Doc Brown's A Level Chemistry Advanced Level Theoretical Physical Chemistry – AS A2 Level Revision Notes – Basic Thermodynamics
GCE Thermodynamics–thermochemistry sub–index links below
Part 1 – ΔH Enthalpy Changes – The thermochemistry of enthalpies of reaction, formation, combustion and neutralisation
Part 1.4 Enthalpy data patterns (a) combustion of alkanes & alcohols, (b) bond enthalpy and bond length
This page describes patterns of enthalpy values for some homologous series of organic compounds e.g. the enthalpies of combustion of alkanes, enthalpy of combustion of alcohols. A second section describes and explains the data patterns for series of bond enthalpies and bond lengths e.g. group 7 hydrides and other sets of 'X–H' bonds and further consideration, single, double and triple carbon–carbon or nitrogen–nitrogen bonds.
Energetics index: GCSE Notes on the basics of chemical energy changes – important to study and know before tackling any of the three Advanced Level Chemistry pages Parts 1–3 here * Part 1a–b ΔH Enthalpy Changes 1.1 Advanced Introduction to enthalpy changes – reaction, formation, combustion : 1.2a & 1.2b(i)–(iii) Thermochemistry – Hess's Law and Enthalpy Calculations – reaction, combustion, formation etc. : 1.2b(iv) Bond Enthalpy Calculations : 1.3a–b Experimental methods for determining enthalpy changes and treatment of results : 1.4 Some enthalpy data patterns : 1.4a The combustion of linear alkanes and linear aliphatic alcohols : 1.4b Some patterns in Bond Enthalpies and Bond Length : 1.4c Enthalpies of Neutralisation : 1.4d Enthalpies of Hydrogenation of unsaturated hydrocarbons and evidence of aromatic ring structure in benzene : Extra Q page A set of practice enthalpy calculations with worked out answers ** Part 2 ΔH Enthalpies of ion hydration, solution, atomisation, lattice energy, electron affinity and the Born–Haber cycle : 2.1a–c What happens when a salt dissolves in water and why? : 2.1d–e Enthalpy cycles involving a salt dissolving : 2.2a–c The Born–Haber Cycle *** Part 3 ΔS Entropy and ΔG Free Energy Changes : 3.1a–g Introduction to Entropy : 3.2 Examples of entropy values and comments * 3.3a ΔS, Entropy and change of state : 3.3b ΔS, Entropy changes and the feasibility of a chemical change : 3.4a–d More on ΔG, Free energy changes, feasibility and applications : 3.5 Calculating Equilibrium Constants : 3.6 Kinetic stability versus thermodynamic feasibility * PLEASE note that delta H/S/G values vary slightly from source to source, so I apologise in advance for any inconsistencies that may arise as I've researched and developed each section.
1.4 Some enthalpy data patterns
1.4a The combustion of linear alkanes and linear aliphatic alcohols
The standard enthalpies of complete combustion (at 298K, 1 atm = 101kPa) are listed below (4 sf)
General formula of these homologous series: Alkanes CnH2n+2 and aliphatic alcohols H(CH2)nOH
and the general equations for complete combustion can be represented as ... (n = 1, 2, 3 etc.)
alkanes: CnH2n+2(g/l) + (11/2n + 1/2)O2(g) ===> nCO2(g) + (n + 1)H2O(l)
alcohols: H(CH2)nOH(l) + 11/2nO2(g) ===> nCO2(g) + (n + 1)H2O(l)
Graph interpretation and comments
The graph of ΔHcomb versus the number of carbon atoms shows an almost linear relationship as the combustion of each extra –CH2– unit usually contributes an extra 632–670kJ to the molar enthalpy of combustion. The first incremental rise in ΔHc from C1 to C2 is slightly anomalous in both homologous series compared to the general trend.
For the first 8 alkanes, this incremental rise ranges from 632 kJ to 670 kJ. For methane ==> ethane the incremental rise is 670 kJ. The increment for butane ==> pentane is 632 kJ and this lesser incremental rise corresponds to a the first change in state involved i.e. some of the energy released on burning pentane must be used to vapourise it and evaporation is an endothermic process. In fact ΔHvap(C5H12) is +36 kJ mol–1. This absorbed energy is not required by methane ==> butane which are already in the gaseous state. Apart from these two small anomalies all the other incremental rises are 653–658 kJ.
In the case of the first 8 alcohols, all liquids at 298K 101kPa, apart from the incremental rise of 641 kJ from methanol to ethanol, all the other incremental rises up this homologous series are 653–656 kJ and these are completely consistent with incremental rises for most alkane.
For the same carbon number (n) the values for alcohols are slightly smaller than those for alkanes because the alcohols are already partially oxidised i.e. the presence of a single oxygen atom in each alcohol molecule.
1.4b Some patterns in Bond Enthalpies and Bond Length
1.4b(i) Examples of single versus multiple bond
For a pair of atoms (similar/dissimilar) the bond length shortens and the bond enthalpy increases in going from a single to double to triple bond (1, 2 and 3 electron pairs involved). This is a rule in chemistry which is always true! The reason is quite simple. A covalent bond results from the sharing of electrons, which is actually the mutual attraction of two positive nuclei for the negative electrons between them. The greater the number of electrons between the two nuclei the stronger the attraction between them. Therefore the nuclei are drawn together more closely to give a shorter bond length and more energy is required to 'pull them apart' i.e. a greater bond enthalpy (bond dissociation energy). The number of bonding electrons for a particular bond divided by 2 is referred to as the bond order.
Since reactions usually involve collision and initiated by bond fission you might think that single bonds would automatically be more reactive i.e. have a lower activation energy due to a smaller bond enthalpy.
BUT, particularly in organic chemistry, the nature of the 'attacking' reagent is a major factor in the feasibility of a reaction. For example, unsaturated alkenes (>C=C< functional group) and alkynes (–CC– functional group) are much more reactive than saturated alkanes with only single C–C bonds. The pi electron clouds of the unsaturated hydrocarbons are very susceptible to attack by electrophilic (electron pair seeking) reagents like bromine Br2, hydrogen bromide HBr etc. The polarised carbonyl group (>Cδ+=Oδ–) in aldehydes and ketones is susceptible to attack at the δ+ carbon by nucleophilic electron pair donors and much more so than the similarly polarised Cδ+–Oδ– bond in alcohols (Cδ+–Oδ––H) or ethers (Cδ+–Oδ––C).
However in the more inorganic situations the expected pattern is observed. Nitrogen, with its triple bond is extremely stable, hence the need for a catalyst and high temperature to make it combine with hydrogen in the Haber Synthesis of ammonia.
1.4b(ii) Some Group VII (Group 7/17) Halogens trends
Some general observations, most of which relate to smaller radii giving shorter stronger bonds:
Halogen molecules X2: From bromine to iodine the bond length increases and bond enthalpy decreases as the radius of the halogen atom increases with increasing number of filled inner electron shells. Fluorine is distinctly anomalous with a much lower than expected bond dissociation energy, though the bond length fits the general trend. This is explained by the close proximity of the small fluorine atoms causing repulsion between them due to the closeness of the outer electron orbitals.
Hydrogen halides HX: From hydrogen fluoride HF(g) to hydrogen iodide HI(g), there is clear trend in increasing bond length and decreasing bond enthalpy. One result is the increasing ease of ionisation from hydrofluoric acid to hydriodic acid so that the HX(aq) acids become stronger down the group. In fact, hydrofluoric acid HF(aq) is a relatively weak acid but hydrochloric, hydrobromic and hydriodic acids are all very strong. The latter three are so strong in aqueous media you don't really see the difference e.g. from pH readings, but in non–aqueous media the differences can be clearly measured.
Halogenoalkanes R3C–X: Based on polarisation of the bond (Cδ+–Xδ–), you might expect the reactivity order with respect to nucleophiles (electron pair donors) attacking the δ+ carbon bond to be R–F > R–Cl > R–Br > R–I as the electronegativity difference decrease from C–F to C–I. However, it is the decreasing bond enthalpies from C–F to C–I that override this polarisation trend giving the reactivity trend R–I > R–Br > R–Cl > R–F.
1.4b(iii) The effect of bond polarity and electronic situation
If the electronegativity difference between two atoms of a covalent bond increases then the polarity of the increases and the bond enthalpy increases with this increased 'ionic' character. I don't know how 'general' this rule is but e.g.
(a) this does also coincide with a decreasing covalent atomic radius across Period 2 which would contribute to the increase in bond enthalpy of X–H from left to right, so I'm not sure which has the greater influence on the trend?
(b) and for the ~non–polar C–H bond, the bond enthalpy is +413 which doesn't fit in with the trend at all, so this would suggest the increasing polarity of the bond does have some effect?
(c) If the polarity of the bond is 'shared out' the bond energy decreases e.g.
(d) These examples also illustrate the difficulties of using average bond enthalpies in theoretical calculations – like it or not, the actual bond enthalpy of an 'A–B' bond is quite dependent on the particular 'electronic' situation even for a particular pair of covalently bonded atoms A and B.
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