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Organic Chemistry I

by: Ryan Kub

Organic Chemistry I CHEM 251

Ryan Kub
GPA 3.77

Mark Brandt

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Mark Brandt
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This 14 page Class Notes was uploaded by Ryan Kub on Monday October 19, 2015. The Class Notes belongs to CHEM 251 at Rose-Hulman Institute of Technology taught by Mark Brandt in Fall. Since its upload, it has received 18 views. For similar materials see /class/225120/chem-251-rose-hulman-institute-of-technology in Chemistry at Rose-Hulman Institute of Technology.


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Date Created: 10/19/15
General Chemistry A Guided Review for Study This review is by no means intended to be a completely comprehensive review of general chemistry This guide is intended to serve as a reminder of selected topics that must be reviewed before beginning organic chemistry You will be quizzed on your understanding and comprehension ofthe topics listed in the quotReview of General Chemistry Topicsquot and this study may help to refresh your memory ofthose topics If you nd that the level of detail here does not refresh your memory ofa topic you should look up that topic in a general chemistry text for more detail Chemical Equilibrium Law of Mass Action For a reaction jA kB k0 mD the law ofmass action is represented by the following equilibrium expression CHD AlBY The reaction quotient Q is obtained by applying the law of mass action but using initial concentrations instead of equilibrium concentrations For 1 Q K system is equilibrium no shift will occur 2 QgtK system shifts left 3 QltK system shifts right where K is the equilibrium constant Significance ofthe Magnitude ofK K K K2 KM K n KC is the equilibrium constant and the subscript 0 indicates that it is in terms of concentration when the reaction occurs in the gas phase the equilibrium constant is Kp where p indicates that it is in terms of pressure refer to a physical chemistry text for a discussion on activities and standard states For dilute solutions KC and Keq are used interchangeably Ka refers to the equilibrium constant for acid dissociation reactions Kb refers to the equilibrium constant for base dissociation and K refers to the autoionization ofwater H20I Haq OH39aq For KW 10X10391 HOH39 then ifH gt 10397 then the solution is acidic and if OH39gt 10397 then the solution is basic By applying LeChatelier39s principle pure water is a system in equilibrium with its ions and adding either acid H or base OH39 will cause the equilibrium to shift to the left toward neutral water that is water ionization is inhibited in acidic or basic solutions We use a logarithmic scale known as the pH scale standing forthe negative ofthe power ofthe Hydrogen ion concentration defined by pH og10H and ifwe take the og10 of both sides ofthe previous equation then we get pH pOH pKW 1400 which is temperature dependent Summa ofe uilibrium constantvalues As the value of As the value of As the value of As the value of KC Kp T the K3 T the the KsJ l the products products acid strength T solubility reac tan ts reac tan ts Copyright 2007 Rebecca DeVasher 1 Relationship between Kand Spontaneous Reactions Gibbs free energy G is a measure ofreaction spontaneity G H TS H enthalpy T temperature and S entropy and AG AH TAS AG 0 condition allows reversible or equilibrium processes to be described As a reaction proceeds the free energythat guarantees its spontaneity is expended and eventually the system will reach equilibrium When equilibrium has been attained AG Gp GR O is the equilibrium condition and AG lt O at constant T P results in a spontaneous process we have AG Gp GR O AGO RTaneq where Qeq K and by substituting Q at equilibrium for K 7A6 AG RTan or K e RT At a fixed temperature K is truly a constant since AGO depends only on T the nature ofthe system and the definition ofthe standard state see quotstandard heats of formationquot For 1 AG 0 K 1 2 AG lt0 K gt 1 3 AG gt O Klt1 AcidBase Equilbria Strong and Weak Acids and Bases Strong acids are strong electrolytes which for practical purposes are assumed to ionize completely in water Most ofthe strong acids are inorganic acids hydrochloric acid nitric acid perchloric acid and sulfuric acid not that H2804 is a diprotic acid thus it will have two dissociation constants Most acids are weak acids which ionize only to a limited extent in water At equilibrium aqueous solutions of weak acids contain a mixture of nonionized acid molecules H3O ions and the conjugate base according to the acid dissociation constant Kan Examples ofweak acids include hydro uoric acid acetic acid and the ammonium ion 80 when calculating the concentration of HF in a solution one has to account forthe following condition HFaq Haq F39aq K 2Ka 70x10394 Notice that the equilibrium constant is a unitless number it is a ratio The limited ionization ofweak acids is related to the equilibrium constant for ionization Like strong acids strong bases are strong electrolytes examples are metal hydroxides NaOH eg and weak bases are weak electrolytes N H3 eg Nearly all weak bases act as bases not by donating OH39 to the solution but by splitting or hydrolyzing waterto produce OH39 Consider ammonia NH3aq H20I NH4aq OH39aq K st179x10395 Chemical Thermodynamics Relationship between Enthalpv Chanqe AHr n and Heat Flow qr n For an isobaric process AH qp The change in enthalpy ofa system has no easily interpreted meaning except at constant pressure where AH heat For a chemical reaction the enthalpy change is given by the equation AH Hproducts Hreacfants Copyright 2007 Rebecca DeVasher 2 At constant pressure exothermic means AH is negative endothermic means AH is positive The quotheat ofreactionquot is the enthalpy change per mole The origin ofthe heat ofa reaction can be traced to the difference in bond energies D0 between the product and reagent molecules Bond breaking is endothermic so bond enthalpies are positive Chemical reactions usually absorb or release heat energy must be absorbed to break a chemical bond amp energy is released when a chemical bond forms see quotbond energiesquot below Comparing Thermochem ical Quantities Sl definition units type temperature hotnesscoldness propertythat controls K intensive direction of heat flows property thermal energy due to molecular motions J extensive energy property heat transfer ofthermal energy due to a J process temperature difference enthalpy adjusted thermal energy J extensive property Standard Heats of Formation AHf Definitions of Standard States For a gas the standard state is a pressure of exactly 1 atm For a substance present in a solution the standard state is a concentration of exactly 1 Mat an applied pressure of1 atm For a pure substance in a condensed state liquid or solid the standard state is the pure liquid or solid For an element the standard state is the form in which the element exists is most stable under conditions of1 atm and the temperature of interest usually 25 C The standard enthalpy offormation AHf ofa compound is de ned as the change in enthalpy that accompanies the formation of1 mole ofa compound from its elements with all substances in their standard states Key Concepts for Doing Enthalpy Calculations When a reaction is reversed the magnitude ofAH remains the same butthe sign changes When the balanced equation for a reaction is multiplied by an integer the value of AH forthat reaction must be multiplied bythe same integer The change in enthalpy for a given reaction can be calculated from the enthalpies offormation ofthe reactants and products AH ZAH f products ZAHfOreactams Elements in their standard states are not included in the AHremon calculations That is HfO for an element in its standard state is zero Hess39s Law Because the energy ofa given element or compound see quotchemical bond energiesquot at a given pressure and temperature is an intrinsic property ofthat substance the Copyright 2007 Rebecca DeVasher 3 heats ofreaction for reactions involving one or more of he same elements or compounds are related to each other as rst shown by Hess in 1840 Hess39s law of constantheatsummation states thatthe heat of any reaction that can be obtained by adding other reactions is given by the same sum ofthe heats ofthose other reactions That is ifreaction a reaction b reaction c then AHa AHb AHC stoichiometric multipliers that adjust the amounts ofreactions to add also apply to the AH relation The heat ofa reaction carried out backward products 9 reactants is the negative of the fonNard heat Hess39s law implies that we do not have to measure the heat of every possible reaction as long as we know the heats of some component reactions The standard reactions that chemists have chosen are the formation reactions those in which a compound is formed from its constituent elements in their most stable states at 1 atm and 298 K Chemical Bond Energies Do We can obtain information about the strength ofa bonding interaction by measuring the energy required to breakthe bond the bond energy Do In an exothermic process q lt 0 the bonds in the products are stronger on average than those ofthe reactants vast majority of reactions That is more energy is released in forming the new bonds in the products than is consumed in breaking the bonds in the reactants The net result is that the quantity ofenergy APE is transferred to the surroundings through heat Many endothermic q gt 0 reactions are also known To systematize the tabulation of experimental reaction heats the results are reported as enthalpy changes at 100 atm pressure and 29815 KNTP per mole ofreaction AHozgs in either kilocalories per mole or kilojoules per mole For example the bond energy of the C H bond D0C H is taken as an average ofthe four bonds in CH4g and differs slightly from that in any individual bond in CH4 or in other molecules possessing that bond ln molecules such as benzene CBHB a further complication occurs because of delocalized 7 bonding which makes the molecule more stable than one would predict based on the bond energies corresponding to a single Lewis structure resonance form The difference between the predicted and experimental heats offormation of molecules such as benzene is called the resonance energy When estimating AH from bond enthalpies one should adopt the following strategy imagine reaction as a dissociation ofreactants into atoms b recombination of atoms into products 1 Add enthalpies for all product bonds 2 Add enthalpies for all reactant bonds 3 AH is approximately the difference between the product and reactant bond enthalpies Summary of Enthalpy of Reaction Reaction type exothermic endothermic heat is released absorbed reaction vessel temperature rises falls enthalpy change is negative positive number or strength of bonds increases decreases Copyright 2007 Rebecca DeVasher 4 Chemical Kinetics Reaction Rate and the Reaction Rate Constant k Rate is expressed in the units of moles per liter per second mol39L39139s391 or molarity per second M39s391 decrease in concentration reactants increase in concentration products rate time time In general for a reaction aA b8 9 cC dD lll i a At b At c At al At rate Factors Affecting Reaction Rate 1 reaction concentration 2 temperature 3 medium 4 catalysts Differential and Integrated Rate Laws For nearly all forward irreversible reactions the rate is proportional to the product of the concentrations ofthe reactants each raised to some power For a general reaction ofthe type aA bB cC dB the rate is proportional to AX By that is rate kAXBy This expression is the rate law forthe general reaction above where k is the rate constant The exponents X and y are called the orders of reaction X is the order with respect to A and y is the order with respect to B These exponents may be integers fractions or zero and mustbe determined experimentally The overall reaction order is the sum ofa exponents in our example xy Experimental Determination of Rate Constants The values of k x and y in the rate law equation rate kAXBy must be determined experimentally One method is to measure the rate initially as a function ofthe initial concentration ofthe reactantsA and B called the initial rate method Another way is to collect concentration data overtime and solve the integrated form ofthe rate equation by setting the general rate law equal to the rate expression as a function of disappearance of product and then applying calculus Arrhenius Energy ofActivation Ea The dependence ofthe rate constant ofa reaction on temperature can be expressed by the following equation known as the Arrhenius equation E k A679 where A represents the collision frequency from the collision theory ofchemical kinetics Ea is the activation energy ofthe reaction kJmol Rthe gas constant 83145JK39mol Tthe absolute temperature and e the base ofthe natural logarithm Copyright 2007 Rebecca DeVasher 5 scale A plot ofln kversus 1T the rate constant at various temperatures must be collected experimentally whose slope is m is equal to EaR and whose intercept b with the ordinate the y axis is In A from the equation 1nk E R In A Chemical Reaction Mechanisms The mechanism ofa reaction is the actual series of steps through which a chemical reaction occurs Knowing the accepted mechanism ofa reaction often helps to explain the reaction39s rate position of equilibrium and thermodynamic characteristics Atheoretical mechanism can be supported by experimental kinetic data not the other way around Kinetic data can support mechanisms but proposed mechanisms must agree with experimental data ifthey are to be a theoretical framework on which to base further explanation of events For an overall reaction ofthe type A2 2B 2AB a possible mechanism would be the following two step process Step 1 A2 B 9 A2B slow Step 2 A2B B 9 2AB fast Note that these two steps add up to the overall net reaction a necessary requirement for a plausible mechanism A2B does not appear in the overall net reaction because it is neither a reactant nor product but is an intermediate Reaction intermediates are often dif cult to detect but can be supported through kinetic data as mentioned earlier For example suppose thatthe rate law ofthe above reaction was rate kA2B the above mechanism would be plausible The slowest step in a reaction mechanism limits the overall rate ofa reaction and this step is often called the ratelimiting step orthe ratedetermining step OxidationReduction Redox Reactions Reactions involvinq transfer ofelectrons between species The law of conservation of charge states that an electrical charge can be neither created nor destroyed Thus an isolated loss or gain of electrons cannot occur oxidation loss ofelectrons OIL and reduction gain ofelectrons RIG must occur simultaneously resulting in an electron transfer called a redox reaction An oxidizing agent causes another atom in a redox reaction to undergo oxidation and is itself reduced A reducing agent causes the other atom to be reduced and is itself oxidized Assigning oxidation numbers 1 The oxidation number in free elements is zero 2 The oxidation number for a monoatomic ion is equal to the charge ofthe ion 3 The oxidation number of each Group lA element in a compound is 1 The oxidation number of each Group A element in a compound is 2 4 The oxidation number of each Group VllA element in a compound is 1 except when combined with an element of higher electronegativity 5 The oxidation number of hydrogen is 1 unless it is in compounds with less electronegative elements than hydrogen and the oxidation Copyright 2007 Rebecca DeVasher 6 state Wm be 6 Th rhost compoundsthe oxroatroh hurhoerot Oxygen T572 The or quot eTectrohegatrvethah oxygen 7 The Sum or the oxroatroh numbers or 5H the atoms ewmo preseht rh a poTyatorhrc roh rs eguaT to the charge otthe roh hr th or r method aTso known as the Tonretectron method rh wmch the eguatroh rs separated mtotwo haTtreactrohs rme oxroatroh part ahothe reductroh part Each hatfrreacuon ba anced redox readton Molecular and lo c Structure Pertodtc Progemes Note the effecuve nudear chargezecah exp am aH perroorctrehos as WeH as cherhrcaT propertres ncvexmv araarsremw ncvexmv Wmquot my omng mp mans 47 KECaSc Tt ncvusmv searcherth ncvusmw xllTananvV chuxmv Elam mans 1 Atormc Radu cerrters oftwo atorhs of that eterherrtthat arejust toughrhg each other H gerreraL the gwen group the atoms wrth the targest atorhrc raoh Wm be tocateo at the bottom or groups and h Grou t The effecwe nudear charge rhcreases steaorty across a penod ahothrs causesthe atorhrc raorus to decrease 2 tohrzatroh Energy The rohrzatroh energy E or sorhetrrhes reterreo to asthe rohrzatroh potehtrat rs he on Copynght 2007 Rebecca DeVasher 7 Removing an electron from an atom always requires an input ofenergy is endothermic The rst ionization energy is the energy required to remove one valence electron from the parent atom the second ionization energy is the energy needed to remove a second electron from the univalent ion to form the divalent ion and so on Successive ionization energies grow increasingly large Ionization energy increases from left to right as the atomic radius decreases Moving down a group the ionization energy decreases as the atomic radius increases Group elements have low ionization energies because the loss ofan electron results in the formation ofa stable con guration octet 3 Electron Affinity Electron af nity is the energy change that occurs when an electron is added to a gaseous atom and it represents the ease with which the atom can accept an electron The strongerthe attractive pull ofthe nucleus for electrons effective nuclear charge or Ze the greaterthe electron af nity will be Two sign conventions 1 the more common one states that a positive electron af nityvalue represents energy release when an electron is added to an atom 2 the other states that a negative electron affinity represents a release ofenergy 4 Electronegativity Electronegativity is a measure ofthe attraction an atom has for electrons in a chemical bond The greater the electronegativity ofan atom the greater its attraction for bonding electrons Electronegativity values are not determined directly they are not experimentally measured they are calculated values The Pauling electronegativity scale is the most commonly used method and ranges from 07 9 40 the most electronegative element fluorine Electronegativities are related to ionization energies in a directly proportional relationship Covalent and lonic Bonding Classification of chemical bonds begins with identifytwo distinct types ionic bonds and covalent bonds In general there is a much more accurate buttime consuming method for distinguishing between ionic and covalent bonds the difference in electronegativity oftwo atoms involved in a bond must be greater than 17 In a covalent bond the electronegativity difference is considerably less than 17 Polari of Bonds and ofMolecules ln ionic bonding an electron from an atom with smaller ionization energy is transferred to an atom with a greater electron af nity and the resulting ions are held together by electrostatic forces In covalent bonding an electron pair is shared between two atoms although not necessarily shared equally on average the majority ofthe electron density will be concentrated on more electronegative ofthe two elements in a polar chemical bond lfthe two atoms involved in a covalent bond have the exact same electronegativity orthe electronegativity values are very close eg C H the bond is considered to be nonpolar For polar bonds the bond can be considered partially ionic and partially covalent due to the electronegativity difference between the two elements Copyright 2007 Rebecca DeVasher 8 Polar bonds are represented in the following way 6quot 539 H Cl A molecule with polar bonds may not be a polar molecule the bond dipole moments may cancel each other out resulting in a nonpolar molecule Although a m9lecule with polar bonds need not be polar a polar molecule must have polar bonds A molecule with a net dipole moment is called polar because it has positive and negative poles H20 for example is a polar molecule See the illustration below for the notation of polarity region of excess electron density corresponds to the arrow end region ofelectron de ciency corresponds to the quotpositivequot portion ofthe arrow 5quotgtlt 0 5 5 H H polarity vectors net dipole for H20 moment for H2O Another important topic intermolecular forces will not be discussed Weakerthan intramolecular chemical bonds intermolecular forces are of considerable importance in understanding the physical properties ofmany substances The following discussions concern covalent bonding Properties of Covalent Bonds 1 Bond Length Bond length is the average distance between the two nuclei ofthe atoms involved in the bond As the number of shared electron pairs increases the two atoms are pulled closertogether leading to a decrease in bond length Thus for a given pair of atoms a triple bond is shorterthan a double bond which is shorterthan a single bond 2 Bond Energy Bond energy is the energy required to separate two bonded atoms For a given pair of atoms the strength ofa bond and therefore the bond energy increases as the number of shared electron pairs increases 3 Bond Order Bond order is the number of bonds between two elements in a covalent bond Bond order is related to bond length and bond energy Bond order can be determined by several methods see discussion on bonding theories In general as the bond order number of bonds increases the bond energy increases and the bond length decreases A single bond has a bond order of 1 double bond a bond order of 2 and a triple bond has a bond order of 3 Copyright 2007 Rebecca DeVasher 9 Covalent Bond Notation The shared valence electrons ofa covalent bond are called the bonding electrons The valence electrons not involved in the covalent bond are called nonbonding electrons The unshared electron pairs can also be called lone electron pairs A convenient notation called a Lewis structure is used to represent the bonding and nonbonding electrons in a molecule facilitating chemical quotbookkeepingquot 1 Lewis Structures Write the skeletal structure ofa compound ie the arrangement of atoms Forthis exercise you must recall the bonding tendencies ofthe elements For example HCN H C N or H CEN Count all valence electrons ofthe atoms H has 1 valence electron C has 4valence electrons N has 5 valence electrons 10 valence electrons Obey the octet rule and minimize formal charge while lling in the remaining valence electrons not accounted for in the skeletal structure H CEN 2 Formal charges The number of electrons officially assigned to an atom in a Lewis structure does not always equal the number ofvalence electrons ofthe free atom The difference between these two numbers is the formal charge ofthe atom Formal charge can be calculated using the following formula 1 Formal charge valence electrons Ebondz39ng electrons nonbonding electrons The sum of all the formal charges equals the total charge on the molecule or ion We calculate formal charges on individual atoms by subtracting the number ofvalence electrons assigned to an atom in its bonded state from the number ofvalence electrons it has as a neutral free atom 3 Resonance For some molecules two or more Lewis structures are needed to illustrate quotsnapshotsquot ofthe actual molecule These quotsnapshotsquot are called resonance structures The combination or weighted average ofthe structures ofthe resonance form more accurately describes the molecule and is known as the resonance hybrid Resonance structures are represented with a double headed arrow between them It is important to note thatthe arrows ofthis type describes equilibrium and arrows ofthis type lt describes resonance a subtle but VERY IMPORTANT difference See the structure of carbonate below Copyright 2007 Rebecca DeVasher 10 All three represent the structure of carbonate Each Lewis structure represents a resonance form and all three since the energies are equivalent averaged together represent the resonance hybrid Resonance structures are not structures forthe actual molecule or ion they exist only on paper Summary Rules for Resonance 1 Resonance structures exist only on paper 2 In writing resonance structures we are only allowed to move electrons 3 All ofthe structures must be proper Lewis structures 4 The energy ofthe actual molecule is lower than the energy that might be estimated for any contributing structure 5 Equivalent resonance structures make equal contributions to the hybrid and a system described by them has a large resonance stabilization 6 The more stable a structure is when taken by itself the greater its contribution to the hybrid a The more covalent bonds a structure has the more stable it is b Structures in which all ofthe atoms have a complete valence shell ofelectrons ie s2p6 are especially stable and make large contributions to the hybrid 0 Charge separation decreases stability d Resonance contributors with negative charge on highly electronegative atoms are more stable than ones with negative charge on less or nonelectronegative atoms Theories of Chemical Bondinq VB VSEPR MO Many molecules contain atoms bonded according to the octet rule Exceptions to this rule include but are not exclusive to hydrogen duet lithium duet beryllium quartet boron sextet and elements below the second period which can have quotexpanded octetsquot 1 Valence Shell Electron Pair Repulsion VSEPR Theory The valence shell electron pair repulsion VSEPR theory uses Lewis structures to predict the molecular geometry of covalently bonded molecules It states that the threedimensional arrangement of atoms surrounding a central atom is determined bythe repulsions between the bonding and the nonbonding electron pairs in the valence shell ofthe central atom These electron pairs arrange themselves as far apart as possible thereby minimizing repulsion 2 Valence Bond VB Theory This theory involves the hybridization of he central metal ion in a transition metal complex The formation ofempty hybrid orbitals on the central metal ion enables a ligand to donate a pair ofelectrons to form a covalent bond with the ligand positioned Copyright 2007 Rebecca DeVasher 1 1 in a definite geometry Geometry magnetic properties and the possibility of color are predicted 3 Molecular Orbital MO Theory Theory describes the arrangement of electrons between two nuclei by considering wave functions 1p and symmetry restrictions Predicts bond orders relative bond energies and magnetic properties 4 Orbital Hybridization The formation of hybrid atomic orbitals on a central atom accounts for directional bonds Bond orders relative bond lengths polarity dipole moment and planarity are predicted This theory is most applicable to compounds containing carbon or nitrogen Molecular Geometrv th 1 Molecular Geometry Geometry is determined by the number of bonding and unshared pairs of electrons on the central atom Shape indicates that the bonding and nonbonding unshared electrons pairs repel each other Rules used to predict molecular geometry lCount the number of valence electrons 2 Place a pair of electrons between the atoms bound together 3 Each noncentral terminal atom is given electrons for its inert gas con guration 4 Any additional electron pairs become lone unshared electrons pairs on the central atom 5 The geometry is determined by the number of bonding and unshared pairs of electrons on the central atom Here are some example structures for molecules and polyatomic ions that consist of a central A atom bonded to two or three B atoms 7 UGL C02 802 NH3 BTF3 COC OSO HN H ClBCl FBrF Bond angle 180 ll93 1067 120 862 Molecular Linear Bent Trlgollal Trlgonal Tshaped pyram1 dal planar geometry This term can be very misleading This structure does not resemble a quotTquot shape Copyright 2007 Rebecca DeVasher 12 Summary of Observed Geometries No of Lone Atoms Bound to the Central Atom Electron Pairs 2 3 4 5 6 0 linear linear mgonal tetrahedral mgona octahedral planar b1pyram1dal trigonal sawhorse square 1 linear bent 120 pyramidal seesaw pyramidal 2 linear bent square 109 T Shaped planar 3 linear linear 2 Bond Angles Bond angles are the angles that result from the atomic arrangement or the geometry of a molecule or ion The bond angle can be predicted from the VSEPR geometry of a molecule or ion In a tetrahedral arrangement for example the bond angles will be measured as follows Copyright 2007 Rebecca DeVasher 39V 5039 where 6 1200 H QH Overview of molecular geometries Electron Pairs 2 3 4 5 6 Molecular BAuB B B Geometry BAB BAB BAQ39B B Bf VB Zero Lone Linear Trigonal B B B 39 Tri onal Palm planar TEtrahedral bipygamidal Octahedral Molecular i E T B1 A B Geometry A AquotB 39 Bv I B One Lone B B B B Q B Pair Bent Trigonal square Vshaped pyramidal Seesaw pyramidal Molecular E B l B G t Au A As eome ry B B B I B B B Two Lone B Pairs Bent Square V39ShaPEd T39Shaped planar Molecular 2quot Geometry BAB Three Lone Pairs LInear Copyright 2007 Rebecca DeVasher


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