Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Chemistry II-AP Chemical Bonding Notes Bonding of atoms - types 1) ionic (electrovalent) 2) covalent (including coordinate, or dative, bonding) 3) metallic [Other types of intermolecular bonding will be discussed later.] The above types of bonding are intramolecular. 1) ionic - transfer of electrons and subsequent union of an anion with a cation - formation of ionic bond. Ions are held together by electrostatic forces - some of the strongest bonds that exist. 2) covalent - sharing of electrons to form molecules. Electrons involved generally are those in the outer energy levels. 1 ex. for NaCl: 3s - loss of 1 electron for Na 5 3p - gain of 1 electron for Cl Ions try to achieve a noble gas configuration. -often called the octet structure - try to fill the s and p subshells. Ionic bonding: occurs when atoms differ greatly in electronegativity; generally with a difference of 1.7 or greater. Typically, ionic bonding occurs between a metal of low ionization energy (Group IA, IIA, or transition metal) and a highly electronegative nonmetal (Group VIA or VIIA). Ionic compounds: characterized as solids with high M.P.; are good conductors of electricity; usually exhibit cleavage. Formation of pure ionic compounds from elements is always exothermic. Bond strength: - also known as the bond dissociation energy - measured in terms of bond energy (amt. of energy required to make or break a bond) - usually limited to gaseous compounds - the more polar the molecule, the stronger the bond. - for nonpolar molecules, the calculated value is close to the observed value, but for polar molecules, the observed energy value is greater than the calculated value. (The difference is greater, too, as the polarity increases.) ex. for H-F Observed: 135 kcal/mole 104 + 37 Calc.: 1/2 D(H2) + 1/2 D(F2) = ----------- = 70. kcal/mole 2 (found by splitting the bond energy of H2 and F2. -2** Linus Pauling made calculations to compensate for the extra strength due to polarity. This was the reason for proposing the E.N scale! 2 Y = 23( E.N.) *** BOND ENERGIES MAY BE USED TO CALCULATE THE H for a gaseous reaction if the Hf of the substances are not known. THESE VALUES ARE ONLY APPROXIMATIONS!! ............................................................... Bond distance: 1. For a nonpolar molecule, the distance is simply the sum of the two atomic radii. 2. For a polar molecule, the distance is slightly less, most probably due to the added strength of the ionic character that shortens the bond. o o ex. F-F = 2(0.64 A) = 1.28 A o o o but for H-F Calc. = 0.37 A + 0.64 A = 1.01 A o yet Observed = 0.92 A 3. With multiple bonds, the bond distance gets shorter for each additional bond ................................................. Multiple Bonds: In general, the bond distance and bond energy for an element is constant regardless of the molecule that is formed. If the observed distance is much less, then a multiple bond probably exists. examples: formula ethane (C2H6) propane (C3H8) butane (C4H10) ethene (ethylene) (C2H4) ethyne (acetylene) (C2H2) bond distance (in Angstroms) 1.54 1.54 1.54 1.33 1.20 for ethylene - a double bond exists bond energy (in kcal/mole) 83 83 83 143 196 for acetylene - a triple bond exists Multiple bonds show extra strength since the atoms are pulled closer together. **(Draw structures of ethylene and acetylene) Comparison of bond strengths for nitrogen: nitrogen gas - N2 vs. hydrazine - N2H4 225 kcal/mole 38 kcal/mole Hydrazine is used as a rocket fuel because upon combustion, it forms nitrogen gas which releases large amounts of energy as it is formed (since it is such a stable molecule). -3Drawing Lewis structures (Lewis Dot Structures or Lewis Diagrams) - used to show bonding configuration for compounds - shows the number of valence electrons and their arrangement general order for an element, 3 6 4 7 E 1 2 "s" subshell electrons 5 8 Locations # 1 and 2 represent the "s" subshell electrons; the rest of the locations represent the six "p" subshell electrons. This single element configuration is not very useful since an element will rearrange its electrons to achieve an octet configuration. +1 for an ionic compound: ex NaCl Na for a molecular substance: ex: HCl H -1 Cl Cl Guidelines for drawing Lewis structures: 1. Add up the total number of valence electrons for all elements. These must all be used. If ions are involved, be sure to add extra electrons for anions, and reduce the number of valence electrons for cations. 2. Draw the symbols for the elements and connect them by single bonds. One line represents a single bond pair. Compounds are often written in the order shown in the formula. If an element has a multiple valence, it is usually the central atom. Some elements - such as hydrogen - only can form single bonds, so they cannot serve as a central atom. 3. If all the valence electrons are used and the octet configuration is not achieved, then multiple bonds must exist. For every two electrons that are needed that do not exist, a multiple bond is needed. If four electrons are needed, either two double bonds may be needed, or a triple bond may be needed. 4. If the octet rule is achieved and extra valence electrons exist, the central atom violates the octet rule. The extra electrons are put around the central atom. -4Special cases when the octet rule is not achieved: 1. He configuration - 1st five elements - often show tendency to form dative bonds. 2. Odd number of valence electrons (often form free radicals) First two cases are rare; third case is more common. 3. Central atom holds more than 8 electrons. Not found in second series elements because they don't have a "d" subshell capability. Found when large central atom holds small, highly electronegative anion. -reason that larger noble gases can form compounds. ex. 1) BF3 2) NO2 3) SF6 RESONANCE: Limitation of Lewis structure - some molecules cannot be represented by just one drawing because electron pair(s) are delocalized. - Need to draw all structures with double arrow between them. ex. SO2 NO3-1 show benzene (aromatics) vs. aliphatics Class problem: draw the Lewis structure for ozone ** WHAT WOULD YOU PREDICT ABOUT BOND STRENGTH AND BOND LENGTH?? .............................................................. FORMAL CHARGE: - used to help determine the correct Lewis structure Formal charge = # of valence electrons - (# of lone electrons) - # of bonded pairs The best Lewis structure is the one that has the fewest number of elements with a formal charge other than zero. (Often compounds are written from the most electropositive element first with the most electronegative being last.) ex. carbon dioxide for O (#1): for O (#2): for C: <C=O=O> 1 2 or 6 - 0 - 4 = +2 6-4-2= 0 4 - 4 - 2 = -2 net: <O=C=O> (remember 4 lone pair electrons) 1 2 6-4-2=0 6-4-2=0 4-4 -0=0 0 0 (preferred because of 3 zeros) - Class problem: do cyanate ion (CNO ) ..................................................................... [N=C=O is best structure] Be sure to review sections on oxidation numbers; and redox reactions. Review section on oxides - acidic, basic, and amphoteric. Know the importance of ionic size for solubility and biological activity. -5HYBRIDIZATION: Rearrangement of electrons in available subshell orbitals to allow for maximum bonding. atomic orbitals: 1. "s" subshell - spherical shape - result in direct overlap "p" subshell - figure-8 configuration - may result in direct overlap if orbitals are lying along the same axis; but may also produce delocalized bonding. direct overlap - sigma () bonding delocalized bonding - pi () bonding When a single bond is formed, a sigma bond is always involved. For multiple bonds, the first bond is a sigma and the others are pi bonds. Specifically: single bond double " triple " 1 sigma 0 pi 1 sigma 1 pi 1 sigma 2 pi Class problem: How many sigma and how many pi bonds exist in a molecule of benzene? Hybridization: involves electrons that directly overlap ex. show carbon as a lone element; then as a central atom 2 2 3 normal configuration s p hybrid sp Instead of 2 bond sites, there are now four bond sites and octet configuration can be achieved. Class problem: do hybridization of the following: of B in BF3 of S in SF6 of P in PCl5 of Be in BeH2 Types of Intermolecular forces (IMF) - discuss each 1. ionic 2. covalent network 3. metallic 4. dipole-dipole 5. hydrogen 6. ion-dipole 7. ion - induced dipole 8. dipole - induced dipole 9. van der Waals (or London dispersion froces) -6Molecular Geometry VSEPR theory (Valence Shell Electron Pair Repulsion theory) - describes the 3-d (geometric) shape of the molecule - based on the idea that electrons tend to repel each other as far as possible - doesn't matter whether they are bonding or nonbonding electron pairs. Classification of shapes: # of e pairs # of e pairs # of e pairs shape (total) (bonding) (nonbonding) ........................................................................... bond angle examples 1 1 0 linear ............................................................................ 180 H-Cl 2 2 0 linear ............................................................................. 180 CO2 120 BF3 < 120 SO2 3 3 0 trigonal planar 3 2 1 4 4 0 tetrahedral 109.5 CH4 4 3 1 trigonal pyramid < 109.5 NH3 bent or V-shaped ............................................................................. (107 for NH3) 4 2 2 bent or V-shaped < 109.5 H2S (104.5 for H2O – important!) 4 1 3 linear 180 H-Cl ............................................................................. 5 5 0 trigonal bipyramid 90 (axial) 120 (equatorial) PCl5 5 4 1 seesaw or teetertotter 90 (a) < 120 (eq) SF4 5 3 2 T-shaped 90 ClF3 5 2 3 linear 180 XeF2 -7Classification of shapes (cont.): - # of e pairs (total) - # of e pairs (bonding) - # of e pairs (nonbonding) shape bond angle examples 6 6 0 octahedral or square bipyramid 90 SF6 6 5 1 square-based pyramid 90 BrF5 6 4 2 square planar 90 XeF4 ????????????????????????????????????????????????????????????????????????????? ?6 ?3 ?3 ? T-shaped 90 ??? ?6 ?2 ?4 ? linear 180 ??? Note #1: if multiple bonds exist, they are treated as single bonds concerning the shape. Note #2: if lone pairs exist, these electrons are found to be the equatorial pairs, not the axial pairs. This is due to the fact that repulsion is minimized with this arrangement. Consider a species with 1 lone pair: if the lone pair were in an axial position, it would encounter electron pair repulsion with three pair of o electrons in a 90 angle; if the lone pair is an equatorial pair, it only encounters two electron pair for repulsion (those closest to it). The other two pair on the equator are less significant that those axial pair closest to it. Note #3: The last two situations with six sets of electron pairs are doubtful because they would be violating the octet rule for no reason - there is not special bonding that is taking place. The first three situations occur because the majority of the electron pairs around the central atom are bonding electron pairs. ........................................................................... Class problem: What about the shape of ions and complexes? ex. ions: nitrate, triiodide, sulfate, sulfite, phosphate, carbonate, chlorate, perchlorate, hydronium, ammonium, amide, hypochlorite ex. complexes: hexachlorostannate (IV), tetrachloroiodo (III), tetrachloroplatinate(II), (I), hexafluorosilicate (IV), hexaaquazinc(II)cation ......................................................................... dichloroargentate Polarity of bonds vs. polarity of molecules 1. You can have polar bonds within a molecule and still not have a polar molecule. Can you have a polar molecule if polar bonds do not exist? 2. Polarity of the molecule deals with the symmetry of the molecule. 3. Polarity of the bond depends on the electronegativity difference. -8Examples: molecule bond type molecular shape methane CCl4 NH3 nonpolar (0.3) polar (0.6) polar (0.8) tetrahedral tetrahedral trig. pyramid polar or nonpolar? nonpolar nonpolar polar (axis of symmetry for ammonia runs from the lone pair, through the N, and through the middle of the H) H2O polar (1.3) bent polar (axis of symmetry runs from between the lone pairs, through the O, and through the middle of the H) HCl HI SO3 SF6 polar (1.0) nonpolar (0.4) polar (1.0) polar (1.5) linear linear trig. planar octahedral polar nonpolar nonpolar nonpolar Depending on the location of lone pair electrons, these electrons can produce highly polar molecules. The strength of the polarity is directly related to the solubility of the molecule, especially in a highly polar solvent like water. Molecular Orbital (M.O.) Theory: Atomic orbitals vs. Molecular orbitals M.O. - describes the distribution of electrons throughout the molecule - not just around the individual atoms - describes the energy of the different electrons Energy - based on quantum mechanics - based on the shape of the region in which the electrons move "orbital" - electron may move around the whole molecule - "molecular orbital" Molecular orbital - still holds a maximum of two electrons (w/ opposite spin) Types of orbitals - sigma and pi actually 3 types - sigma-s, sigma-p, and pi-p Bonding vs. anti-bonding orbitals Bonding orbitals: electrons are generally found between the nuclei - they are attracted by both nuclei which results in a lower energy level (more stable) than when they were isolated. -serve to stabilize the molecule. Anti-bonding * * * sigma-s , sigma-p , and pi-p