Atoms and Molecules: 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 elements form compounds with unique properties.

• 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, exhibiting emergent properties distinct from its constituent elements.

• Emergent Properties: Characteristics of a compound that are different from those of its individual elements.

Example: Table salt (NaCl) is a compound formed from sodium (a reactive metal) and chlorine (a toxic gas), but together they form a safe, edible substance.

The Elements of Life

Of the 92 naturally occurring elements, only a small fraction are essential for life. Four elements—carbon, hydrogen, oxygen, and nitrogen—make up the majority of living matter.

• Essential Elements: Elements required for an organism to survive, grow, and reproduce.

• Trace Elements: Elements required in minute quantities, such as iron and iodine.

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

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

Case Study: Evolution of Tolerance to Toxic Elements

Some elements are toxic, but certain species can adapt to environments containing these elements. For example, plants in serpentine soils have evolved mechanisms to tolerate high levels of toxic metals.

• Adaptation: Evolutionary process by which organisms become better suited to their environment.

• Serpentine Soil: Soil rich in magnesium and heavy metals, toxic to most plants.

Concept 2.2: An Element’s Properties Depend on the Structure of Its Atoms

Atomic Structure and Subatomic Particles

Atoms are the smallest units of matter that retain the properties of an element. They are composed of subatomic particles: protons, neutrons, and electrons.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle orbiting the nucleus.

• Dalton: Unit of atomic mass (1 dalton ≈ mass of 1 proton or 1 neutron).

Example: Helium atom contains 2 protons, 2 neutrons, and 2 electrons.

Atomic Number and Atomic Mass

The atomic number is the number of protons in an atom’s nucleus, while the mass number is the sum of protons and neutrons. Atomic mass is the total mass of an atom, approximated by the mass number.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

• Atomic Mass: Total mass, measured in daltons.

Formula: (where N = number of neutrons)

Isotopes and Radioactivity

Isotopes are atoms of the same element with different numbers of neutrons. Some isotopes are radioactive and decay spontaneously, emitting energy and particles.

• Isotope: Variant of an element with different neutron number.

• Radioactive Isotope: Unstable isotope that decays, used in medical diagnostics and radiometric dating.

• Half-life: Time required for half the atoms of a radioactive isotope to decay.

Example: PET scans use radioactive tracers to monitor metabolism in the body.

Radiometric Dating

Scientists use the decay of radioactive isotopes to date fossils and rocks. The ratio of parent to daughter isotopes and the known half-life allow calculation of the sample’s age.

• Radiometric Dating: Method for determining the age of objects using isotope ratios.

• Half-life Equation: (where = remaining atoms, = initial atoms, = time elapsed, = half-life)

Concept 2.2: The Energy Levels of Electrons

Electron Shells and Potential Energy

Electrons occupy specific energy levels called shells. The potential energy of an electron depends on its distance from the nucleus; electrons in higher shells have more energy.

• Potential Energy: Energy due to position or structure.

• Electron Shell: Energy level where electrons reside.

• Valence Shell: Outermost electron shell, determines chemical reactivity.

Electron Distribution and Chemical Properties

The arrangement of electrons in shells determines an atom’s chemical behavior. The periodic table reflects electron distribution, with elements in each row having sequentially filled shells.

• Valence Electrons: Electrons in the outermost shell.

• Chemically Inert: Atoms with full valence shells (e.g., noble gases).

Electron Orbitals

Orbitals are three-dimensional regions where electrons are likely to be found. Each shell contains a specific number of orbitals, and each orbital holds up to two electrons.

• Orbital: 3D space where an electron is found 90% of the time.

• Shell Structure: First shell: 1 orbital (2 electrons); Second shell: 4 orbitals (8 electrons).

Concept 2.3: Chemical Bonding Between Atoms

Covalent Bonds

Covalent bonds involve the sharing of valence electrons between atoms. These bonds are fundamental to the structure of biological molecules.

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

• Single Bond: One pair of shared electrons.

• Double Bond: Two pairs of shared electrons.

• Valence: Bonding capacity of an atom.

Electronegativity and Bond Polarity

Electronegativity is an atom’s ability to attract electrons. If atoms share electrons unequally, the bond is polar; if equally, it is nonpolar.

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

• Polar Covalent Bond: Unequal sharing, resulting in partial charges.

• Nonpolar Covalent Bond: Equal sharing of electrons.

Ionic Bonds

Ionic bonds form when atoms transfer electrons, resulting in charged ions that attract each other. Ionic compounds, or salts, are stable and often form crystals.

• Ionic Bond: Attraction between oppositely charged ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic Compound: Compound formed by ionic bonds (e.g., NaCl).

Weak Chemical Interactions

Hydrogen Bonds

Hydrogen bonds occur when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom. These bonds are crucial for the structure of DNA and proteins.

• Hydrogen Bond: Weak attraction between hydrogen and electronegative atoms (O, N).

Van der Waals Interactions

Van der Waals interactions are weak attractions between molecules due to temporary charge differences. Collectively, these can be significant, as seen in gecko adhesion.

• Van der Waals Interaction: Weak, transient attraction between molecules.

Molecular Shape and Function

Molecular Shape

The shape of a molecule is determined by the positions of its atoms’ orbitals. Shape is critical for biological function, such as enzyme-substrate recognition.

• Hybrid Orbitals: Orbitals formed by mixing s and p orbitals, creating specific shapes.

• Molecular Mimicry: Molecules with similar shapes can bind the same biological receptors.

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

Chemical reactions involve the making and breaking of chemical bonds. Reactants are transformed into products, and all reactions are reversible.

• Reactant: Starting molecule in a reaction.

• Product: Resulting molecule after a reaction.

• Reversible Reaction: Reaction that can proceed in both directions.

Chemical Equilibrium

Equilibrium is reached when the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant.

• Chemical Equilibrium: State where reaction rates are equal and concentrations are stable.

Photosynthesis: An Example of a Chemical Reaction

Photosynthesis is a key biological reaction powered by sunlight, converting carbon dioxide and water into glucose and oxygen.

• Photosynthesis Equation:

Example: Plants use photosynthesis to produce food and oxygen for life on Earth.