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Unit 2
7–9% of examThis unit is about how atoms hold together. You'll compare ionic, covalent and metallic bonding, draw Lewis diagrams, and use them with VSEPR theory to predict a molecule's shape, bond angles and polarity. Those structures explain the properties you study in Unit 3.
Longer videos that cover the whole unit. Good for a first pass or a final review.
Bonds fall on a range from nonpolar covalent (electrons shared about equally) through polar covalent to ionic, depending mostly on the difference in electronegativity between the two atoms. As a rule of thumb, a metal bonded to a nonmetal is ionic, two nonmetals are covalent, and in a metal the valence electrons are spread out (delocalized) across the whole solid.
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A graph of potential energy against the distance between two nuclei dips to a lowest point: that distance is the bond length, and the depth of the dip is the bond energy needed to pull the atoms apart. Between the same two atoms, double and triple bonds are shorter and stronger than single bonds, and Coulomb's law shows that ions with bigger charges and smaller sizes attract each other more strongly.
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In an ionic solid, cations and anions pack into a repeating 3-D pattern (a lattice) that keeps opposite charges close and like charges apart. You should be able to sketch this as a particle model with sensible ion sizes; in NaCl, for example, each Cl⁻ ion is larger than each Na⁺ ion.
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Metals are modeled as positive metal ions in a 'sea' of delocalized valence electrons. In a substitutional alloy such as brass, atoms of similar size (zinc) take the place of some copper atoms; in an interstitial alloy such as steel, much smaller carbon atoms fit into the gaps between the larger iron atoms.
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A Lewis diagram shows a molecule's valence electrons as bonding pairs and lone pairs. To draw one, count the total valence electrons, connect the atoms with single bonds, then place the remaining electrons so atoms usually end up with an octet, using double or triple bonds when you run short.
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When you can draw more than one equivalent Lewis diagram, as for ozone (O₃) or nitrate (NO₃⁻), the real particle is a blend called a resonance hybrid, with bonds partway between single and double. When several different diagrams are possible, formal charge helps you pick the best: choose the one with formal charges closest to zero, with any negative formal charge on the more electronegative atom.
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VSEPR theory says the electron domains around a central atom spread as far apart as possible, which predicts shapes such as linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal pyramidal and bent. From the shape you can decide whether a molecule is polar, name the central atom's hybridization (sp, sp² or sp³), and count sigma and pi bonds: a single bond is one sigma bond, a double bond is one sigma plus one pi, and a triple bond is one sigma plus two pi.
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