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8Chapter 01s - Chapter 1 Structure Determines Properties...

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Unformatted text preview: Chapter 1 Structure Determines Properties Dr. Wolf's CHM 201 & 202 1- 1 1.1 Atoms, Electrons, and Orbitals Dr. Wolf's CHM 201 & 202 1- 2 Atoms are composed of Atoms are composed of Protons + positively charged mass = 1.6726 X 10-27 kg Neutrons neutral mass = 1.6750 X 10-27 kg • – Dr. Wolf's CHM 201 & 202 Electrons negatively charged mass = 9.1096 X 10-31 kg 1- 3 Atomic Number and Mass Number Atomic Number and Mass Number A X Z Atomic number (Z) = number of protons in number nucleus nucleus (this must also equal the number of electrons (this in neutral atom) in Mass number (A) = sum of number of protons sum + neutrons in nucleus neutrons Dr. Wolf's CHM 201 & 202 1- 4 Schrödinger Equation Schrödinger Equation Schrödinger combined the idea that an Schrödinger electron has wave properties with classical equations of wave motion to give a wave equation for the energy of an electron in an atom. Wave equation (Schrödinger equation) gives a Wave series of solutions called wave functions ( ψ ). Dr. Wolf's CHM 201 & 202 1- 5 Wave Functions Wave Functions Wave Only certain values of ψ are allowed. Only Each ψ corresponds to a certain energy. Each The probability of finding an electron at a The particular point with respect to the nucleus is particular given by ψ 2. given Each energy state corresponds to an orbital. Dr. Wolf's CHM 201 & 202 1- 6 Figure 1.1 Probability distribution (ψ 2) for an for electron in a 1s orbital. electron Dr. Wolf's CHM 201 & 202 1- 7 A boundary surface encloses the region where the probability of finding an electron is high—on the order of 90-95% 1s 2s Figure 1.2 Boundary surfaces of a 1s orbital orbital and a 2s orbital. orbital. Dr. Wolf's CHM 201 & 202 1- 8 Quantum Numbers Quantum Numbers Each orbital is characterized by a unique Each set of four quantum numbers. set The principal quantum number (n) is a whole The principal is number (integer, 1, 2, etc.) that specifies the shell number shell and is related to the energy of the orbital. and The angular momentum quantum number (l) is The angular (l) usually designated by a letter (s, p, d, f, etc) usually etc) and describes the shape of the orbital. Values for l range from 0 to (n-1) Dr. Wolf's CHM 201 & 202 1- 9 Quantum Numbers Quantum Numbers Quantum Each orbital is characterized by a unique Each set of four quantum numbers. set The magnetic quantum number m is a whole The is number ranging from –l through 0 to +l. This gives the orientation of the orbital. The spin quantum number is either + or - ½ The spin is for the electron spin clockwise or counter for clockwise Dr. Wolf's CHM 201 & 202 1- 10 s Orbitals s Orbitals s Orbitals are spherically symmetric. Orbitals The energy of an s orbital increases with the The number of nodal surfaces it has. number it A nodal surface is a region where the nodal probability probability of finding an electron is zero. A 1s orbital has no nodes; a 2s orbital has one; one; a 3s orbital has two, etc. Dr. Wolf's CHM 201 & 202 1- 11 The Pauli Exclusion Principle The Pauli Exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers. Two electrons can occupy the same orbital only when they have opposite spins. There is a maximum of two electrons per orbital. Dr. Wolf's CHM 201 & 202 1- 12 First Period First Period Principal quantum number (n) = 1 Hydrogen Helium Z=1 1s 1 1s Z=2 1s 2 2s 2p H He Dr. Wolf's CHM 201 & 202 1- 13 p Orbitals p Orbitals p Orbitals are shaped like dumbells. Orbitals Are not possible for n = 1. Are Are possible for n = 2 and higher. Are Dr. Wolf's CHM 201 & 202 1- 14 p Orbitals p Orbitals p Orbitals are shaped like dumbells. Orbitals Are not possible for n = 1. Are Are possible for n = 2 and higher. Are There are three p orbitals for each value There of n (when n is greater than 1). Dr. Wolf's CHM 201 & 202 1- 15 p Orbitals p Orbitals p Orbitals are shaped like dumbells. Orbitals Are not possible for n = 1. Are Are possible for n = 2 and higher. Are There are three p orbitals for each value There of n (when n is greater than 1). Dr. Wolf's CHM 201 & 202 1- 16 p Orbitals p Orbitals p Orbitals are shaped like dumbells. Orbitals Are not possible for n = 1. Are Are possible for n = 2 and higher. Are There are three p orbitals for each value There of n (when n is greater than 1). Dr. Wolf's CHM 201 & 202 1- 17 Second Period Second Period Principal quantum number (n) = 2 Z 1s 2s 2p Li 3 Be 4 B5 C6 Dr. Wolf's CHM 201 & 202 1- 18 Second Period Second Period Z N 2p 8 F 2s 7 O 1s 9 Ne 10 Dr. Wolf's CHM 201 & 202 1- 19 Electron Configurations Electron Configurations Electrons fill the lowest energy levels first Electrons (calcium shown) (calcium E N E R G Y Dr. Wolf's CHM 201 & 202 5s 4s 3s 2s 1s 4d 4p 3d 3p 2p 1.2 Ionic Bonds Positively charged ions called CATIONS X OGIONS Negatively charged ions called DXX IONS ANIONS Dr. Wolf's CHM 201 & 202 1- 21 Ionic Bonding Ionic Bonding An ionic bond is the force of electrostatic attraction between oppositely charged ions Na+ (cation) (cation) Cl– (anion) Dr. Wolf's CHM 201 & 202 1- 22 Ionic Bonding Ionic Bonding Ionic bonds are common in inorganic Ionic chemistry chemistry but rare in organic chemistry. Carbon shows less of a tendency to form Carbon cations cations than metals do, and less of a tendency to than form form anions than nonmetals. Dr. Wolf's CHM 201 & 202 1- 23 1.3 Covalent Bonds and the Octet Rule Dr. Wolf's CHM 201 & 202 1- 24 The Lewis Model of Chemical Bonding The Lewis Model of Chemical Bonding In 1916 G. N. Lewis proposed that atoms In atoms combine in order to achieve a more stable combine more electron configuration. Maximum stability results when an atom is isoelectronic with a noble gas. An electron pair that is shared between An two atoms constitutes a covalent bond. two Dr. Wolf's CHM 201 & 202 1- 25 Covalent Bonding in H22 Covalent Bonding in H Two hydrogen atoms, each with 1 electron, H. .H can share those electrons in a covalent bond. H: H Sharing the electron pair gives each hydrogen Sharing an electron configuration analogous to helium. an Dr. Wolf's CHM 201 & 202 1- 26 Covalent Bonding in F22 Covalent Bonding in F Two fluorine atoms, each with 7 valence electrons, .. .. . F: : .. . F .. can share those electrons in a covalent bond. .. .. : .. : .. : FF Sharing the electron pair gives each fluorine an Sharing electron configuration analogous to neon. electron Dr. Wolf's CHM 201 & 202 1- 27 The Octet Rule The Octet Rule In forming compounds, atoms gain, lose, or In share electrons to give a stable electron configuration characterized by 8 valence electrons. electrons. .. .. : .. : .. : FF The octet rule is the most useful in cases The involving covalent bonds to C, N, O, and F. involving Dr. Wolf's CHM 201 & 202 1- 28 Example Example Combine carbon (4 valence electrons) and four fluorines (7 valence electrons each) . . C. . .. : .. . F to write a Lewis structure for CF4. .. F .. : .. : .. : .. : C: .. : F .. F : .. : F The octet rule is satisfied for carbon and The each fluorine. each Dr. Wolf's CHM 201 & 202 1- 29 Example Example It is common practice to represent a covalent bond by a line. We can rewrite .. F .. : .. : .. : .. : C: .. : F .. F : .. : F as .. : .. F .. : F: C .. F .. : : .. : F Dr. Wolf's CHM 201 & 202 1- 30 1.4 Double Bonds and Triple Bonds Dr. Wolf's CHM 201 & 202 1- 31 Inorganic Examples Inorganic Examples .. .. .. : O: : C : : O: :O C Carbon dioxide H : C : :: N: H C .. O: N: Hydrogen cyanide Dr. Wolf's CHM 201 & 202 1- 32 Organic Examples Organic Examples HH .. .. H: C: : C:H H Ethylene H C H H H : C : :: C:H Dr. Wolf's CHM 201 & 202 Acetylene H C C CH 1- 33 1.5 Polar Covalent Bonds Polar and Electronegativity and Dr. Wolf's CHM 201 & 202 1- 34 Electronegativity Electronegativity Electronegativity is a measure of the ability of an element to attract electrons toward of itself when bonded to another element. itself An electronegative element attracts electrons. An electronegative An electropositive element releases electrons. An electropositive Dr. Wolf's CHM 201 & 202 1- 35 Pauling Electronegativity Scale Pauling Electronegativity Scale Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Electronegativity increases from left to right in the periodic table. Electronegativity decreases going down a group. Dr. Wolf's CHM 201 & 202 1- 36 Generalization Generalization The greater the difference in The electronegativity electronegativity between two bonded atoms; the more polar the between .. .. : N N: : .. .. : FF bond. H—H bond. nonpolar bonds connect atoms of the same electronegativity Dr. Wolf's CHM 201 & 202 1- 37 Generalization Generalization The greater the difference in The electronegativity electronegativity between two bonded atoms; the more polar the between bond. .. δ− bond. δ+ H .. : F δ+ H .. δ− δ+ : δ− δ+ :δ− OCO .. .. OH .. polar bonds connect atoms of different electronegativity Dr. Wolf's CHM 201 & 202 1- 38 Electrostatic Potential Maps Electrostatic Potential Maps Electrostatic potential maps show the charge distribution within a molecule. δ+ H .. δ− F .. : Solid surface Red is negative charge; blue is positive. Dr. Wolf's CHM 201 & 202 1- 39 Electrostatic Potential Maps Electrostatic Potential Maps Electrostatic potential maps show the charge distribution within a molecule. δ+ H .. δ− F .. : Transparent surface Red is negative charge; blue is positive. Dr. Wolf's CHM 201 & 202 1- 40 Electrostatic Potential Maps Electrostatic Potential Maps Electrostatic potential maps show the charge distribution within a molecule. δ+ δ+ .. δ− δ− H .. : H Li F Red is negative charge; blue is positive. Dr. Wolf's CHM 201 & 202 1- 41 1.6 Formal Charge Formal charge is the charge calculated Formal for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs. electron Dr. Wolf's CHM 201 & 202 1- 42 Nitric acid Nitric acid Formal charge of H H .. O .. .. O: N : O: .. We will calculate the formal charge for each We atom in this Lewis structure. atom Dr. Wolf's CHM 201 & 202 1- 43 Nitric acid Nitric acid Formal charge of H H .. O .. .. O: N : O: .. •Hydrogen shares 2 electrons with oxygen. •Assign 1 electron to H and 1 to O. •A neutral hydrogen atom has 1 electron. •Therefore, the formal charge of H in nitric acid Therefore, is 0. is Dr. Wolf's CHM 201 & 202 1- 44 Nitric acid Nitric acid Formal charge of O H .. O .. .. O: N : O: .. •Oxygen has 4 electrons in covalent bonds. •Assign 2 of these 4 electrons to O. •Oxygen has 2 unshared pairs. Assign all 4 of Oxygen these electrons to O. these •Therefore, the total number of electrons Therefore, assigned to O is 2 + 4 = 6. assigned Dr. Wolf's CHM 201 & 202 1- 45 Nitric acid Nitric acid Formal charge of O H .. O .. .. O: N : O: .. •Electron count of O is 6. •A neutral oxygen has 6 electrons. •Therefore, the formal charge of O is 0. Dr. Wolf's CHM 201 & 202 1- 46 Nitric acid Nitric acid Formal charge of O H .. O .. .. O: N : O: .. •Electron count of O is 6 (4 electrons from Electron unshared pairs + half of 4 bonded electrons). unshared •A neutral oxygen has 6 electrons. •Therefore, the formal charge of O is 0. Dr. Wolf's CHM 201 & 202 1- 47 Nitric acid Nitric acid Formal charge of O H .. O .. .. O: N : O: .. •Electron count of O is 7 (6 electrons from Electron unshared pairs + half of 2 bonded electrons). unshared •A neutral oxygen has 6 electrons. •Therefore, the formal charge of O is -1. Dr. Wolf's CHM 201 & 202 1- 48 Nitric acid Nitric acid Formal charge of N H .. O .. .. O: N – : O: .. •Electron count of N is 4 (half of 8 electrons in Electron covalent bonds). covalent •A neutral nitrogen has 5 electrons. •Therefore, the formal charge of N is +1. Dr. Wolf's CHM 201 & 202 1- 49 Nitric acid Nitric acid Formal charges .. O: H .. O+ .. N – : O: .. •A Lewis structure is not complete unless formal Lewis charges (if any) are shown. charges Dr. Wolf's CHM 201 & 202 1- 50 Formal Charge Formal Charge An arithmetic formula for calculating formal charge. Formal charge = Formal group number group – in periodic table Dr. Wolf's CHM 201 & 202 number of – bonds number of unshared electrons 1- 51 Formal Charge Formal Charge "Electron counts" and formal charges in NH4+ and BF4charges and 1 H H 4 Dr. Wolf's CHM 201 & 202 N H + H .. : .. F .. : F: 7 –.. B .. : F : .. : F 4 1- 52 1.7 Structural Formulas of Organic Molecules Dr. Wolf's CHM 201 & 202 1- 53 Constitution Constitution The order in which the atoms of a The molecule are connected is called its constitution or connectivity. constitution connectivity The constitution of a molecule must be The determined in order to write a Lewis structure. structure. Dr. Wolf's CHM 201 & 202 1- 54 Condensed structural formulas Condensed structural formulas Lewis structures in which many (or all) Lewis covalent bonds and electron pairs are omitted. omitted. H H H H C C C H : O: H H Dr. Wolf's CHM 201 & 202 H can be condensed to: CH3CHCH3 or (CH3)2CHOH OH 1- 55 Bond-line formulas Bond-line formulas CH3CH2CH2CH3 is shown as is CH3CH2CH2CH2OH is shown as CH is Omit atom symbols. Represent Omit structure by showing bonds between carbons and atoms other than hydrogen. hydrogen. Atoms other than carbon and hydrogen Atoms are called heteroatoms. heteroatoms Dr. Wolf's CHM 201 & 202 OH OH 1- 56 Bond-line formulas Bond-line formulas H H2C H2C Cl C C Cl CH2 CH2 is shown as H Omit atom symbols. Represent Omit structure by showing bonds between carbons and atoms other than hydrogen. hydrogen. Atoms other than carbon and hydrogen Atoms are called heteroatoms. heteroatoms Dr. Wolf's CHM 201 & 202 H 1- 57 Constitutional Isomers Dr. Wolf's CHM 201 & 202 1- 58 Constitutional isomers Constitutional isomers Isomers are different compounds that Isomers have the same molecular formula. have Constitutional isomers are isomers Constitutional that differ in the order in which the atoms are connected. atoms An older term for constitutional An isomers is “structural isomers.” isomers Dr. Wolf's CHM 201 & 202 1- 59 A Historical Note A Historical Note O NH4OCN NH H2NCNH2 Ammonium cyanate Urea In 1823 Friedrich Wöhler discovered that In when ammonium cyanate was dissolved in hot water, it was converted to urea. water, Ammonium cyanate and urea are Ammonium constitutional isomers of CH4N2O. constitutional Ammonium cyanate is “inorganic.” Urea is Ammonium “organic.” Wöhler is credited with an important early contribution that helped overturn the theory of “vitalism.” Dr. Wolf's CHM 201 & 202 1- 60 Examples of constitutional isomers Examples of constitutional isomers H H .. O: N C+ H –: :O .. Nitromethane H H C .. O .. N .. .. O: H Methyl nitrite Both have the molecular formula CH3NO2 but the atoms are connected in a different order. the Dr. Wolf's CHM 201 & 202 1- 61 1.8 Resonance Dr. Wolf's CHM 201 & 202 1- 62 Resonance Resonance two or more Lewis structures may be two written for certain compounds (or ions) written Recall from Table 1.5 Dr. Wolf's CHM 201 & 202 1- 63 Table 1.5 How to Write Lewis Structures Table If an atom lacks an octet, use electron If pairs on an adjacent atom to form a double or triple bond. double Example: Nitrogen has only 6 electrons in the Nitrogen structure shown. H .. .. H C O N O: .. .. .. Dr. Wolf's CHM 201 & 202 H 1- 64 Table 1.5 How to Write Lewis Structures Table If an atom lacks an octet, use electron If pairs on an adjacent atom to form a double or triple bond. double Example: All the atoms have octets in this Lewis All structure. H .. .. H C O N O: .. .. Dr. Wolf's CHM 201 & 202 H 1- 65 Table 1.5 How to Write Lewis Structures Calculate formal charges. Example: None of the atoms possess a formal None charge in this Lewis structure. charge H .. .. H C O N O: .. .. Dr. Wolf's CHM 201 & 202 H 1- 66 Table 1.5 How to Write Lewis Structures Calculate formal charges. Example: This structure has formal charges; is This less stable Lewis structure. less H .. – + H C O N O: .. .. .. Dr. Wolf's CHM 201 & 202 H 1- 67 Resonance Structures of Methyl Nitrite Resonance Structures of Methyl Nitrite same atomic positions differ in electron positions H C .. O .. N .. .. O: H H H C + O .. N .. .. – O: .. H more stable more Lewis structure structure Dr. Wolf's CHM 201 & 202 lless stable ess Lewis structure structure 1- 68 Resonance Structures of Methyl Nitrite Resonance Structures of Methyl Nitrite Resonance same atomic positions H C .. O .. N .. differ in electron positions H .. + O: HCO .. H N .. .. – O: .. H more stable more Lewis structure structure Dr. Wolf's CHM 201 & 202 lless stable ess Lewis structure structure 1- 69 Why Write Resonance Structures? Why Write Resonance Structures? Electrons in molecules are often delocalized between two or more atoms. Electrons in a single Lewis structure are Electrons assigned to specific atoms-a single Lewis structure is insufficient to show electron delocalization. is Composite of resonance forms more accurately depicts electron distribution. depicts Dr. Wolf's CHM 201 & 202 1- 70 Example Example Ozone (O3) Lewis structure of Lewis ozone shows one double bond and one single bond one •• •O • + O •• •• •– O• •• Expect: one short bond and one Expect: long bond long Reality: bonds are of equal length Reality: (128 pm) (128 Dr. Wolf's CHM 201 & 202 1- 71 Example Example Ozone (O3) Lewis structure of Lewis ozone shows one double bond and one single bond one •• •O • + O •• •– O• •• •• Resonance: •• •O • + O Dr. Wolf's CHM 201 & 202 •• •• •– O• •• – •• •O • •• + O •• O• • •• 1- 72 Example Example Ozone (O3) Electrostatic potential map shows both end carbons are equivalent with respect to negative charge. Middle atom is positive. •• •O • + O Dr. Wolf's CHM 201 & 202 •• •• •– O• •• – •• •O • •• + O •• O• • •• 1- 73 1.10 The Shapes of Some Simple Molecules Dr. Wolf's CHM 201 & 202 1- 74 Valence Shell Electron Pair Repulsions Valence Shell Electron Pair Repulsions Valence The most stable arrangement of groups The attached to a central atom is the one that has the maximum separation of electron pairs the (bonded or nonbonded). (bonded Dr. Wolf's CHM 201 & 202 1- 75 Table 1.6 Methane Table 1.6 Methane Table Methane Table Methane tetrahedral geometry H—C—H angle = 109.5° Dr. Wolf's CHM 201 & 202 1- 76 Table 1.6 Methane Table 1.6 Methane Table Methane Table Methane tetrahedral geometry each H—C—H angle = each 109.5° 109.5° Dr. Wolf's CHM 201 & 202 1- 77 Table 1.6 Water Table 1.6 Water Table Water Table Water bent geometry H—O—H angle = 105° H H O : .. but notice the tetrahedral arrangement but of electron pairs of Dr. Wolf's CHM 201 & 202 1- 78 Table 1.6 Ammonia Table 1.6 Ammonia Table Ammonia Ammonia trigonal pyramidal geometry H—N—H angle = 107° H H N : H but notice the tetrahedral arrangement but of electron pairs of Dr. Wolf's CHM 201 & 202 1- 79 Table 1.6 Boron Trifluoride Table 1.6 Boron Trifluoride Table Boron Table Boron F—B—F angle = 120° trigonal planar geometry trigonal allows for maximum separation separation of three electron pairs Dr. Wolf's CHM 201 & 202 1- 80 Multiple Bonds Multiple Bonds Multiple Four-electron double bonds and six-electron Four-electron triple bonds are considered to be similar to a two-electron single bond in terms of their spatial two-electron requirements. requirements. Dr. Wolf's CHM 201 & 202 1- 81 Table 1.6: Formaldehyde Table 1.6: Formaldehyde Table Formaldehyde Table Formaldehyde H—C—H and H—C—O angles are close to 120° trigonal planar geometry trigonal H C O H Dr. Wolf's CHM 201 & 202 1- 82 Table 1.6 Carbon Dioxide Table 1.6 Carbon Dioxide Table Carbon Table Carbon O—C—O angle = 180° linear geometry O Dr. Wolf's CHM 201 & 202 C O 1- 83 1.11 Molecular Dipole Moments Dr. Wolf's CHM 201 & 202 1- 84 Dipole Moment Dipole Moment A substance possesses a dipole moment substance if its centers of positive and negative charge do not coincide. do µ=exd (expressed in Debye units) + — not polar Dr. Wolf's CHM 201 & 202 1- 85 Dipole Moment Dipole Moment A substance possesses a dipole moment substance if its centers of positive and negative charge do not coincide. do µ=exd (expressed in Debye units) — + polar Dr. Wolf's CHM 201 & 202 1- 86 Molecular Dipole Moments Molecular Dipole Moments δ- O δ+ C O δ- molecule must have polar bonds necessary, but not sufficient need to know molecular shape because individual bond dipoles can cancel because Dr. Wolf's CHM 201 & 202 1- 87 Molecular Dipole Moments Molecular Dipole Moments O C O Carbon dioxide has no dipole moment; µ = 0 D Carbon Dr. Wolf's CHM 201 & 202 1- 88 Figure 1.7 Figure 1.7 Carbon tetrachloride µ=0D Dr. Wolf's CHM 201 & 202 Dichloromethane µ = 1.62 D 1.62 1- 89 Figure 1.7 Figure 1.7 Resultant of these two bond dipoles is Resultant of these two bond dipoles is µ=0D Carbon tetrachloride has no dipole moment because all of the individual bond dipoles cancel. Dr. Wolf's CHM 201 & 202 1- 90 Figure 1.7 Figure 1.7 Resultant of these two bond dipoles is Resultant of these two bond dipoles is µ = 1.62 D 1.62 The individual bond dipoles do not cancel in dichloromethane; it has a dipole moment. Dr. Wolf's CHM 201 & 202 1- 91 1.13 Acids and Bases: The Arrhenius View Dr. Wolf's CHM 201 & 202 1- 92 Definitions Definitions Arrhenius An acid ionizes in water to give protons. A An base ionizes in water to give hydroxide ions. base Brønsted-Lowry An acid is a proton donor. A base is a proton An acceptor. acceptor. Lewis An acid is an electron pair acceptor. A base An is an electron pair donor. is Dr. Wolf's CHM 201 & 202 1- 93 Arrhenius Acids and Bases Arrhenius Acids and Bases An acid is a substance that ionizes to give An protons when dissolved in water. protons H A . H + + . A– A base is a substance that ionizes to give base hydroxide ions when dissolved in water. hydroxide M Dr. Wolf's CHM 201 & 202 . OH . . –. .. M+ + . OH . 1- 94 Arrhenius Acids and Bases Arrhenius Acids and Bases Strong acids dissociate completely in water. Strong Weak acids dissociate only partially. Weak H A . H + + . A– Strong bases dissociate completely in water. Strong Weak bases dissociate only partially. Weak M Dr. Wolf's CHM 201 & 202 . OH . . –. .. M+ + . OH . 1- 95 Acid Strength is Measured by pKa H . H + + . A– A Ka = [H+][A–] [HA] pKa = – log10Ka log Dr. Wolf's CHM 201 & 202 1- 96 1.14 Acids and Bases: The Brønsted-Lowry View Brønsted-Lowry definition an acid is a proton donor a base is a proton acceptor Dr. Wolf's CHM 201 & 202 1- 97 A Brønsted Acid-Base Reaction A Brønsted Acid-Base Reaction A proton is transferred from the acid to the proton base. base. B. + H A . base Dr. Wolf's CHM 201 & 202 + B H + . A– . acid 1- 98 A Brønsted Acid-Base Reaction A Brønsted Acid-Base Reaction A proton is transferred from the acid to the proton base. base. B. + H A . base Dr. Wolf's CHM 201 & 202 acid + B H + . A– . conjugate conjugate conjugate conjugate acid base acid base 1- 99 Proton Transfer from HBr to Water Proton Transfer from HBr to Water hydronium ion H .. . O . + H Br . .. .. . H base Dr. Wolf's CHM 201 & 202 acid H + ..– . O H + . Br . . .. . . H conjugate conjugate conjugate acid base 1- 100 Equilibrium Constant for Proton Transfer Equilibrium Constant for Proton Transfer H .. – + . O H + . Br . .. . . . H H .. . O . + H Br . .. .. . H Ka = [H3O+][Br–] [HBr] Takes the same form as for Arrhenius Ka, but Takes but H3O+ replaces H+. H3O+ and H+ are considered equivalent, and there is no difference in Ka values for Arrhenius and Brønsted acidity. Brønsted Dr. Wolf's CHM 201 & 202 1- 101 Equilibrium Constant for Proton Transfer Equilibrium Constant for Proton Transfer H .. – + . O H + . Br . .. . . . H H .. . O . + H Br . .. .. . H Ka = [H3O+][Br–] [HBr] pKa = – log10 Ka log Dr. Wolf's CHM 201 & 202 1- 102 Water as a Brønsted Acid Water as a Brønsted Acid H .. – . N . + H OH .. . . H base Dr. Wolf's CHM 201 & 202 acid H .N . H – .. H + . OH .. . conjugate conjugate conjugate acid base 1- 103 Dissociation Constants (pKaa) of Acids* Dissociation Constants (pK ) of Acids* stronger stronger acid acid pKa -10.4 Conj. Base HI HBr weaker weaker acid acid Acid -5.8 Br– H2SO4 HCl H3O+ -4.8 – HSO4 -3.9 Cl– -1.7 H2O I– strong acids are stronger than hydronium ion Dr. Wolf's CHM 201 & 202 1- 104 Important Generalization! Important Generalization! stronger stronger acid acid Acid HI pKa -10.4 Conj. Base – I– – -5.8 Br – H2SO4 -4.8 – HSO4 HCl + H3O -3.9 Cl -1.7 H2O HBr weaker weaker acid acid – The stronger the acid, the weaker the conjugate The base. base. Dr. Wolf's CHM 201 & 202 1- 105 Dissociation Constants (pKaa) of Acids* Dissociation Constants (pK ) of Acids* Acid H3O+ HF pKa –1.7 3.5 CH3CO2H NH4+ 4.6 H2O 15.7 9.2 Conj. Base H2O F– CH3 O2– C NH3 HO– weak acids are weaker than hydronium ion Dr. Wolf's CHM 201 & 202 1- 106 Dissociation Constants (pKaa) of Acids* Dissociation Constants (pK ) of Acids* Acid CH3OH H2O CH3CH2 H O (CH3)2CHOH pKa 15.2 15.7 ~16 ~17 Conj. Base CH3O– HO– CH3CH2O– (CH3)2CHO– (CH3)3COH ~18 (CH3)3CO– alcohols resemble water in acidity; their conjugate bases are comparable to hydroxide ion in basicity bases Dr. Wolf's CHM 201 & 202 1- 107 Dissociation Constants (pKaa) of Acids* Dissociation Constants (pK ) of Acids* Acid NH3 (CH3)2NH pKa ~36 ~36 Conj. Base NH2– (CH3)2N– ammonia and amines are very weak acids; their conjugate bases are very strong bases Dr. Wolf's CHM 201 & 202 1- 108 Dissociation Constants (pKaa) of Acids* Dissociation Constants (pK ) of Acids* Acid pKa Conj. Base HC H 26 H H H H H2C CH2 43 CH3CH3 HC C– H CH H H H – H 45 62 H – – CH H2C – CH3CH2 Most hydrocarbons are extremely weak acids. Dr. Wolf's CHM 201 & 202 1- 109 1.15 What Happened to pKb? Dr. Wolf's CHM 201 & 202 1- 110 About pKaa and pKbb About pK and pK A separate “basicity constant” Kb is not separate necessary. Because of the conjugate relationships in the Brønsted-Lowry approach, we can examine acid-base reactions by relying exclusively on pKa values. Dr. Wolf's CHM 201 & 202 1- 111 H Example Example H N• • N H •• •• Which is the stronger base, ammonia (left) or Which pyridine (right)? pyridine Recall that the stronger the acid, the weaker the Recall conjugate base. conjugate Therefore, the stronger base is the conjugate of Therefore, the weaker acid. the Look up the pKa values of the conjugate acids of ammonia and pyridine in Table 1.7. ammonia Dr. Wolf's CHM 201 & 202 1- 112 Example Example H H + N H pKa = 9.3 weaker acid pKa = 5.2 stronger acid H + N H Dr. Wolf's CHM 201 & 202 Therefore, ammonia is a Therefore, stronger base than pyridine stronger 1- 113 1.16 How Structure Affects Acid Strength Dr. Wolf's CHM 201 & 202 1- 114 The Main Ways Structure Affects Acid Strength The Main Ways Structure Affects Acid Strength The strength of the bond to the atom from The strength which the proton is lost. The electronegativity of the atom from which electronegativity the proton is lost. the Changes in electron delocalization on Changes electron ionization. ionization. Dr. Wolf's CHM 201 & 202 1- 115 Bond Strength Bond Strength Bond strength is controlling factor when Bond comparing acidity of hydrogen halides. hydrogen HF pKa HCl HBr HI 3.1 -3.9 -5.8 -10.4 weakest acid weakest strongest H—X bond Dr. Wolf's CHM 201 & 202 strongest acid weakest H—X bond 1- 116 Bond Strength Bond Strength Recall that bond strength decreases in a group in going down the periodic table. in Generalization: Bond strength is most important factor when considering acidity of protons bonded to atoms in same group of same periodic table (as in HF, HCl, HBr, and HI). Another example: H2S (pKa = 7.0) is a Another stronger acid than H2O (pKa = 15.7). stronger Dr. Wolf's CHM 201 & 202 1- 117 The Main Ways Structure Affects Acid Strength The Main Ways Structure Affects Acid Strength The strength of the bond to the atom from The which the proton is lost. which The electronegativity of the atom from which electronegativity the proton is lost. the Changes in electron delocalization on Changes electron ionization. ionization. Dr. Wolf's CHM 201 & 202 1- 118 Electronegativity Electronegativity Electronegativity is controlling factor when Electronegativity comparing acidity of protons bonded to atoms in the same row of the periodic table. same Dr. Wolf's CHM 201 & 202 1- 119 Electronegativity Electronegativity pKa CH4 60 NH3 36 weakest acid weakest least electronegative Dr. Wolf's CHM 201 & 202 H2O 15.7 HF 3.1 strongest acid most electronegative 1- 120 Electronegativity Electronegativity R . O. + H A .. R + . O H + . A– . . H H The equilibrium becomes more favorable as A The becomes better able to bear a negative charge. becomes Another way of looking at it is that H becomes Another more positive as the atom to which it is attached becomes more electronegative. Dr. Wolf's CHM 201 & 202 1- 121 Bond strength versus Electronegativity Bond strength versus Electronegativity Bond strength is more important when comparing acids in which the proton that is lost is bonded to atoms in the same group of the periodic table. periodic Electronegativity is more important when comparing acids in which the proton that is lost is bonded to atoms in the same row of the periodic table. periodic Dr. Wolf's CHM 201 & 202 1- 122 Acidity of Alcohols Acidity of Alcohols In many acids In the acidic proton acidic is bonded to oxygen. oxygen. Alcohols (RO—H) Alcohols resemble water (HO—H) in their acidity. acidity. Dr. Wolf's CHM 201 & 202 pKa HO—H 15.7 CH3O—H 15.2 CH3CH2O—H 16 (CH3)2CHO—H 17 (CH3)3CO—H 18 1- 123 Acidity of Alcohols Acidity of Alcohols Electronegative substituents can increase the Electronegative acidity of alcohols by drawing electrons away from the —OH group. from pKa CH3CH2OH 16 weaker Dr. Wolf's CHM 201 & 202 CF3CH2OH 11.3 stronger 1- 124 Inductive Effect Inductive Effect F H C C F F O δ+ H H The greater acidity of CF3CH2OH compared to OH CH3CH2OH is an example of an inductive effect. CH OH inductive Inductive effects arise by polarization of the Inductive electron distribution in the bonds between atoms. atoms. Dr. Wolf's CHM 201 & 202 1- 125 Electrostatic Potential Maps Electrostatic Potential Maps The greater positive character of the proton of The the OH group of CF3CH2OH compared to the OH CH3CH2OH is apparent in the more blue color CH OH in its electrostatic potential map. CH CH2OH Dr. Wolf's CHM 201 & 202 3 CF3CH2OH 1- 126 Another example of the inductive effect Another example of the inductive effect O CH3C O H pKa Dr. Wolf's CHM 201 & 202 4.7 weaker O CF3C OH 0.50 stronger 1- 127 The Main Ways Structure Affects Acid Strength The Main Ways Structure Affects Acid Strength The strength of the bond to the atom from The which the proton is lost. The electronegativity of the atom from which the proton is lost. the Changes in electron delocalization on Changes electron ionization. ionization. Dr. Wolf's CHM 201 & 202 1- 128 Electron Delocalization Electron Delocalization R . O. + H A .. R + . O H + . A– . . H H Ionization becomes more favorable if electron Ionization delocalization increases in going from right to left in the equation. left Resonance is a convenient way to show Resonance electron delocalization. electron Dr. Wolf's CHM 201 & 202 1- 129 Nitric Acid Nitric Acid •• H O• •O• + H H• • •• O •• N+ • •• – O• •• • pKa = -1.4 •• H + •O H + H• Dr. Wolf's CHM 201 & 202 – •• •O • •• O• N+ • •• – O• •• • 1- 130 Nitric Acid Nitric Acid • Nitrate ion is stabilized by Nitrate electron delocalization. electron Dr. Wolf's CHM 201 & 202 – •• •O • •• O• N+ • •• – O• •• • 1- 131 Nitric Acid Nitric Acid – •• •O • •• •• – •O• •• Negative charge is shared equally by all three oxygens. three N+ O• •• • •O •• • Dr. Wolf's CHM 201 & 202 •• – •• •O• •• N+ •• – O• •• • •• – •• •O • •• O• N+ • •• – O• •• • 1- 132 Acetic Acid Acetic Acid •• H O• •O• + H H• • •• O •• C pKa = 4.7 • CH3 •• H + •O H + H• Dr. Wolf's CHM 201 & 202 – •• •O • •• O• C • CH3 1- 133 Acetic Acid Acetic Acid •• Acetate ion is stabilized by Acetate electron delocalization. electron – •• •O • •• O• C • CH3 Dr. Wolf's CHM 201 & 202 1- 134 Acetic Acid Acetic Acid Negative charge is shared equally by both oxygens. both •O •• • C •• – •O• •• – •• •O • CH3 Dr. Wolf's CHM 201 & 202 •• •• O• C • CH3 1- 135 1.17 Acid-Base Equilibria Dr. Wolf's CHM 201 & 202 1- 136 Generalization Generalization The equilibrium in an acid-base reaction is The favorable if the stronger acid is on the left and the weaker acid is on the right. the Stronger acid + Stronger base Dr. Wolf's CHM 201 & 202 Weaker acid + Weaker base 1- 137 Example of a strong acid Example of a strong acid H .. . O . + H Br . .. .. . H pKa = -5.8 stronger acid H + .. . O H + . Br . – . .. . . H pKa = -1.7 weaker acid The equilibrium lies to the side of The the weaker acid. (To the right) right Dr. Wolf's CHM 201 & 202 1- 138 Example of a weak acid Example of a weak acid •• O• H •• • • O • + H—OCCH3 •• •• H pKa = 4.7 weaker acid •• H O• – •• • + OCCH3 •O—H + • •• • • H pKa = -1.7 stronger acid The equilibrium lies to the side of The the weaker acid. (To the left) the Dr. Wolf's CHM 201 & 202 1- 139 Important Points Important Points A strong acid is one that is stronger than H3O+. A weak acid is one that is weaker than H3O+. A strong base is one that is stronger than HO–. A weak base is one that is weaker than HO–. The strongest acid present in significant The quantities when a strong acid is dissolved in water is H3O+. water The strongest acid present in significant quantities when a weak acid is dissolved in water is the weak acid itself. water Dr. Wolf's CHM 201 & 202 1- 140 Predicting the Direction of Acid-Base Reactions Predicting the Direction of Acid-Base Reactions – •• •• H—O • + H—OC6H5 •• •• • Phenol Phenol pKa = 10 stronger acid – •• •• H—O—H + • OC6H5 •• •• • Water pKa = 15.7 weaker acid The equilibrium lies to the side of the weaker The acid. (To the right) Phenol is converted to right Phenol phenoxide ion by reaction with NaOH. phenoxide Dr. Wolf's CHM 201 & 202 1- 141 Predicting the Direction of Acid-Base Reactions Predicting the Direction of Acid-Base Reactions O – •• •• HOCO• + H—OC6H5 •• •• • Phenol pKa = 10 weaker acid O – •• •• HOCO—H + • OC6H5 •• •• • Carbonic acid pKa = 6.4 stronger acid The equilibrium lies to the side of the weaker The acid. (To the left) Phenol is not converted to phenoxide ion by reaction with NaHCO3. phenoxide Dr. Wolf's CHM 201 & 202 1- 142 1.18 Lewis Acids and Lewis Bases Dr. Wolf's CHM 201 & 202 1- 143 Definitions Definitions Arrhenius An acid ionizes in water to give protons. A An base ionizes in water to give hydroxide ions. base Brønsted-Lowry An acid is a proton donor. A base is a proton An acceptor. acceptor. Lewis An acid is an electron pair acceptor. A base An is an electron pair donor. is Dr. Wolf's CHM 201 & 202 1- 144 Lewis Acid-Lewis Base Reactions Lewis Acid-Lewis Base Reactions The Lewis acid and the Lewis base can be The either a neutral molecule or an ion. either Lewis acid A+ + Lewis base + A + A+ + A + Dr. Wolf's CHM 201 & 202 – •B • – B • • •B • •B A—B – A—B A—B + – A—B + 1- 145 Example: Two Neutral Molecules Example: Two Neutral Molecules CH2CH3 F3B + •O • Lewis acid Lewis • • CH2CH3 – F3B CH2CH3 + O• • CH2CH3 Lewis base Product is a stable substance. It is a liquid with a boiling point of 126°C. Of the two reactants, BF3 is a gas and CH3CH2OCH2CH3 with a boiling point of 34°C. boiling Dr. Wolf's CHM 201 & 202 1- 146 Example: Ion + Neutral molecule Example: Ion + Neutral molecule •• – H—O• + •• • Lewis base •• H3C—Br • •• • •• H—OCH3 •• + •• – • Br• •• •• Lewis acid Reaction is classified as a substitution. But notice how much it resembles a Brønsted acid-base reaction. •• – H—O• + •• • Dr. Wolf's CHM 201 & 202 •• H—Br • •• • •• H—O—H •• + •• – • Br• •• •• 1- 147 Example: Ion + Neutral molecule Example: Ion + Neutral molecule •• – H—O• + •• • Lewis base •• H3C—Br • •• • •• H—OCH3 •• + •• – • Br• •• •• Lewis acid Brønsted acid-base reactions are a subcategory of Lewis acid-Lewis base reactions. •• – H—O• + •• • Dr. Wolf's CHM 201 & 202 •• H—Br • •• • •• H—O—H •• + •• – • Br• •• •• 1- 148 End of Chapter 1 Dr. Wolf's CHM 201 & 202 1- 149 ...
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