(a) Introduction to hydrogen bonding

(b) Evidence for hydrogen bonding - spectra, X-ray crystallography and bond lengths

(c) The evidence and theory of hydrogen bonding in simple covalent hydrides (lots of data tables)

(d) Hydrogen bonding in mixtures of water, alcohols and ammonia

(e) Hydrogen bonding - more on the structure of ice, anomalous density of water and consequences

(f) Hydrogen bonding in carboxylic acid dimers

(g) Some of the exceptions to the usual N–H, O–H and H–F hydrogen bonding situations.

(h) Hydrogen bonding in types of protein structures

(i) Hydrogen bonding in the alpha helix of DNA and RNA

(j) Examples of intermolecular AND intramolecular hydrogen bonding for the same aromatic molecule and contrasted them with their positional isomers

(k) Hydrogen bonding in sugars (e.g. glucose) and polysaccharides (e.g. starch)

(l) Hydrogen bonding in polyamides and other polymers

(m) The effect of hydrogen bonding on the dynamic viscosity of liquids

(n) Why some aromatic compounds dissolve in water and others do not

Other evidence is discussed here

(c) More on the evidence for, and the theory of, the intermolecular hydrogen bonding in simple covalent hydrides See also (b) Evidence for hydrogen bonding - spectra, X-ray crystallography and bond lengths Reminders and questions to be answered Hydrides are simple binary compounds formed by combining hydrogen with another element. Hydrogen bonding is the strongest of the intermolecular attractive forces between polar covalent molecules. What evidence is there for hydrogen bonding in covalent hydrides? Why do hydrogen fluoride, water and ammonia have much higher boiling points than expected? Graphs of boiling point versus hydride formula/period are presented and discussed with reference to the intermolecular forces contributing to the overall intermolecular bonding. The evidence and theory of hydrogen bonding in simple covalent hydrides Some data provides evidence of an extra strong intermolecular force due to hydrogen bonding and some data does not, so I've discussed lots of data tables in this context - evidence for hydrogen bonding or not? Abbreviations used in data tables: n/a = not applicable X = element of that group; Mpt/Bpt = melting point/boiling point in oC/K, Pol'y = polarizability in 10-24 cm3 (ignore units, just think of it as a relative value) Electrons = sum of atomic numbers of atoms in the molecular formula, ? = unknown data ΔHvap = enthalpy of vaporization in kJ mol-1,  D = Debye for dipole moments I've found some unreliable, but I've done my best to sort it from internet research. Data table of permanent dipole moments of Group 14-17 (4-7) hydride molecules and noble gases (D in Debye units)   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 D XH3 D H2X D HX D atom D 2 CH4 0.00 NH3 1.48 H2O 1.84 HF 1.91 Ne n/a 3 SiH4 0.00 PH3 0.55 H2S 0.92 HCl 1.05 Ar n/a 4 GeH4 0.00 AsH3 0.16 H2Se 0.40 HBr 0.80 Kr n/a 5 SnH4 0.00 SbH3 0.12 H2Te 0.20 HI 0.42 Xe n/a 6 PbH4 0.00 BiH3 ? H2Po ? HAt ? Rn n/a Comments on dipole moments (X = element of that particular group. Group 4/14: Symmetrical tetrahedral molecules, no net dipole moment. Group 5/15: Ammonia has a relatively large dipole moment from a highly polar bonds and hence a permanently highly polar molecule. However, this isn't enough to predict a particularly high intermolecular forces directly increased by hydrogen bonding. However, the high dipole moment would suggest that there will be a relatively strong intermolecular bond between ammonia molecules. The other dipole moments decrease down the group as expected with deceasing difference in electronegativity of the X-H bonds Group 6/16: Comments as above noting the relatively high dipole moment of water. Group 7/17: Comments as above noting the relatively high dipole moment of hydrogen fluoride. Group 0/18: Dipole moments not applicable to single atoms.   Data table of electron numbers of Group 14-17 (4-7) hydride molecules and noble gases AND data table of polarizability of Group 14-17 (4-7) hydride molecules and noble gases   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 electrons XH3 electrons H2X electrons HX electrons atom electrons 2 CH4 10 NH3 10 H2O 10 HF 10 Ne 10 3 SiH4 18 PH3 18 H2S 18 HCl 18 Ar 18 4 GeH4 36 AsH3 36 H2Se 36 HBr 36 Kr 36 5 SnH4 54 SbH3 54 H2Te 54 HI 54 Xe 54 6 PbH4 86 BiH3 86 H2Po 86 HAt 86 Rn 86   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 Pol'y XH3 Pol'y H2X Pol'y HX Pol'y atom Pol'y 2 CH4 2.45 NH3 2.10 H2O 1.50 HF 0.80 Ne 0.38 3 SiH4 4.78 PH3 4.24 H2S 3.63 HCl 2.52 Ar 1.66 4 GeH4 4.77 AsH3 5.47 H2Se 4.77 HBr 3.62 Kr 2.50 5 SnH4 ? SbH3 ? H2Te ? HI 5.45 Xe 4.01 6 PbH4 ? BiH3 ? H2Po ? HAt ? Rn ? Comments on numbers of electrons in hydride molecules and noble gases All of these are calculated from adding together the atomic numbers of the molecular formula giving 10, 18, 36, 54 and 86 for the hydride molecules of periods 2, 3, 4, 5 and 6. This electron number has been added to the graph of hydride boiling points versus period number (further down the page). AND comments on polarizability In all cases down the group, as expected, the polarizability increases with increased numbers of electrons as extra shells of electrons are added down from one period to another. However, the head elements in each group of period 2, C, N, O, F (plus Ne) do seem to resist being polarized more than expected. Non of this data provides direct evidence for hydrogen bonding, but when you look at the melting/boiling points of ammonia, water and hydrogen fluoride, you should realise they do NOT fit in with the polarizability trend.   Data table of the melting points (oC) for Group 13-18 hydrides and Group 18 noble gases (old group numbers 3 to 7 and 0)   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 mpt/oC XH3 mpt/oC H2X mpt/oC HX mpt/oC atom mpt/oC 2 CH4 -183 NH3 -78 H2O 0 HF -84 Ne -246 3 SiH4 -185 PH3 -113 H2S -86 HCl -115 Ar -189 4 GeH4 -165 AsH3 -116 H2Se -66 HBr -89 Kr -157 5 SnH4 -146 SbH3 -88 H2Te -49 HI -51 Xe -111 6 PbH4 ? BiH3 ? H2Po ? HAt ? Rn -71 Comments on melting points of group 4/14 to 7/17 hydrides and noble gases Group 4/14: Generally a steady increase in melting point, with increased numbers of electrons, increased polarizability, increasing strength of intermolecular bonding from instantaneous dipole - induced dipole attractive forces (Van der Waals, London Forces). Group 5/15: Apart from ammonia, explanation as for group 4/14. Ammonia's melting point is far higher than expected (anomaly) and this is explained by the extra intermolecular force of hydrogen bonding contributing to the total intermolecular bond. Group 6/16: Argument as for group 5/15, but here it is water having an anomalously higher melting point - attributed to the extra contribution of hydrogen bonding to the overall intermolecular bond. Group 7/17: Again, as for group 5/15 and group 6/16, hydrogen fluoride has an anomalously high melting point - attributed to the extra contribution of hydrogen bonding to the overall intermolecular bond. Group 0/18: Though lower values than for Group4/14, generally as expected, a steady increase in melting point, with increased numbers of electrons, increased polarizability, increasing strength of intermolecular bonding (thinking of them as monatomic molecules).   Data table of the enthalpies of vaporisation ΔHvap in kJ mol-1 for Group 13-18 hydrides and Group 18 noble gases (old group numbers 3 to 7 and 0)   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 ΔHvap XH3 ΔHvap H2X ΔHvap HX ΔHvap atom ΔHvap 2 CH4 8.2 NH3 23.4 H2O 41.1 HF 30.2 Ne 1.8 3 SiH4 13.0 PH3 14.6 H2S 18.7 HCl 16.2 Ar 6.5 4 GeH4 23.0 AsH3 16.7 H2Se 19.9 HBr 17.6 Kr 9.0 5 SnH4 32..6 SbH3 21.1 H2Te 23.8 HI 19.8 Xe 12.6 6 PbH4 ? BiH3 ? H2Po ? HAt ? Rn 16.4 Comments on enthalpies of vaporization ΔHvap Group 4/14: Down the group the ΔHvap steadily increases in an expected pattern i.e. the increasing in the electrons in the molecule produces an increase in the instantaneous dipole - induced dipole forces, increasing the strength of the intermolecular bond, so the molecules need a greater kinetic energy to change state. They are symmetrical non-polar molecules. Group 5/15: Beyond NH3, the expected steady rise in ΔHvap is expected, more electrons in the molecule, greater polarizability, greater intermolecular bonding, greater KE required to vaporize. Ammonia's ΔHvap is far higher than expected (anomaly) and this is explained by the extra intermolecular force of hydrogen bonding contributing to the total intermolecular bond. Group 6/16: Argument the same as for group 5/15, but here water has the much higher than expected ΔHvap due to the extra contribution of hydrogen bonding. Group 7/17: Argument the same as for group 5/15 and group 6/16, but here hydrogen fluoride has the much higher than expected ΔHvap due to the extra contribution of hydrogen bonding. Group 0/18: The same argument applies as for Group 4/14 - just think of the noble gases as monatomic molecules, but not as polarizable as the group 4/14 hydrides on the same period, so the values are smaller.   Data table of the boiling points (K) for Group 13-18 hydrides and Group 18 noble gases (old group numbers 3 to 7 and 0) See also the graph after the table and initial accompanying notes.   Group 4/14 hydride Group 5/15 hydride Group 6/16 hydride Group 7/17 halogen hydride Group 0/18 Noble Gas Period XH4 Bpt/K XH3 Bpt/K H2X Bpt/K HX Bpt/K atom Bpt/K 2 CH4 112 NH3 240 H2O 373 HF 293 Ne 27 3 SiH4 161 PH3 185 H2S 212 HCl 188 Ar 87 4 GeH4 184 AsH3 218 H2Se 232 HBr 206 Kr 121 5 SnH4 221 SbH3 256 H2Te 271 HI 238 Xe 166 6 PbH4 260 BiH3 295 H2Po 310 HAt 270 Rn 211 Group 4/14: Down the group the boiling point steadily increases in an expected pattern i.e. the increasing in the electrons in the molecule produces an increase in the instantaneous dipole - induced dipole forces, increasing the strength of the intermolecular bond, so the molecules need a greater kinetic energy to change state. The molecules are very weakly polar, which doesn't seem to have any significant effect on the value of the boiling point. This data set (and graph line) provides a useful baseline when comparing the boiling points of group 5/15, group 6/16 and group 7/17 hydrides. Group 5/15: Beyond NH3, the expected steady rise in boiling point is expected, more electrons in the molecule, greater polarizability, greater intermolecular bonding, greater KE required to vaporize the hydride. However, ammonia's boiling point is far higher than expected (anomaly) and this is explained by the extra intermolecular force of hydrogen bonding contributing to the total intermolecular bond. Group 6/16: Argument the same as for group 5/15, but here water has the much higher than expected boiling point due to the extra contribution of hydrogen bonding. Group 7/17: Argument the same as for group 5/15 and group 6/16, but here hydrogen fluoride has the much higher than expected boiling point due to the extra contribution of hydrogen bonding. Group 0/18: The same argument applies as for Group 4/14 - just think of the noble gases as monatomic molecules, but not as polarizable as the group 4/14 hydrides on the same period. Further discussion on the evidence for hydrogen bonding after the graph. Graph of the boiling points (K) of the hydrides of Groups (13-17) plus those of noble gases for comparison Three graph lines show anomalous behaviour and two do not. X is the non-metal other than hydrogen. The graphs showing the very anomalous, 'higher than expected', boiling points of NH3, H2O and HF i.e. the hydrides from the period 2 (group head elements N, O and F). Other graphs show similar anomalies e.g. by plotting enthalpy of vaporization (ΔHvap) or melting point (mpt) versus the period number or electron number (it isn't just boiling point higher than expected anomalies for NH3, H2O and HF). The hydrides considered here are from the combination of a non-metal from groups 4-7 (groups 14-17) and their boiling points are compared with each other and the noble gases group 0 (group 18) as a sort of base-line trend. The graphs of boiling point versus period for the Group 4 hydrides and the noble gases all show the expected gradual rise in boiling point due to the greater number of electrons in the bigger molecule, facilitating a greater number of transient dipole – induced dipole interactions (more polarizable), therefore increasing the intermolecular forces, apart from three molecules! You get a similar series of graph lines if you plot the enthalpy of vaporization on the y axis. However, ammonia NH3, water H2O and hydrogen fluoride HF show considerably higher boiling points than 'expected'. These anomalous boiling points are accounted for by the phenomena of hydrogen bonding., the strongest of the intermolecular forces (intermolecular bonding). Hydrogen bonding is the strongest of the permanent dipole – permanent dipole intermolecular force, though it is not a true ionic or covalent bond but hydrogen bonding is the only intermolecular force in which the intermolecular bond is directional in space. (see the diagrams of the hydrogen bonding in water and alcohols). Generally speaking, it only occurs where hydrogen is bonded to one of the three most electronegative elements, namely nitrogen, oxygen and fluorine. These three elements can pull electrons towards themselves the most in a covalent bonding situation and will be more partially negative than other elements. Hydrogen only has one electron, so if that electron is pulled away, there is a just a minute proton left behind that will be particularly partially positive. Molecules with this type of highly polar bond with have much stronger permanent dipole – permanent dipole forces, and these are significant enough to have their own category which we call 'hydrogen bonding'. BUT this is still a type of intermolecular force of attraction between molecules which results from a type of bond within the molecule, namely N–H, O–H and H–F. In the hydrogen bond A–Hδ+ llll :Bδ––H, A and B are both very electronegative giving the partial charge distribution as shown, and strong electrostatic attraction between H and B (or H and A). A and B can be the same or different, but for this section we are dealing with .. N–Hδ+llllδ–:N–H ... etc. hydrogen bond between ammonia molecules O–Hδ+llllδ–:O–H ... etc. hydrogen bond between water molecules H–F:δ–llll δ+H–F ... etc. hydrogen bond between hydrogen fluoride molecules   Apart from section (d) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (d) onwards Sub-index of notes on hydrogen bonding (d) Hydrogen bonding in mixtures of water, alcohols and ammonia HO–Hδ+llllδ–:NH3 ... the hydrogen bond between water and ammonia molecules, to explain why ammonia is very soluble in water, they are both highly polar molecules with permanent dipoles. R-O–Hδ+llllδ–:O(H)-R ...the hydrogen bond between alcohol molecules R-O–Hδ+llllδ–:O-H2 ... the hydrogen bond between alcohol and water molecules, why water and ethanol are miscible. δ+llllδ– represents the directional hydrogen bond and : the lone pair of electrons on the electronegative O, N or F atom)   Hydrogen bonding diagrams for pure water and aqueous ammonia solution     Hydrogen bonding diagrams for a pure alcohol and an aqueous solution of an alcohol Let the two liquids or solutes that can be hydrogen bonded be A an B. A and B are the same in pure ammonia, water and hydrogen fluoride, but different in e.g. solutions of ammonia, hydrogen fluoride or a soluble alcohol in water where there would be three possible permutations of hydrogen bonding (AllllA, BllllB and AllllB). If A and B are different, but can form hydrogen bonds between them, then a solution of the mixture can form. Ammonia is very soluble in water because it readily forms multiple hydrogen bonds with water, so in terms of hydrogen bond interactions H2OllllH2O ~ H2OllllNH3  ~ NH3llllNH3. Methanol, ethanol, propan-1-ol and propan-2-ol are very soluble and miscible in all proportions, again because the formation of hydrogen bonds aiding solvation, again with the not too bulky alcohols in terms of interactions, on face value we have H2OllllH2O ~ H2OllllHOR  ~ ROHllllHOR. BUT, in reality, in terms of strength of intermolecular hydrogen bonding, it is more like H2OllllH2O > H2OllllHOR  > ROHllllHOR, but not so different to prevent miscibilty. From the butanols onwards, the they are not miscible, and the solubility rapidly decreases as more of the (stronger)  water hydrogen bonds are disrupted. Although all alcohols have the hydrophilic -OH group, the increasing length (bulk) of the hydrophobic 'hydrocarbon' alkyl group disrupts (breaks) more and more of the hydrogen bonds between the water molecules. Consequently, thermodynamically, it becomes more and more less favourable for the alcohol to dissolve in water as the H2OllllH2O bonds become >> H2OllllHOR bonds. An important exam note: You must clearly show (if it is) the directional linearity of the Xδ--Hδ+ǁǁǁ:Xδ- atomic arrangement of the hydrogen bond including the single X-H covalent bond and the lone pair on the other X atom too! (X is usually O, N or F) You must do this accurately in exams when drawing intermolecular bonding diagrams of water or alcohols because it is the only spatially directed intermolecular force, all the rest of the other types of intermolecular bonding forces are randomised - the δ+ and δ- electric fields acting in all directions. For more on solubility see Solubility of covalent compounds, miscible and immiscible liquids and ... Physical properties (boiling points, solubility) and intermolecular forces (including isomeric ethers) Apart from section (e) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (e) onwards Sub-index of notes on hydrogen bonding (e) Hydrogen bonding - more on the structure of ice, anomalous density of water and consequences - some useful others not so desirable! The structure of ice and the anomalous density of water and important examples of hydrogen bonded molecules in biochemistry are also discussed. Before discussing the structure of ice we need to look at the anomalous density behaviour of water. Left graph - density versus temperature for a typical liquid: (1) Increase in temperature of solid, increase in thermal vibration, molecules move increasing a little more apart, density falls. (2) Melting occurs when intermolecular forces weakened and increased freedom of movement moves the molecules a little bit apart decreasing the density. (3) Increasing the temperature increases the KE of the liquid molecules, more energetic collisions, increasing with increase in temperature, steadily lowers the density of the liquid as the molecules bash each other a bit further apart. Right graph -density versus temperature for  water: You also need to refer to the diagrams of ice structure below, as well as the graphs above. (1) Increase in temperature of solid, increase in thermal vibration, molecules move increasing a little more apart, density of ice falls - normal behaviour. (2) Ice melts, but instead of a decrease in density, you get an increase in density of liquid water compared to ice - anomalous behaviour. This is due to the partial breakdown of the open crystal structure of ice and the liquid molecules, despite their greater KE of movement, they actually have the freedom to get closer together (about 10% less volume) and so the density increases - use a little imagination when looking at the ice diagrams - imagine when the open crystal structure breaks down, the water molecules can actually get closer together. (3a) From 0oC to 100oC there is a continuous breakdown of the hydrogen bonding on 'ice-like' structures in water. YES! the ice structure does not completely break down at 0oC on melting. Clumps of water molecules persist and gradually get broken down with increase in temperature - increase in KE of molecules. At the same time normal thermal expansion is going on! From 0oC to 4oC the effect of ice structure breakdown outweighs the normal thermal expansion, so you get a 2nd anomaly of the maximum density at 4oC. (3b) The maximum density at 4oC is because the breakdown of ice-like structures in liquid water is exactly balanced by the effect of normal thermal expansion. (3c) From 4oC the increasing KE of the molecules and more energetic collisions outweighs the break down of hydrogen bonded clumps of water molecules and normal thermal expansion takes place. BUT, even at 100oC, there is still a low concentration of small clumps of water molecules held together by hydrogen bonds AND the anomalously high boiling point of liquid water is due to the continuous attraction between the very polar water molecules. ice diagram (i) ice diagram (ii) Ice diagram (i): A rough sketch of the arrangement of hydrogen bonds to give the open crystal structure of ice.  In ice there is a tetrahedral 3D network of two intramolecular bonds (-O-H) and two intermolecular hydrogen bonds (-O–Hδ+llllδ–O<). The linear nature of the configuration of the hydrogen bond plus covalent bond leads to a 'extended linear bond' the arranges the atoms further apart than a normal close packing arrangement, δ–O–Hδ+llllδ–:O  but don't forget the linkages are extended ... δ–O–Hδ+llllδ–:O–Hδ+llllδ–O–Hδ+llllδ–:O–Hδ+llll  and occur in water too, albeit on a more transient basis! Some extra notes concerning ice formation and clathrates Water has a very high specific heat capacity due to the thermal energy absorbed in weakening the hydrogen bonds, that's why it is of good use in central heating systems or a coolant liquid in a heat exchanger. The caged structure of ice (illustrated above) allows the formation of clathrates, which are 'cage' structures trapping molecules as gas hydrates e.g. with chlorine you get ratios like Cl2.6H2O frozen under pressure Other gas-ice clathrates include hydrates of the gases CH4, CO2, N2O and CH3F BUT, these are NOT true compounds, the trapped gas molecules are held in place by the 'cage' structure of ice and intermolecular bonding. The methane clathrate is of great significance in these days of climate change and global warming. Methane is trapped in ice in arctic/Antarctic regions and tundra at lower latitudes too as well as deep in oceans. These methane hydrates are unstable and break down on warming to release the powerful greenhouse methane. Ice causes burst pipes! The expansion of water when freezing explains why fossil fuel engines of road vehicles and ships must be protected by “antifreeze” in water and why unprotected pipes in houses break if they are allowed to freeze because of the great expanding force created in ice formation. The formation of the open crystal structure of ice creates a very high pressure, even though it is caused by the formation of supposedly weak intermolecular hydrogen bonds. Another important consequence of hydrogen bonding in biology The anomalous density behaviour of ice has really important implications for aquatic life. Being less dense than liquid water, ice forms and floats on the surface of water, life can go on as normal in the liquid water below the ice, otherwise sinking ice would cause freezing from bottom up and extinguish aquatic life.. Ice acts as a protective layer and actually provides some insulation from the cold atmosphere, and in deeper ponds, rivers and lakes, most aquatic life can go on as normal because the ice forms from the top down. Apart from section (f) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (f) onwards Sub-index of notes on hydrogen bonding (f) Hydrogen bonding in carboxylic acid dimers See also (j) Examples of intermolecular AND intramolecular hydrogen bonding for the same aromatic molecule and its positional isomers Both aliphatic and aromatic carboxylic acids are often isolated dimers (via two hydrogen bonds) in the vapour phase and when dissolved in non-polar organic solvents but free individual molecules in aqueous solution - solvation via the carboxylic acid group and hydrogen bonding. Evidence molecular mass measurements shows that Mr is twice expected. In the diagram below, when R = CH3, we have ethanoic acid and C6H5 gives benzoic acid. In the liquid and solid states the situation for carboxylic acids is more complex, they can be dimers, or multi-linked intermolecular bonding of various complex structures, including linked dimers and the involvement of other intermolecular forces - instantaneous dipole - induced dipole and permanent dipole - permanent dipole attractions (Van der Waals forces). The melting points and boiling points are usually higher than expected because of the effective doubling of the molecular mass i.e. effectively doubling the instantaneous dipole - induced dipole intermolecular bond (London dispersion forces) and these dimers are themselves connected in chains throughout the crystal structure by permanent dipole - permanent dipole attractions including hydrogen bonding. So, the situation overall is very complex with various configurations via hydrogen bonds that will increase the strength of the intermolecular bond and raise melting points and boiling points above what might be expected without hydrogen bonds. Enthalpies of melting or boiling will also be increased as a greater vibrational or translational kinetic energy is needed to overcome the attractive forces of the intermolecular bond. As an 'electron' point of comparison, octan–1–ene (1–octene), C8H16, has 64 electrons and boils at 121oC close to the 118oC bpt of ethanoic acid (the dimer also has 64 electrons). Apart from section (g) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (g) onwards Sub-index of notes on hydrogen bonding (g) Some of the exceptions to the usual N–H, O–H and HF hydrogen bonding situations. Trichloromethane will hydrogen bond with propanone Trichloromethane will hydrogen bond with propanone in a mixture of the liquids because: the combined effect of the three quite electronegative chlorine atoms on one carbon atom makes the -CCl3 grouping behave like a very electronegative atom, and (ii) the >C=O: carbonyl bond is quite polar too. Cl3C–Hδ+llllδ–:O=C(CH3)2 permanent dipole – permanent dipole interactions of a hydrogen bond nature (llll). When you mix the two liquids you get a temperature rise, exothermic change, which is from the thermal energy released as the hydrogen bonds form. Good experimental evidence of a stronger intermolecular bond formation. Hydrogen bonding in potassium difluoride This is a real curiosity and blurs the distinction between an intramolecular covalent bond and an intermolecular hydrogen bond. For potassium bifluoride (potassium hydrogen fluoride) KHF2 the linear shaped ion in aqueous solution and pure crystals is the hydrogen bifluoride ion [FlllH-F]- or is it? because it doesn't conform to usual single covalent H-F bond and a FlllH hydrogen bond. X-ray diffraction analysis has shown both F-H bonds are identical lengths at 0.114 nm yet the ion is considered to contain the strongest known hydrogen bond. It is described as having a 3-centre 4 electron bond, so do both H-F bonds have partial covalent bond character and partial hydrogen bond character i.e. a sort of hybrid bond. The hydrogen bond would be expected to be ~0.24 and the H-F bond length is 0.917 nm, so the bond length in the ion (0.114 nm) is, as expected from the theory above, an intermediate value, but it seems both bonds are much 'closer' to being covalent bonds than one of them being a hydrogen bond. Apart from section (h) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (h) onwards Sub-index of notes on hydrogen bonding (h) Hydrogen bonding in types of protein structures Hydrogen bonding is widespread in many homologous series of organic compounds such as alcohols, carboxylic acids, acid amides (including synthetic polymers), amino acid derivatives like proteins (including enzymes), and RNA and DNA nucleic acid derivatives. Of great biochemistry importance are the hydrogen bonds between the bases in RNA and DNA. The directional nature of the hydrogen bond partly explains why proteins including enzymes and DNA/RNA can have quite stable complex 3D structures. The hydrogen bond is therefore partly responsible for holding together e.g. the specific 3D folded shapes of enzyme proteins, which can only function if they are in the 'correct' folded 3D shape, in particular the active site that accepts the substrate molecule. Descriptors of types of protein structure Primary structure: The amino acid sequence in the polypeptide chain Secondary structure: How the protein chains are folded to give a 3D shape e.g. a helix or sheet configuration. Tertiary structure: The full 'extra' folded 3D structural shape of an individual protein of whatever type i.e. the folding of the secondary structure into the final stable and functional shape - can be referred to as the 'monomer' if a quaternary structure is possible. Quaternary structure: The structure formed by joining identical monomers together, often via covalent bonding e.g. insulin can be made with 2, 4 or 6 protein units, haemoglobin is made from four protein molecules each with an O2 accepting site. The above diagram illustrates a section of the primary and secondary structure of protein. Intermolecular hydrogen bonding can allow folding into sheets (shown above or an alpha helix The hydrogen bond involved is usually the >N-Hδ+llllδ-O=C< intermolecular bond, and both groups are derived from, and linked between, the peptide linkages (peptide bonds) of the primary protein structure of the polypeptide.   A few degrees above 37oC, many enzymes in your body start to become denatured and this includes the breaking of thermally weakened intramolecular hydrogen bonds which are important in holding together the complex folded structure of a protein enzyme. If the 'active site' is damaged (i.e. the active part of the 3D folded structure), where the substrate molecules is changed to a desired useful product, then the enzyme cannot fulfil its biochemical catalytic function This is because the protein molecule no longer has the correct configuration (correct folded shape) and its 'stereochemistry' disrupted - problems for the organism! A change in pH can also denature a protein e.g. protonation of an amino group might break a hydrogen bond. See also in organic chemistry notes sections ... 14.4 Protein-enzyme structure, function and inhibition 14.5 Protein analysis & synthesis AND combinatorial chemistry and autosynthesis Apart from section (i) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see and comparing the boiling pints of different organic molecules and homologous series boiling point comparison and explanation of trends and other case studies of boiling points of covalent molecules related to intermolecular forces otherwise read on from section (i) onwards Sub-index of notes on hydrogen bonding (i) Hydrogen bonding in the alpha helix of DNA and RNA See also organic chemistry part 8.7 Amine bases and the hydrogen bonding in DNA and RNA Hydrogen bonding is the principal intermolecular forces that holds the double helix together in DNA and RNA, that is linking the single strands of DNA into the folded double helix structure. Illustrated above and below are the hydrogen bonds holding the two strands of the nucleotide polymer together in the alpha helix structure of DNA You can see, that because of the shapes of the bases (cytosine and guanine or thymine and adenine), the hydrogen bonds neatly dovetail the two strands of DNA together. hydrogen bonding between the >N-H group of one base and the O=C< group of the other paired base linking the DNA strands together in the alpha helix formation. >N–Hδ+llllδ–O=C<   or  >N–Hδ+llllδ–N<   (for both DNA and RNA) Hydrogen bonding and other intermolecular forces are partly responsible for old (and even prehistoric) DNA retaining its compact alpha helix structure (but usually fragmented), enabling forensic scientists and archaeologists to carry out genetic analysis from long dead animals (usually of humans of human interest).   RNA usually single stranded, but can form a double helix structure like DNA and again you can see because of the shapes of the two bases (uracil and adenine), the hydrogen bonds neatly dovetail the two strands of RNA together.

See also organic chemistry part 8.7 Amine bases and the hydrogen bonding in DNA and RNA

• Apart from section (j) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see

  • and comparing the boiling pints of different organic molecules

  • and homologous series boiling point comparison and explanation of trends

  • and other case studies of boiling points of covalent molecules related to intermolecular forces

  • otherwise read on from section (j) onwards

Sub-index of notes on hydrogen bonding

(j) Examples of intermolecular AND intramolecular hydrogen bonding for the same aromatic molecule and contrasted them with their positional isomers

Some molecules only display intramolecular hydrogen bonding or intermolecular hydrogen bonding and others both.

As already mentioned in section (f), carboxylic acids are often dimers in the vapour phase and when dissolved in non-polar organic solvents (measured M_r is twice expected), but free individual molecules in aqueous solution.

In the liquid and solid states the hydrogen bonding situation for carboxylic acids is more complex, they can be dimers or exhibit intramolecular hydrogen bonding and extended structures via intermolecular hydrogen bonding including the involvement of other intermolecular forces including instantaneous dipole- induced dipole and permanent dipole - permanent dipole forces (Van der Waals forces).

When isomers are compared, you can see the effect of the different hydrogen bonds on e.g. melting point and solubility

To illustrate this, I've described the situation for a variety of aromatic molecules and looked at their melting points and boiling points.

Intramolecular hydrogen bonding tends to reduce the melting/boiling points of the molecules and also tends to reduce their solubility in water.

Case studies of hydrogen bonding

1. nitrophenols

2. aminophenols 

3.hydroxybenzoic acids

4. methyl esters of hydroxybenzoic acids

5. aminobenzoic acids

6. hydroxybenzaldehydes

Intermolecular AND intramolecular hydrogen bonding in nitrophenols and aminophenols

Name of 'phenol'	Structure (hydrogen bonding diagrams below)	Mpt/oC	Bpt/oC	Comments
2-nitrophenol		45	216	2-nitrophenol exhibits both intramolecular and intermolecular hydrogen bonding. All three isomers can display intermolecular hydrogen bonding.
3-nitrophenol		97	high, dec.
4-nitrophenol		114	dec. 279
2-aminophenol		174	dec.?	2-aminophenol displays both intramolecular and intermolecular bonding. All three isomers can display intermolecular hydrogen bonding.
3-aminophenol		122	dec.?
4-aminophenol		187	284

1. nitrophenols (see data table above)

All these three nitrophenols exhibit intermolecular hydrogen bonding, but only 2-nitrophenol can form an intramolecular hydrogen bond and, as with many cases, it seems to significantly lower the melting point.

Is this because the other isomers can form several hydrogen bonds per molecule (diagram above).

Note for this example of 2-nitrophenol, because of the C-O and C-N bond angles from the benzene ring, the 'hexagonal ring' does not allow a linear hydrogen bond >O–H^δ+llll^δ–O-N and also notes the significantly lower melting point and boiling point of 2-nitrophenol.

2-nitrophenol also has a much lower solubility in water (data below in g/100 cm³ at 25^oC)

2-nitrophenol (0.2),  3-nitrophenol (1.35),  4-nitrophenol (1.24)

The reason is due to the intramolecular hydrogen bond in 2-nitrophenol inhibiting forming hydrogen bonds with water to aid solvation. The other two isomeric nitrophenols can readily form hydrogen bonds with water.

 

2. aminophenols (see data table above)

All three aminophenols can display intermolecular hydrogen bonding, but only 2-aminophenol can form a five membered ring intramolecular hydrogen bond involving >O–H^δ+llll^δ–N and all three isomers can form intermolecular hydrogen bonds (no diagram yet).

All three isomers OH hydrogen bonds

Can other intermolecular hydrogen bonds form from the protons of the amine group rather than the nitrogen atom? i.e. >N–H^δ+llll^δ–N  or  >N–H^δ+llll^δ–O

Intermolecular AND intramolecular hydrogen bonding in hydroxybenzoic acids, methyl hydroxybenzoic acids and aminobenzoic acids

The hydrogen bonding in carboxylic acids is complicated and there are up to three possibilities depending on the structure and physical state of the acid.

(i) I've already covered the formation of a carboxylic dimer via hydrogen bonding in section (f) Hydrogen bonding in carboxylic acid dimers

(ii) All carboxylic acids can exhibit intermolecular hydrogen bonding.

(iii) As well as (ii), a the third possibility is an intramolecular hydrogen bond e.g. aromatic carboxylic acids with an -OH or -NH₂ substituent in the 2- position (these are described first in this section).

Names of aromatic acid	Structure (hydrogen bonding diagrams below)	mpt/oC	Comments
2-hydroxybenzoic acid		158	2-hydroxybenzoic acid exhibits intramolecular hydrogen bonding. All three isomers can display intermolecular hydrogen bonding.
3-hydroxybenzoic acid		202
4-hydroxybenzoic acid		214
2-aminobenzoic acid		147	2-aminobenzoic acid exhibits intramolecular hydrogen bonding. All three isomers can display intermolecular hydrogen bonding.
3-aminobenzoic acid		173
4-aminobenzoic acid		188

3. Hydroxybenzoic acids

HO-C₆H₄-COOH  (three positional isomers)

You get the same situation with 2-hydroxybenzoic acid in which the intramolecular hydrogen bonding dominates (diagram below).

All three isomeric hydroxybenzoic acids exhibit intermolecular hydrogen bonding, but only 2-hydroxybenzoic acid can  give an intramolecular hydrogen bond.

Note for 2-hydroxybenzoic acid, because of the C-O and C-C bond angles from the benzene ring, it does not display a linear hydrogen bond O–H^δ+llll^δ–O-H and note the significantly lower melting point of 2-hydroxybenzoic acid.

Other intermolecular hydrogen bonds can be formed between the carbonyl group of the carboxylic acid group and the phenol group  -O–H^δ+llll^δ–O=C<

4. Extra data: The three isomeric methyl esters of hydroxybenzoic acid

HO-C₆H₄-COOCH₃  (three positional isomers)

methyl 2-hydroxybenzoate: mpt -9^oC and bpt 220^oC, exhibits both intramolecular and intermolecular hydrogen bonds

methyl 3-hydroxybenzoate: mpt 72^oC and bpt >281^oC, only intermolecular hydrogen bonds

methyl 4-hydroxybenzoate: mpt 126^oC and bpt 265^oC, only intermolecular hydrogen bonds

All three methyl esters of the hydroxybenzoic acids exhibit intermolecular hydrogen bonding, but only methyl 2-hydroxybenzoate can display intramolecular hydrogen bonding (diagram above).

Note for this example of methyl 2-hydroxybenzoate, because of the C-O and C-C bond angles from the benzene ring, it does not display a linear hydrogen bond O–H^δ+llll^δ–O-H and note the significantly lower melting point and boiling point of methyl 2-hydroxybenzoate.

Other intermolecular hydrogen bonds can be formed between the carbonyl group of the ester linkage and the phenol group  -O–H^δ+llll^δ–O=C<

 

5. Aminobenzoic acids

H₂N-C₆H₄-COOH  (three positional isomers)

All three isomeric aminobenzoic acids can exhibit intermolecular hydrogen bonding, but only 2-aminobenzoic acid can also form an intramolecular hydrogen bond and notes much lower melting point, presumably the intramolecular hydrogen bond inhibits intermolecular hydrogen bonding.

Note for this example of 2-aminobenzoic acid, because of the C-O and C-N bond angles from the benzene ring, it does not display a linear hydrogen bond O–H^δ+llll^δ–N

Other intermolecular hydrogen bonds can be formed between the carbonyl group of the ester linkage and the phenol group  -O–H^δ+llll^δ–O=C<  or  -O–H^δ+llll^δ–O=C< and probably other possibilities too.

 

6. Extra data: Hydroxybenzaldehydes (diagram below)

Intramolecular and intermolecular hydrogen bonds in hydroxybenzaldehyde molecules

HO-C₆H₄-CHO  (three positional isomers)

2-hydroxybenzaldehyde: mpt -7^oC and bpt 197^oC

3-hydroxybenzaldehyde: mpt 101^oC and bpt 191^oC

4-hydroxybenzaldehyde: mpt 115^oC and bpt 310^oC

All three isomeric hydroxybenzaldehydes exhibit intermolecular hydrogen bonding, but only 2-hydroxybenzaldehyde can  give an intramolecular hydrogen bond (diagram above).

Note for 2-hydroxybenzaldehyde, because of the C-O and C-C bond angles from the benzene ring, it does not display a linear hydrogen bond O–H^δ+llll^δ–O-H and note the significantly lower melting point of 2-hydroxybenaldehyde.

Other intermolecular hydrogen bonds can be formed between the carbonyl group of the aldehyde group and the phenol group  -O–H^δ+llll^δ–O=C<

 

• Apart from section (k) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see

  • and comparing the boiling pints of different organic molecules

  • and homologous series boiling point comparison and explanation of trends

  • and other case studies of boiling points of covalent molecules related to intermolecular forces

  • otherwise read on from section (k) onwards

Sub-index of notes on hydrogen bonding

(k) Hydrogen bonding in sugars (e.g. glucose) and polysaccharides (glycogen, starch and cellulose)

All sugars and polysaccharide molecules have hydroxyl (OH) groups which can engage in hydrogen bonding which explains why sugars are very soluble in water.

The diagram above shows the intermolecular hydrogen bonding between glucose molecules themselves and high solubility in water via hydrogen bonding between glucose and water molecules which aids solvation.

A non-polar hydrocarbon - highly polar sugar comparison

I've chosen to discuss two molecules of similar size, molecular mass and number of electrons, which theoretically exhibit similar instantaneous dipole - induced dipole intermolecular bonding.

A considerable contrast when comparing the melting point and solubility in water of glucose C₆H₁₂O₆ (M_r = 180, 96 electrons) mpt ~146^oC and the cycloalkane cyclotridecane C₁₃H₂₆ (M_r = 182, 104 electrons) mpt ~-57^oC.

The cyclo-hydrocarbon has no polar bonds, melts at a much lower temperature and boils intact without thermal decomposition. There are only instantaneous dipole - induced dipole intermolecular bonds in cyclotridecane, but glucose has additional intermolecular bonding from permanent dipole - permanent dipole forces and hydrogen bonding.

Glucose is obviously a crystalline solid at room temperature and only begins to melt at a much higher temperature than cyclotridecane. However, before it can boil, because of the strong intermolecular bonding, it readily decomposes and browning occurs ('caramel' initially) and all sorts of compounds are formed, eventually carbonisation from water loss leaves are dark coloured carbonaceous residue!

 

You get hydrogen bonding in polysaccharide carbohydrates e.g. starch and cellulose

Cellulose, the main structural component of wood, forms long chains from glucose linked by glycoside bonds plus cross-linked by covalent and hydrogen bonds

n C₆H₁₂O₆ ===>  (C₅H₁₀O₅)_n + nH₂O

M_r of polymers is 10⁵ to 10⁷ (number of units 617 to 61700 per molecule)

Glycogen, starch and cellulose are natural polymers formed from glucose molecule.

There are many possible sites for hydrogen bonding in the complex structures of polysaccharide and cellulose molecules.

The diagram just points out one possible set of hydrogen bonding interactions attracting two polysaccharide molecules together.

Unlike poly(alkene) polymers, the intermolecular hydrogen bonding is so strong in polysaccharides that they do NOT melt and at higher temperatures they begin to thermally degrade.  In other words, covalent bonds are broken before all the hydrogen bonds are!

The strong intermolecular hydrogen bonding between the polymer chains accounts for why cellulose and starch are insoluble in water (starch is slightly soluble in warm water).

The intermolecular hydrogen bonding inhibits solvation to dissolve in water, whose molecules cannot penetrate the tightly bonded polysaccharide structure.

• Apart from section (l) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see

  • and comparing the boiling pints of different organic molecules

  • and homologous series boiling point comparison and explanation of trends

  • and other case studies of boiling points of covalent molecules related to intermolecular forces

  • otherwise read on from section (l) onwards

Sub-index of notes on hydrogen bonding

(l) Examples of hydrogen bonding in synthetic polymers (polyamides and poly(ethenol)

Polyamides

Nylon-6,6

The hydrogen bonding between adjacent polymer molecules via the >C=O and H-N groups.

Hydrogen bonding between polymer chains increases the tensile strength of plastic materials.

Repeating unit of Nylon-6,6

In the manufacturing process the strength of Nylon fibres are increased by a technique called cold-drawing in which the threads are tensioned to help the polymer molecules to line up and therefore maximising the surface-surface contact and increasing the intermolecular bonding forces of attraction.

So the effect is to increase the strength of the instantaneous dipole - induced dipole forces and the permanent dipole - permanent dipole attraction AND the symmetrical molecular alignment maximises the hydrogen bonding (llll) between the C=O and the NH groups of adjacent polymer molecules i.e. >C=O^δ-llll^δ+H-N hydrogen bond.

 

Kevlar - hydrogen bonding between adjacent polymer molecules >C=O and H-N

The Kevlar polymer molecules are aligned in sheets to maximise the instantaneous dipole - induced dipole forces, permanent dipole - permanent dipole between them.

The molecular sheet arrangement allows the maximum hydrogen bonding (llll) between the C=O and the NH groups of adjacent Kevlar molecules i.e. >-C=O^δ-llll^δ+H-N< creating beautiful and strong sheets of polymer.

Overall the strong intermolecular bonding gives Kevlar considerable tensile strength and is used in bullet armour protection clothing

Other polymers exhibiting hydrogen bonds

Poly(ethenol) is unusual for a poly(alkene) because it dissolves in water, the hydrogen bonding with water allows solvation (diagram above).

The polymer chain has lots of hydroxy groups to hydrogen bond with water  e.g. CO-H^δ+llll^δ-O-H₂  (see above diagram).

There will of course be hydrogen bonds between poly(ethenol) molecules themselves, if their combined strength exceeds the solvation forces described above, then the polymer becomes insoluble.

• Apart from section (m) onwards - see other detailed case studies of intermolecular bonding including hydrogen bonding see

  • and comparing the boiling pints of different organic molecules

  • and homologous series boiling point comparison and explanation of trends

  • and other case studies of boiling points of covalent molecules related to intermolecular forces

  • otherwise read on from section (m) onwards

Sub-index of notes on hydrogen bonding

(m) The effect of hydrogen bonding on the dynamic viscosity of liquids

Viscosity is a measure of the resistance of a liquid to flow. It generally increases the longer the molecule due to increased intermolecular bonding. For long molecules the viscosity also increases due to greater entanglement.

However, here, we are comparing three molecules of similar size and similar number of electrons, in terms of the effect of their intermolecular bonding on their properties including viscosity.

Their viscosity is quoted in units of dynamic viscosity at 20^oC

Units: centipoise (cP) = millipascal seconds (mPa s) i.e. 1 cP = 1 mPa s

They must be compared at the same temperature because the viscosity of a liquid decreases with increasing temperature.

All have five carbon or oxygen atoms and the same number of electrons in the molecule.

The viscosity data are at 20^oC.

Comparison comments

I've used three 'acronym' shorthands for the intermolecular bonding forces (ID-ID, Pd-PD and HB).

In terms of electron clouds, they are similar sized molecules with equal numbers of electrons, so it is relatively easy to compare the contributing inter molecular bonding forces.

The polarizabilities are quite similar, though the least polar molecule, pentane, has the highest polarizabilty.

Down the table the intermolecular bonding gets stronger increasing all of the following: the dynamic viscosity, boiling point and the enthalpy of vaporisation of the compound.

This is a pretty clear pattern for this set of organic compounds, that is easy to explain in terms of the strength of the total intermolecular bond.

Pentane - a non-polar molecule

... δ⁺=> attraction <=δ^–

The only force contributing to the intermolecular bonding is the instantaneous dipole - induced dipole force (ID-ID), the weakest of the intermolecular attractive forces.

For liquids, the dynamic viscosity decreases with increase in temperature e.g. pentane falls to 0.17 at 50^oC.

butanone - polar molecule

There are two principal contributions to the intermolecular bond - ID-ID attractive forces and permanent dipole - permanent dipole attractions (PD-PD). The extra PD-PD attractive force (illustrated with propanone below) significantly increases the strength of the intermolecular bond e.g. it raises the boiling point by 44^oC compared to pentane.

••••permanent dipole - permanent dipole attractions

butan-1-ol (1-butanol) - polar molecule

Here there is the ID-ID contribution, plus the stronger permanent dipole - permanent dipole attraction (PD-PD) and hydrogen bonding (HB) both via the hydroxyl group. The extra HB contribution from RO–H^δ+llll^δ–:O(H)R bonding between neighbouring alcohol molecules to the intermolecular bond, considerably raises both the boiling point and viscosity

Propane-1,3-diol - polar molecule

As for butan-1-ol in terms of intermolecular bonding, BUT, there are two hydroxyl groups which will increase the 'sites' for hydrogen bonding and so further increasing the total strength of the intermolecular bonding. The double hydroxyl grouping hydrogen bonding considerably increases the boiling point and viscosity even more.

• See other detailed case studies of intermolecular bonding including hydrogen bonding see

  • and comparing the boiling pints of different organic molecules

  • and homologous series boiling point comparison and explanation of trends

  • and other case studies of boiling points of covalent molecules related to intermolecular forces

First, consider the fact that hydrogen bonding explains why the melting point of highly polar phenol (40^oC) is so much higher than for methylbenzene (-95^oC), even though they are similar sized and shaped molecules with the same number of electrons and similar contributions of instantaneous dipole - induced dipole forces contribution to the intermolecular bond.

The extra contribution to the intermolecular bond from the hydrogen bonding raises the melting point of phenol by 135^oC compared to the non-polar hydrocarbon methylbenzene, which cannot exhibit any permanent dipole - permanent dipole attractive forces.

The hydrogen bonding described for phenol also explains why phenols can dissolve in water (some very soluble) and aromatic hydrocarbons are almost insoluble.

Polar phenol molecules can hydrogen bond with water and dissolve (solvation) to a greater or lesser degree but relatively non-polar methylbenzene cannot be solvated by the highly polar water solvent.

The non-polar hydrocarbon disrupts the hydrogen bonding water with no compensating hydrogen bonding or even permanent dipole - permanent dipole attractive forces.

Conversely, aromatic hydrocarbons will dissolve in non-polar organic solvents in which polar phenols are much less soluble.

This sort of argument for similar sized organic molecules can be generalised to predict the likely solubility of an organic compound in water.

See section 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Some learning objectives for hydrogen bonding

(a) Learning objectives for the origin of the hydrogen bond

a1 Know that a hydrogen bond is the strongest of the intermolecular forces contributing to the total overall force of an intermolecular bond.

a2 Know that the hydrogen bond is a special case of a permanent dipole - permanent dipole attractive force.

a3 Know that, unlike other intermolecular forces that act in all directions, the hydrogen bond is directed at a particular angle between the two molecules joined by a hydrogen bond..

a4 Know that a hydrogen intermolecular bond is usually formed between a hydrogen atom (proton) and an electronegative element such as nitrogen, oxygen and fluorine and takes the form of a Y–H^δ+llll:X^δ– link where usually X = N, O or F and Y = N, O and F.

a5 Do NOT confuse a strong intramolecular bond with a much weaker intermolecular bond.

a6 Take care in how you present a diagram of a hydrogen bond i.e. show the directional linearity of the X^δ--H^δ+ǁǁǁ:X^δ- where X is usually O, N or F.

(b) Learning objectives for the spectral and X-ray diffraction evidence for the intermolecular hydrogen bond (see also (c) about the properties of hydrides)

b1 Know that the O-H and N-H chemical shifts are changed if these groups in the molecule are hydrogen bonded.

b2 Know that the wavenumber of O-H infrared spectral line is shifted and broadened if the hydroxy molecule is hydrogen bonded.

b3 Know that from X-ray crystallography the spatial positions of the hydrogen bond can be determined and also show that intermolecular hydrogen bonds are much longer than full covalent bonds.

(c) Learning objectives for evidence from the state of change properties of the hydrides of groups 5/15 to 7/17 of the periodic table.

c1 Know that ammonia, water and hydrogen fluoride have relatively high dipole moments which suggest that there will be a relatively strong intermolecular bond between these molecules.

c2 Know that the melting/boiling points of ammonia, water and hydrogen fluoride do NOT fit in with the polarizability and general mpt/bpt trends for the hydrides of group 4/14 to 7/17..

c3 Be familiar with graphs that show the anomalous behaviour of ammonia, water and hydrogen fluoride e.g. graphs of boiling point versus period number or electron number.

c4 Know, and be able to explain, why the melting/boiling points and enthalpies of state changes for ammonia, water and hydrogen fluoride are higher than expected due to hydrogen bonding.

c5 Be able to draw diagrams to illustrate the hydrogen bonding in ammonia, water and hydrogen fluoride.

(d) Learning objectives for liquid ammonia, water and alcohols, including solution mixtures

d1 Using your knowledge of the nature of the intermolecular hydrogen bond, be able to explain why ammonia is very soluble in water.

d2  Similarly, be able to explain why lower alcohols are very soluble in water (miscible), but solubility rapidly falls as the carbon chain gets longer.

d3 Know that in an alcohol molecule, the hydroxyl group is hydrophilic and solvated by water via hydrogen bonding, but the 'alkyl' chain is hydrophobic and disrupts the hydrogen bonding between water molecules reducing the solubility of the alcohol.

d4 Be able to draw diagrams illustrating the hydrogen bonds in water, alcohols and aqueous ethanol.

(e) Hydrogen bonding - more on the structure of ice, anomalous density of water and consequences

e1 Be able to explain with particle theory of state changes and hydrogen bonding why ice is less dense than water and be able to draw hydrogen bonding diagrams showing the structure of liquid water and ice.

e2 Know that ice as an open cage like structure that can trap small molecules including the greenhouse gas methane and these so-called clathrates are unstable and readily break down to release the gas.

e3. Know and be able to describe the biological significance and importance to aquatic life of ice forming on the surface of water, rather than sinking.

e4 Be able to explain why ice formation can burst pipes and why water using cooling systems need ant-freeze adding to the water!

(f) Hydrogen bonding in carboxylic acid dimers

f1 Know that carboxylic acids can form dimers via two hydrogen bonds and be able to draw a diagram showing the dimer structure.

f2 Know that the measured doubled molecular mass provides evidence of a dimer structure.

f3 Know that the dimer structure increases the strength of the intermolecular bond raising the melting point and boiling point above what might be expected on the basis non-hydrogen bonding forces (Van der Waals forces).

f4 Appreciate the situation is complex with various configurations possible within the hydrogen bonding networks all of which tend to increase the kinetic energy needed to change state i.e. melting or boiling.

(g) Some of the exceptions to the usual N–H, O–H and H–F hydrogen bonding situations.

g1 Appreciate that hydrogen bonds can be formed between molecules that at first sight might not fit in with the usual N O F molecules.

g2 Multi-halogenated (e.g. with electronegative Cl) molecules can be sufficiently polar to form a hydrogen bond with another highly polar molecule.

g3 Know that there is some evidence that the strongest of hydrogen bonds may have  partial covalent character.

(h) Hydrogen bonding in types of protein structures

h1 Know that hydrogen bonding is very important in the 3D structure of proteins e.g. it helps hold the folded structure of a protein which maybe tissue or an enzyme.

h2 Recognise and be able to draw the hydrogen bond >N-H^δ+llll^δ-O=C< system that holds the secondary structure of proteins at the peptide linkage.

h3 Know that increased temperature or change in pH can weaken hydrogen bonds and denature an enzyme inhibiting its function because the protein molecule no longer has the correct configuration (correct folded shape).

(i) Hydrogen bonding in the alpha helix of DNA and RNA

i1 Know that hydrogen bonding is the most important intermolecular bond hold the 3D structure of the double helix structure of DNA and RNA.

i2 Be able to recognise from DNA and RNA structure where the hydrogen bonds might form between the bases.

i3 Be able to draw a diagram illustrating the hydrogen bond between the single strands of DNA or RNA i.e. >N–H^δ+llll^δ–O=C<   or  >N–H^δ+llll^δ–N<   (for both DNA and RNA)

(j) Examples of intermolecular AND intramolecular hydrogen bonding for the same aromatic molecule and contrasted them with their positional isomers

j1 Know and be able to recognise and describe the difference between an intermolecular hydrogen bond and an intramolecular bond hydrogen bond.

j2 Know that some molecules can display both an intermolecular hydrogen bond and an intramolecular hydrogen bond.

j3 Be able to draw diagrams illustrating an intramolecular hydrogen bond and a linked system of intermolecular hydrogen bonds.

j4 In terms of isomers, appreciate that the formation of an intramolecular bond affects physical properties such as melting point and solubility in solvents.

(k) Hydrogen bonding in sugars and polysaccharides

k1 know that sugars like glucose and polysaccharides like starch have hydroxyl groups that allow intermolecular hydrogen bonding between the molecules.

k2 Know that sugar molecules readily dissolve in water, the solvation process aided by hydrogen bonding between the water and sugar molecules.

k3 Be able to argue that the hydrogen bonding between the polymer chains of starch and cellulose is too strong to allow solvation to dissolve in water.

(l) Hydrogen bonding in polyamides and other polymers

j1 Be able to recognise and explain the hydrogen bonds that give strength to polyamide plastic materials like Nylon and Kevlar.

j2 Appreciate that the manufacturing process maximises the hydrogen bonding (llll) between the C=O and the NH groups of adjacent polymer molecules i.e. >C=O^δ-llll^δ+H-N hydrogen bond to maximise the tensile strength of the polymer.

j3 Be able to explain why poly(ethenol) can dissolve in water with the aid of hydrogen bonding to water molecules aiding the solvation process.

(m) The effect of hydrogen bonding on the dynamic viscosity of liquids

m1 Be able to predict the boiling point order of a group of similar sized molecules with a similar number of electrons.

m2 You must be able to argue this order on the basis of which attractive forces may be exhibited by a molecule - predictions you should be able to make on the basis of its molecular structure.

m3 From the same arguments, be able to predict the increasing order of the values of enthalpy of vaporisation and dynamic viscosity.

(n) Why some aromatic compounds dissolve in water and others do not

n1 Be able to argue why a polar compound like phenol is much more soluble in water than methylbenzene due to the hydrogen bonding between the phenol and water molecules.

n2 Be able to contrast polar molecules and non-polar molecules in their ability to dissolve in highly polar solvents like water or relatively non-polar hydrocarbon solvents.

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Index of ALL my chemical equilibrium context revision notes Index