Bonding General Concepts

Bonding: General Concepts 8.1 Types of Chemical Bonds. There are 3 basic categories of bonds: ionic, molecular and metallic. • Ionic: the attraction due to electrostatic charge between oppositely charged particles. Usually formed by a metal with a nonmetal. • Covalent: result of sharing e- between nonmetallic atoms. Frequently bonds between nonmetal and metalloid atoms qualify as covalent as well. • Metallic: solid metals. An atom is bonded to several neighbors, but electrons are relatively free to move about in the 3-D structure. (Causes electrical conductivity and luster.) Lewis Structures and the Octet Rule Valence: a measure of an atom's ability to form ionic or covalent bonds. Valence e- are outer shell electrons only. A Lewis or e- dot structure shows valence e- around the element symbol. 1 dot = 1 e-. Ex. Sulfur [Ne]3s23p4 6 valence e- Draw the Lewis diagram for sulfur here: For an active metal or representative element, the number of valence e- = column number. Many atoms undergo reactions that leave them with 8 or an octet of valence e- like a noble gas (except those that end up like He with only 2 valence e-. There are many exceptions to the octet rule but it is useful for many bonding concepts. Ionic Bonding 2Na(s) + Cl2(g) ???aCl(s) The NaCl is a crystal of regularly arranged Na+ and Cl- ions in a 3-D structure. The NaCl(s) is a crystal of regularly arranged Na+ and Cl- ions in a 3-D structure. Na has lost 1 e- and Cl has gained one. Draw the e- dot diagrams of the electron transfer here: Na+ now has 2s2 2p6: an octet. Cl is now 3s2 3p6: also an octet. Coulomb’s Law: Calculates bond energy. Ionic bond energy energy depends on the charges of the ions and the distance between them. Epotential = (2.31 x 10-19 Jnm)(Q1Q2) where Q1 & Q2 = numerical ionic charges (+1, -2 etc) r r = distance between ionic centers (in nm) What this means is that large ions with a low charge form weaker crystals with low melting and boiling points. Small, highly charged ions form hard crystals with high m.p. and b.p. Ex. CsI has a m.p. of only 621°C while Al2O3 has a m.p. of 2027°C. 8.2 Electronegativity Identical atoms share e- exactly equally resulting in a non-polar covalent bond. The other extreme involves one atom completely removing one or more e- from another, forming ions. Most often, shared e- are not shared equally. A polar covalent bond develops, that is, the shared e- spend more time around one of the atoms than the other. **The ability of an atom to attract shared e- to itself is called electronegativity. This concept was developed by Linus Pauling, a 20th century American chemist. (Pauling's scheme is not the only one. There are a couple of others, and some of the values vary significantly.) The most electronegative element is F with a value of about 4.0. The scale has Paulings as units and a range of 0.0 to 4.0. The least is Fr with a value of about 0.7. Electronegativity values are relative and approximate. They undoubtedly vary with situation. 8.3 Bond Polarity and Dipole Moment Molecular shape and bond polarity determine charge distribution in molecules. If the centers of + and - charge are not in the same place, the molecule is polar. The degree of polarity is measured by the molecule's dipole moment. **A molecule may have polar bonds and yet be non-polar overall because of its geometry. Ex. CO2 O = C = O The bonds are polar. Electron density is greater around O + + atoms, but the centers of both + and - charge are on carbon. Result, no overall molecular polarity and no dipole moment. This helps to explain many of carbon dioxide's physical properties. Your text book has several nice images of electron density models on pages 336-7 that demonstrate polar and non-polar molecules. Both bond polarity and molecular shape are taken into consideration. Because it is bent rather than linear, a water molecule is polar. Draw the water molecule and other examples here: *Unshared e- prs. can and often do contribute to polarity. 8.4 Electron Configurations and Sizes of Ions Important in terms of lattice energy, as discussed earlier. Removal of ions by water softening systems depends largely on ionic size. Many biological systems are very ion specific. Size is important. Ex. K vs Na vs Ca. Review basic electron configuration with your instructor if necessary. • Metal atoms lose outer e- and become smaller as cations. • Nonmetals become larger as anions as they achieve noble gas configuration. 8 outer e- leads to maximum e- repulsion. Outer E level expands. Isoelectric ions: ions with the same noble gas configurations, show the effect of nuclear charge on ion size. O2- F- Ne Na+ Mg2+ Al3+ All have 10 e- , the same as Ne. 4502159144000largest smallest 354965141605008 p+ 13 p+ The particle radius decreases as effective nuclear charge (Zeff) increases. 8.5 Energy Effects in Binary Ionic Compounds and the Born-Haber Cycle. Energy of Bonding: Ionic compound formation is always exothermic. • I is endothermic. • e- affinity is usually exothermic. • Lattice energy is always quite highly exothermic. Technically, lattice energy is the energy required to take a crystal apart, but an equal amount of energy is released when a crystal of a particular material is formed. Lattice E for NaCl = -785kJ/mol. The result is a solid substance with a high m.p. In 3-D, each Na+ ion is surrounded by 6 Cl- ions, and each Cl- by 6 Na+ ions. Strong attraction exists between oppositely charged particles. The arrangement of ions and geometry of the crystal depends on the ratio of ions in the compound. Attractive forces are maximized and repulsive forces are minimized producing brittle solids with high melting points. **ALL ionic compounds are solids at room temperature (unless dissolved in a solvent.) Born-Haber Cycles The energy of ionic bonding involves several phases that are brought together in a series of steps called a Born-Haber Cycle. The steps of the cycle follow: Step Endo- or Exothermic Example 1. Sublimation of solid metal Endothermic Li(s) Li(g) 2. Ionization of metal atoms to cations Endothermic Li(g) Li+(g) + 1 e- 3. Sublimation of nonmetal* Endothermic ½ Br2(l) ½ Br2(g) 4. Dissociation of nonmetal molecules Endothermic ½ Br2(g) Br(g) 5. Electron affinity of nonmetal atom to anions Exothermic Br(g) + 1e- Br-(g) 6. Formation of ionic solid: lattice energy Exothermic Li+(g) + Br-(g) LiBr(s) *May not be necessary if nonmetal is a gas to start. These types of problems make use of Hess’s Law to derive a final balanced equation. Calculating lattice energy involves a version of Coulomb’s Law: E = k(Q1Q2/r) where k is a constant that depends on the structure of the solid. As charges increase, energy increases, and as distance between ionic centers increases, energy decreases. With larger radius ions, distance increases. MgCl2 has a higher m.p. than CaCl2 because the Ca ion has a larger radius than Mg and therefore forms "weaker" crystals (as distance increases,? lattice energy decreases.) Question: If energy increases as Q increases, why doesn't Na form +2 ions? Answer: Side Note: Transition Metal Ions: For some elements, the octet rule must be modified or ignored completely. Ex. Metals of Group IB: Cu, Ag, Au. Ag is 5s14d10 and nearly always becomes Ag+ n = 4 now has 18 e- and is filled. For many transition metals, noble gas configuration is impossible because, below their ns1 or ns2 outer e-, lies a partially filled (n-1)d sublevel. Since some of these "d" e- may be lost in addition to the ns e-, multiple ionic charges are often possible for many of the transition metals. Ex. Cu may be Cu+ or Cu2+. There are no simple rules or easy explanations as to what charges such a metal may have. *Review ions and their charges including the polyatomic ions. Ex. Hg22+ mercury(I), VO2+ vanadyl, O22- peroxide, CN- cyanide and OH- hydroxide 8.6 Partial Ionic Character of Covalent Bonds The difference between the electronegativities for 2 atoms that are bonded together is a measure of polarity. Ex. HCl Cl = 3.2 and H = 2.2. Difference = 1.0. The bond is polar. Electronegativity difference can be translated into % ionic character. The scale is roughly this: Electronegativity Difference 0-0.4 0.5-2.0 >2.0 % Ionic Character 0-5% 6-50% >50% Bond Type Non-polar covalent Polar covalent Ionic Since Cl is more electronegative than H, the shared e- spend more of their time around Cl than H. The result is that Cl has a partial negative charge (?-) and H a partial positive charge (?+). ?+ H--Cl ?- e- are more attracted to Cl than to H. Another way to show this is as follows: H--Cl The arrow points to the more electronegative element in the bond with + a small plus sign at the other end of the arrow. 8.7 The Covalent Chemical Bond Why do atoms form covalent bonds? As with most natural processes, thermodynamics is the deciding factor. Systems seek the lowest possible energy (and become as thermodynamically stable as possible in the process.) The ?Hf° of methane, CH4, is -1652 kJ/mol. In other words, 1652 kJ/4 = 413 kJ of energy is released every time C covalently bonds to H. The bond energies of a great many covalent bonds between different kinds of atoms are known. We can actually calculate the ?H°rxn for most any chemical reaction by comparing the energy required to break existing bonds and the energy released in the formation of the new covalent bonds. Molecular modeling: We are about to undertake the construction of various molecular models. But please remember: Models help us understand what we cannot see, but they often do not reflect reality. Models are often wrong. When models are wrong and we discover the flaws, we learn a lot. 8.8 Covalent bond Energies and Chemical Reactions When we look at bond energies, for example in Table 8.4 on p. 351, we are seeing average energies for the bonds between 2 atoms. Be aware that the true energy of the bond varies with circumstance. We can calculate the enthalpy of a reaction by using the following equation: ?H = ? D (bonds broken) – ? D (bonds formed) Energy required - energy released D = bond dissociation energy: the energy required to break 1 mole of covalent bonds between atoms. Do Example problem 54a in the homework. 8.9 The Localized Electron Bonding Model The localized electron (LE) model assumes that molecules are atoms bound tightly together by sharing pairs of electrons that are a blend of the atomic orbitals of the atoms. Localized electrons “belonging” to an atom may be of 2 types: lone pairs (which do NOT bond 2 atoms together.) bonding pairs, found in the space between 2 atomic nuclei that DO bond atoms together. The LE model has 3 parts: Lewis structures: electron dot models of molecules that involve valence electrons. Prediction of the molecular geometry using the VSEPR model (discussed later). Description of the atomic and molecular orbitals used by atoms to share electrons or hold lone pairs (Chapter 9). 8.10 Lewis Structures Lewis structures (named after American chemist G. N. Lewis) are a 2-D way of representing molecules. They show us a generally correct picture of each atoms bonding and lone pair electrons. Its chief weakness is that it does not show 3-D geometry. Below are the steps to follow in order to construct a correct Lewis diagram. This is an important skill that requires practice. Count the total number of valence electrons in the molecule. If the particle is an ion, add or subtract the appropriate number of electrons based on the particle’s charge. Determine which the central atom is and arrange the other atoms around it. The central atom is usually the singleton in the group. If you still cannot tell, you may need to check formal charges (Section 8.12). Connect bonded atoms with a single covalent bond (1 pair of electrons.) Add the remaining electrons so as to satisfy the duet rule (2 e-) for hydrogen and the octet rule (8 e-) for other atoms in the structure. This is sometimes tricky. Start by giving all the atoms an octet of electrons (except H) and see how many electrons you have. If there are too many, remove a pair from each of 2 bonded atoms and replace with a double bond between them. Triple bonds are also possible. It is also possible for 1 atom to donate both of the electrons necessary to make a bond. Such a bond is called a coordinate covalent or dative bond. Do a final count on the electrons to be sure there is the correct number. Practice several examples in class. 8.11 Exceptions to the Octet Rule Exceptions with fewer than an octet of electrons. Some smaller atoms (H, Be) emulate He with only a duet of electrons. Boron is a special case and is reasonably stable with only 6 electrons. Ex. BF3. Boron only appears to have a sextet of electrons, although the octet rule can be satisfied of there is a double bond in the molecule. F donates a pair of electrons for the second bond. Exceptions with more than an octet of electrons. Nonmetals in Row 3 or lower of the periodic table are large enough central atoms to accommodate more than 8 electrons. This is called an expanded valence shell. Notable examples include PCl5 and SF6. This is possible because there is a d-sublevel available to participate in the hybridization process (Chpt. 9). 8.12 Resonance Sometimes it is possible to have more than one Lewis diagram for a molecule that satisfies all the structural requirements. A good example is the nitrate ion, NO3-. Draw the nitrate ion below. Resonance is common where there alternating single and double bonds, and you are not sure where to put the double bond. Another type of resonance involves molecules in which the central atom has an odd number of electrons, such as NO2 or ClO2. Draw NO2’s resonance structures below. Formal Charge Formal charge is used to help determine the most likely structure for an atom or ion when several structures are possible. This may involve Structures where the central atom is not certain. Ex. SCN- Molecules or ions with non-equivalent Lewis structures (those that contain differing number of single and multiple bonds). Ex. SO42- Formal charge = (number of valence electrons on the free atom) – (number of valence electrons assigned to the atom in the molecule) The first part is easy. The number of electrons the atom is assigned in the molecule is calculated as follows: Unbonded electron pairs belong solely to the atom. The atom counts ½ of any shared electrons. Or, shared electrons are divided equally between the 2 atoms sharing them. Deciding which structure is best. Negative formal charges are assigned to the most electronegative element. The total of the formal charges in the particle must add up the charge of the particle. The set of formal charges that is closest to zero is the most likely structure. Example: The sulfate ion can be drawn with all single bonds or with 2 double bonds. Calculate the formal charge of the atoms 8.13 Molecular Geometries The VSEPR Model Valence Shell Electron Pair Repulsion e- are attracted to the nucleus, but are repelled by each other. They stay as far apart as possible while maintaining their distance from the nucleus. An atom with an octet of 4 pairs of bonding electrons would probably take on the shape of a tetrahedron, with the nucleus at the center. See p. 371-2 in the text for examples of geometries. A tetrahedron is the geometric shape allowing greatest distance between e- pairs that are equidistant from the nucleus. CH4 is an example of such a molecule. The model also works for differing numbers of e- pairs about the nucleus. Examples: Molecule Shape Bond Angle BeCl2 linear 180º BF3 trigonal planar 120º CH4 tetrahedral 109.5º PCl5 trigonal bipyramidal 90º & 120º SF6 octahedral all 90º *Virtually all molecular geometries are variations of these 5 basic shapes. ALL OF THE POSSIBLE SHAPES ARE SHOWN IN THE TABLES ON PP. 371-374. Electron pairs surrounding the central atom of a molecule may or may not be shared e- pairs. H2O is a bent rather than a linear molecule, even though we might expect it to be the same as BeCl2. Oxygen in a water molecule possesses 2 unshared pairs of electrons that influence the shape of the molecule. If we examine the e- pair geometry alone, we see that O has an octet of e- which takes on a tetrahedral geometry. We must be aware of unshared e- pairs before we can use the VSEPR model to predict molecular shape. 1. We take all e- pairs into account when describing e- pair geometry. 2. We describe only the atoms in the molecule when stating molecular shape. Unshared e- pairs influence geometry and shape. Example: CH4, NH3, H2O and HF are all tetrahedral if we only look at e- pair geometry, however the shapes vary in their descriptions. The shapes are tetrahedral, pyramidal, bent and linear respectively. The bond angles change slightly as unshared pairs appear. Unshared e- pairs take up more space than shared pairs which are held between 2 nuclear centers. So the bond angle in ammonia is only 107º rather than the predicted 109.5º associated with tetrahedral geometry. And likewise, the bond angle of a water molecule is affected even more and measures only 104.5º. SF4 is an interesting molecule. e- pr. geometry is trigonal bipyramidal, and S has 1 unbonded e- pair (making it an exception to the octet rule.) The one lone pair of unbonded electrons may occupy one of two positions: axial or equatorial (radial). 1. If the e- pr. is in an axial position, it will experience three 90º interactions with the pairs bonded in the equatorial positions. 2. If the unbonded pair is in an equatorial position, it will experience only two 90º interactions with the 2 bonded pairs in the axial positions. The unbonded pair is in reality found in this radial position because the repulsion of only 2 pairs at 90º is less than that of 3 pairs at 90º. Also, the lone unbonded e- pr. demonstrates the same effect on bond angles as seen in previous examples. The axial bond angle is 186º rather than the predicted 180º and the atoms in the equatorial positions are only 116º apart rather than the expected 120º. The VSEPR model can be used to predict just about any molecular geometry, whether for molecules obeying the octet rule or for exceptions. ** When applying the model, treat multiple bonds like single bonding e- prs. For example, CH2O has 4 bonds, 2 C-H bonds and a double bond C=O. The molecule is trigonal planar geometry. To use the model: 1. Count valence e- and draw the Lewis structure. 2. Determine the number of bonds (counting multiple bonds as singles) and unshared e- pairs . 3. Determine the e- pr. geometry for this number of bonds 2 = linear 3 = trig. planar 4 = tetrahedral 5 = trig. bipyramidal 6 = octahedral and put unshared pairs in positions that minimize e- pr. repulsions. 4. Describe the shape of the molecule in terms of the positions of the bonded atoms. (Unbonded pairs are NOT described as part of the molecular shape.)