Enduring Understanding 2.C.2, 2D: Ionic and Metallic Bonding
- Ionic Bonding results from the net Coulombic attraction of positively and negatively charged anions packed together in a regular crystal lattice.
- Coulombic force is proportional to charge, so higher charges result in stronger interactions.
- Coulombic force is inversely proportional to (the square of) distance, so smaller ions that can pack more closely together will have stronger interactions.
- Example: Which of the following would have the more exothermic lattice energy, NaF or KBr?
- NaF would have a more exothermic lattice energy (-922 kJ/mol vs. -688 kJ/mol) because it is composed of smaller ions that can pack more tightly together.
- In ionic compounds, electrons are tightly held by the ions, and the ions cannot move translationally relative to each other.
- This explains many properties of ionic solids. They are hard and brittle, they are not malleable or ductile (i.e. cannot be shaped without cracking/breaking), and they do not conduct electricity.
- Metallic bonding describes a lattice of positively charged ions, surrounded by a mobile 'sea' of valence electrons. In contrast with ionic bonding, the valence orbitals are delocalized over the entire metal lattice, electrons are free to move and are not associated with individual cations.
- The 'free valence electrons' model explains several properties of metals: they conduct electricity, are malleable and ductile (can have their shape changed without breaking) and are not volatile.
- As mentioned above, they type of bonding observed in the solid state determine the properties of solids.
- Molecular solids:
- Consist of nonmetals bonded covalently to each other.
- Are composed of distinct molecules of covalently bonded atoms, that are attracted to each other by relatively weak (London and dipole) forces
- Usually have low melting and boiling points.
- Electrons are tightly bound in well-defined bonds, so they do not conduct electricity as a solid or in solution.
- Examples: CO2, I2, S8
- Ionic solids:
- have low vapor pressure (strong Coulombic attractions between ions)
- are brittle and cannot be deformed (ions in lattice not free to slide over each other)
- Solids do not conduct electricity (electrons are tightly bound to ions)
- In aqueous solution, or when melted to a liquid, ionic compounds do conduct electricity (ions are now free to move). This is often an identifying feature of an ionic solid.
- Tend to be soluble in polar solvents and insoluble in nonpolar solvents.
- Examples: NaCl, Fe2O3
- Metallic Solids:
- Conduct heat and electricity well (electrons are delocalized and free to move)
- Are malleable and ductile (the cations are more free to move relative to each other than in ionic solids)
- Are shiny ('lustrous') and good conductors of heat.
- Examples: all pure metals: Na, Fe, Al, Au, Ag...
- Metals can also exist as mixtures called alloys, where atoms either substitute for the metal atoms in the lattice, or fill in empty spaces in the lattice. Different atoms in the metal lattice can change the properties of the pure metal.
- Examples: Carbon atoms (about 2%) mixed with iron form steel, which is much stronger (less malleable) than pure iron. Brass is another alloy, composed of 70% copper and 30% zinc.
- Network Covalent solids form large 2D or 3D networks of covalently bonded atoms.
- They are only formed by nonmetals, that can form covalent bonds
- Because all atoms are covalently bonded, they have extremely high melting points.
- Three-dimensional network covalent solids are extremely hard and brittle. (e.g. diamond)
- Two-dimensional network covalent solids have layers than can slide past each other more easily (e.g. graphite)
- Examples: Diamond, graphite (both carbon), silicon dioxide, silicon carbide.
- Sample Question: An unknown substance is a colorless crystalline solid. It melts at 801°C, its crystals are brittle and break, and it dissolves in water to form a conducting solution. Which of the following is the most likely formula for this compound? PCl5, NaCl, Cu, SiC?
- Answer: NaCl. The properties indicate the compound must be an ionic solid; the other three choices are not ionic solids.
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Related Links: Chemistry Chemistry Quizzes AP Chemistry Notes Covalent Bonding |
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