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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 participate in chemical bonding. • Configurations that result in inert gas cores or half-filled or fully-filled subshells are particularly stable. Types of Chemical Bonds Types of Atoms Type of Bond Characteristics of Bond Metal & nonmetal Ionic Electrons transferred Nonmetal & nonmetal Covalent Electrons shared Metal & Metal Metallic Electrons pooled Some pointers... • Ionic bonds: Cations and anions form and are held together electrostatically. • Covalent bonds: Electrons are shared by atoms (they travel back and forth between the atoms). The most stable configurations have the electrons between the nuclei (draw on board, Johnston!). • Metallic bonds: Electrons move through the entire crystal lattice of the metal. Attraction & Repulsion (Explain in detail!) Electrostatic Repulsion Electrostatic Attraction Some examples of bonding types... We can represent valence electrons as dots... Details—ionic bonds... • Atoms either lose electrons to form cations • or • gain electrons to form anions. • They try to attain an inert gas core, if possible! For instance, consider potassium... 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 exceptions: – 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 & Cations Combine to Form a Crystal Lattice A quick aside... • We can break up the process in easy to visualize steps. • The energies of the steps are additive. • 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! The Born-Haber Cycle is the vizualization! Lattice Energy • This is the main provider of energy to stabilize the crystal! • However, note that the electron affinity also sneaks in her! Trends in Lattice Energy (Size) Ion Size Examples Metal Chloride Lattice Energy (kJ/mol) LiCl -834 NaCl -788 KCl -701 CsCl -657 Trends in Lattice Energy (Charges) Ion Charge Examples Compound Lattice Energy (kJ/mol) NaF -910 CaO -3414 Just for pHun... (The release of lattice energy) • Let’s look at the reaction of sodium and chlorine. • All the steps occur together, of course. • Be, here, you get to see lattice energy in all its glory! • Here is the link...http://www.youtube.com/watch?v=Mx5JJWI2aaw This is in accordance with Coulomb’s Law... Some Effects: Ions in Solution A little demonstration... Ionic Compounds by Themselves Don’t Conduct Electricity Ionic Compounds in Aqueous Solution DO Conduct! Ionic Compounds are Called “Strong Electrolytes” • Let’s look at a presentation by a kindly old professor (as contrasted to your mean old professor)... • Here is the link: http://www.youtube.com/watch?v=1XWnovm6JLs A Short Critique of the Ionic Bond Model... • A great deal can be explained by assuming simple electrostatic attractions/repulsions. • These ideas allow one to develop methods to calculate lattice 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... Drawing Lewis Structures—Some Rules • Here, we stick 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 – Exceptions to these will be discussed later We construct a few on the board—I shall construct as we go Some terminology 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 directionality. • Works very well with organic compounds. • Exceptions are interesting and will be discussed later! Electronegativity & Polar Bonds • In all covalent bonds, electrons are shared. • However, some elements attract electrons more readily than others. • This property is called electronegativity. • Bonds formed this way are called polar bonds. • Polar bonds have dipole moments. HF: An example of a polar bond (explain picture as we go) How we can detect this property... Trends in electronegativity (back to the good old periodic table) Note the values... • 𝝌 = 4.0 for F • 𝝌 = 3.5 for O • These two are the most electronegative • 𝝌 = 0.7 for Cs • This is the least electronegative. • Note that 𝝌 = 0 does not happen. Bond Polarities • Bond polarity depends on the electronegativity differences between the bonded atoms. • We shall look at three cases on the next three slides. Identical Atoms: Always nonpolar Ionic Bonds: The most polar A very polar covalent bond This table is useful... 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”) A Clarification... • 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 multiply the bond length in pm (divided by 100) times 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! 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... Now the real pHun! Writing 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 NO2CN- Resonance & Formal Charge • Sometimes there are more than one possible equivalent Lewis structures. • In this part of the lecture we shall discuss the phenomenon of resonance. What is meant by “equivalent” Two structures are equivalent if they can be converted from one to the other by a simple rotation of the entire molecule or a simple reflection of the entire molecule. The relative positions of the electrons are left un changed! A very simple example: Ozone The true structure is a “resonance hybrid” (explain verbally) More about this... • Neither structure exists independently. • The true structure is a linear combination of the separate structures. • This is an example electron delocalization. Multiple 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? →→→→ Formal Charge! (Rules below) Draw the structure. Be sure to take into account anions and cations. 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... Selecting 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, negative formal charge should reside on the most electronegative 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 function. • A bonus: We analyze also OCN-. Exceptions to the Octet Rule ①Odd-Electron Species (“free radicals”) ②Incomplete octets (forced by formal charge) ③Expanded octets (when extra electrons can be accommodated by d-orbitals) 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! Some Verbal Examples i. ii. iii. iv. v. 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 distinct entity. • It is a very powerful concept! • We can assign properties to bonds. 1) 2) 3) 4) Bond Energies Bond Lengths Bond Angles (next chapter) Bond dipole moments (discussed in passim) Bond Energies • Chemical reactions (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. 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 reaction is favored. • The next slide shows the equation... The Equation 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 variations 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). 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.” • Sometimes, this is called a “Fermi gas.” • The electrons are delocalized over the enter metal chunk. • Paired electrons can be very far apart!