Lecture Outline Bonding Extremes

Name ______________________________________ Pre-lecture exercises. Remember that you may be asked to turn this page in at the beginning of Friday's lecture. To make grading easier, please print and work directly on this page. 1 Chapter 1 of the text discusses covalent and ionic bonding. With this in mind, see if you can figure out the Lewis structures of the following compounds. Show all lone pairs and formal charges clearly. KOH FCH3 NH4Br NaNH2 2 In general chem you learned to use VSEPR to predict the geometries of simple molecules. Ch 1 contains a refresher. Using VSEPR ideas, sketch the following compounds to show their 3dimensional structures as clearly as you can. Note whether the central atom in each is linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral. H3C: – H2S=O (add lone pairs) Lecture outline Bonding extremes — 1. Covalent — e–-pair is shared between two atoms this is what we mean when we draw a "line" to represent a bond 2. Ionic (this perhaps should not be thought of as "bonding" in the usual sense — it's an electrostatic attraction between opposite charges) Ionic "bonding" often occurs between atoms of very different electronegativities (e.g., Na and Cl), but electronegativity alone is not a reliable predictor of when two atoms are likely to form a polarized covalent bond as opposed to existing as a pair of ions. There are two common situations where we really need to recognize the presence of ions (1) Alkali metals — (except that C–Li bonds are really polar-covalent) (2) Octet required — Covalent bonds are polar if the atoms have different electronegativities, though they remain covalent, not ionic, i.e. the e–-pair is still shared, but unequally, resulting in partial charges on the two bonded atoms. Note that in these cases the formal charges are all zero — – these are based on the simple Lewis structure and the approximation that all bonding e s are equally shared. Of course that's badly oversimplified, but ya gotta start somewhere. Denoting partial charges is an attempt to refine the Lewis structure to bring it more in line with reality. e.g., water has polar covalent bonds H–Cl gas is another example. (Note that in water, this is ionized — that's hydrochloric acid; but here, we're talking about H–Cl gas — that's polar covalent.) How important do you think it will be to instantly recognize the difference between HOCH3 and NaOCH3? (hint: no, it's waaaaay more important than that) Geometry — Predict shapes of simple molecules with VSEPR model (valence shell e–-pair repulsion) — e–-pairs repel, so they "want" to get as far apart as possible (although badly oversimplified, this provides reasonable predictions for geometries of 2nd-row atoms in simple bonding situations. And even where it fails it serves as a decent starting point for understanding the actual geometries.) + Ok, let's chunk these to the table below: HCN, CH4, H2O, H2C=NH, CH3 , NH3, CO2 # of atoms + lone pairs around central atom: 2 3 4 examples: "electron geometry": linear trigonal planar tetrahedral bond angles: 180° (about) 120° (about) 109.5° Model-building exercises for you to do after class — In your model set, you will find atoms (the small “jacks”) representing these three geometrical groups — ignore the smooth posts and focus on the ribbed ones — note how these are arranged in space. Use the connectors to put these together and — voilá (that’s French for touché)— you’ve made a molecule. Look at a tetrahedral C. Which two of the following are correct representations of this C? What's wrong with the others? Explain how they've been mangled. C C C C C C Here are some applications of VSEPR geometries that are more applicable to organic chem... Add Hs — Fix this — Cl Cl OH More of these As time permits.... N Molecular polarity — overall polarity affects properties like boiling point, melting point, solubility, etc. Molecular dipole is vector sum of bond dipoles (Ignore lone pairs and non-polar bonds like C–H bonds) Even if a molecule has polar bonds, molecular symmetry can cause cancellation of bond dipoles and result in a molecular dipole moment of zero. Some examples are CO2, BF3, and CCl4.