Chapter 1 Book Notes
Chapter 1 Book Notes CHM 2210
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Popular in Chemistry
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This 6 page Class Notes was uploaded by Amanda Weaver on Sunday August 30, 2015. The Class Notes belongs to CHM 2210 at University of Florida taught by Dr. Laura Beth Peterson in Summer 2015. Since its upload, it has received 50 views. For similar materials see Organic Chemistry 1 in Chemistry at University of Florida.
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Date Created: 08/30/15
CHAPTER 1 COVALENT BONDING AND SHAPES OF MOLECULES o The most common elements for compounds are C H O and N 0 Organic chemistry the study of compounds of carbon 11 Electron Structure of Atoms Electrons exist outside the nucleus in shells There can be 2n2 electrons per shell Energy levels are quantized which means they occur discretely Electrons are delocalized which means that they are spread out over a larger volume of space Subshells include s p d and f o p subshells exist in px py and pz Each can hold up to two electrons A Electron Configuration of Atoms 0 We determine the electron configuration of atoms in their groundstate electron configuration or the lowest energy level as they appear on the periodic table The rules include 1 Auf39bau Principle Orbitals ll in the order of increasing energy ls 2s 2p 3s 3p 4s 4p 3d 2 Pauli Exclusion Principle only two electrons may occupy an orbital and their spins must be paired opposite Thus you can only have Ti and not TT or ii 3 Hund s Rule one electron is added to each orbital before any are filled a Ex 2p4 Ti T T B The Concept of Energy Energy the ability to do work Potential energy energy that is stored and ready to be released and do work Ground state standard electron configuration of an atom Excited state electron configuration that differs from the ground state 9 ie a higher energy state than naturally occurs for a given element 0 Ionization energy the energy it takes to remove and electron from an atom or molecule 0 The further an electron is from the nucleus the easier it is to remove the lower its IE and the sooner it is ionized I Ex 2p orbital for C is the furthest out so these electrons have the first ionization energy C Lewis Dot Structure 0 Valence electrons electrons in the outermost shell highest energy level 0 Valence shell the shell in which the outermost electrons are located 0 EX Valence electrons are in 4p therefore the valence shell is 4 0 Lewis Dot Structure uses a series of dots to depict the valence electrons of a given element 12 Lewis Model of Bonding 0 Octet rule the tendency of Group 1A7A elements to strive to achieve PNG configurations 8 valence A Formation of Chemical Bonds 1 An atom may become ionic a Anion atom that gains electrons b Cation atoms that loses electrons c Ionic interaction the interaction between and anion and cation where two ions are held together by primarily the attraction of opposite forces also called and ionic bond NaCl 2 An atom may share electrons with one or more other atoms to complete its valence shell a Covalent bond a chemical bond formed by sharing electrons AlCl3 3 Bonds may be partially ionic and partially covalent which are called polar covalent bonds B Electronegativity and Chemical Bonds 0 Electronegativity a measure of an atom s attraction for electrons that it shares in a chemical bond with another atom H F 19 0 Periodic trend of EN up and to the right 0 Electron af nity the energy that is released upon addition of an electron 0 Periodic trend of EA up and to the right Formation of Ions 0 Ions are formed by the transfer of an electron from a low EN atom to a high EN atom 0 An ion is formed if the difference in EN is 19 or higher Covalent Bonds 0 Formed when two atoms share electrons o The energy that is released upon the covalent bond forming is the bond dissociation energy bond dissociation enthalpy BDE and it must be absorbed for the bond to break 0 H2 has a high BDE which is why it occurs diatomically 0 Bond length the distance between nuclei participating in a bond 0 Lone pair a pair of electrons that is not participating in a bond 0 There can be up to four bond with the second row atoms minus one for every lone pair Polar Covalent Bonds 0 Polarity o lt 05 Nonpolar covalent o 05 19 Polar covalent o gt19 Ionic bond 0 To find the polarity of the bond take the absolute value of the difference in EN of the atoms 0 Bond dipole moment a measure of the polarity of a covalent bond The product of the charge of either atom times the distance between the atoms C Lewis Structures for Molecules and Polyatomic Ions 1 Determine the number of valence electrons in the molecule or ion add the number of valence electrons contributed by each atom 2 Determine the connectivity arrangement of atoms in the molecule or ion a Isomers difference compounds with the same formula Connect the atoms with single bonds then fill the rest according to the octet rule 4 Each bonding pair of electrons is shown by a single line between atoms Lone pairs are shown by two dots Two electrons single bond four electrons double bond six electrons triple bond 6 Table 18 pg 12 has common compounds practice drawing these 9 U D Formal Charge 0 Formal Charge the charge on an atom in a molecule or polyatomic ion 1 Write a correct Lewis structure for the molecule or ion 2 Assign each atom all its unshared nonbonding electrons and half of its shared electrons 3 Find the difference of the atom s valence electron number and the number of electrons it has in the molecule or polyatomic ion 4 The sum of all formal charges is the charge on the molecule or polyatomic ion 5 Make sure that there are no extra electrons ie Period 2 atoms cannot hold more than 8 electrons in their outer shell E Exceptions to the Octet Rule 0 Elements in the 3A Group have incomplete valence shells Therefore they can only be surrounded by up to six 6 electrons Ex BF3 AlCl3 0 Other elements such as Sulfur can have an expanded octet because of its six valence electrons 13 Functional Groups 0 Functional groups 0 Carboncontaining compound 0 Units by which organic compounds are divided 0 Sites of characteristic chemical reactions a functional group always reacts the same 0 Serve as a basis for naming organic compounds A Alcohols 0 Alcohol an OH hydroxyl bonded to a tetrahedral C atom o Hydroxyl group an OH group 0 Condensed structural formula CH3 CH2 OH a way of shortening Lewis structures We will primarily use line angle structures though covered in class notes 0 Alcohols are classified as primary secondary or tertiary depending on how many C atoms are bonded to the Carbon atom with the OH group 0 Primary 1 a C OH where the C atom is only bonded to one other C atom and two H 0 Secondary 2 a C OH where the C is bonded to two other C atoms and one H o Tertiary 3 a C OH where the C is bonded the three other C atoms and no H B Amines 0 Amino group a Nitrogen atom bonded to one two or three C atoms by single bonds 0 Primary 1 amine an N C bond where the N atom is bonded to one C atom Secondary 2 amine an N C bond where the N atom is bonded to two C atoms Tertiary 3 amine an N C bond where the N atom is bonded to three C atoms C Aldehydes and Ketones 0 Carbonyl group the functional group of both aldehydes and ketones C o Aldehyde a compound containing CHO group CH20 o Ketone a compound containing a carbonyl group bonded to two carbons CH3 2C 0 D Carboxylic Acids 0 Carboxyl COOH the functional group of a carboxylic acid 0 Carboxyl carbonyl hydroxyl E Carboxylic Esters a derivative of a carboxylic acid in which H of the carboxyl group is replaced by C F Carboxylic Amides a derivative of a carboxylic acid in which the OH is replaced by an amine 14 Bond Angles and Shapes of Molecules 0 VSEPR valenceshell electronpair repulsion a method of predicting bond angles based on knowledge that electrons will repel one another Tetrahedral bond angles of 1095 Pyramidal bond angles of 1073 Triangular bond angles of 120 Linear bond angles of 180 Octahedral bond angles of 90 15 Polar and Nonpolar Molecules 0 Molecular dipole moment mu the sum of each individual bond dipoles 0 Compare each of the individual polar bonds and if it is zero then the molecule is nonpolar 0 Asking Tanvir about how to do mu 16 Quantum Wave Mechanics A Moving Particles Exhibit Properties of a Wave 0 Energy E of a photon is proportional to the frequency v 0 De Broglie equation 7 hmv B Shapes of Atomic s and p Orbitals 0 s sphere p peanut d dumbbell f ower 17 A Combined Valence Bond and Molecular Orbital theory Approach to Covalent Bonding A Molecular Orbital Theory Formation of Molecular Orbitals 0 Molecular Orbital MO Theory e39 in molecules occupy molecular orbitals that extend over the entire molecule and are formed by the combination of the atomic orbitals that make up the molecule 1 The number of orbitals that is formed is equal to the number that is combinedb 2 Molecular orbitals are arranged in order of increasing energy 3 The molecular orbitals are filled following the same guidelines as atomic orbitals o Bonding molecular orbital e39 have a lower energy than they would in the isolated atomic orbital 0 Sigma bonding molecular orbital e39 density lies between the two nuclei along the axis joining them and is cylindrically symmetric about the axis 0 Antibonding molecular orbital e39 have a higher energy than they would have in the isolated atomic orbitals 0 Ground state an atom in its lowest energy o Excited state any electron state other than the ground state 0 All of the orbitals of all of the atoms take part in the molecular orbitals and thus the orbitals extend over the entire molecule B Valence Bond Theory Hybridization of Atomic Orbitals 0 Valence Bond VB Theory bonds are created by the overlap of atomic orbitals on adjacent atoms 0 Hybrid Orbital an orbital formed by the combination of two or more orbitals o Hybridization the combination of atomic orbitals of different types 0 sp3 Hybrid Orbitals Bond Angles of Approximately 1095 0 sp3 hybrid orbitals formed by one s orbital and three p orbitals 0 sp3 hybridized an atom with the hybridization state sp3 0 sp2 Hybrid Orbitals Bond Angles of Approximately 120 0 sp2 hybrid orbital formed by the combination of one 5 atomic orbital and two p orbitals 0 sp2 hybridized an atom with the hybridization state sp2 0 sp Hybrid Orbitals Bond Angles of Approximately 180 0 sp hybrid orbital formed by the combination of one s and one p orbital 0 sp hybridized an atom with the hybridization state sp C An Analysis of S and P 0 S and P are exceptions to the octet rule so they are typically drawn with more than eight surrounding electrons However there are ongoing tests to investigate whether the sp3 hybridizations are held and the e39 density changes to better conform to the sp3 hybridization D Combining Valence Bond VB Theory and Molecular Orbital MO Theory The Creation of Sigma and Pi Bonding and Antibonding Orbitals 0 Combines both theories and allows the visualization of overlapping bonding or nonbonding sigma and pi bonds 0 Antibonding bonds that have more energy than the separate orbitals o Bonding bonds that have less energy than the separate orbitals 0 If a molecule has double bonds sp2 should be considered If it has triple bonds sp should be considered 0 Pi bonding molecular orbital the combination of two parallel pi bonds 18 Resonance A Theory of Resonance o Resonance a theory that many molecules are best described by multiple structures 0 ContributingResonance structuresResonance contributors individual Lewis structures 0 Double headedarrows connect each contributing structure 0 Curved arrows indicate the shift of electrons on Lewis structures 0 Moves from high density electron areas to low density areas 0 Electron pushing the movement of electrons using curved arrows B Rules for Writing Acceptable Contributing Structures PPM All structures must have the same number of valence electrons All structures must obey the rules of covalent bonding see Lewis structures All nuclei must be in the same position Contributing structures must have the same amounts of paired and unpaired electrons C Estimating the Relative Importance of Contributing Structures There is a preference system to determine the resonance structure that is dominant 0 Preference 1 Filled valence shells 0 Preference 2 Maximum number of covalent bonds 0 Preference 3 Least separation of unlike charges 0 Preference 4 Negative charge on a more electronegative atom 19 Molecular Orbitals for Delocalized System A Resonance Revisited Pi bonds are shown in hybrid resonance structures as dashed lines Charges are split evenly by atoms that hold formal charges Delocalization spreading of electron density over a larger space has a stabilizing effect Conjugation lack of an intervening atom between pi bonds and lonepair electrons B A Greater Reliance on Molecular Orbital Theory Since MO theory tells us to add and subtract all of the atomic orbitals it is more useful when discussing delocalized molecules C Hybridization Considerations in Light of Resonance and MO Theory When a resonance structure has a pi bond it will most likely be sp2 never 5193 110 Bond Lengths and Bond Strengths in Alkanes Alkynes P9P Order of bond length single bond gt double bond gt triple bond For C C bonds the shorter that the C C bond gets the shorter that the C H bond gets The shorter the bond the stronger the bond A triple bond is not three times as strong as a single bond nor is a double bond twice as strong