounds with unique properties.

• Matter: Anything that takes up space and has mass.

• Element: A substance that cannot be broken down into other substances by chemical reactions.

• Compound: A substance consisting of two or more elements in a fixed ratio. Compounds exhibit emergent properties distinct from their constituent elements.

• Emergent Properties: New characteristics that arise when elements combine to form compounds.

• Example: Table salt (NaCl) is a compound with properties different from sodium and chlorine.

The Elements of Life

Only a subset of the 92 naturally occurring elements are essential for life. The majority of living matter is composed of a few key elements, while others are required in trace amounts.

• Essential Elements: About 20–25% of elements are required for life.

• Major Elements: Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter.

• Minor Elements: Calcium, phosphorus, potassium, and sulfur constitute most of the remaining 4%.

• Trace Elements: Required in minute quantities for proper biological function.

Concept 2.2: An element’s properties depend on the structure of its atoms

Atoms are the smallest units of matter that retain the properties of an element. The structure of atoms, including their subatomic particles, determines the element’s characteristics.

• Atom: Smallest unit of matter retaining element properties.

• Subatomic Particles: Neutrons (no charge), protons (positive charge), electrons (negative charge).

• Atomic Nucleus: Contains protons and neutrons; electrons form a cloud around the nucleus.

• Dalton: Unit of mass for atoms and subatomic particles.

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in the nucleus.

• Mass Number: Sum of protons and neutrons.

• Atomic Mass: Total mass of an atom, approximated by mass number.

Isotopes

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

• Radioactive Isotopes: Decay spontaneously, emitting particles and energy.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

The Energy Levels of Electrons

• Energy: Capacity to cause change.

• Potential Energy: Energy due to position or structure.

• Electron Shells: Electrons occupy shells with characteristic energy levels and distances from the nucleus.

• Valence Electrons: Electrons in the outermost shell; determine chemical behavior.

• Inert Elements: Elements with full valence shells are chemically unreactive.

Electron Orbitals

• Orbital: Three-dimensional space where an electron is found 90% of the time.

• Shell Structure: Each shell contains a specific number of orbitals; each orbital holds up to 2 electrons.

• Atoms interact to complete their valence shells.

Concept 2.3: The formation and function of molecules and ionic compounds depend on chemical bonding between atoms

Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds that hold atoms together in molecules and compounds.

Covalent Bonds

• Covalent Bond: Sharing of a pair of valence electrons between two atoms.

• Molecule: Two or more atoms held together by covalent bonds.

• Single Bond: Sharing one pair of electrons.

• Double Bond: Sharing two pairs of electrons.

• Valence: Bonding capacity of an atom.

• Electronegativity: Atom’s attraction for electrons in a covalent bond.

• Nonpolar Covalent Bond: Electrons shared equally.

• Polar Covalent Bond: Electrons shared unequally, resulting in partial charges.

• Example: Water (H2O) has polar covalent bonds.

Ionic Bonds

• Ionic Bond: Attraction between oppositely charged ions formed by electron transfer.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic Compound: Compound formed by ionic bonds, such as salts.

• Example: Sodium chloride (NaCl).

Weak Chemical Interactions

• Hydrogen Bond: Attraction between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom (usually oxygen or nitrogen).

• Van der Waals Interactions: Weak attractions between molecules due to transient local partial charges.

• Importance: Weak bonds help maintain the functional form of large biological molecules.

Molecular Shape and Function

• Molecular Shape: Determined by the positions of atom orbitals; crucial for biological function.

• Hybrid Orbitals: s and p orbitals may hybridize, creating specific shapes.

• Recognition: Molecular shape determines how molecules interact and respond to each other.

• Example: Opiates and endorphins have similar shapes and bind the same brain receptors.

Concept 2.4: Chemical reactions make and break chemical bonds

Chemical reactions involve the making and breaking of chemical bonds, transforming reactants into products. All reactions are reversible, and equilibrium is reached when forward and reverse reactions occur at the same rate.

• Reactants: Starting molecules in a chemical reaction.

• Products: Resulting molecules from a chemical reaction.

• Reversibility: Products can become reactants in the reverse reaction.

• Chemical Equilibrium: State where forward and reverse reactions occur at the same rate; concentrations of reactants and products remain constant.

• Example: Photosynthesis: Sunlight powers the conversion of carbon dioxide and water to glucose and oxygen.

  • Equation: