Doc Brown's Chemistry Advanced A Level Notes - Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 8

Part 8.6 The evidence for, and the theory of, the intermolecular hydrogen bonding in simple covalent hydrides and its importance in other molecules

email doc brown - comments - query?

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.

Sub-index for this page

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

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

8.6.3 Other examples of hydrogen bonding

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

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'.

• 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.

• 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 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–H^δ+ llll :B^δ––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).

  • ^δ–N–H^δ+llll^δ–:N–H^δ+    hydrogen bond between ammonia molecules

  • ^δ–O–H^δ+llll^δ–:O–H^δ+    hydrogen bond between water molecules

  • ^δ+H–F:^δ–llll ^δ+H–F^δ–      hydrogen bond between hydrogen fluoride molecules

    • (llll represents the directional hydrogen bond)

    • Intermolecular hydrogen bonding in water

  • 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)

  • 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. ^δ–A–H^δ+llll^δ–B,

    • in ice this is  ^δ–O–H^δ+llll^δ–:O, 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

  • An important exam note:

    • You must clearly show the directional linearity of the X^δ--H^δ+ǁǁǁ:X^δ- 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.

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

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 0^oC to 100^oC 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 0^oC 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 0^oC to 4^oC the effect of ice structure breakdown outweighs the normal thermal expansion, so you get a 2nd anomaly of the maximum density at 4^oC.

(3b) The maximum density at 4^oC is because the breakdown of ice-like structures in liquid water is exactly balanced by the effect of normal thermal expansion.

(3c) From 4^oC 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 100^oC, 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,

^δ–O–H^δ+llll^δ–:O, but don't forget that the linkages

^δ–O–H^δ+llll^δ–:O–H^δ+llll^δ–O–H^δ+llll^δ–:O–H^δ+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!

Hydrogen bonding is the strongest type of intermolecular force.

Its the same intermolecular forces that holds the double helix together in DNA and RNA, and partly responsible for holding together the specific 3D protein shape of enzymes.

Another important consequence of hydrogen bonding in biology

The anomalous density behaviour of ice has really important implications for aquatic life.

Because ice forms and floats on the surface of water, life can go on as normal in the liquid water below the ice.

The ice actually provides some insulation from the cold atmosphere, and in deeper ponds, rivers and lakes, most aquatic life can go on as normal.

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8.6.3 Other examples of hydrogen bonding

• 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,

    • and (ii) the >C=O carbonyl bond is quite polar too.

  • Cl₃^δ–C–H^δ+llll^δ–:O=C^δ+(CH₃)₂ permanent dipole – permanent dipole interactions of a hydrogen bond nature (llll).

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

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

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 Index: Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. K_c and K_p 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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