Atoms, Elements and Compounds 1. Basic Definitions Elements, Compounds and Mixtures Element: Pure substance consisting of only one type of atom. Compound: Pure substance consisting of two or more different elements chemically bonded together. Mixture: Two or more substances physically blended but not chemically bonded. 2. Atomic Structure The Atom Structure: Nucleus (containing protons and neutrons) surrounded by electrons in shells. Subatomic Particles: Particle Relative Charge Relative Mass Location Proton +1 1 Nucleus Neutron 0 1 Nucleus Electron -1 1/1840 Shells Atomic and Mass Numbers Proton Number (Atomic Number): Number of protons in the nucleus of an atom. Mass Number (Nucleon Number): Total number of protons and neutrons in the nucleus. Electronic Configuration Electrons occupy shells around the nucleus. Configuration for proton numbers 1-20: 2, 8, 8, 2. Periodic Table Relation: Group Number: Equals the number of electrons in the outer shell (Group I-VII). Period Number: Equals the number of occupied shells. Group VIII (Noble Gases): Have full outer shells, making them unreactive. Isotopes Definition: Atoms of the same element with the same number of protons but different numbers of neutrons. Chemical Properties: Isotopes have identical chemical properties because they have the same electronic configuration. Relative Atomic Mass: Calculated using the weighted average of isotope abundances: $\text{Relative Atomic Mass} = \frac{\sum (\text{isotope mass} \times \text{abundance})}{\text{total abundance}}$ 3. Ions and Bonding Ion Formation Cations: Positive ions formed when an atom loses electrons. Anions: Negative ions formed when an atom gains electrons. Ionic Bonding Definition: Strong electrostatic attraction between oppositely charged ions. Formation: Typically between Group I (metal) and Group VII (non-metal). Formation: Occurs between any metallic and non-metallic elements. Representation: Use dot-and-cross diagrams to show electron transfer. Structure: Exists as a giant ionic lattice. Properties: High melting and boiling points due to strong electrostatic forces throughout the lattice. Conduct electricity when molten or aqueous (ions are free to move). Poor conductors when solid (ions fixed in position). Covalent Bonding Definition: A pair of shared electrons between two atoms to achieve noble gas configurations. Simple Molecules: Examples: $\text{H}_2$, $\text{Cl}_2$, $\text{H}_2\text{O}$, $\text{CH}_4$, $\text{NH}_3$, $\text{HCl}$. Representation: Use dot-and-cross diagrams. Simple Molecules: Examples: $\text{CH}_3\text{OH}$, $\text{C}_2\text{H}_4$, $\text{O}_2$, $\text{CO}_2$, $\text{N}_2$. Properties: Low melting and boiling points due to weak intermolecular forces, despite strong covalent bonds within the molecule. Poor electrical conductivity (no free ions or electrons). Giant Covalent Structures Diamond: Each carbon atom bonded to four others in a tetrahedral lattice. Extremely hard, used in cutting tools. Graphite: Each carbon atom bonded to three others in hexagonal layers. Layers slide (lubricant) and delocalised electrons allow conductivity (electrode). Silicon(IV) Oxide ($\text{SiO}_2$): Structure similar to diamond (Si bonded to 4 O atoms). High melting point and hard. Metallic Bonding Definition: Electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons. Properties: Electrical Conductivity: Delocalised electrons are free to move and carry charge. Malleability and Ductility: Layers of ions can slide over each other without breaking the metallic bond.