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10/12/12 Chemical Bonding I: Lewis Theory How Valence Electrons Arrange Themselves to Give Chemical Bonds Basic Ideas from the Previous Chapter... •  Valence electrons are the electrons that parGcipate in chemical bonding. •  ConfiguraGons that result in inert gas cores or half-­‐filled or fully-­‐filled subshells are parGcularly stable. Types of Chemical Bonds Types of Atoms Type of Bond Characteris5cs of Bond Metal & nonmetal Ionic Electrons transferred Nonmetal & nonmetal Covalent Electrons shared Metal & Metal Metallic Electrons pooled 1 10/12/12 Some pointers... •  Ionic bonds: CaGons and anions form and are held together electrostaGcally. •  Covalent bonds: Electrons are shared by atoms (they travel back and forth between the atoms). The most stable configuraGons have the electrons between the nuclei (draw on board, Johnston!). •  Metallic bonds: Electrons move through the enGre crystal laUce of the metal. AVracGon & Repulsion (Explain in detail!) Electrosta5c Repulsion Electrosta5c A:rac5on Some examples of bonding types... 2 10/12/12 We can represent valence electrons as dots... Details—ionic bonds... •  Atoms either lose electrons to form ca1ons •  or •  gain electrons to form anions. •  They try to aVain an inert gas core, if possible! For instance, consider potassium... 3 10/12/12 Octets... •  Except for H and He, most atoms try to get an octet of electrons. (H makes “duets.”) •  This is called the “octet rule.” •  It holds true with just a few excepGons: –  Elements such as Be and B –  Elements with available empty d-­‐orbitals can take on more. –  We shall discuss all this soon when we get to Lewis dot structures! Anions & CaGons Combine to Form a Crystal LaUce A quick aside... •  We can break up the process in easy to visualize steps. •  The energies of the steps are addiGve. •  This is a sneak peak at the First Law of Thermodynamics. •  We shall not say anything more here other than energies add the same way as masses or lengths! 4 10/12/12 The Born-­‐Haber Cycle is the vizualizaGon! LaUce Energy •  This is the main provider of energy to stabilize the crystal! •  However, note that the electron affinity also sneaks in her! Trends in LaUce Energy (Size) Ion Size Examples Metal Chloride LaAce Energy (kJ/mol) LiCl -­‐834 NaCl -­‐788 KCl -­‐701 CsCl -­‐657 5 10/12/12 Trends in LaUce Energy (Charges) Ion Charge Examples Compound LaAce Energy (kJ/mol) NaF -­‐910 CaO -­‐3414 Just for pHun... (The release of laUce energy) •  Let’s look at the reacGon of sodium and chlorine. •  All the steps occur together, of course. •  Be, here, you get to see laUce energy in all its glory! •  Here is the link...
h"p://www.youtube.com/watch?v=Mx5JJWI2aaw This is in accordance with Coulomb’s Law... 6 10/12/12 Some Effects: Ions in SoluGon A liVle demonstraGon... Ionic Compounds by Themselves Don’t Conduct Electricity Ionic Compounds in Aqueous Solu5on DO Conduct! Ionic Compounds are Called “Strong Electrolytes” •  Let’s look at a presentaGon by a kindly old professor (as contrasted to your mean old professor)... •  Here is the link: h"p://www.youtube.com/watch?v=1XWnovm6JLs 7 10/12/12 A Short CriGque of the Ionic Bond Model... •  A great deal can be explained by assuming simple electrostaGc aVracGons/repulsions. •  These ideas allow one to develop methods to calculate laUce energies of many types of ionic compounds. •  The concept of electrolytes is introduced. •  In reality, however, there is no such thing as a pure ionic bond! We shall discuss this shortly. Covalent Bonding •  Here, we shall start with the Lewis model. •  We shall look first a just simple octets (and duets) as done in organic chemistry. •  The model will have to be extended for some compounds of elements in the 3rd row (and further) in the periodic table. •  But, for now, we note that organic chemistry is different than other endeavors→→→ Yes, organic chemistry is different... 8 10/12/12 Drawing Lewis Structures—Some Rules •  Here, we sGck to just octets and duets. •  Extended structures will come later. •  Steps to follow: –  Count the valence electrons –  If you have ions, make allowances for this –  Distribute electrons so that H has duets and other atoms have octets –  ExcepGons to these will be discussed later We construct a few on the board—I shall construct as we go Some terminology 9 10/12/12 Double & Triple Bonds We shall construct some of these on the board... O2 N2 CO C2H2 Acetone Benzene Power of the Lewis Model •  Very good at showing what can and cannot form! •  Predicts direcGonality. •  Works very well with organic compounds. •  ExcepGons are interesGng and will be discussed later! ElectronegaGvity & Polar Bonds •  In all covalent bonds, electrons are shared. •  However, some elements aVract electrons more readily than others. •  This property is called electronega1vity. •  Bonds formed this way are called polar bonds. •  Polar bonds have dipole moments. 10 10/12/12 HF: An example of a polar bond (explain picture as we go) How we can detect this property... Trends in electronegaGvity (back to the good old periodic table) 11 10/12/12 Note the values... •  𝝌 = 4.0 for F •  𝝌 = 3.5 for O •  These two are the most electronegaGve •  𝝌 = 0.7 for Cs •  This is the least electronegaGve. •  Note that 𝝌 = 0 does not happen. Bond PolariGes •  Bond polarity depends on the electronegaGvity differences between the bonded atoms. •  We shall look at three cases on the next three slides. IdenGcal Atoms: Always nonpolar 12 10/12/12 Ionic Bonds: The most polar A very polar covalent bond This table is useful... 13 10/12/12 Another way to look at these... The dipole moment... •  When two atoms are separated by a distance, r, equal and opposite charges of size q are present on each atom. •  We usually represent this as the dipole moment, μ. •  μ is in units of Debyes. •  Details are on the next slide... Details about μ (“mu”) 14 10/12/12 A ClarificaGon... •  Suppose we have two ions with q = +e and -­‐e. •  e = 1.60217653 x 10-­‐19C. •  That is, these are ions of charge +1 and -­‐1. •  Let them be separated by 100 pm. •  Then→→→→... To get μ for a pure ionic bond... Just mulGply the bond length in pm (divided by 100) Gmes the value in the previous slide! This is easy to do (I hope) since I have rounded things! We show this now→→→ We can thus define the % ionic character of a bond! 15 10/12/12 Example from the book... r = 130pm ∴ μionic = 4.803 x 130/100 = 6.24D In the example, μobs = 3.5D So............ Some typical examples... (Note the ∆𝝌 values!) Some more examples... 16 10/12/12 Now the real pHun! WriGng Lewis Structures!!! 1)  Write the correct skeletal structure for the molecule. 2)  Calculate the total number of electrons by summing the valence electrons of each atom. (Be sure to take ions into account!) 3)  Distribute the electrons among the atoms giving octets to all atoms other than H (duet for it). 4)  If any atoms lack an octet, form double or triple bonds as necessary. On-­‐the-­‐board examples (“Explain as we go” mode!) CO2 NH3 NH4+ C2H6 C2H4 C2H2 N2O NO2-­‐ CN-­‐ Resonance & Formal Charge •  SomeGmes there are more than one possible equivalent Lewis structures. •  In this part of the lecture we shall discuss the phenomenon of resonance. 17 10/12/12 What is meant by “equivalent” Two structures are equivalent if they can be converted from one to the other by a simple rotaGon of the enGre molecule or a simple reflecGon of the enGre molecule. The relaGve posiGons of the electrons are lev un changed! A very simple example: Ozone The true structure is a “resonance hybrid” (explain verbally) 18 10/12/12 More about this... •  Neither structure exists independently. •  The true structure is a linear combina1on of the separate structures. •  This is an example electron delocaliza1on. MulGple Equivalent Structures are Possible (e.g., nitrate) Are nonequivalent Lewis structures possible? •  Yes! •  It is possible to draw alternate structures that obey the rules but are not equivalent. •  I shall draw CO2 structures as an example. •  This means that we have to choose a best Lewis structure (or best set of resonance structures). •  What is the key to this? →→→→ 19 10/12/12 Formal Charge! (Rules below) Draw the structure. Be sure to take into account anions and caGons. Start with the atom’s group #. Subtract 1 for each electron completely “owned by the atom.” Subtract 1 for each pair of electrons shared by an atom. This procedure gives the formal charges of each atom. A simple first example: HF Same for hydrogen. Here is the summary... 20 10/12/12 SelecGng the best Lewis structure 1)  The sum of all formal charges on the atoms must equal the total charge of the species. 2)  The best structure always has the differences in the formal charges minimized. 3)  When formal charge cannot be zero, negaGve formal charge should reside on the most electronegaGve atom. 4)  Now, we are ready for CO2! Three CO2 structures! We draw these on the board. All the structures are “legal.” But, one is obviously the best. The minor structures, however, do contribute to the calculated wave funcGon. •  A bonus: We analyze also OCN-­‐. • 
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ExcepGons to the Octet Rule ①  Odd-­‐Electron Species (“free radicals”) ②  Incomplete octets (forced by formal charge) ③  Expanded octets (when extra electrons can be accommodated by d-­‐orbitals) 21 10/12/12 Odd-­‐electron species •  These do occur in nature but tend to be unstable. •  We look at a few examples –  NO –  NO2 (compare this to N2O4) –  The t-­‐butyl free radical Incomplete Octets •  Be & B are strange! •  We explain BeCl2. •  BF3; why F cannot have a double bond and, thus, B is forced into an incomplete octet. •  The very strange and curious case of BH3! Expanded Octets •  Molecules such as AsF5, SF6, PCl5, and many similar such exist. •  How do we accommodate these? •  We put electrons into vacant d-­‐orbitals! •  We shall show how to do this with several examples momentarily. •  In some cases, we shall also have to invoke formal charges and resonance structures! 22 10/12/12 Some Verbal Examples i. 
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SF6 AsF5 SF4 XeF2 The sulfate anion (lots of resonance structures and a BIG surprise!) Bond Energies & Bond Lengths •  The chemical bond can be treated as a disGnct enGty. •  It is a very powerful concept! •  We can assign properGes to bonds. 1) 
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Bond Energies Bond Lengths Bond Angles (next chapter) Bond dipole moments (discussed in passim) Bond Energies •  Chemical reacGons (we have to restrict ourselves to the gas phase here) can be thought as occurring by the breaking of old bonds and the forming of new bonds. •  Typical bond energies are shown in the next slide. •  Note that these bond energies are averages since they can—and do—vary slightly in different molecules. 23 10/12/12 Humongous Table! General Rules •  Bond breaking is endothermic; it takes energy to break a bond! •  Bond forming is exothermic; you get energy back if a bond is formed. •  Usually, if a net process is exothermic, the reacGon is favored. •  The next slide shows the equaGon... The EquaGon 24 10/12/12 Two ways to do this... •  Brute force: Break ALL the bonds in the reactant(s) and then form ALL the bonds in the product(s). •  Finesse: Just look at the actual old bonds broken and the actual new bonds formed. •  Use whichever way works best for you! We shall now give YOU the opportunity to do some by yourself. I shall then explain them! Bond Lengths •  These are defined as the distance between the nuclei of the bonded atoms. •  As with bond energies, these are averages since there are slight variaGons according to the molecular structure. •  The next few slides give some typical values. •  Nowadays, we use pm and the length unit. •  Before that, we used the Ångstrom (1Å = 10-­‐10m). 25 10/12/12 Another humongous table... A quick note on trends... •  For a given atom pair, single bonds are longer than double bonds. •  And, of course, triple bonds are shorter than double bonds. Bonding in Metals •  A common model is the “electron sea model.” •  SomeGmes, this is called a “Fermi gas.” •  The electrons are delocalized over the enter metal chunk. •  Paired electrons can be very far apart! 26