2.4: Simple and Giant Structures

• Non-metals can join together by covalent bonds to form small molecules like chlorine or ammonia, known as simple molecular substances
• Non-metals can also bond covalently to form giant covalent structures like diamond, graphite, or silicon dioxide
• Covalent bonds are strong and require a lot of energy to break apart
• In simple molecular substances, the atoms within each molecule are strongly bonded together, but only weak intermolecular forces need to be broken to melt or boil the substance
• As molecules get larger in simple molecular substances, more intermolecular forces are present, leading to higher melting and boiling points
• Simple molecular substances do not conduct electricity because they lack free electrons and have no electric charge
• Giant covalent structures are made of non-metal atoms bonded together in regular repeating lattices, such as diamond, graphite, and silicon dioxide
• Giant covalent structures have high melting and boiling points due to the need to break all the strong covalent bonds
• Giant covalent structures generally do not conduct electricity because they lack charged particles, except for graphite which is an exception
• Silicon dioxide, also known as silica, is made of silicon and oxygen atoms in a ratio of 1:2 and is the main component of sand
• Simple molecular substances are made of small molecules with weak intermolecular forces, while giant covalent structures have all atoms covalently bonded in regular repeating lattices, making them stronger with higher melting and boiling points
• Graphite is made of 100% carbon
• Carbon atoms in graphite have three covalent bonds per carbon atom
• Carbon atoms in graphite are arranged in hexagons
• Weak intermolecular forces exist between layers of carbon atoms in graphite
• Graphite is soft and slippery due to the ability of layers to slide over each other
• Graphite has a giant covalent structure with a rigid 3D network of atoms
• The weak intermolecular forces between layers allow them to easily slide over each other
• Graphite has one delocalized electron per carbon atom
• Graphite can be used as lubricants and in pencils
• Graphite has a high melting point due to the strong covalent bonds between carbon atoms
• Graphite can carry an electric current
• Delocalized electrons move between layers in graphite, allowing it to conduct electricity
• Graphite: Graphite is a form of carbon consisting of layers of carbon atoms arranged in a hexagonal lattice structure. It is a soft, black, slippery solid, commonly used as a lubricant and in pencil leads.
• Diamond: Diamond is another allotrope of carbon that consists of tetrahedrally bonded carbon atoms forming a three-dimensional network. Diamonds are hard, transparent crystals with excellent thermal conductivity and electrical insulation properties.
• Intermolecular forces: Intermolecular forces are the attractive or repulsive forces between molecules. These forces include hydrogen bonding, dipole-dipole interactions, and van der Waals forces. They play a crucial role in determining the physical properties of substances, such as melting and boiling points, solubility, and viscosity.
• Hydrogen bonding: Hydrogen bonding occurs when a hydrogen atom in one molecule forms a weak attraction to an electron pair on another nearby molecule. This interaction results from the partial positive charge on the hydrogen atom and the partial negative charge on the other molecule. The strength of hydrogen bonding depends on factors like the distance between the two molecules and their relative orientations.
• Lattice: In chemistry and materials science, a lattice refers to the regular, repeating arrangement of atoms, ions, or molecules in a crystalline solid. The lattice structure determines many of the material's properties, including its strength, conductivity, and optical behavior.
• 1. Layers: Layers refer to the parallel sheets or planes of atoms, molecules, or ions arranged within a material. In the context of graphite, the layers consist of hexagonal arrays of carbon atoms bonded covalently to form flat sheets that stack on top of each other. These layers can slide over each other easily, giving graphite its lubricating properties.
• Nonmetal: Nonmetals do not exhibit metallic characteristics and often exist as gases or liquids under standard conditions. Some nonmetals are brittle solids, while others are volatile liquids or gases. Examples of nonmetals include chlorine gas, sulfur, and iodine.
• Melting point: Melting point is the temperature at which a solid substance transitions into a liquid state. Different materials have different melting points due to differences in intermolecular forces and molecular structures.
• Boiling point: Boiling point is the temperature at which a liquid turns into a gas. Like melting point, it varies depending on the type of substance and the strength of intermolecular forces.
• Metalloid: Metalloids have properties intermediate between metals and nonmetals. They may be classified as semiconductors due to their ability to conduct electricity but less efficiently than metals. Silicon and germanium are examples of metalloids.
• Non-metals can join together by covalent bonds to form small molecules like chlorine or ammonia, known as simple molecular substances
• Non-metals can also bond covalently to form giant covalent structures like diamond, graphite, or silicon dioxide
• Covalent bonds are strong, requiring a lot of energy to break apart atoms bonded together
• In simple molecular substances, like chlorine, intermolecular forces need to be broken to melt or boil the substance, not the strong covalent bonds within the molecules
• As molecules get larger in simple molecular substances, more intermolecular forces exist, leading to higher melting and boiling points
• Simple molecular substances do not conduct electricity due to the lack of free electrons or electric charge within the molecules
• Giant covalent structures consist of non-metal atoms bonded in regular repeating lattices, such as diamond, graphite, and silicon dioxide