Class Notes for ALL OF CHAPTER 6!!!
Class Notes for ALL OF CHAPTER 6!!! CH 101
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This 23 page Class Notes was uploaded by Allie Newman on Monday October 19, 2015. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Professor John McDuffie in Fall 2015. Since its upload, it has received 91 views. For similar materials see General Chemistry in Chemistry at University of Alabama - Tuscaloosa.
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CH 101 Class Notes for Chapter 6 Morphine l A Molecular lmposter Morphine is a natural product from the sap of the opium poppy that has pain relieving attributes Morphine binds to nerve opioid receptors by altering the transmitted nerve signals resulting in less pain and inducing feelings of euphoria and tranquility Morphine works in a body similar to endorphins binding to the opioid receptor s active site Like a key in a lock or lockinkey Morphine is a molecular imposter mimicking the action of endorphins because of similarities in molecular shape To understand drug interactions researchers use bonding theories to simulate the shape of potential drug molecules and how they would interact Bonding Theories Explain how and why atoms attach together to form molecules Explain why some combinations of atoms are stable and others are not Why is water H20 not H0 or H3O Can be used to predict the shapes of molecules Can be used to predict the chemical and physical properties of compounds Lewis Bonding Theory One of the simplest bonding theories is called Lewis theory Lewis theory Emphasizes valence electrons to explain bonding Predicts many properties of molecules such as molecular stability shape size and polarity Why Do Atoms Bond Chemical bonds form because they lower the potential energy between the charged particles that compose atoms A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms To calculate this potential energy you need to consider the following interactions Nucleustonucleus repulsions Electrontoelectron repulsions Nucleustoelectron attractions Types of Bonds Bonds between atoms are classi ed based on the types of atoms that are bonded together Atom Types Bond Types Characteristic of Bond Metal and Ionic Electrons are transferred Nionmetal between atoms cation anion Monmetal and Covalent Electrons are shared Monmetal between atoms Metal and Metal Metallic Electrons are pooled between atoms com mu nal sharing Bond Polarity Most bonds have some degree of sharing and some degree of ion formation to them Bonds are classi ed as covalent if the amount of electron transfer is insuf cient for the material to display the classic properties of ionic compounds If the sharing is unequal enough to produce a dipole in the bond the bond is classi ed as polar covalent Polar Covalent Bonding Covalent bonding between unlike atoms results in unequal sharing of the electrons One atom pulls the electrons in the bond closer to its side One end of the bond has larger electron density than the other The result is a polar covalent bond Bond polarity The end with the larger electron density gets a partial negative charge The end that is electron de cient gets a partial positive charge Electronegativity EN Increases across a period left to right and decreases down a group top to bottom Fluorine is the most electronegative element Francium is the least electronegative element Noble gas atoms are not assigned EN values Opposite of atomic size trend The ability of an atom to attract bonding electrons to itself is called electronegativity The larger the difference in electronegativity the more polar the bond Negative end toward more electronegative atom Electronegativity Difference and Bond Type o If the difference in electronegativity between bonded atoms is 0 the bond is pure covalent 0 Equal sharing of the atoms in the bond 0 If the difference in electronegativity between bonded atoms is 01 to 04 the bond is nonpolar covalent If the difference in electronegativity between bonded atoms is 05 to 19 the bond is polar covalent o Unequal sharing of electrons between the atoms in the bond o If the difference in electronegativity between bonded atoms is larger than or equal to 20 the bond is ionic TABLE 61 The Effect of Electronegativity Difference on Bond 39I39ype Electronegati vity Difference AIEN Bond Type Small 0 04 Covalent Clz intermediate 04 20 Polar covalent HCI Large 20 Ionic NaCl The Continuum of Bond Types WWW Pom m covalent bond 10mg bond i i 5 3 Electrons shared Electrons shared Electrons equally unequally transferred A 4A V A I I I 1390 1394 1390 33 Electronegativity difference AEN Bond Dipole Moments Dipole moment u is a measure of bond polarity A dipole is a material with a and end It is directly proportional to the size of the partial charges q and directly proportional to the distance r between them 1 dipole moment qr Measured in Debyes D Generally the more electrons two atoms share andl the larger the atoms are the larger the dipole moment TERIE 62 Dipole Moments of Several Molecules in the Gas Phase 7 7 Dipole Moment Molecule AEN 7 D ClF 10 088 HlF 19 182 LiF 30 633 Percent Ionic Character The percent ionic characteris the percentage of a bond s measured dipole moment compared to what it would be if the electrons were completely transferred The percent ionic character indicates the degree to which the electron is transferred 100 KBr KCl Lil H 39 OKF 4 75 K1 C31 3 l g LiCll CSCI 3 L 39 NaCl 393 50 g p LlBr i a HF 8 l r 2 l a a 5 HI HCl 2 1C1 0 E1 39 HBr l 0 ll 2 3 Electronegativity difference Lewis Structures of Molecules and Bonding Lewis theory on structures Predicts the distribution of valence electrons in a molecule Useful for understanding the bonding in many compounds Can be used to predict shapes of molecules Can be used to predict properties of molecules and how they will interact together Lewis theory on bonding Implies that another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms Shared electrons would then count toward each atom s octet Covalent bonding is the sharing of electrons between atoms in a bond 0 Lewis theory predicts that atoms will be most stable when they have their octet of valence electrons o It does not require that atoms have the same number of lone pair electrons they had before bonding Covalent Bonding Model versus Reality Lewis theory of covalent bonding Implies that the attractions between atoms are directional The shared electrons are most stable between the bonding atoms Predicts covalently bonded compounds will be found as individual molecules Rather than an array like ionic compounds 0 Compounds of nonmetals are made of individual molecule units Octet Rule A Guideline for Molecule Formation When atoms bond they tend to gain lose or share electrons to result in a quotnoble gaslike con guration l nszng 5 Nonmetals 2 39 period elements must obey the octet rule eg eight valence electrons around each atom in the molecule Exceptions to the octet rule expanded octets Involve the nonmetal elements located in the 3rd period and below Nonmetals 3rOI period down in the periodic table follow the octet rule when they are not the center atom The center atom is the atom in the molecule where the other elements individually bond to attach When they are the center atom they can accommodate more than eight electrons Using empty valence dorbitals that are predicted by quantum theory Writing Lewis Structures for Molecular Compounds Write the correct skeletal structure for the molecule The less electronegative atom in the molecule is usually the center atom Simple molecules have a center atom to which all of the other atoms in the molecules are attached bonded to For example for H20 the center atom is 0 so both H atoms are attached to the oxygen atom H O H The more electronegative atoms are usually terminal attached to the center atom Hydrogen atoms are always in the terminal position Determine the total number of valence electrons each atom is bringing in to form the molecule Examples For the molecule H Br H atom brings in 1 electron and Br atoms brings in 7 electrons for a total 8 electrons ln polyatomic ions the charge on the ion also must be accounted For the polyatomic ion NOZ a total of 18 electrons are brought in 5 electrons from N total 12 from O 2 oxygen atoms x 2 and 2 from the 1 charge Distribute the electrons among the atoms in the molecule giving octets or duets in the case of hydrogen to as many atoms as possible The best practice is to place two electrons around an atom at a time Bonding pairs electrons between two atoms Nonbonding or lone pairs electrons not participating in bonding but complete the atom s octet The total number of electrons brought in must be accounted in the Lewis structure and must not violate any criteria ie H atoms can only have single bonds or two electrons total If any atoms lack an octet form double or triple bonds as necessary to give them octets Atoms that can multiple bond with each other or to themselves are as follows Double bond 4 electrons or two pairs of electrons between atoms C O N S amp P Triple bond 6 electrons or three pairs of electrons between atoms CO N amp S Resonance and Formal Charges Two additional concepts to write the best possible Lewis structures for a large number of compounds The concepts are Resonance used when two or more valid Lewis structures can be drawn for the same compound Formal charge an electron bookkeeping system that allows us to discriminate between alternative Lewis structures Resonance Lewis theory localizes the electrons between the atoms that are bonding together Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons the concept is called resonance Delocalization of charge helps to stabilize the molecule When there is more than one Lewis structure for a molecule that differ only in the position of the electrons they are called resonance structures The actual molecule is a combination of the resonance forms a resonance hybrid The molecule does not resonate between the two forms though we often draw it that way Examplelj O3 molecule sz lt gt Q bzb Resonance hybrid structure a b o Resonance hybrid Just as the offspring of two different dog breeds is a hybrid that is intermediate between the two breeds a the structure of a resonance hybrid is intermediate between that of the contributing resonance structures b Rules of Resonance Structures Resonance structures must have the same connectivity Only electron positions can change Resonance structures must have the same number of electrons Second row elements have a maximum of eight electrons Bonding and nonbonding Third row can have expanded octet Formal charges must total the same Better structures have fewer formal charges Better structures have smaller formal charges Better structures have the negative formal charge on the more electronegative atom Formal Charge 0 The concept of formal charge is useful because it helps distinguish between competing skeletal structures or competing resonance structures 0 In general these four rules apply 1 The sum of all formal charges in a neutral molecule must be zero 2 The sum of all formal charges in an ion must equal the charge of the ion 3 Small or zero formal charges on individual atoms are better than large ones 4 When formal charge cannot be avoided negative formal charge should reside on the most electronegative atom 0 Formal chargg of valence electrons of nonbonding electrons 12 x of bonding electrons HIFI Example The molecule HF has 0 zero formal charge o The formal charge on H atom Formal charge 1 O 12 2 O o The formal charge on F atom Formal charge 7 6 12 2 0 Example Formal Charge 502 During bonding atoms may end with more or fewer electrons than the valence electrons they brought in order to ful ll octets This results in atoms having a formal charge Formal Charge FC valence e nonbonding e 12 bonding e I OS O The Lewns structure for 502 u left 0 FC6 4 124O S FC6 2 1261 RightO FC6 6 122 1 Sum of all the formal charges in a molecule O o In an ion total equals the charge Drawing Resonance Structures and Formal Charge N0339 1 Draw the rst Lewis structure that maximizes octets 2 Assign formal charges 3 Move electron pairs from atoms with formal charge toward atoms with formal charge 4 If formal charge atoms are in the second row only move in electrons if you can move out electron pairs from multiple bonds 5 If formal charge atoms are in the third row or below keep bringing in electron pairs to reduce the formal charge even if you get an expanded octet i ll 39l I ll ON O O NO Expanded Octets OddElectron and Other Species The Exceptions to the Octet Rule 0 The exceptions o Oddelectron species free radicals or radicals Molecules or ions with an odd number of electrons Incomplete octets molecules or ions with fewer than eight electrons around an atom Legitimate Lewis structures cannot be written for they do not meet the octet rulequot as required by the Lewis model Example N0 0 Has 11 valence electrons Distribution of 11 electrons cannot meet the criteria under the Lewis model 0 NO does exist as a molecule 0 The Lewis model is not sophisticated enough to work for an odd number of electron compounds 0 Incomplete octets Elements speci cally metalloids and H atom whose tendency is not to have a complete octet H can only accompany two electrons duet Boron metalloid Prefer 6 electrons than 8 electrons O O H IFIBIFI o Expanded octets Molecules or ions with more than eight electrons around an atom Involve the nonmetal elements located in the 3rd period and below Nonmetals 3rOI period down in the periodic table follow the octet rule when they are not the center atom The center atom is the atom in the molecule where the other elements individually bond to attach When they are the center atom they can accommodate more than eight electrons Using empty valence dorbitals that are predicted by quantum theory Bond Energies Chemical reactions involve breaking bonds in reactant molecules and making new bonds to create the products The AH reaction can be estimated by comparing the cost of breaking old bonds to the income from making new bonds The amount of energy in the gaseous state that it takes to break one mole of a bond in a compound is called the bond energy Trends in Bond Energies In general the more electrons two atoms share the stronger the covalent bond Must be comparing bonds between like atoms CEC 837 kJ gt CC 611 kJ gt C C 347 kJ CEN 891 kJ gt C N 615 kJ gt C N 305 kJ o In general the shorter the covalent bond the stronger the bond Must be comparing similar types of bonds Br F 237 kJ gt Br Cl 218 kJ gt Br Br 193 kJ Bonds get weaker down the column Bonds get stronger across the period Average Bond Energies TABLE 63 Average Bond Energies 7 Bond Energy Bond Energyquot Bond Energy Bond klmol klmo klmo H H 436 C c 347 NEN 946 H C 414 CC 611 0 0 142 HiN 389 CEC 837 00 498 H O 464 C 0 360 F F 159 HiF 565 CO 736 CIC 243 H Cl 431 C Cl 339 Br Br 193 H Br 364 N N 183 l l 151 H l 297 NN 418 quot47991 in C02 Using Bond Energies to Estimate AH rxn The actual bond energy depends on the surrounding atoms and other factors We often use average bond energies to estimate the Aern Works best when all reactants and products in gas state Bond breaking is endothermic AHbreaking is positive Bond making is exothermic AHmaking is negative Aern Z AHbonds broken Z AHbonds formed Covalent Bonding Model versus Reality for Bond Strength 0 Lewis theory predicts that the more electrons two atoms share the stronger the bond Single bond lt Double bond lt Triple bond 0 Lewis theory would predict that double bonds are twice as strong as single bonds but the reality is they are less than twice as strong 0 Bond strength is measured by how much energy must be added into the bond to break it in half Covalent Bonding Model versus Reality for Bond Length Lewis theory predicts that the more electrons two atoms share the shorter the bond should be When comparing bonds to like atoms Bond length is determined by measuring the distance between the nuclei of bonded atoms In general tripe bonds are shorter than double bonds and double bonds are shorter than single bonds E Bond Lengths c12 12 Bond Lengths The distance between the nuclei of bonded atoms is called the bond length Because the actual bond length depends on the other atoms around the bond we often use the average bond length Averaged for similar bonds from many compounds T EBLE 64 verage Bond Lengths Bond Length Bond Length Bond Length Bond 13m 7 pm pm H H 74 0 D 154 NEN 1 10 H C 110 CC 134 0 0 145 H N 100 CEC 120 00 121 H 0 97 3 0 143 F F 143 H F 92 CC 120 Cl Cl 189 H Cl 127 C Cl 178 Br Br 228 H Br 141 N N 145 l l 266 H l 161 NN 123 Trends in Bond Lengths In general the more electrons two atoms share the shorter the covalent bond Must be comparing bonds between like atoms cac 120 pm lt CC 134 pm lt C C 154 pm CEN 116 pm lt CN 128 pm lt C N 147 pm 0 Generally bond length decreases from left to right across a period C C 154 pm gt C N 147 pm gt C 0 143 pm Generally bond length increases down the column F F 144 pm gt Cl Cl 198 pm gt Br Br 228 pm 0 In general as bonds get longer they also get weaker VSEPR Theory and Molecular Geometries Structure Determine Properties Structure of Compounds Determines Their Properties Properties of molecular substances depend on the structure of the molecule The structure includes many factors The skeletal arrangement of the atoms The kind of bonding between the atoms Ionic polar covalent or covalent The shape of the molecule Bonding theory predicts the shapes of molecules What is Molecular Geometry Molecules are threedimensional objects Molecular geometries Describe the shape of a molecule with terms that relate to geometric gures Have characteristic quotcornersquot that indicate the positions of the surrounding atoms around a central atom in the center of the geometric gure Have characteristic angles that are called bond angles Lewis Theory Predicts Electron Groups Lewis theory predicts there are regions of electrons in an atom These regions of electron groups should repel each other because the regions are negatively charged Some regions result from placing shared pairs of valence electrons between bonding nuclei Other regions result from placing unshared valence electrons on a single nuclei This idea can then be extended to predict the shapes of the molecules The position of atoms surrounding a central atom will be determined by where the bonding electron groups are The positions of the electron groups will be determined by trying to minimize repulsions between them VSEPR Valence Shell Electron Pair Repulsion Theory Electron groups around the central atom will be most stable when they are as far apart as possible This is the basis for VSEPR valence shell electron pair repulsion theory Because electrons are negatively charged they should be most stable when they are separated as much as possible The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule Electron 033 Electron Groups The Lewis structure predicts the number of valence electron pairs around the central atoms Each lone pair of electrons constitutes one electron group on a central atom Each bond constitutes one electron group on a central atom regardless of whether it is single double or triple Example N02 u on There are three electron groups on N o a Three lone pairs 39 O N O 39 One single bond One double bond Electron Group Geometry There are ve basic arrangements of electron groups around a central atom Arrangements are as follows Linear trigonal planar tetrahedral bipyramidal and octahedral Based on a maximum of six bonding electron groups Though there may be more than six on very large atoms it is very rare 0 Each of these ve basic arrangements results in ve different basic electron geometries In order for the molecular shape and bond angles to be a perfect geometric gure all the electron groups must be bonds and all the bonds must be equivalent 0 For molecules that exhibit resonance it doesn t matter which resonance form you use as the electron geometry will be the same Two Electron Groups Linear Electron Geometry When there are two electron groups around the central atom they will occupy positions on opposite sides of the central atom This results in the electron groups taking a linear geometry The bond angle is 180 Linear geometry Linear geom eny 39 180 rm Three Electron Groups Trigonal Planar Electron Geometry When there are three electron groups around the central atom they will occupy positions in the shape of a triangle around the central atom This results in the electron groups taking a trigonal planar geometry The bond angle is 120 J EszrE Four Electron Groups Tetrahedral Electron Geometry When there are four electron groups around the central atom they will occupy positions in the shape of a tetrahedron around the central atom This results in the electron groups taking a tetrahedral geometry The bond angle is 1095 a Linear geometry b Trigonal planar geometry 1095 Tetrahedral geometry Five Electron Groups Trigonal Bipyramidal Electron Geometry When there are ve electron groups around the central atom they will occupy positions in the shape of two tetrahedral that are base to base with the central atom in the center of the shared bases This results in the electron groups taking a trigonal bipyramidal geometry The positions above and below the central atom are called the axial positions The bond angle between axial and equatorial positions is 90 The positions in the same base plane as the central atom are called the equatorial positions The bond angle between equatorial positions is 120 L Trigoml bipymmidal geometry 1 Octahedral Electron Geometry When there are six electron groups around the central atom they will occupy positions in the shape of two squarebase pyramids that are basetobase with the central atom in the center of the shared bases This results in the electron groups taking an octahedral geometry It is called octahedral because the geometric gure has eight sides All positions are equivalent The bond angle is 90 Octahedral geometry j Electron Pair Geometry versus Molecular Geometry The actual geometry of the molecule may be different from the electron geometry When the electron groups are attached to atoms of different size or when the bonding to one atom is different than the bonding to another this will affect the molecular geometry around the central atom Lone pairs also affect the molecular geometry Lone pair groups Occupy space on the central atom but are not seen as points on the molecular geometry Take up more spacequot on the central atom because their electron density is exclusively on the central atom rather than shared like bonding electron groups This affects the bond angles making the bonding pair angles smaller than expected Effect of Lone Pairs on Molecular Structure 0 The bonding electrons are shared by two atoms so some of the negative charge is removed from the central atom Relative sizes of repulsive force interactions are as follows 0 Lone pair to Lone pair gt Lone pair to Bonding pair gt Bonding pair to Bonding pair Bonding electron Pair L j t 7 Nuclei Nucleus Bent Molecular Geometry Derivative of Trigonal Planar Electron Geometry When there are three electron groups around the central atom and one of them is a lone pair the resulting shape of the molecule is called a trigonal planarbent shape The bond angle is less than 120 because the lone pair takes up more space lone pair Pyramidal and Bent Molecular Geometries l Derivatives of Tetrahedral Electron Geometry When there are four electron groups around the central atom and one is a lone pair the result is called a pyramidal shape because it is a triangularbase pyramid with the central atom at the apex The bond angle is less 1095 Lone pair Electron geometry Molecular geometry tetrahedral trigonal pyramidal Pyramidal and Bent Molecular Geometries l Derivatives of Tetrahedral Electron Geometry When there are four electron groups around the central atom and two are lone pairs the result is called a tetrahedralbent shape It is planar It looks similar to the trigonal planarbent shape except the angles are smaller The bond angle is less than 1095 l Lone pair Electron geometry Molecular geometry tetrahedral bent Bond Angle Distortion from Lone Pairs Ellfect of Lone Pairs on Molecular Geometry Derivatives of the Trigonal Bipyramidal Electron Geometry When there are ve electron groups around the central atom and some are lone pairs they will occupy the equatorial positions because there is more room The bond angles between equatorial positions are less than 120 The bond angles between axial and equatorial positions are less than 90 Linear 180 axial to axial 0 When there are ve electron groups around the central atom and one is a lone pair the result is called the seesaw shape or sawhorse aka distorted tetrahedron Three 90 lone pairbonding Two 90 lone pair bonding pair replulsiolns pair repulsions r Molecular geometry Axial lone pair Equatorial seesaw Does not occur lone pair Derivatives of the Trigonal Bipyramidal Electron Geometry When there are ve electron groups around the central atom and some are lone pairs they will occupy the equatorial positions because there is more room When there are ve electron groups around the central atom and two are lone pairs the result is Tshaped 2F I l Br F IE 39 Electron geometry Molecular geometry trigonal bipyramidal 39ll shaped Derivatives of the Trigonal Bipyramidal Electron Geometry When there are ve electron groups around the central atom and some are lone pairs they will occupy the equatorial positions because there is more room When there are ve electron groups around the central atom and three are lone pairs the result is Linear r 2 Electron geometry Molecular geometry trigonal bip yramidal linear Lu Derivatives of the Octahedral Geometry When there are six electron groups around the central atom and some are lone pairs each even number lone pair will take a position opposite the previous lone pair When there are six electron groups around the central atom and one is a lone pair the result is called a square pyramid shape The bond angles between axial and equatorial positions are less than 90 51393quot ili Br lj O Electron geometry l L Molecular geometry J F F I I O octahedral square pyramidal o a or o Derivatives of the Octahedral Geometry When there are six electron groups around the central atom and some are lone pairs each even number lone pair will take a position opposite the previous lone pair When there are six electron groups around the central atom and two are lone pairs the result is called a square planar shape The bond angles between equatorial positions are 90 Electron geometry Molecular geometry octahedral square planar Using VSEPR to Predict Molecular Geometries The steps 1 Draw the Lewis structure 2 Determine the number of electron groups around the central atom 3 Classify each electron group as a bonding or lone pair and count each type Remember multiple bonds count as one group 4 Use Table 65 to determine the shape and bond angles Representing ThreeDimensional Shapes on Paper Drawing molecules to show their dimensionality on paper 2D is dif cult How to draw a 3D representation of molecule on paper By convention the central atom is put in the plane of the paper Put as many other atoms as possible in the same plane and indicate with a straight line For atoms in front of the plane use a solid wedge For atoms behind the plane use a hashed wedge ilillllllll 1 Straight line Hatched wedge Solid wedge Bond in plane of paper Bond going into the page Bond coming out of the page Illustrations of Molecular Geometries of Molecules Using 3D Notations 2 e A A Q r r X X a a X X Linear Trigenal planar Bent x 6 m r X La x A 39 39 a x A39WwX WX I X X W X e Tetrahedral Trigonal pyramidal Trlglonal blpyrarniclal X X m ax g Xe I ax X X A v 16 33 fr A lX Xi X X X 23 X X W Seesaw Octahedral Square planar Multiple Central Atoms and Their Geometries Many molecules have larger structures with many interior atoms Think of them as having multiple central atoms For multiple center molecules Each center atom has a designated a shape Example Glycine The shape around the o N atom is trigonal pyramidal Left C is tetrahedral Right C is trigonal planar O is bent Trigonal pyramidal O H ixi C C o H Wquot lFour interior atoms Bent Tetra hedral Glycine l Ball andStick Modiel of Glycine Polarity of Molecules For a molecule to be polar it must have the following Polar bonds Electronegativity difference Bond dipole moments measured An unsymmetrical shape Vector addition Polarity affects the intermolecular forces of attraction Example Boiling points and solubilities Like dissolves like Nonbonding pairs affect molecular polarity a strong pull in its direction Molecule Polarity The bond between the H atom and Cl atom in HCI molecule is polar The bonding electrons are pulled toward the Cl end of the molecule because Cl is more electronegative EN than the H atom The net result is a polar molecule Net dipole moment 6 5 m 77 I M Low electron High electron density density Polar bond 0 The bond between the C atom and O atoms in C02 molecule is polar o The bonding electrons are pulled toward the 0 ends of the molecule equally because 0 is more electronegative EN than the C atom o The net result is a nonpolar molecule No net dipole moment o The bond between the O atom and H atoms in H20 molecule is polar Both sets bonding electrons are pulled toward the 0 end of the molecule equally because 0 is more electronegative EN than the H atom Because of its geometry the lone pairs do not negate the diploe moment that result from the each of the 0 H bonds 0 The net result is a polar molecule Net dipole moment 5 or 39 fi 5f Summarizing Determining Molecular Shape and Polarity 0 Draw the Lewis structure for the molecule and determine its molecular geometry 0 Determine if the molecule contains polar bonds 0 A bond is polar if the two bonding atoms have suf ciently different electro negativities o If the molecule contains polar bonds superimpose a vector pointing toward the more electronegative atom on each bond 0 Make the length of the vector proportional to the electronegativity difference between the bonding atoms 0 Determine if the polar bonds add together to form a net dipole moment 0 Sum the vectors corresponding to the polar bonds together If the vectors sum to zero the molecule is nonpolar o If the vectors sum to a net vector the molecule is polar Vector Addition THEME Ei i Gunmen Eeeee ef Adding Dipele Elements te Z etemine whether a Heieml Is Linear Nl l lp li llr The dieelle me meets eil twee i elehtiieel pellet beetle eeihting ih epeeeite direeliiehe will ee meet The meleeule ie hehpelen Eeht Feller The elieeile me mettie ei twe eeler her e with an angle eli ieee that 1 EU between them will het C l i ei The reeuilteht eieelle me merit neuter ie ehewn in reel The meleeule ie eel eh Triiglenell pile her Nenpeler The tilipelle memente eif three identieell pelleir heth at quotl 2W frern eeehl either will eeneel The melleeule ie nehpelet Vector Addition One Dimension 0 One Dimension Tettehedr el Wheeler The e p elle me m ehte ef fee r the htieel eel er it E e in e tetteh ed re ll ewe nge merit ii irerh eeeh then will eeinaeel The eleeui e lie ne 1 tie l err Trigenei eeir emieel ller The dipele memente ef three El i lf eehee in e trigehel pererhieel arrangement heee eeeh etheri Willi net eeneel The reeeitent ellipelle merrieht eeeter ie ehewn in reel The meleeeie ie pellet Hie it ellll eeeee in which the eieelee ell twe err metre peler beetle ea heel the hentie ere eeeemed te he idemieell If ene Lr mere tell the lhshtle are different tram the etherlelr the dipellee willll net eeeeell and the mleeulle 39eiliilll lee elem 0 To add two vectors that lie on the same line assign one direction as positive 0 Vectors pointing in that direction have positive magnitudes 0 Consider vectors pointing in the opposite direction to have negative magnitudes Example l 5 5 r B el Example 2 10 shill tel Example 3 5 5 1 l E E 10 71 ll 4 eel 5 71 Il Del Del 0 i the vectors exactly cancel Vector Addition Two or More Dimensions 0 Two or More Dimensions 0 To add two vectors draw a parallelogram in which the two vectors form two adjacent sides 0 Draw the other two sides of the parallelogram parallel to and the same length as the two original vectors 0 Draw the resultant vector beginning at the origin and extending to the far corner of the parallelogram as shown in Examples 4 and 5 0 To add three or more vectors add two of them together rst and then add the third vector to the result Examples 6 and 7 Example 4 Example 5 D Zero r if C i i731 lmi c mime vectors exactly cancel up EU 1 Molecular Polarity Affects Solubility in Water Polar molecules are attracted to other polar molecules Water is a polar molecule therefore Other polar molecules will dissolve in water As well as ionic compounds 0 Some molecules have both polar and nonpolar parts Example Soap molecule attract one another Opposite magnetic poles Opposite partial charges on molecules attract one another