INTRODUCTION. heat. ammonium cyanate

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1 1 INTRDUTIN N N ammonium cyanate heat N urea N I can produce urea, without having the need for kidneys or an animal at all, be it human being or a dog. The cyanogen acidic ammonia is urea. Letter from F. Wöhler to J. Berzelius, 1828 ISMERS have identical molecular formulas but have different arrangements of atoms in their molecules, that is, different structures. onstitutional Isomers (Structural Isomers) differ in their "connectivity", that is, in the order of attachment of the atoms in the molecule. These isomers have different physical and chemical properties.

2 2 Structure Name dimethyl ether ethanol n-butane isobutane b.pt. ( ) density (g/ml) The properties of a substance depend upon its structure. structure physical & chemical properties Structure hemical transformations Bonding Electrons Sugar, salicin, and morphium will be produced artificially. f course, we do not know the way yet by which the end result may be reached since the prerequisite links are unknown to us from which these materials will develop however, we will get to know them. F. Wöhler and J. von Liebig

3 3 ELETRNI STRUTURE F TMS tom consists of a dense inner core -- the nucleus (protons, neutrons) -- and electrons that surround the nucleus. Particle Relative Mass harge proton (p) 1 +1 neutron (n) 1 0 electron (e ) 1/ The number of protons in the nucleus (atomic number, Z) uniquely determines the atom's identity. In a neutral atom, #p = #e; #n is variable (atomic isotopes). { } { D (deuterium) } isotopes of carbon isotopes of hydrogen Electrons have particle-like properties (mass) and wave-like properties (diffraction). Solve WVE EQUTIN > WVE FUNTINS and quantum numbers Wave Functions express the energies and the positions of electrons in an atom. Quantum Numbers n is the designation for the principle, or main, energy level (or energy shell) in the atom that may be populated with electrons. The closer an electron is to the nucleus, the less energy it has and the more stable it is.

4 4 l is an energy sublevel (or subshell) within a given main energy level, n. [ s p d f sublevels]. There are restrictions on the number of sublevels within a given main level. ml is an orbital within a given sublevel. n orbital is a region of space in which an electron is most likely to be found. [ one s three p five d seven f ] rbitals have unique "shapes". Maximum of two electrons per orbital. ms designates the quantum number for the electron spin. The opposite nature of their spins differentiates between the two electrons that may occupy a given orbital (thus having the same n, l, and ml). Spins may be designated as +1/2, 1/2; clockwise, counterclockwise ; Pauli Exclusion Principle Each electron in an atom has a unique set of quantum numbers (n, l, ml, ms)

5 5 tomic orbital chart Main level (n) Sublevel (l) rbital ( m l ) Spin ( m s ) Max. # electrons per sublevel Max. # electrons per main level 1 s s ± 1/ s s ± 1/2 2 p p x p y p z ± 1/ s s ± 1/2 2 p p x p y p z ± 1/2 6 d five d's ± 1/ s s ± 1/2 2 p p x p y p z ± 1/2 6 d five d's ± 1/2 10 f seven f's ± 1/ * * lthough each main energy level contains a number of sublevels 6 equal to n (verify for n = 4, above), main levels 5-7 do not * require more than 4 sublevels to accommodate the electrons in all 7 * known atoms (due to sublevel energy overlap). The sublevels, orbitals, and electron-occupancy for n = 5-7 will be the same as for n = 4. Electron onfiguration is the description of the distribution of electrons within an atom. In the ground state, electrons will occupy the atomic orbitals of lowest energy. n l #e or n ml #e 1 = 1s 1 Li 3 = 1s 2 2s 1 B 5 = 1s 2 2s 2 2p 1

6 6

7 7 Valence Electrons are those electrons that occupy the outermost (highest) main energy level (n) of an atom. They are the most loosely held and engage in chemical reactions. Electron Dot Symbols designate the valence electrons (dots) surrounding the inner core electrons and nucleus (atomic symbol) of an atom. MIN BLK TMS I II III IV V VI VII VIII und's Rule In filling degenerate orbitals with electrons, each orbital is occupied singly with electrons of the same spin; subsequently, electrons of opposite spin are added. ENERGY LEVEL DIGRM 2s 2px 2py 2pz Inc. Energy 1s SUGGESTED PRBLEMS: 1. Write electron configurations for all atoms Z = Select a few atoms from problem 1. and show the application of und s Rule by drawing energy diagrams as shown above.

8 8 Rule of Electronic Stability Electronic stability is greatest when atoms have an n(s 2 p 6 ) configuration in their valence shell and resemble the closest Noble Gas configuration. The exceptions are elements that can attain the 1s 2 configuration more easily than the s 2 p 6 configuration, such as and Li. toms gain, lose, or share electrons in order to achieve this "octet" of electrons. Mg = 1s22s22p63s2 Mg +2 = 1s22s22p6 l = 1s22s22p63s23p5 l 1 = 1s22s22p63s23p6 Electronegativity is the tendency of an atom to attract electrons toward itself (to the positively charged nucleus). E.N. is a periodic property that increases up a family and increases across a row in the periodic table. Electronegativity B < < Electronegativity increases across a row The positive charge on the nucleus increases as protons are added to the nucleus. Negatively charged valence electrons are added to the same main energy level. The effective positive charge of the nucleus increases for all the electrons that are in the same valence shell and so the attraction between the electrons and the nucleus increases across the row. Electronegativity decreases down a column The positive charge on the nucleus increases as protons are added to the nucleus. Negatively charged valence electrons are added to main energy levels further from the nucleus. lthough there is an increase in the number of positive protons in the nucleus, the increased number of inner electrons screens the effective positive charge on the nucleus from attracting the outer valence electrons.

9 - EMIL BNDING hemical bond -- the force of attraction that holds atoms together in a molecule. The force is the attraction between the negatively charged electrons and the positively charged nuclei. The drive for bond formation is to achieve a filled valence shell. Ionic bond Results from the electrostatic attraction between ions of opposite charge. K K + e - F + e - F ovalent Bond -- results from sharing of two electrons between two atoms. The shared electrons are electrostatically attracted to both nuclei. There are several ways to represent covalent bonding between atoms in a molecule. Before going into detail about how to represent bonding using the atomic orbitals of atoms, we will construct structural formulas using the same electron dots as for individual atoms. WRITING ELETRN DT STRUTURES (1) From the molecular formula, calculate the number of valence electrons contributed by all the atoms in the molecule. djust the total number of electrons for net charge, if any. (2) rrange the atomic symbols to correspond with the known connectivity. Generally, more electronegative atoms and 's tend to be terminal atoms in a molecule. Ignore molecular geometry for now. (3) Distribute the valence electrons in pairs so that each atom is bonded to at least one other atom. Distribute the remaining valence electrons among the atoms, giving them multiple bonds and/or non-bonded electron pairs. Try to give each atom an octet (except ).

10 -2- (4) heck your drawing to make sure that all atoms have no more electrons than their valence level can accommodate: a first row element can have no more than two; elements in the second row of the periodic chart can have no more than eight electrons; elements in the third and higher rows may have more than eight electrons because of the availability of the unfilled d sublevel. P and S are notable examples of atoms that may have more than eight electrons in their valence shells they are in the 3 rd main energy level. Example: ommon bonding for atoms without formal charges: I II III IV V VI VII

11 -3- Formal harge on an atom is a method of electron bookkeeping that indicates the difference in electron "ownership" by an atom in its atomic vs. its bonded molecular state. It does not indicate the actual charge distributions on the atoms in a molecule. (1) Find the atom's valence Group number (I-VIII). (2) Subtract from (1) the number of valence electrons the atom owns in the molecule (all of its non-bonding electrons and half of its shared bonding electrons). (3) Total: zero, there is no formal charge on that atom; nonzero, there is a negative (excess electrons) or positive (deficiency of electrons) formal charge on the atom. The sum of the formal charges on the atoms in a molecule equals the net charge of the molecule or ion. There is a net ionic charge on the polyatomic ion. There is no formal charge on hydrogen. There is a formal +1 charge on oxygen. N SUGGESTED PRBLEM: There are 7 structural isomers with the molecular formula that contain one double bond, including acetic acid, above. Draw as many of them as you can. Include all non-bonding electrons in your drawings. None of the isomers should have formal charges. The answer will be posted later on the class website with other hapter 1 problems. The next page will show you how to begin working this problem in a systematic way.

12 -4- pproach to solving the isomer problem: In this and all other problems, do not use bond line formulas: use atomic symbols for atoms. Understand the meaning of "isomer": same molecular formula, but different connectivity. To simplify the drawings initially, do not draw in the ydrogen atoms. Draw the different possible connectivities using only the arbon and xygen atoms. Where to start? nywhere is usually fine, but the simplest or most obvious might be the best way. For example, try putting the and atoms in a continuous chain (no branching). Since there are various combinations, start with one order of atoms and complete the drawing. 1. Perhaps start with this connectivity: Now choose where to put the double bond: Now put in the hydrogen atoms and non-bonding electrons: = (int: this will be a valid structure) 1B. See if there is another place to put the double bond: = Now put in the hydrogen atoms and non-bonding electrons: (int: this structure will have formal charges) 1. nd then another: = Now put in the hydrogen atoms and non-bonding electrons: (int: this structure will have formal charges) This now exhausts the possibilities of placement of a double bond in this connectivity. fter you complete each drawing, check that it fits the requirements for a correct answer: it contains the requisite number of atoms and electrons; there is only one double bond; there are no formal charges. 2. hange the connectivity slightly: Repeat the steps above by choosing a place for the double bond and then drawing in the 's and non-bonding electrons. heck each drawing for conformity for a correct answer. 3. ontinue to change the connectivity and repeat the drawing steps until you have exhausted all possibilities for a continuous chain of 's and 's and for a branched structure of 's and 's. ints: There are four 2 2 connectivities two of them are shown above. There are two remaining with continuous chain connectivity and one with branched connectivity. You know there are seven correct answers. Many of the drawings will not conform to a correct answer because they require formal charges on some of the atoms.

13 -5- Resonance is a term that indicates there is more than one way to reasonably assign electrons in a molecule without moving or rearranging the nuclei or changing the connectivity. The molecule's real electronic structure (the resonance hybrid) is a blend of the individual resonance contributors. The resonance contributors do not necessarily contribute equally to the overall hybrid structure; rather the major contributor(s) will be those that are the most plausible Lewis structures. (See the guidelines below on better resonance structures.) ere are some factors to consider when judging which resonance contributors are better/more stable/more important than others: toms that have an octet are better. Structures that have more covalent bonds are usually better. For neutral molecules, structures in which there are no formal charges on atoms are usually better than those where formal charges are required. Where formal charges are necessary, these are generally of as small a magnitude as possible. mong alternative electronic structures, the most stable is that in which negative formal charges are placed on the more electronegative atoms. structure in which like formal charges (both positive or both negative) reside on adjacent atoms is less stable. Some of the statements above are redundant in some cases. For example, look at the formaldehyde resonance contributors. The first-drawn structure is the best on the basis of the first, second, and third statements, which are essentially equivalent here.

14 -6- Example: In the Lewis structure for acetic acid drawn on a preceding page, it was explicitly stated that we would consider only that one resonance structure. Draw another resonance structure for acetic acid in which all atoms have an octet. Include formal charges and all non-bonding electrons. Solution: Take a bonding pair of electrons from the = and make it a non-bonding pair on the (exactly as was done for formaldehyde on the preceding page). [You should draw the result of this operation to note that the now does not have an octet. The structure is a resonance contributor but not a very good one.] To give the an octet, take a non-bonding pair from the other and make it a second bonding pair between the and. Note that all the atoms have an octet, which is good. The number of covalent bonds in the two structures is the same. owever, the second structure has formal charges which renders it less important than the one that has no formal charges.

15 -7- SUGGESTED PRBLEMS 1. Draw resonance contributors (include formal charges) for urea that resemble the ones for acetic acid. Remember, none of these atoms can have more than 8 electrons. Begin by drawing the complete electron dot structure for urea and proceed as with the example for acetic acid. 2. There are three connectivities (structures) for the anion composed of carbon, nitrogen and oxygen: N, N, N. Some resonance contributors for the N structure are shown below. onfirm that each resonance contributor accounts for all the valence electrons and that all of them show the same number of valence electrons. onfirm the assignment of formal charges. N N N N 2 Explain why the one best resonance representation (this is not the hybrid) for this anion is the one shown below, taking into account the guidelines for better resonance structures as given on a previous page. N 3. Peroxynitrous acid ( the connectivity is N) is useful for fast oxidations; however, the compound rapidly isomerizes to nitric acid. Draw electron-dot structures for peroxynitrous acid and nitric acid. Explain why these two species are isomers. Draw all resonance contributors for peroxynitrous acid, stating which contributors are equivalent and which are not. ssign formal charges to all atoms in the drawings. Repeat for nitric acid. 4. The molecule I below has an important resonance contributing structure II. Explain why this might be an important contributor and explain why III is not a valid resonance contributor. [int: count electrons.]

16 -8- N N N I II III Some selected answers to above problems 1. Notice the similar relationship of the = and : in acetic acid and the = and N: in urea. lso, you can draw yet another resonance contributor for urea that uses the non-bonding pair on the other N. These two resonance contributors are equivalent. Both of them contribute equally to the resonance hybrid structure, but both of them contribute less than the structure that has no formal charges on N or. N N 3. N N N None of these are equivalent resonance contributors. The middle structure probably contributes the least because N does not have an octet. Notice that although in the third structure has a formal positive charge, has an octet. aving an octet is always more important than anything else.

17 -9-5. Draw resonance contributors for the polyatomic ions carbonate and nitrate. Include formal charges. re the contributors equivalent or non-equivalent? The answers are in your text. These are excellent first problems to solve. nother question in your text asks you to write a structure for sulfate, (S 4 ) 2. owever, this problem is similar to the ones above, that is, there are several resonance contributors, no one of which is adequate to show the true structure. notable difference between the central atoms N,, and S in these three problems is that S is in the 3 rd row of the Periodic Table and so can have >8 electrons in its valence shell. You can try to complete the full set of resonance structures below. There are 32 valence electrons in sulfate. S S (The double bond can be between S and any of the 's) S 1+ S (The two double bonds can be between S and any of the 's. Draw the other structures with two double bonds in the space below. There should be 5 more.) ere is one of them: 1+ S S S 1+ You can repeat this process using three double bonds (total 4) and four double bonds (1, the answer shown in your textbook). S The resonance structures in each column are equivalent to each other. The sulfate ion has a tetrahedral shape, with 4 equivalent S bonds.

18 -10- SYMBLI REPRESENTTINS F VLENT BND FRMTIN IN 2 1. Lewis Electron Dot Structures 2. /M Pictures energy released -435 kj/mol 1 s atomic orbitals molecular orbital Sigma () covalent bonds result from a "head to head" overlap of atomic orbitals along the internuclear axis. ny single bond is necessarily a sigma bond (as shown above). Each of the two atomic orbitals that form the bond can be occupied by one electron or two electrons can occupy one orbital and the other orbital is empty. (Examples shown later.) The resulting bonding molecular orbital is occupied by the two electrons. 3. Energy Level Diagrams Number of tomic rbitals = Number of Molecular rbitals [onservation of rbitals] * (antibonding M) 1s 1s (bonding M)

19 -1 Unoccupied molecular orbital (antibonding, *) tomic orbitals ccupied molecular orbital (bonding, ) Before looking at the representation of bonding in larger molecules, we need to reconsider the atomic orbitals involved in covalent bonding.

20 -12- DRWING TMI ND MLEULR RBITL PITURES Sigma () covalent bonds result from a "head to head" overlap of atomic orbitals along the internuclear axis. The table below shows how a sigma covalent bond can be formed from combinations of s+s orbitals or s+p orbitals or p+p orbitals. The identity of the orbital is irrelevant the bond is formed head-to-head along the internuclear axis. The nucleus of each atom is shown as a single dot. The electrons are likely to be found in the volume enclosed by the orbital shape. TMI RBITLS from two participating atoms BNDING MLEULR RBITLS two atoms are bonded together s s s p p p ny single bond is necessarily a sigma bond (as shown above). Each of the two atomic orbitals that form the bond can be occupied by one electron or two electrons can occupy one orbital and the other orbital is empty. (Examples shown later.) The resulting bonding molecular orbital is occupied by the total of two electrons. Later, we discuss molecules of the type = or, containing a double bond or a triple bond. We will see that in a set of multiple bonds (double or triple bond), only one can be a sigma bond; the others are pi () covalent bonds.