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Acids and Bases Section 4 (Chapter 6, M,F&T) Acid-Base and Donor-Acceptor Chemistry There are a variety of definitions for acids and bases. Some definitions are quite specific (and limited) A few of the more popular acid-base definitions are: Arrhenius Brønsted-Lowry Solvent system Lewis Arrhenius Acids and Bases Arrhenius acids are defined as substances which increase the concentration of H3O+ ions when added to water (e.g. H2SO4) H2SO4 + H2O HSO4- + H3O+ Arrhenius bases are substances that increase the concentration of OH- ions when added to water (e.g. NH3) NH3 + H2O D NH4+ + OHIt is a definition that is limited to aqueous solutions Brønsted Acids and Bases A more general definition of acids and bases that is defined as follows: Brønsted acids are proton (H+) donors Brønsted bases are proton acceptors The definition applies to all Arrhenius cases, and even in non-aqueous solutions HCl + H2O H3O+ + ClNH3 + H2O D NH4+ + OHNH3 + HCl NH4+ + Cl- Structure and Brønsted Acidity The ability of a Brønsted acid to donate a proton will depend on the polarity of the H-X bond (in most Brønsted acids, X = N, O, or a halogen) Electron-withdrawing groups attached to X will increase the quantity of partial positive charge on the H-atom, making it more susceptible to nucleophilic attack by a solvent (inductive effect) O O H3C O acetic acid H F3C O H Acids and Bases There are a variety of definitions for acids and bases. Some definitions are quite specific (and limited) A few of the more popular acid-base definitions are: Arrhenius Brønsted-Lowry Solvent system Lewis trifluoroacetic acid O-H bond which is broken to yield H+ ion 1 Pauling’s Rules for Oxyacids Structure and Brønsted Acidity For oxyacids, acid strength increases with the number of oxygen atoms: more O-atoms, greater inductive effect the stability of the conjugate base may also be the driving force behind dissociation (resonance structures) - O O O Cl O H O O Cl O + + H + H - O O O Cl O H O Cl O - O Cl O H + O Cl O H + + Cl Ka = big O O H Cl O + H + Ka = big Ka = 1.1 x 10-2 Ka = 3.0 x 10-8 Solvent System Definition The solvent system definition of acids and bases is one that evolves from the autodissociation reaction: 2 H2O D H3O+ + OHBy this definition, an acid is anything added to the solvent that increases the concentration of the cation of the autodissociation reaction (e.g. H3O+). For example: H2SO4 + H2O H3O+ + HSO4- Solvent System Definition An acid is a substance that increases the concentration of the cation of the autodissociation reaction When SbF5 is added to a BrF3 solvent, the following reaction occurs: SbF5 + BrF3 D SbF6- + BrF2+ Thus SbF5 is an acid (in BrF3) by the solvent system definition A base is a substance that increases the concentration of the anion of the autodissociation reaction When KF is added to BrF3, the following reaction takes place: KF + BrF3 D K+ + BrF4Thus KF is a base by the solvent system definition The solvent definition is also fairly general, since many solvents are capable of autodissociation: 2 NH3 D NH4+ + NH22 H2SO4 D H3SO4+ + HSO42 OPCl3 D OPCl2+ + OPCl42 BrF3 D BrF2+ + BrF4The last two equations don’t involve H+ ions Solvent System Definition Solvent System Definition To predict the pKa of an oxyacid whose formula can be written OpE(OH)q pKa = 8 – 5p hydrogen-free oxygens Where p is the number of hydrogen-free oxygen atoms. O For polyprotic acids (for which q > 1), there will be an increase in pKa of 5 units for H S successive proton transfers O O O Sulfuric acid (O2S(OH)2, p = 2 and q = 2. H pKa1 ~ -2; pKa2 ~ +3 Even for protic solvents, this definition is more useful than the Brønsted definition, since it treats acidity not as an absolute property of the solute, but must be specified in relation to the solvent used. Example, for acetic acid (CH3COOH) in water: CH3COOH + H2O D CH3COO- + H3O+ solute For acetic acid in H2SO4: CH3COOH + H2SO4 D CH3COOH2+ + HSO4solute Thus acetic acid is an acid in water, but a base in H2SO4 2 Lewis Acids and Bases The definition proposed by Lewis is the most general, and can be summarized by: Lewis Acids and Bases Lewis acids are electron-pair acceptors Lewis bases are electron-pair donors The following are examples: H3O+ + OH- 2H2O BF3 + Et2O BF3OEt2 4NH3 + Cu2+ [Cu(NH3)4]2+ The Lewis acid-base reaction is driven by the base’s ability to donate electrons to the acid Recognizing Lewis acids vs. Lewis bases is not always easy, but bases typically have lone pairs or negative charges, while acids are often cations or may have empty (acceptor) orbitals Lewis Acids Lewis Bases Molecules possessing nitrogen atoms (amines, imines, etc.) (e.g. ammonia, pyridine) H Cations (e.g. carbocations; electrophiles are thus Lewis acids) + N N H H Molecules having oxygen atoms (e.g. water) Includes metal ions (e.g. Fe3+) 3+ Fe O H H Anions (F-, C6H5COO-) - O - O F F Molecules (or ions) that have complete octets, but can rearrange to accept more electrons Molecules that can handle expanded octets (3rd period elements and heavier) and can accept additional e-’s Closed-shell systems that can accommodate more electrons through p* orbitals O - Ge F O OH O H 1. Adduct formation (base donated e- pair to acid) F B F F + NC 2. Displacement reaction F F F N CN CN CN substituents (cyano) are electron-withdrawing, and lower the energy of the p* MO in this molecule F F H N H H + H B F F F F Ge F F 2F NC 2- F F O F F F adduct formed with neutral base indicated with arrow - C + O - B Lewis Acid-Base Reaction Types Lewis Acids Molecules with empty (acceptor) orbitals (e.g. BF3) and incomplete octets F F H H H B N F H H H B + N N F H H N CH3 3. Double displacement H3C Si Br CH3 CH3 + AgCl H3C Si Cl CH3 + AgBr adduct formed with anionic Lewis base indicated with line 3 Lewis Acids and Bases The Acid-Base Interaction The Lewis acid-base reaction is driven by the base’s ability to donate electrons to the acid Recognizing Lewis acids vs. Lewis bases is not always easy, but bases typically have lone pairs or negative charges, while acids are often cations or may have empty (acceptor) orbitals Electronic Factors Hard Soft Acid-Base Concepts Factors Influencing Acid-Base Reactions There are four basic things which must be considered in acid-base (donor-acceptor) reactions: 1. The strength of the A-B bond (electronics) 2. The energy change involved in structural rearrangements 3. Steric contributions 4. Solvent effects Electron donors and acceptors tend to react in ways that favor hard-hard and soft-soft interactions, proposed by Pearson Hard acids are small in size and/or highly charged (e.g. Li+, Ti3+, BF3) (or whose d-electrons are relatively unavailable for bonding) and bind preferentially to small/light basic species F- >> Cl- > Br- > IR2O >> R2S R3N >> R3P Soft acid species are polarizable, and are large, have low charge if ionic (e.g. Ag+, BH3, Hg2+) F- << Cl- < Br- < IR2O << R2S R3N << R3P Soft and Borderline Lewis Acids - low or zero oxidation states, availability of d-electrons for p-bonding 4 Electronic Factors Hard Soft Acid-Base Concepts Hard-Hard HSAB Guidelines Soft-Soft There is a greater separation between the frontier orbitals in a hard species than in a soft species. Hard-hard interactions have more ionic character, while soft-soft have more covalent character. Pearson’s Hardness Parameter 1 I A 2 Hard-hard and soft-soft interactions tend to be favorable Hard-hard creates strong interaction because of ionic component Hard-hard and soft-soft interactions are favored over hard-soft Soft-soft interaction creates bonding MO that is significantly more stable (lower energy) than MO of base (HOMO) or acid (LUMO) Hardness 1 I A 2 Hard-Soft Acid Base Model Pearson: favourable interactions: Hard acid and hard base: ionic interactions dominant Soft acid and soft base: covalent interactions dominant Drago: Quantitative treatment including parameters for electrostatic and covalent contributions A + B AB Hreaction (gas phase or in inert solvent) -H = EAEB + CACB 5 Electronic Factors Electronic Factors HSAB Concepts HSAB Concepts Using HSAB guidelines, reactions between acids and bases can be often be predicted successfully (though not always) Q: Why is AgI(s) very water-insoluble, but LiI very water-soluble? A: AgI is a soft acid-soft base combination, while LiI is hard-soft. The interaction between Li+ and Iions is not strong. Q: Is OH- or S2- more likely to form an insoluble salt with a +3 transition metal ion? A: The harder species will bind more strongly. Between OH- or S2-, OH- is the harder species. AgI(s) + H2O(l) essentially no reaction LiI(s) + H2O(l) Li+(aq) + I-(aq) Qualitative Analysis Separation of Cations In the separation of the group cations carried out this year, HSAB rules are used to separate classes of cations based on different hard and soft interactions soft and Group II: Hg2+, Cd2+, Cu2+, Sn2+, Sb3+, Bi3+ borderline acids Group III: Mn2+, Fe2+, Co2+, Ni2+, Zn2+, Al3+, Cr2+ H2S(g) D 2H+(aq) + S2-(aq) Even at low S2- concentrations, the group II ions precipitate (stronger interactions with the soft base, S2-) Raising the pH increases the S2- concentration, which allows the precipitation of group III ions The group IV are then precipitated as hydroxides. These cations are harder and prefer the hard base OH-. borderline Group IV: Ca2+, Mg2+, Ba2+, K+, NH4+ hard acids Ambidentate Bases GENERAL UNKNOWN SCN- (thiocyanate) can interact through either its S or N atom with Lewis acids. It can donate an electron pair through more than one atom. Interaction will be through the S-atom with a soft acid, or through the N-atom when interacting with hard acids. Cr(III) interacts as Cr-NCS, while Pt(II) does so as Pt-SCN ACIDIC CONDITIONS Decanted Solution (Contains Group III & IV) Precipitate containing Group II Cations BASIC CONDITIONS Decanted Solution Containing Group IV Cations Precipitate containing Group III Cations The soft and borderline cations are separated through reaction with the soft base sulfide, S2-. Group II sulfides are less soluble than group III, so in order to selectively remove group II ions, a low pH is used: 6 Electronic Factors Inductive Effects Electron donating substituents enhance base strength and electron-withdrawing groups enhance electron acceptor (acid) strength gas-phase base strengths Factors Influencing Acid-Base Reactions P Me Me Me H P H H 1. The strength of the A-B bond (electronics) 2. The energy change involved in structural rearrangements 3. Steric contributions 4. Solvent effects PMe3 stronger base than PH3 NMe3 > NHMe2 > NH2Me > NH3 strongest base There are four basic things which must be considered in acid-base (donor-acceptor) reactions: weakest base This plays a role in bond lengths also Me = methyl; alkyl, aryl groups are electron donating; F, CF3, CN, etc. are e- withdrawing Structural Factors Structural Rearrangement Factors Influencing Acid-Base Reactions In some cases, a center must adjust its hybridization in order to accommodate the formation of a new bond F B F F F H N H + F B N F H opposite order to what is expected for inductive effect sp2 H H H 1. The strength of the A-B bond (electronics) 2. The energy change involved in structural rearrangements 3. Steric contributions 4. Solvent effects sp3 Order of Lewis acid strength for BX3 (X = halides) is BF3 < BCl3 < BBr3 This is due to better p-orbital overlap in BF3 than in BCl3, which is better than BBr3 (B-F bonds are shortest). Thus more energy is needed to change from the sp2-hybridized form of BF3. Steric Factors Influence of Solvent Size/Bulk of Lewis Acid/Base Solvent Properties Bulky and/or large groups may interfere with interaction between the donor and acceptor sites of the base and acid steric effect N H3C CH 3 H3C N H3C > > N N > III H 3C > H 3C N > N CH3 > H3C Since nearly all acid-base reactions occur in solution, the properties of a solvent are critical to the success or failure of a reaction. There are five features of solvents that are influential in acid-base reactions: IV reactions with H+ ions (inductive effect of alkyl donor enhances base strength; bulkiness of t-butyl group in III offsets inductive effect N CH3 H 3C II I There are four basic things which must be considered in acid-base (donor-acceptor) reactions: CH3 N H 3C reactions with BF3 shows behavior that is influenced significantly by steric effects of substituents Usable temperature range Large temperature range desirable Dielectric constant, e Important: ability to reduce attraction between ions Solvent’s donor-acceptor properties Affect energies of reactants, products Solvent’s protic acidity/basicity Will it protonate the reactants? Nature and extent of autodissociation Solutes encounter not only solvent molecules, but also cations and anions of autodissociation 7 Solvent Properties Solvation Effects Although in the gas phase, the amine bases exhibit the following trend in base strength: NMe3 > NHMe2 > NH2Me > NH3 In aqueous solution, the trend is NHMe2 > NH2Me > NMe3 > NH3 and NHEt2 > NH2Et ~ NEt3 > NH3 There are four basic things which must be considered in acid-base (donor-acceptor) reactions: 1. The strength of the A-B bond (electronics) 2. The energy change involved in structural rearrangements 3. Steric contributions 4. Solvent effects When the base reacts with water, the ammonium-type conjugate acid produced is charged. The presence of three methyl groups in NMe3 hampers the solvent’s ability to solvate the charged ion (more H-atoms, more H-bonding), making it less stable Aquated Metal Ions Factors Influencing Acid-Base Reactions Me = methyl Et = ethyl The interaction of water molecules with metal ions of high charge and small size (or having a high charge density) can lower the pH of a solution, even though there appears to be no proton donor present The base-acid interaction weakens the O-H bond in associated water molecules, enabling H+ ions to be released into solution M H n+ + O H + [M(H2O)6]n+(aq) + H2O(l) ⇌ H3O+(aq) + [M(H2O)5(OH)](n-1)+(aq) Aquated Metal Ions Smaller and highly charged cations (hard) like Al3+, Fe3+, and Ti3+ are better at pulling away electron density from water molecules than larger ions, thus these aquated ions would be expected to be quite acidic: [Al(H2O)6]3+(aq) + H2O(l) ⇌ H3O+(aq) + [Al(H2O)5(OH)]2+(aq) pKa = 5.0 [Ti(H2O)6]3+(aq) + H2O(l) ⇌ H3O+(aq) + [Ti(H2O)5(OH)]2+(aq) pKa = 3.9 coordination complexes For comparison, pKa for acetic acid is 4.74 8 Coordination Complexes Aquated Ions: Interesting Cases For [Cr(H2O)6]3+, formation of a dinuclear complex is observed in basic solution (this also happens for Fe3+) [Cr(H2O)6]3+(aq) + H2O ⇌ [Cr(H2O)5(OH)]2+ + H3O+(l) 2 Cr(H2O)5(OH)]2+(aq) ⇌ [(H2O)4Cr(mOH)2Cr(H2O)4]4+(aq) + 2H2O(l) m denotes a “bridging” molecule. When bases (“ligands”) interact with metal ions, a coordination complex results. This interaction is created by a neutral or anionic molecule (ligand) donating at least one electron pair to the Lewis acid metal ion The bonding in these complexes results from donation of an electron pair by the ligand to the metal ion (coordinate covalent bond) Metal ions commonly coordinate four, six, or more ligands. These types of complexes bridge the domains of inorganic chemistry, organometallic chemistry (M-C bonds) and biochemistry (porphyrins, proteins, etc.) Bridging molecules (or bridging “ligands”) link Lewis acids ferrocene [Co(NH3)6]3+ Aquated Metal Ions HEMOGLOBIN HbO2 + CO ⇌ HbCO + O2 Lewis basicity: O O C K = 200 O C N - S2- The interaction of water molecules with metal ions of high charge and small size (or having a high charge density) can lower the pH of a solution, even though there appears to be no proton donor present The base-acid interaction weakens the O-H bond in associated water molecules, enabling H+ ions to be released into solution M H n+ + O H + [M(H2O)6]n+(aq) + H2O(l) ⇌ H3O+(aq) + [M(H2O)5(OH)](n-1)+(aq) 9 Coordination Complexes When bases (“ligands”) interact with metal ions, a coordination complex results. This interaction is created by a neutral or anionic molecule (ligand) donating at least one electron pair to the Lewis acid metal ion The bonding in these complexes results from donation of an electron pair by the ligand to the metal ion (coordinate covalent bond) Metal ions commonly coordinate four, six, or more ligands. These types of complexes bridge the domains of inorganic chemistry, organometallic chemistry (M-C bonds) and biochemistry (porphyrins, proteins, etc.) Geometrical Isomerism Two species having the same molecular formula and the same structural framework, but having different spatial arrangements of atoms around a central atom or double bond Exists in Square planar species: Pt(PPh3)2Cl2 Octahedral species: SnMe2F4, SH3F3 Trigonal bipyramidal species: Fe(CO)4PPh3 Double bonds (cis-, trans-): 2-butene “LEWIS BASES” mer- and fac- Isomerism cis-, trans- Isomerism X X cis-[Co(NH3)4Cl2]+ NH3 Co NH3 M Y Cl NH3 Y Y M X NH3 Y Y NH3 fac-MX3Y3 mer-MX3Y3 Cl NH3 Co Cl NH3 trans-[Co(NH3)4Cl2]+ NH3 Cl Cl NH3 Cl Pt Pt NH3 X Y X Cl X NH3 cis-Pt(NH3)2Cl2 NH3 Cl trans-Pt(NH3)2Cl2 10 Chelating Ligands Some molecules/ions are capable of donating electron pairs through more than one atom at once. This interaction results in the formation of a chelate (pronounced: key-late) ring Chelating ligands tend to form very stable complexes with metal ions. Some ligands are even capable of forming more than one chelate ring (example EDTA: ethylene diamine tetraacetic acid) Nice to know: five- and six membered rings tend to be the most stable, and more chelate rings means more stable How many chelate rings in this structure? From Harris, Quantitative Chemical Analysis, 6 th Ed. The Chelate Effect Polydentate ligands form more stable complexes with transition metal ions than monodentate ligands. They can easily replace monodentate ligands in displacement reactions For example, ethylene diamine (en) will replace ammonia in [Cd(NH3)4]2+ [Cd(NH3)4]2+(aq) + 2en(aq) D [Cd(en)2]2+(aq) + 4NH3(aq) The additional stability of a chelate complex over a monodentate one is known as the chelate effect, and is thermodynamic in origin a bidentate ligand H2N : H2N en = : NH2 Chelate Effect atoms in a ligand Optical Isomerism The chelate effect is a result of an entropy increase, and is not so much an enthalpic effect: G = H - TS Cd2+(aq) + 4NH3(aq) [Cd(NH3)4]2+(aq) Ho = -52.5kJ/mol; So = -41.9 J/K.mol Cd2+(aq) + 2en(aq) [Cd(en)2]2+(aq) Ho = -55.7 kJ/mol; So = +10.4 J/K.mol NH2 denticity = # of donor Mz+ Similar to carbon compounds, tetrahedral complexes will also exhibit optical isomerism (chiral complexes). Octahedral complexes incorporating at least two bidentate ligands are also chiral. It is seen in the reaction below that four monodentate ligands are displaced by two bidentate ligands, resulting in a greater degree of disorder (So = +52.3 J/K.mol): [Cd(NH3)4]2+(aq) + 2en(aq) [Cd(en)2]2+(aq) + 4NH3(aq) ENANTIOMERS 11 Optical Isomerism Optical Isomerism cis-complexes of this type exhibit this type of isomerism, but not transmirror plane N1 N1 Cl N2 Cl N2 Co N3 Co Cl N3 Cl N4 N4 Optical Activity rotate 180o A solution of one optical isomer will rotate plane-polarized light by +° N1 Cl N3 A solution of the other optical isomer will rotate it by -° Co N2 Cl An equimolar mixture of the two isomers (racemic) will show no rotation N4 What types of isomers can exist for the following complexes? “propeller complexes” N N N N N N N M M N N N N N (no relation) Lewis Acids and Bases The Lewis acid-base reaction is driven by the base’s ability to donate electrons to the acid Recognizing Lewis acids vs. Lewis bases is not always easy, but bases typically have lone pairs or negative charges, while acids are often cations or may have empty (acceptor) orbitals [Ru(NH3)3(OH2)3]2+ Fe(CO)4Cl2 Ru(bpy)3 Ru(bpy)2Cl2 Ni(CO)2Br2 Cu(NH3)(OH2)BrCl [Ru(tpy)2]2+ bpy N N N N N tpy Polydentate Ligands Other interesting polydentate ligands come from the crown ether class of compounds M+ 12 Stability Constants of Coordination Complexes Crown Ethers Consider the formation of ML6 (where L is a neutral ligand) by the addition of L to an aqueous solution of the cation: [M(H2O)6]z+(aq) + 6L(aq) D [ML6]z+(aq) + 6H2O(l) We can describe this formation reaction with a constant (like K): is the cumulative formation constant (here, 6 ligands in one step) We should break down the formation of this complex step-by-step, since the coordination of each ligand involves 1. displacement of a water molecule stepwise 2. coordination of the new ligand molecule formation For a metal cation of charge z+, constants [M(H2O)6]z+(aq) + L(aq) [M(H2O)5L]z+(aq) + H2O(l) [M(H2O)5L]z+(aq) + L(aq) [M(H2O)4L2]z+(aq) + H2O(l) [M(H2O)4L2]z+(aq) + L(aq) [M(H2O)3L3]z+(aq) + H2O(l) [M(H2O)3L3]z+(aq) + L(aq) [M(H2O)2L4]z+(aq) + H2O(l) [M(H2O)2L4]z+(aq) + L(aq) [M(H2O)L5]z+(aq) + H2O(l) [M(H2O) L5]z+(aq) + L(aq) [ML6]z+(aq) + H2O(l) K1 K2 K3 K4 K5 K6 6 ML z M ( H O ) L 6 2 z 6 6 We call K the stepwise stability (or formation) constant. β is the cumulative stability (or formation) constant In contrast to solubility product constants and acid dissociation constants, K is usually quite large Thus, for [M(H2O)6]n+(aq) + 6L(aq) D [ML6]n+(aq) + 6H2O(l) β6 = K 1 K 2 K 3 K 4 K 5 K 6 or log β6 = logK1 + logK2 + logK3 + logK4 + logK5 + logK6 where each K is calculated as Kn = éëM (H 2O)6-n Lnz+ ùû éë M (H 2O)z+ ù 7-nû [ L ] Stability Constants of F- Complexation Stepwise stability constants for [Al(H2O)6-xFx](3-x)+ (x = 1 to 6) A Possible Exam Question? O O 2- C C O O Consider the formation of a tris(oxalato)iron (III) salt from [Fe(H2O)6]3+(aq). (oxalate = C2O42-) Give expressions for the stepwise equilibria for the formation of [Fe(ox)3]3- from Fe3+(aq) and ox2(log β1 = 7.54, log β2 = 14.59, log β3 = 20.00). What are the numeric values of K1, K2, and K3? Propose a reason for why K decreases in this series? ox2- 13 Answers a) Fe3+(aq) = [Fe(H2O)6]3+(aq) oxalate is a bidentate dianion (ox2-) Stepwise formation of [Fe(ox)3]3-: [Fe(H2O)6]3+(aq) + ox2-(aq) D [Fe(H2O)4(ox)]1+(aq) + 2H2O(l) [Fe(H2O)4(ox)]1+(aq) + ox2-(aq) D [Fe(H2O)2(ox)2]1-(aq) + 2H2O(l) [Fe(H2O)2(ox)2]1-(aq) + ox2-(aq) D [Fe(ox)3]3-(aq) + 2H2O(l) b) β3 = K1K2K3, β2 = K1K2, and β1 = K1. So K1 = 107.54 = 3.5 x 107 K2 = β2/K1 = 1.1 x 107 K3 = β3/K1K2 = 2.6 x 105 c) K will decrease as the charge of the reactant complex decreases, since electrostatic interaction will be less. K1 K2 K3 The Hydrogen Bond – Donor-Acceptor Complex Hydrogen Bonding in H2O - 2d d+ d+ H O H O Caused by: d+H H i) High POLARITY of the O-H bond ii) Availability of unshared electrons on oxygen Limited to H and O? NO! But need high electronegativities and unshared electron pairs H with N, O, F, (S, Cl) Hydrogen Bonding Do not confuse the phenomenon of hydrogen bonding between molecules with the bonds between O and H within a molecule! Hydrogen Bonding 14 The Hydrogen Bond Definition of a ‘hydrogen bond’ is a moving target A hydrogen bond is formed between an H atom attached to an electronegative atom, and another electronegative atom that possesses a lone pair of electrons. The Hydrogen Bond X−HB Hydrogen bond formation has varying contributions from three components: An X−HB interaction is called a hydrogen bond if it constitutes a local bond, and if X−H acts as a proton donor towards Y. 1. An electrostatic component, from the polarity of the XH bond. The hydrogen bond is an attractive interaction between the hydrogen from a group X−H and an atom or a group of atoms B, in the same or different molecule(s), where there is evidence of bond formation. 3. (London) dispersion forces. 2. A partial covalent character, and transfer of charge from B to XH, from a donor-acceptor interaction. Evidence for a Hydrogen Bond XHB XHB linear angle indicative of relatively strong H-bond, short HB distance. Increased deviation from linearity, with longer HB distances, indicates weaker H-bond. Weakening, lengthening of XH bond, decreasing vibrational frequency, formation of a new HB vibrational mode (IR, Raman spectroscopies). Deshielded H nucleus, strong downfield shift in 1H NMR spectrum. 15 Predicting H-Bond Strengths Electrostatic Potential Map for Molecular Iodine I2 XHB ⇌ XHB ⇌ X¯HB+ - Competition between two acids, XH and HB+ pKa(XHB) = pKa(HX) - pKa(BH+) Molecular Orbitals of I2 The Halogen Bond Out-of-phase combination of p-orbitals: * antibonding LUMO Near linear F-Cl-O due to alignment of acceptor * LUMO In-phase combination of p-orbitals: -bonding Lengthening of F-Cl bond 16