Principles of Chemistry I
Principles of Chemistry I CHEM 1307
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Bonding amp Molecular Structure CH3 0 CH3 Hg 11 N CN II N HC CC Cc MIL 2H C 1 0 T TH2 HZCCNCH Chemical Bonding A chemical bond is formed when the valence electrons between atoms reorganize Ionic Bond One or more valence electrons is transferred from one atom to another creating a positive and negative ion The bond is the attractive force between the ions NM ct a Namjet Na ttz Covalent Bond Valence electrons are shared between atoms No ions are formed in a covalent bond Covalent Bonding and Valence Electrons Covalent bonds are typically formed between nonmetals while ionic bonds are formed between a metal and a nonmetal Chemical reactions result in the gain loss or rearrangement of valence electrons The core electrons are not involved in bonding Element Periodic Group Core Electrons Valence Electrons Total Configuration Main Group Elements Na 1A Si 4A As 5A Transition Elements Ti 48 Co 88 Mo 68 15 22522p6 Ne 1522522p6 Ne 1522522p63523p63d1 Ar3d10 1522522pquot3523p6 Ar Ar Kr 351 3523p2 4524p3 3d24s2 3d7452 4d5551 Ne3s1 Ne3523p2 Ar3d 4524p3 Ar3d2452 Ar3d7452 Kr4d5551 Nominated for the Nobel Prize 34 times but never won Mysterious circumstances surround his death in his lab at UC Berkeley While working with HCN he collapsed under a work bench and died Suicide accident or heart attack The Life and Death of G N Lewis v Proposed the concept of the covalent bond after drawing cubical structures to represent the position of electrons in specific atoms 439 lg My wmw C J39 g 9 r 70 3 5 g M a F g m 41 a Lewis Dot Structures Between the s and p orbitals in a main group element there can be eight valence electrons 4 electron pairs We can visually depict the number of valence electrons for any element and this allows us to predict reactivity TABLE 82 Lewis Dot Symbols for Main Group Atoms 1A 2A 3A 4A 5A 6A 7A 8A ns1 ns2 nsznp1 nsznp2 nsznp3 nsznpquot nsznp5 nsznp5 Li Be 3 C 91 F1 we Na Mg 39Al Si 39P39 23957 d A One electron is placed on each of the four sides of the symbol before any electrons are paired Hund s Rule 539 Lewis Dot Structures When drawing a Lewis dot structure a bond results when one or more electron pairs is shared between two atoms Lone pair of u II n IFFZ gtZFZFI ZF Fi Shared or bonding electron pair We count each bond as two electrons Each atom is surrounded by eight electrons or an octet so each atom has obtained a noble gas configuration Lewis Dot Structures The more electrons that are shared between atoms the more bonds are formed In dinitrogen N2 there are a total of 10 valence electrons In order to form an octet around each N a triple bond is made NN Each N atom is surrounded by 8 electrons Each N atom donated 5 valence electrons to the structure In carbon dioxide there are a total of 16 valence electrons To form octets around each atom double bonds are made Double and triple bonds are common for C N and O Drawing Lewis Dot Structures A set of rules can make drawing Lewis clot structures easy Determine the arrangement of atoms in a molecule The central atom typically has the lowest electron affinity Hydrogen and halogens are typically terminal atoms preferring to form only one bond each Determine the total number of valence electrons Count up the valence e39 for each atom in the molecule For an anion add the number of e39 equal to the charge For a cation subtract the number of e39 equal to the charge Make a single bond between each pair of atoms If any e39 remain add them as lone pairs to terminal atoms If the central atom does not have an octet change lone pairs into bonding pairs to form multiple bonds Drawing Lewis Dot Structures Draw the Lewis structure for N02 Central Atom N Number of valence e39 5 N 2 x 6 O 1 charge 16 Single bond between each atom O N O 2 bonds 4 electrons Add lone pairs of e39 to terminal atoms l G u 0 2 bonds 4 electrons I O to N O I 6 lone pairs 12 electrons co co Make central atom have an octet 0 I G O N O 4 bonds 8 electrons 4 lone pairs 8 electrons Predicting Lewis Structures Knowing the number of valence e39 in an atom we can predict the number of bonds that will form in an uncharged molecule Carbon has 4 valence electrons and therefore needs to make four bond in order to obtain an octet of electrons Group 4A Group 5A Group 6A Group 7A C N 0 IF Nitrogen 5 valence electrons 3 bonds Oxygen 6 valence electrons 2 bonds Fluorine 7 valence electrons 1 bond Predicting Lewis Structures What is the Lewis structure of C2F4 Central Atoms C Number of valence e39 2 X 4 C 4 X 7 F 36 Single bond between each atom F F 5 bonds 10 electrons F I C F Add lone paIrs of e39 to termInal atoms FI F 5 bonds 10 electrons 12 lone pairs 24 electrons F C C FI Make central atoms have an octet I F F C C 6 bonds 12 electrons F F 12 lone pairs 24 electrons Oxyacids and Their Anions In the absence of water oxyacids are molecular compounds Lewis structures for the acidsanions have the same number of e39 o H u H N O N O l I a O H O TABLE 84 Lewis Structures of Common Oxoacids and Their Anions HNo3 H o N6 H3P04 z39o H H2504 6 H nitric acid l phosphoric acid I sulfuric acid I 39 IO f O H Q H 19 H r V 3 3 H 39 quotgit 9 T 0 Pg ht 393 H50quot o H m r3 em 0 p 03p a emquot b p b39 hydrogen sulfate ion l V I Q T QV I 27 2 HClOL CIJ H HOCl H g g 504 v 4quot perchloric acid 10Cll0 hypochlorous acid sul telon 9 0 cm 30 ocr perchlorate ion I hypochlorite ion Isoelectronic Species Molecules or ions that have the same number of valence electrons and comparable Lewis structures 2N20 NEN CEQ CEN CO and CM are very toxic but behave differently in water N2 is an inert gas that makes up 79 of air TABLE 85 Some Common Isoelectronic Molecules and Inns Representative Representative Formulas Lewis Structure Formulas Lewis Structure 8H CH NH H C039N0 o llio 7 H H39NiH o H N 30 Heir P0315047 no N H 01 MN SCN39 N20 oo NOZ 0E5 C52 Isoelectronic Species Is the acetylide ion C223 isoelectronic with N2 Acetylide ion 10 electrons Dinitrogen 10 electrons ecCe N N What common polyatomic ion is isoelectronic with HF HF 8 electrons OH39 8 electrons H H quot9 O Formal Charges The charge on an atom in a molecule or polyatomic ion The formal charges should add to the overall charge on the ion or should add to zero for a neutral compound Formal charge group number lone pair e39 12 bonding e39 The group number gives the number of valence electrons The lone pairs belong to the atom that they surround Bond pairs are divided equally between the bonded atoms Where does the charge reside in GO Number of valence e39 7 CI 6 O 1 charge 14 0 Formal charge 0 6 6 12 2 391 O Formal charge CI 7 6 12 2 0 ZCI c1 oe Formal Charges Where does the charge reside in CN39 Number of valence e39 4 C 5 N 1 charge 10 Formal charge C 4 2 12 6 1 C N Formal charge N 5 2 12 6 0 zc N Where does the charge reside in NF4 Number of valence e39 5 N 4 x 7 F 1 charge 32 F o o l F Iv F Formal charge N 5 0 12 8 1 F39 Formal charge F 7 6 12 2 0 Resonance Structures The possible structures of a molecule where more than one Lewis structure can be written differing in the number of bond pairs between a given set of atoms 201020 lt gt 020 0 The structure for ozone is the combination of the two resonance structures drawn above Since the central oxygen has one single bond and one double bond drawn to it the average bond length is 15 Resonance structures have equal energy The molecule alternates between the different forms Resonance Structures Benzene is a classic example of resonance Benzene is composed of six identical carbon atoms arranged in a ring i i i HCC H H C H HCCCH i ll lt gt ii i Cl l H Cl H H CI H H E H H H H The CC bond is intermediate between a CC double bond and a CC single bond Resonance Structures How many resonance structures exist for CO3 Drawthem 2 2 2 IOCAOI IOCOZ IOgCOI H I 9 0 9 No single structure accurately describes the carbonate ion The actual structure is a combination of the three resonance forms There is an equal distribution of electrons over all the atoms involved Where do the formal charges belong Exceptions to the Octet Rule Some atoms are surrounded by fewer than four pairs of electrons boron trifluoride 3 B 3 X 7 F 24 F Il3 F I F I The boron atom is short one pair of electrons so it is reactive Another atom can donate an electron pair to the boron coordinate covalent bond T 39 B 3939 Exceptions to the Octet Rule Some atoms are surrounded by more than four pairs of electrons This can occur in elements in the third period and higher Typically the central atom is bonded to F CI or O sulfur hexafluoride 6 S 6 x 7 F 48 Exceptions to the Octet Rule If there are more than four groups bonded to the central atom this indicates the octet has been exceeded The octet can also be exceeded with fewer than four bonded atoms due the to placement of lone pairs TABLE 8G Lewis Structures in Which the Central Atom Exceeds an Octet Group 4A Group SA Gmup 6A Group 7A Group 3 Sir Pri sr Xer 39F39 5Fquot 39 F 5F F 5 4f F F C l 39 XI 7 l 7 7 7 e F VF F I F39 r r 39 F SiFJ PF w M F F39 r 5L F P F FBl r r x dr39 r e F1F39 r r r r r r 39F l r F F H H The extra orbitals in the valence shell of elements located in period three and higher accounts for this observation Molecules with an Odd Number of Electrons When a molecule has an odd number of electrons it is impossible to obey the octet rule and at least one e39 must be unpaired How do you draw a Lewis structure for N02 Number of valence e39 N 5 2 x 60 17 Free Radicals Chemical species with unpaired electrons Free radicals are typically very reactive Examples include H Cl NO and N02 VSEPR Model Lewis structures allow us to predict the threedimensional geometry of molecules and ions The valence shell electronpair repulsion theory states that molecular geometry is determined based on the repulsion between electron pairs Molecules adopt the shape that minimizes the electron pair repulsions ie keeps electron pairs farthest apart The positions of the valence electrons define the angles between bonds VSEPR theory does not work well for compounds containing transition metals Central Atoms with Single Bond Pairs The simplest structures involve central atoms with only single covalent bonds no lone pairs 180 z a no 1095 90 120 90 M2 Ax m Ax5 Ax5 Exampie BeF Example BF3 Example39 CF Example39 PF Example ch Central atom does not Expanded octet have an octet only 339 period or higher Central Atoms w Single Bond Prs amp Lone Prs Electron Pair Geometry The geometry of a molecule or ion when all electrons lone and bonding are considered Molecular Geometry The geometry of a molecule or ion when only bonding electrons are considered Lone pairs of electrons still occupy space but are not included in the shape description of the molecule or ion Let s consider ammonia Number of valence e39 N 5 3 x 1 H 8 H lT H HN WH V Moleculargeometry H H trigonal pyramidal Lewis structure Electron pair Actual HNH geometry tetrahedral angle 10750 Effect of Lone Pairs on Bond Angles The electron pair geometry for ammonia is tetrahedral so we would expect the bond angles to be 1095 However the experimentally determined angles are 1075 This is because lone pairs of electrons take up more space than bonding pairs of electrons and force the bonding pairs closer together lone pairlone pair gt lone pairbond pair gt bond pairbond pair Four electron pairs tetrahedral electron pair geometry Methane CH4 Ammonia NH3 Water H20 4 bond pairs 3 bond pairs 2 bond pairs no lone pairs 1 lone pair 2 lone pairs Central Atoms with Expanded Octets Five electron pairs trigonal bipyramidal electron pair geometry 2 bond pairs 3 lone pairs 3 bond pairs 2 lone pairs There are two inequivalent positions the electrons can occupy Equatorial The positions in the trigonal plane that lie in the equator of an imaginary sphere around the central atom Two axial neighbors 90 apart Axial The north and south poles of the molecule Three equatorial neighbors 90 apart Central Atoms with Expanded Octets Six electron pairs octahedral electron pair geometry gtIlt 3 sra BrF5 Xer 6 bond pairs 5 bond pairs 4 bond pairs No lone pairs 1 one pair 2 one pairs There are no distinct positions so one lone pair of electrons can be placed in any position If a molecule has a second lone pair of electrons the pairs want to be as far apart as possible so they are placed 180 apart Molecular Geometry Draw the Lewis structure for ICIZ39 What is the geometry of the ion Central Atom I Number of valence e39 7 I 2 x 7 CI 1 charge 22 Lewis Structure Electron pair Geometry trigonal bipyramidal Molecular Geometry linear Multiple Bonds and Molecular Geometw Double and triple bonds involve more electrons than single bonds but do not change the overall molecular geometry 180 0921 Lewis structure Molecular structure linear electron pair geometry linear H H101 Cysteine HSCHZCHNH2COZH C1 electron pair geometry trigonal planar C2 C3 S N O electron pair geometry tetrahedral Bond Polarity amp Electronegativity Although we represent covalent bonding as equal sharing of electrons between two atoms the electrons are actually shared unequally when two different atoms are connected 6 6 6 partial charge Polar Covalent Bond e39 are closer to one atom which gets a 839 charge e39 are farther from the other atom which get a 8 charge The bond therefore has a positive and negative end Bond Polarity amp Electronegativity In ionic compounds the electron is transferred completely from one atom to the other so we use a instead of 5839 Bonds can be purely covalent purely ionic or anywhere in between Electronegativity x Proposed by Linus Pauling in the 1930s A measure of the ability of an atom in a molecule to attract electrons to itself F has the highest electronegativity while Cs has the lowest 1954 Nobel Prize 1962 Nobel Peace Prize Electronegativity Electronegativity increases going up a column and across a period to the right V Cr Mn Fe Co 16 17 15 18 19 l Nb Mo Tc Ru Rh A 16 22 19 22 23 Ta W Re 05 Ir 15 24 19 22 22 an 15719 E12519 10 14 20 24 30 40 Metals range from around 08 2 Metalloids are values around 2 Nonmetals have values greater than 2 Electronegativity What is the difference in electronegativity Ax for the following molecules 3A 4A 5A 6A 7A B 20 Al 5139 P 16 19 22 Ga Ge As 18 20 22 In Sn Sh Te 18 20 19 21 Tl Pb Bi Po At 16 23 20 20 22 AX 40 22 18 AX 35 22 13 AX 30 22 08 Electronegativity For the following bond pairs decide which is the more polar pair and indicate the negative and positive poles BF and BCl B is a metalloid value 2 and F is the MOST electronegative atom Cl lies below F in the same group therefore Cl is less electronegative 8BF5 is the more polar bond I SiO and PP Si is a Grp 14 atom while 0 is a Grp 16 atom therefore 0 is more electronegative A bond between two atoms of the same kind is nonpolar ESSi08 is the more polar bond I Electroneutrality Principle Sometimes formal charges and electronegativity are at odds To resolve this problem Pauling suggested that electrons will be distributed to give the lowest charges on each atom Which structure is favored Why Formal charges 0 0 0 1 0 1 Resonance structures OC0 IOEC Oi A B Which structure is favored Why Formalcharges 1 0 O O D 1 1 0 2 Resonance structures id CENC39i gt lt gt EOEC39 N37 A B C Molecular Polarity When a polar molecule is placed in an electric field the molecules align with the field Electric Field OFF Dipole Moment u The magnitude and direction of the poles in a molecule Units of debye named after Peter Debye 1936 Nobel Prize in Chemistry Dipole Moments Why are some molecules polar and others are not A molecule will be polar if 1 It has polar bonds 2 The e density is distributed r 3380 unevenly in the molecule No net dipole moment 5 7 A 5 H 20 J Terminal atoms same distance Molecule is bent dipoles from C atoms are planar do not cancel 5 Dipole Moments Determine whether the following molecules are polar 5 5 A 5 5 ctzco NH3 BF3 N0 net dipole M 147D u E 147D moment 7 Net dipole Net dipole J i r y 399 a 1 XO gt XC gt XC If one position is occupied by a lone pair or there are different terminal atoms the molecule will be polar Bond Order The number of bonds between two atoms Bond order 1 Bond order I 2 39Esc39gtoO 3915 Bond order ozone 3 Qairs of electrons 2 00 bonds Bond Length The distance between the nuclei of two bonded atoms Depends on size of atoms larger atom longer bond H F H CI H l Depends on bond order higher BO shorter bond Bond Dissociation Enthalpy The enthalpy change for breaking a bond in a molecule with the reactants and products in the gas phase gt CH3 CH3 ARH gt CH2 CH2 ARH kJmol HCECH g gt CH g CH g ARH 962 kJmol Breaking bonds is always endothermic Forming bonds is always exothermic Let s calculate the enthalpy change for the reaction H2 9 Clz g gt 2 HCI 9 Aern 2 AH bonds broken 2 AH bonds formed Bond Dissociation Enthalpy Let s calculate the enthalpy change for the reaction H2 9 Clz g gt 2 HCI 9 Which bonds are broken HH and ClCl Which bonds are formed 2 HCl Bonds BDE kJ mol HH 436 ClCl 242 HCl 432 Aern 2 AH bonds broken 2 AH bonds formed Aern 436 242 2 x 432 186 kJ Bond Dissociation Enthalpy Let s calculate the enthalpy change for the reaction N2 9 3 I39Iz 9 gt 2 NI39I3 9 Which bonds are broken NN and 3 HH Which bonds are formed 6 NH Bonds BDE kJmol NN 945 HH 436 NH 391 Aern 2 AH bonds broken 2 AH bonds formed Aern 945 3 x 436 6 x 391 93 kJ Bond Dissociation Enthalpy Let s calculate the enthalpy change for the reaction H o 0 ll H3C C CH3 9 H Hg gt H3C C CH3 9 H Bonds BDE kJmol CO 745 HH 436 C H 413 C0 358 OH 463 Which bonds are broken CO HH Which bonds are formed C H C O OH Aern 2 AH bonds broken 2 AH bonds formed Aern 745 436 413 358 463 53 kJ Bond Dissociation Enthalpy Use tables of information to help you TABLE 89 Some Average Bond Dissociation Enthalpies kJmnl Single Bands H C 316 305 353 A85 7 7 272 339 285 N 163 201 283 7 7 7 192 7 U 156 7 452 335 7 213 201 F 155 565 490 284 253 249 Si 222 293 331 310 P 201 326 S 225 255 7 l 242 216 Bi 193 Multiple Bonds NN 418 CC 610 NEN 945 LEE 835 CN 615 0 715 EN 837 CED 1046 Bond Dissociation Enthalpy Rank the following in order of increasing bond length and bond strength SiF SiC SiO Si constant rF lt r0 lt rC Bond length sum of radii SiF lt SiO lt SiC Bond strength SiC lt SiO lt SiF NNN NNEN Bond order N N NN NEN Bond length NEN lt NN lt N N Bond strength N N lt NN lt NEN