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Theoretical–Physical Advanced Level Chemistry Revision Notes PART 8.6

8.6 The evidence for, and the theory of, the intermolecular hydrogen bonding in simple covalent hydrides and its importance in other 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 (intermolecular bonding involved). The structure of ice and the anomalous density of water and important examples of hydrogen bonded molecules in biochemistry are also discussed.

(c) doc b KS4 Science GCSE/IGCSE Notes on reversible reactions and chemical equilibrium

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibrium and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series


8.6.1 The evidence and theory of hydrogen bonding in simple covalent hydrides

The anomalous boiling points of NH3, H2O and HF from the period 2 (group head) elements N, O and F

  Group 4 (14) hydride Group 5 (15) hydride Group 6 (16) hydride Group 7 (17) 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 277 Rn 211

X is the non-metal other than hydrogen

  • 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, 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'.

  • TOP and LINKSThese 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.

  • Generally speaking, it only occurs where hydrogen is bonded to one of the three most electronegative elements, namely nitrogen, oxygen and fluorine (but there are some exceptions)

  • These elements can pull electrons 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 strong 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 , A and B are both very electronegative giving the partial charge distribution as shown, and strong electrostatic attraction between H and B.

    • δN–Hδ+ llll δN–Hδ+    hydrogen bond between ammonia molecules

    • δO–Hδ+ llll δO–Hδ+    hydrogen bond between water molecules

    • δ+H–Fδ– llll δ+H–Fδ–      hydrogen bond between hydrogen fluoride molecules

    • A and B are the same in ammonia, water and hydrogen fluoride, but different in e.g. solutions of ammonia or hydrogen fluoride in water where there would be two permutations of hydrogen bonding (A llll A and A llll B)

    • Extra notes: (llll represents the directional hydrogen bond)

    • There is electrostatic repulsion between the two highly electronegative elements A and B which tends to keep the hydrogen bond linear – the delta minus is effectively on the lone pair on the highly electronegative atom.

    • Because A is very electronegative, the electron density on the H atom is relatively small, and the repulsion between B and H is not very strong so the A B distance can be relatively short in the strongest hydrogen bonds and causes repulsion between the atoms A and B, which keeps the bond system linear i.e. δA–Hδ+ llll δB,

      • in ice this is  δO–Hδ+ llll δO, which gives it a very open crystal structure (fully discussed in the next section 8.6.2)

    • Check out the zig–zag shape of (HF)n (n = 2 to 7?) in the vapour just above the boiling point of hydrogen fluoride


8.6.2 Hydrogen bonding - the structure of ice and the anomalous density of water

Before discussing the structure of ice we need to look at the anomalous density behaviour of water.

Left graph - typical liquid:

(1) Increase in temperature of solid, increase in thermal vibration, molecules move increasing a little more apart, density falls.

(2) Melting occurs and increased freedom of movement moves the molecules 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.

Right graph - water: Need to refer to the diagrams of ice structure below.

(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, they actually have the freedom to get closer together (about 10% less volume) and so the density increases.

(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.

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,

δO–Hδ+ llll δO, but don't forget that the linkages

δO–Hδ+ llll δO–Hδ+ llll δO–Hδ+ llll δO–Hδ+ llll  occur in water too, albeit on a more transient basis!

Ice diagram (ii): The roughly tetrahedral arrangement of the oxygen atoms of the water molecule which are linked by a combination of an O-H covalent bond (O-H) and a hydrogen bond ().

This diagram gives a much more accurate 3D picture of the lattice of ice crystals pinched from an old textbook - I'm afraid I'm no great artist!


8.6.3 Other examples of hydrogen bonding (link to examples already discussed)

  • Hydrogen bonding is widespread in many homologous series of organic compounds (as discussed in section 8.2.2) such as alcohols, carboxylic acids, amino acid derivatives like proteins, RNA and DNA and acid amides etc.

    • Of great biochemistry importance are the hydrogen bonds between the bases in RNA and DNA.

  • The directional nature of the hydrogen bond helps explain why ice is quite a hard material and proteins, enzymes and DNA etc. can have quite stable complex 3D structures.

  • Even though hydrogen chloride is a polar molecule, the permanent dipole from is not sufficient to give hydrogen bonding.

  • Some of the exceptions to the N–H, O–H and H–F hydrogen bonding situations.

    • 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.

    • Cl3δ–C–Hδ+ llll δ–O=Cδ+(CH3)2 permanent dipole – permanent dipole interactions of a hydrogen bond nature (llll).

  • 8.2.1 A summary of Van der Waals forces and an introduction to intermolecular forces

  • See also Hydrogen bonding in protein structure, enzyme structure and function

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WHAT NEXT?

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibrium and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series

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