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Theoretical-Physical Advanced Level Chemistry - Equilibria - Chemical Equilibrium Revision Notes PART 8.6

 8.6 The evidence for, and the theory of, for  hydrogen bonding in simple covalent hydrides

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).

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. 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

8.6 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

• 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!

• However, ammonia NH₃, water H₂O and hydrogen fluoride HF show considerably higher boiling points than 'expected'.

• This is accounted by the phenomena of hydrogen 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 the three most electronegative elements, namely nitrogen, oxygen and fluorine (* there are some exceptions)

• These elements can pull electrons the most 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-H^δ+ llll :B^δ- , A and B are both very electronegative giving the partial charge distribution as shown, and strong electrostatic attraction between H and B.

  • Extra note:

  • 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 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 tends to keep the bond linear.

  • This 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 3D structures.

  • Check out the zig-zag shape of (HF)_n (n = 2 to 7?) in the vapour just above the boiling point of hydrogen fluoride and the hydrogen bonds between the bases in RNA and DNA.

• 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.

• 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.

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

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. 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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