Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

What is Matter?

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

• All organisms are composed of matter.

Elements and Compounds

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

• Compound: A substance consisting of two or more elements in a fixed ratio.

• Compounds have emergent properties that are different from those of their constituent elements.

The Elements of Life

• Of the 92 natural elements, about 20–25% are essential elements required for life.

• Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) make up about 96% of living matter.

• The remaining 4% is mostly calcium (Ca), phosphorus (P), potassium (K), and sulfur (S).

• Trace elements are required in minute quantities.

Case Study: Evolution of Tolerance to Toxic Elements

• Some elements can be toxic to organisms.

• Species can adapt to environments with toxic elements (e.g., some plants adapted to serpentine soils).

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

Atoms and Subatomic Particles

• Atom: The smallest unit of matter that retains the properties of an element.

• Composed of subatomic particles:

  • Neutrons: No electrical charge

  • Protons: Positive charge

  • Electrons: Negative charge

• Neutrons and protons form the atomic nucleus.

• Electrons form a "cloud" around the nucleus.

• Neutron and proton mass are nearly identical and measured in daltons.

• Electron mass is negligible in atomic mass calculations.

Atomic Number and Atomic Mass

• Atomic number: Number of protons in the nucleus.

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

• Atomic mass: Total mass of an atom, approximately equal to the mass number.

Isotopes and Radioactivity

• Atoms of the same element can have different numbers of neutrons, called isotopes.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

• Used as radioactive tracers in medicine (e.g., PET scans to monitor cancer).

Radiometric Dating

• A "parent" isotope decays into a "daughter" isotope at a fixed rate (half-life).

• Scientists use the ratio of isotopes to date fossils and rocks.

• Half-lives can range from seconds to billions of years.

Energy Levels of Electrons

• Energy: The capacity to cause change.

• Potential energy: Energy due to location or structure.

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

• Electrons occupy electron shells with characteristic energy levels.

Electron Distribution and Chemical Properties

• The chemical behavior of an atom is determined by the distribution of electrons in its shells.

• The periodic table arranges elements by electron configuration.

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

• Elements with full valence shells are chemically inert.

Electron Orbitals

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

• Each shell has 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

Chemical Bonds

• Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds.

Covalent Bonds

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

• A molecule consists of two or more atoms held together by covalent bonds.

• Single bond: Sharing one pair of electrons (e.g., H—H).

• Double bond: Sharing two pairs of electrons (e.g., O ═ O).

• Structural formula: Shows arrangement of atoms and bonds (e.g., H—H).

• Molecular formula: Shows the number of each type of atom (e.g., H2).

• Valence: Bonding capacity of an atom.

Electronegativity and Bond Polarity

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

• Nonpolar covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally, leading to partial charges.

Ionic Bonds

• Atoms may strip electrons from their partners, forming ions:

  • Cation: Positively charged ion.

  • Anion: Negatively charged ion.

• Ionic bond: Attraction between cations and anions.

• Ionic compounds (salts): Compounds formed by ionic bonds (e.g., NaCl).

• Salts are often crystalline and dissociate easily in water.

Weak Chemical Interactions

• Most strong bonds in organisms are covalent, but weak bonds (hydrogen bonds, van der Waals interactions) are crucial for biological function.

• Weak bonds are reversible, allowing dynamic interactions.

Hydrogen Bonds

• Form when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (usually O or N).

Van der Waals Interactions

• Occur when transient, uneven electron distribution creates temporary charges, leading to attractions between molecules.

• Collectively, these can be strong (e.g., gecko toe hairs adhering to surfaces).

Molecular Shape and Function

• A molecule’s shape is determined by the positions of its atoms’ orbitals.

• Shape is crucial for molecular recognition and function (e.g., morphine and endorphins bind the same brain receptors due to similar shapes).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Chemical reactions involve the making and breaking of chemical bonds.

• Reactants: Starting molecules.

• Products: Resulting molecules.

Example: Photosynthesis

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

Chemical Equilibrium

• All chemical reactions are reversible.

• At chemical equilibrium, the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant.

Summary Table: Types of Chemical Bonds and Interactions

Additional info: This summary expands on the provided notes with definitions, examples, and a comparative table for clarity and completeness.