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) examples of bond enthalpy and bond length patterns
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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.
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)
Note for reference:
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.
Note on global warming and carbon dioxide emissions
In the debate on fossil fuels its often quoted that natural gas (mainly methane) is 'greener' than heavier fuel oils, so I thought I'd put a data input to justify the statement or otherwise.
(a) In the table of alkane combustion data I've worked out the energy released per gram of pure alkane fuel.
(b) Also, in the table of data, I've worked out the energy released per gram of carbon dioxide formed.
I'm not using these figures to argue the case for continuing to use large quantities of fossil fuels for power generation or heating, but methane is definitely 'greener' than other hydrocarbon fuels.
(c) Burning 'pure' coal i.e. burning pure carbon
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).
Notes: (i) In carbon dioxide, bond enthalpy of C=O is +805 kJ/mol, bond length is 0.116 nm
(ii) The triple bond of carbon monoxide comprises a double covalent bond plus a dative covalent bond.
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 fluorine to iodine the bond length increases and, except for fluorine, the 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 aqueous 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 decreases 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 bond increases but does the bond enthalpy increases with this increased 'ionic' character?
From N-H to F-H there is an increase in bond energy as the bond polarity increases with an increasing difference in electronegativity, BUT ...
(a) This does also coincide with a decreasing covalent atomic radius across Period 2 which would contribute to a decrease in bond length and the increase in bond enthalpy of X–H from left to right - which is a general expected trend.
(b) For the ~non–polar C–H bond, the bond enthalpy is +413 which doesn't quite fit in with the trend.
(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.
enthalpy patterns for combustion
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data patterns for the combustion of alkanes,
data patterns for the combustion of alcohols, patterns in bond
enthalpy (bond energies), relating bond length to bond enthalpy,
graphs of enthalpies of combustion for linear alcohols, enthalpy of
combustion graph for linear alkanes, explaining variation of bond
enthalpy with bond length for single, double and triple bonds, bond
lengths and bond enthalpies for halogen molecules, bond lengths and
bond enthalpies for hydrogen halides, bond lengths and bond
enthalpies for halogenoalkanes (haloalkanes, alkyl halides)
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