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Organic Chemistry I

by: Houston Kovacek

Organic Chemistry I CHM 355

Marketplace > Marshall University > Chemistry > CHM 355 > Organic Chemistry I
Houston Kovacek
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This 0 page Class Notes was uploaded by Houston Kovacek on Sunday November 1, 2015. The Class Notes belongs to CHM 355 at Marshall University taught by Staff in Fall. Since its upload, it has received 67 views. For similar materials see /class/233286/chm-355-marshall-university in Chemistry at Marshall University.

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Date Created: 11/01/15
Chapter 24 Chemistry of Coordination Compounds We see color all around us but what is color When we see white we see all of the wavelengths of light in the visible region of the electromagnetic spectrum ca 400 7 700 nm transmitted or re ected in roughly equal amounts Black is the opposite it is the transmission or re ection of no light We see color when one or more wavelengths of light are transmitted or re ected to a greater extent than the others or exclusively Broadly stated there are inorganic and organic sources of color In CHM 355356 you will learn how organic compounds can produce color Examples of colored organic compounds include the dyes that color your clothing the ink on a newspaper and the indicators used in the titrations you did in CHM 217 We will discuss the origin of inorganic color in one type of compound coordination compounds in this chapter but in reality the discussion applies to the large majority of inorganic compounds Inorganic sources of color include many paints gemstones and hemoglobin and chlorophyll although strictly speaking these last two also have an organic component as well 241 Metal Complexes A complex or complex ion is an assembly of a metal ion and bound Lewis bases Compounds that contain complex ions are called coordination complexes The Lewis bases that bind coordinate to the metal atom or ion are called hm Ligands are usually anions or polar neutral molecules Examples include halide ions water ammonia cyanide ion carbonate ion hydroxide ion and ethylene diamine H2NCH2CH2NH2 The metal is most usually a transition metal although main group metals can form coordination complexes e g AlH2063 as seen in Chapter 17 The reason for this is transition metals generally have empty or partly lled valence d orbitals The result is that the ligands bind to the metal through a coordinate covalent bond Chap 16 notes p 20 The gure below shows a typical coordination complex 131 13 2 Cu2 4NH Cu ltan 3an HSNxf NH3 NH3 ltan The ligands and metal combine to form the coordination sphere The species that comprise the coordination sphere should be written in square brackets as follows CuNH342 or CuNH34C12 The brackets indicate that the species generally behaves as a single unit We will discuss this in greater detail shortly Finally the arrows represent bonds but also show that the ligands are supplying all of the electrons Charges Coordination Numbers and Geometries The charges on coordination complexes are obtained by treating each constituent ion or group of the complex as if it were a free species Thus for AgNH32Cl The C1 is a chloride ion so the complex ion is AgNH32 Ammonia is a neutral molecule so the silver must be in the 1 oxidation state Ex What is the oxidation number of cobalt in CoNH35ClNO32 Let the oxidation number of Co be x NH3 is neutral Cl is chloride Cl39 and N03 is nitrate NO339 Thus 0 x 50 1 21 x 3 ie a Co3 ion The atoms in a ligand that actually bind to the metal are called donor atoms In the rst example the nitrogens in the ammonia are donor atoms In the second example there are ve nitrogen donor atoms and one chlorine donor atom The number of donor atoms bound to the central atom or ion is called the coordination number CN In the previous examples the coordination numbers are 2 and 6 respectively Several factors in uence coordination number The sizes of the metal and ligands are generally most important The most common coordination numbers are 4 and 6 Fourcoordinate complexes are almost always either tetrahedral or square planar and 6 coordinate complexes are almost always octahedral We ll come back to why in a few pages 242 Ligands with More than One Donor Atom Both of the ligands used in the examples thus far have had only one donor atom ie one site of attachment These ligands are called monodentate Some ligands contain more than one donor atom Ligands binding to a metal using two or more donor atoms are called polydentate Bidentate ligands bind through 2 donor atoms T1i m3 1amp3 and hexadentate ligands bind through 3 4 5 and 6 donor atoms respectively Common polydentate ligands include Name Formula abbr ethylene diamine H2NCH2CH2NH2 en 2 carbonate C03239 2 oxalate 39OZCCOZ39 ox 2 H I H 5 IH H C H bipyridine IC C bpy 2 H CNC lNC H H ii a e I I o cCH2 HZCc o ethylene d1amme tetraacetate EDTA 6 4 Polydentate ligands are commonly called chelates or chelating ligands The name derives for the Greek word for claw because these ligands grasp at the metal at multiple places Chelates are an important class of ligands because they bind significantly more strongly to metals than to monodentate ligands For example NiH2062aq 6 NH3 aq NiNH362aq Kf 12 x 109 3 enaq Nien32aq Kf 68 x 1017 If one assumes the ligand concentration is signi cantly larger than the Ni2 concentrations then the ethylene diamine complex is about 100 million times more stable than the amine complex Why is this so If you think about the reverse reaction you see that it is much more difficult to remove ethylene diamine ligands than ammonia molecules Both reactions demonstrate that Ni2 has a signi cant preference for amine type ligands over water thus the likelihood of an N donor atom letting go at any given time is quite small nonetheless there is a difference Let s assume that we have one of each of the above complexes and that one of the N donor atoms has let go in each complex In the case of NiNH362 the resultant complex ion will be NiNH35H202 Once the released ammonia molecule may dri away if this happens the complex must wait until it encounters another ammonia molecule before it can reform the original complex When Nien32 releases an N donor atom the product complex has the formula Nien2en39H202 where en39 is a sineg bound ethylene diamine At first glance this appears little different from the ammonia case but there is an important difference While one end of the en is loose the other is bound so the unbound end cannot dri away like the ammonia Thus there is a very high likelihood that it will rebind rapidly to reform the original complex ion The more points of attachment for a ligand the less likely that all will release from the metal Thus formation constants tend to increase with increasing numbers of donor atoms on chelates It is for this reason that EDTA is added to many foods eg mayonnaise and salad dressings to preserve freshness If metal ions that would catalyze spoilage get into the food from a spoon for example the EDTA binds to them so they are effectively deactivated Metals and Chelates in Living Svstems Read on your own 243 V 39 A r nnrdimtinn Chemistry The system of naming coordination complexes is in some ways similar to and in other ways different from naming simple inorganic salts The rules for nomenclature are 1 Name the cation first and anion second 2 Within a complex ion the ligands are named alphabetical order Numbering prefixes e g di tri are not used in alphabetizing 3 Anionic ligands end in the letter o while neutral molecules with a few exceptions retain their names 4 The prefixes di tri tetra penta and hexa are used to indicate the number of each ligand If the ligand name includes such a prefix the ligand name should be placed in parentheses and preceded by bis 2 tris 3 tetrakis 4 pentakis 5 and hexakis 6 UI V Ifa complex is an anion it should end in ate O V Place the metal oxidation number in parentheses as a Roman numeral following the metal name Table 242 lists some common ligands and their names as ligands Notably water and ammonia have significant name changes to aqua and ammine respectively A few examples that demonstrate these rules appear below Ex NiH206C12 hexaaquanickelII chloride Na3FeCN6 sodium hexacyanoferrateIII PtNH32C12 diamminedichloroplatinumII Coen3Br2 trisethylene diaminecobaltII bromide In these examples it is important to note that the alkali metal and halide counterions don t have numbering pre xes It is presumed that you can calculate their number from the other information in the name 244 Isomerism Compounds with the same composition but different structures are called isomers Isomers are very important in both coordination chemistry and in the organic chemistry that many of you will study next year Two broad categories of isomers are structural isomers which have different atoms bound to each other and stereoisomers which have the same atoms bound to each other but differ in their spatial arrangement It is easiest to learn this by seeing examples Structural Isomerism One type of structural isomerism is linkage isomerism This may arise when a ligand has more than one chemically distinct donor atom For example CoNH35N022 and CoNH35ONO2 are linkage isomers The blue letter indicates the donor atom O N I NH35C0N NH35Co Q 9 Q pentaamminenitrocobaltIH pentaamminenitrito cobaltIII Note the difference in the names Frequently linkage isomer ligands will have a dilTerent name for each coordination mode The other structural isomerism we will cover is coordination sphere isomerism This occurs when different ligands that are part of the overall formula bind to the metal For example three compounds exist with the general formula CrH206Cl3 They are CrHZO6Cl3 violet CrH205ClClZHZO green CrHZO4C12Cl2HZO green The last two complexes possess m solvent molecules These are molecules of solvent that occupy spaces in the lattice but are not chemically bound to the complex ions What kind of experiment might distinguish these compounds from one another if all you knew was the generic formula and possible alternative structures Stereoisomerism This is the most common and most important class of isomerism Geometrical isomerism occurs when the same ligands bind to different sites on the metal There are several types of geometrical isomerism In square planar complexes systems with one or two pairs of different ligands can usually exist in two different forms All other formulations eg 4 of the same ligand are limited to only one isomer In one isomer the same ligands lie directly opposite one another trans In the other isomer they lie in adjacent positions cis trans diamminedichloroplatinumII cis diamminedichloroplatinumII In octahedral complexes cislrcms isomerism is also possible as is an isomerism associated with pairs of the same ligands C transtetraamminedichlorocobaltIII cistetraamminedichlorocobaltIII In this case we assume the ligand appearing only twice is the ligand described by the prefix Where multiple pairs of ligands appear multiple pre xes are permitted Facial fac isomerism describes 3 of the same ligands lying on one face of an octahedron while meridional mer isomerism refers to 3 of the same ligands lying on a plane passing through the center of the compleX a meridian 0L factriamminetrichlorocobalt11 mer triamminetrichlorocobaltIII Optical Isomerism This type of isomerism occurs when mirror images of a molecule cannot be superimposed on each other The individual isomers are called enantiomers and molecules that eXhibit optical isomerism are said to be M Chirality is a very important property in biological systems because many biologically molecules are present as only one enantiomer The other enantiomer is either biologically inactive or in some cases hazardous Most biological optically active molecules are organic chemicals composed solely of C and H along with some or all of Cl N 9 O and F We can see chirality by using Coen32 as an example Figure 1 shows a Coen32 ion A and it s mirror image B The second shows ion A as it is rotated around the zaxis As you can see it is not the same as ion B No amount of rotating will get these molecules to appear identical to one another N 2 2 NIII I N NIIn I N C l m 390 N N ll A B Figure 1 2 2 2 2 K N KT N N N NIIC0N A IIIIC0N NIIC0 NI11C0N Nil N N I M Nil NV I N N N J N Figure 2 Almost all chemical physical properties of optical isomers are identical For example enantiomers have identical boiling points melting points color density and reactivity with non chiral molecules However they do differ from each other in two important ways 1 Enantiomers rotate polarized light in opposite directions by an equal amount In polarized light the waves are all aligned 2 They react differently with other chiral molecules Since many biologically active molecules exhibit optical isomerism this has important implications for living systems When chiral molecules form from achiral starting reagents each enantiomer forms in equal proportion The mixture of products is called a w mixture When they form from chiral reagents one enantiomer is usually preferred sometimes exclusively 245 Color and Magnetism Almost all colored substances fall into one of two categories 1 transition metals with partially lled d orbitals and 2 organic molecules with extended delocalized systems of TE bonding These molecules include the dyes used to color your clothing and will be discussed next year in your organic chemistry class In this chapter we discuss case 1 Metals with no d electrons an or 10 d electrons d10 are ususually colorless Examples include Sc3 Ti4 and Zn2 examples include Ti02 the base in house paint and ZnO zinc oxide which you ve seen in sunblock Why do these ions yield colorless compounds To answer this we must rst answer the question what is color and why do we see it at all The human eye can only see wavelengths of light between about 400 and 700 nm Your eyes cannot detect wavelengths signi cantly outside this range Seeing occurs when light moves directly from a source to our eyes passes through a substance on the way to our eyes or re ects off of an object on its way to our eyes We ll consider the nal two methods as they are most relevant here If an object absorbs all of the light passing through it or striking it the object appears black because no re ected light makes it to our eyes Assuming the incoming light is white an object that absorbs little or none of the light appears white opaque objects or colorless transparent objects Gray occurs when some but not all of the light is absorbed and all wavelengths are absorbed equally If an object absorbs all but one wavelength of light we see the color of that wavelength If it absorbed only one wavelength of light we would see the complementary color to that wavelength Do you remember the color wheel from art class The primary colors are red yellow and blue The secondary colors are their complements green purple and orange l 1 respectively In actuality there will be different shades of these colors and many compounds absorb at more than one wavelength but the general idea still holds We determine the wavelength of maximum absorbance kmax by passing a light beam containing each wavelength of light through a sample and measuring how much light is absorbed at each wavelength Figure 2426 p 967 provides a view of a relatively simple visible light spectrophotometer In some respects it is similar to the Spec20 you use in CHM 218 Magnetism If one thinks of an electron as a spinning particle then it will generate a magnetic eld Since each unpaired electron will generate a similar magnetic eld they won t be exactly equal because they interact with one another knowing the size of a magnetic eld for a complex tells us the number ofunpaired electrons it possesses One might wonder why this is necessary since if you know the number of electrons in d orbitals the number of unpaired electrons should be easy to determine It turns out it isn t so easy For example if an octahedral complex ion has 5 d electrons it may have 5 unpaired electrons or only 1 We will shortly see why this occurs 246 CrystalField Theory The previous observation and others led to the realization that the electronic structure of transition metal complexes must be more complicated than originally believed A theory that does a very good job of predicting the electronic behavior of transition metal complexes is called crystal eld theory We will begin by considering an octahedral complex 6 donor atoms Remember there are 5 d orbitals at each energy level principle quantum number n dxy dxz dyz dzz and dxzyz Herea er I will refer to them as xy xz yz 22 and x2y2 respectively gtx gt 9 dxy It is easiest to approach this theory by considering a all metal cation a cation with only 1 electron in its d orbitals and extrapolating to the other cases In this situation there are two possible ways of placing the electron In possibility 1 one in ve complexes will have an electron in the xy orbital where it remains constantly Likewise 20 of the complexes will have the electron in the xz yz 22 and x2y2 orbitals respectively Again the electrons are locked into their respective orbitals The other possibility is that the electron is free to roam from orbital to orbital spending a statistical amount of time in each orbital What this means is that the electron spends 20 of its time in the xy 20 in ya etc Which is the correct view If you think of the electron as a wave each orbital has a node at the nucleus In other words the wave equations that describe the 5 orbitals all equal the same value at this point Thus an electron at the nucleus is equally likely to exit the node in any of the ve d orbitals Thus this description is correct Now consider the 6 ligands that will attach to the metal when the complex forms They begin at in nite distance from the metal and begin to approach it To minimize their interaction zaxis L L yaxis I L L xaxis 13 with each other VSEPR theory they will approach along opposite ends of the 3 coordinate axes ie one each from the x x y y 2 and 72 directions The lone pairs of electrons are attracted to the metal cation and are pulled towards it As the distance between the metal and ligands drops below in nity the ligand electrons are not only attracted to the metal nucleus they repel the metal d electron The alignment of the d orbitals now becomes important Two orbitals line up along the coordinate axes z2 and x2 2 and 3 inbetween xy xz and yz Ligand electronidelectron repulsion reaches a maximum along the coordinate axes so as ligands approach the energy of the 22 and x2y2 orbitals increases relative to the xy x2 and yz orbitals This can be shown pictorially as When the d orbitals split the 3 that drop in energy lower by 25 A0 while those that increase in energy do so by 3 5 A0 Thus there is no net change in energy for the orbitals in total Since we only have one d electron it goes into the lower energy set Now it will spend onethird of the time in each of the three orbitals When light with an energy equal to the gap A0 is absorbed the electron jumps from the lower set to the upper set Light not equal to this gap or any other gap in the compound passes through or re ects off A0 is related to wavelength by E hV and c 7N T 22 22 22 x22 xy hv y W77 WFyT It just happens that the energy gap between the split d orbitals corresponds to the visible region of the electromagnetic spectrum We must now deal with the different colors observed for complexes Why is it that different metals yield different colored complexes even though the 14 same ligands are bound to each e g CoH206C12 red vs NiH206C12 green Likewise why do different ligands cause complexes to exhibit dilTerent colors NiH206C12 green vs NiNH36C12 lavender The oxygen in water is more electronegative than the nitrogen in ammonia For this reason its lone pair orbitals don t project out into space as far ie it holds its electrons more tightly Thus as a water molecule approaches the metal its lone pairs interact less strongly with the metal than does nitrogen s lone pair So one expects A0 to be smaller for water since the gap is proportional to the level of interaction A general ordering of ligands has been determined Cl39 lt F39 lt H20 lt NH3 lt en lt CN39 A mcreasmg In a similar manner a higher charge on the metal draws ligands in closer and increases both interaction and A0 In the complexes above Ni2 is smaller than C02 and this should lead to greater metalligand interaction and a larger A0 We now return to the topic that ended the last section why does the number of unpaired electrons vary from complex to complex In 611 a0 and a3 complexes there is no ambiguity about lling the orbitals they go in the lower orbital set each into separate orbitals so that none is paired In a 614 metal things are a little different The fourth electron may either go into the upper set of orbitals or it may pair with an electron in the lower set In the latter case energy must be supplied to overcome the repulsion of the electrons occupying the same orbital If A0 is larger than the energy required to pair the electrons the electron goes into the lower orbitals If A0 is smaller than the pairing energy it goes into the upper set of orbitals L 22 x2y2 z x2y2 i L T LL i T xy xz yz 30 W W Epalr gt A0 Eparr lt A0 A complex with the larger number of unpaired electrons is called high spin while the one with the smaller number is called low spin Skip the section on Tetrahedral and Sguare Planar Complexes February 12 2005


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