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standard enthalpy of formation
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The standard enthalpy of formation or standard heat of formation of a compound is the change of enthalpy during the formation of 1 mole of the substance from its constituent elements, with all substances in their standard states. The standard pressure value p⦵ = 105 Pa (= 100 kPa = 1 bar) is recommended by IUPAC, although prior to 1982 the value 1.00 atm (101.325 kPa) was used.{{GoldBookRef| file=S05921 | title = standard pressure}} There is no standard temperature. Its symbol is ΔfH⦵. The superscript Plimsoll on this symbol indicates that the process has occurred under standard conditions at the specified temperature (usually 25 Â°C or 298.15 K). Standard states are as follows:
  1. For a gas: the hypothetical state it would have assuming it obeyed the ideal gas equation at a pressure of 1 bar
  2. For a solute present in an ideal solution: a concentration of exactly one mole per liter (1 M) at a pressure of 1 bar
  3. For a pure substance or a solvent in a condensed state (a liquid or a solid): the standard state is the pure liquid or solid under a pressure of 1 bar
  4. For an element: the form in which the element is most stable under 1 bar of pressure. One exception is phosphorus, for which the most stable form at 1 bar is black phosphorus, but white phosphorus is chosen as the standard reference state for zero enthalpy of formation.BOOK,weblink Principles of Modern Chemistry, 978-0-8400-4931-5, 547, Oxtoby, David W, Pat Gillis, H, Campion, Alan, 2011,
For example, the standard enthalpy of formation of carbon dioxide would be the enthalpy of the following reaction under the above conditions:
C(s, graphite) + O2(g) → CO2(g)
All elements are written in their standard states, and one mole of product is formed. This is true for all enthalpies of formation.The standard enthalpy of formation is measured in units of energy per amount of substance, usually stated in kilojoule per mole (kJ mol−1), but also in kilocalorie per mole, joule per mole or kilocalorie per gram (any combination of these units conforming to the energy per mass or amount guideline).All elements in their standard states (oxygen gas, solid carbon in the form of graphite, etc.) have a standard enthalpy of formation of zero, as there is no change involved in their formation.The formation reaction is a constant pressure and constant temperature process. Since the pressure of the standard formation reaction is fixed at 1 bar, the standard formation enthalpy or reaction heat is a function of temperature. For tabulation purposes, standard formation enthalpies are all given at a single temperature: 298 K, represented by the symbol ΔfH{{su|b=298 K|p=⦵}}.

Hess's law

For many substances, the formation reaction may be considered as the sum of a number of simpler reactions, either real or fictitious. The enthalpy of reaction can then be analyzed by applying Hess's Law, which states that the sum of the enthalpy changes for a number of individual reaction steps equals the enthalpy change of the overall reaction. This is true because enthalpy is a state function, whose value for an overall process depends only on the initial and final states and not on any intermediate states. Examples are given in the following sections.

Ionic compounds: Born–Haber cycle

File:Born-haber cycle LiF.svg|450px|thumb|Standard enthalpy change of formation in Born–Haber diagram for lithium fluoridelithium fluorideFor ionic compounds, the standard enthalpy of formation is equivalent to the sum of several terms included in the Born–Haber cycle. For example, the formation of lithium fluoride,
Li(s) + {{1/2}} F2(g) → LiF(s)
may be considered as the sum of several steps, each with its own enthalpy (or energy, approximately):
  1. The standard enthalpy of atomization (or sublimation) of solid lithium.
  2. The first ionization energy of gaseous lithium.
  3. The standard enthalpy of atomization (or bond energy) of fluorine gas.
  4. The electron affinity of a fluorine atom.
  5. The lattice energy of lithium fluoride.
The sum of all these enthalpies will give the standard enthalpy of formation of lithium fluoride.
Delta H_text{f} = Delta H_text{sub} + text{IE}_text{Li} + frac{1}{2}text{B(F–F)} - text{EA}_text{F} - text{U}_text{L}.
In practice, the enthalpy of formation of lithium fluoride can be determined experimentally, but the lattice energy cannot be measured directly. The equation is therefore rearranged in order to evaluate the lattice energy.Moore, Stanitski, and Jurs. Chemistry: The Molecular Science. 3rd edition. 2008. {{ISBN|0-495-10521-X}}. pages 320–321.
U_text{L} = Delta H_text{sub} + text{IE}_text{Li} + frac{1}{2}text{B(F–F)} - text{EA}_text{F} + Delta H_text{f}.

Organic compounds

The formation reactions for most organic compounds are hypothetical. For instance, carbon and hydrogen will not directly react to form methane (CH4), so that the standard enthalpy of formation cannot be measured directly. However the standard enthalpy of combustion is readily mesurable using bomb calorimetry. The standard enthalpy of formation is then determined using Hess's law. The combustion of methane (CH4 + 2 O2 → CO2 + 2 H2O) is equivalent to the sum of the hypothetical decomposition into elements followed by the combustion of the elements to form carbon dioxide and water:
CH4 → C + 2 H2 C + O2 → CO2 2 H2 + O2 → 2 H2O
Applying Hess's law,
ΔcombH⦵(CH4) = [ΔfH⦵(CO2) + 2 ΔfH⦵(H2O)] − ΔfH⦵(CH4).
Solving for the standard of enthalpy of formation,
ΔfH⦵(CH4) = [ΔfH⦵(CO2) + 2 ΔfH⦵(H2O)] − ΔcombH⦵(CH4).
The value of ΔfH⦵(CH4) is determined to be −74.8 kJ/mol. The negative sign shows that the reaction, if it were to proceed, would be exothermic; that is, methane is enthalpically more stable than hydrogen gas and carbon.It is possible to predict heats of formation for simple unstrained organic compounds with the heat of formation group additivity method.

Use in calculation for other reactions

The standard enthalpy change of any reaction can be calculated from the standard enthalpies of formation of reactants and products using Hess's law. A given reaction is considered as the decomposition of all reactants into elements in their standard states, followed by the formation of all products. The heat of reaction is then minus the sum of the standard enthalpies of formation of the reactants (each being multiplied by its respective stoichiometric coefficient, ν) plus the sum of the standard enthalpies of formation of the products (each also multiplied by its respective stoichiometric coefficient), as shown in the equation below:WEB,weblink Enthalpies of Reaction, www.science.uwaterloo.ca, 2 May 2018, no,weblink" title="web.archive.org/web/20171025201240weblink">weblink 25 October 2017,
ΔrH⦵ = Σν Î”fH⦵(products) − Σν Î”fH⦵(reactants).
If the standard enthalpy of the products is less than the standard enthalpy of the reactants, the standard enthalpy of reaction is negative. This implies that the reaction is exothermic. The converse is also true; the standard enthalpy of reaction is positive for an endothermic reaction. This calculation has a tacit assumption of ideal solution between reactants and products where the enthalpy of mixing is zero.For example, for the combustion of methane, CH4 + 2 O2 → CO2 + 2 H2O:
ΔrH⦵ = [ΔfH⦵(CO2) + 2 ΔfH⦵(H2O)] − [ΔfH⦵(CH4) + 2 ΔfH⦵(O2)].
However O2 is an element in its standard state, so that ΔfH⦵(O2) = 0, and the heat of reaction is simplified to
ΔrH⦵ = [ΔfH⦵(CO2) + 2 ΔfH⦵(H2O)] − ΔfH⦵(CH4),
which is the equation in the previous section for the enthalpy of combustion ΔcombH⦵.

Key concepts for doing enthalpy calculations

  1. When a reaction is reversed, the magnitude of ΔH stays the same, but the sign changes.
  2. When the balanced equation for a reaction is multiplied by an integer, the corresponding value of ΔH must be multiplied by that integer as well.
  3. The change in enthalpy for a reaction can be calculated from the enthalpies of formation of the reactants and the products
  4. Elements in their standard states make no contribution to the enthalpy calculations for the reaction, since the enthalpy of an element in its standard state is zero. Allotropes of an element other than the standard state generally have non-zero standard enthalpies of formation.

Examples: standard enthalpies of formation at 25 Â°C

Thermochemical properties of selected substances at 298 K and 1 atm">

Inorganic substances{| class"wikitable sortable Wikipedia is not a sure reference, any one can edit and say hat water is dry "

! Species! Phase! Chemical formula! ΔfH⦵ /(kJ/mol)
Aluminium
| Aluminium| Solid| Al| 0
| Aluminium chloride| Solid| AlCl3| −705.63
| Aluminium oxide| Solid| Al2O3| −1675.5
| Aluminium hydroxide| Solid| Al(OH)3| −1277
| Aluminium sulphate| Solid| Al2(SO4)3| −3440
Ammonia (ammonium hydroxide) >| Aqueous NH3 (NH4OH) −80.8
| Gas NH3 −46.1
Ammonium nitrate >| Solid NH4NO3 −365.6
Barium
| Barium chloride| Solid| BaCl2| −858.6
| Barium carbonate| Solid| BaCO3| −1213
| Barium hydroxide| Solid| Ba(OH)2| −944.7
| Barium oxide| Solid| BaO| −548.1
| Barium sulfate| Solid| BaSO4| −1473.2
Beryllium
| Beryllium| Solid| Be| 0
| Beryllium hydroxide| Solid| Be(OH)2| −902.9999
| Beryllium oxide| Solid| BeO| −609.4(25)
Boron
| Boron trichloride| Solid| BCl3| −402.96
Bromine
| Bromine| Liquid| Br2| 0
| Bromide ion| Aqueous| Br−| −121
| Bromine| Gas| Br| 111.884
| Bromine| Gas| Br2| 30.91
| Bromine trifluoride| Gas| BrF3| −255.60
| Hydrogen bromide| Gas| HBr| −36.29
Cadmium
| Cadmium| Solid| Cd| 0
| Cadmium oxide| Solid| CdO| −258
| Cadmium hydroxide| Solid| Cd(OH)2| −561
| Cadmium sulfide| Solid| CdS| −162
| Cadmium sulfate| Solid| CdSO4| −935
Calcium
| Calcium| Solid| Ca| 0
| Calcium| Gas| Ca| 178.2
| Calcium(II) ion| Gas| Ca2+| 1925.90
| Calcium carbide| Solid| CaC2| −59.8
| Calcium carbonate (Calcite)| Solid| CaCO3| −1206.9
| Calcium chloride| Solid| CaCl2| −795.8
| Calcium chloride| Aqueous| CaCl2| −877.3
| Calcium phosphate| Solid| Ca3(PO4)2| −4132
| Calcium fluoride| Solid| CaF2| −1219.6
| Calcium hydride| Solid| CaH2| −186.2
| Calcium hydroxide| Solid| Ca(OH)2| −986.09
| Calcium hydroxide| Aqueous| Ca(OH)2| −1002.82
| Calcium oxide| Solid| CaO| −635.09
| Calcium sulfate| Solid| CaSO4| −1434.52
| Calcium sulfide| Solid| CaS| −482.4
| Wollastonite| Solid| CaSiO3| −1630
Caesium
| Caesium| Solid| Cs| 0
| Caesium| Gas| Cs| 76.50
| Caesium| Liquid| Cs| 2.09
| Caesium(I) ion| Gas| Cs+| 457.964
| Caesium chloride| Solid| CsCl| −443.04
Carbon
| Carbon (Graphite)| Solid| C| 0
| Carbon (Diamond)| Solid| C| 1.9
| Carbon| Gas| C| 716.67
| Carbon dioxide| Gas| CO2| −393.509
| Carbon disulfide| Liquid| CS2| 89.41
| Carbon disulfide| Gas| CS2| 116.7
| Carbon monoxide| Gas| CO| −110.525
| Carbonyl chloride (Phosgene)| Gas| COCl2| −218.8
| Carbon dioxide (un–ionized)| Aqueous| CO2(aq)| −419.26
| Bicarbonate ion| Aqueous|HCO3–| −689.93
| Carbonate ion| Aqueous|CO32–| −675.23
Chlorine
| Monatomic chlorine| Gas| Cl| 121.7
| Chloride ion| Aqueous| Cl−| −167.2
| Chlorine| Gas| Cl2| 0
Chromium
| Chromium| Solid| Cr| 0
Copper
| Copper| Solid| Cu| 0
| Copper(II) oxide| Solid| CuO| −155.2
Copper(II) sulfate >| Aqueous CuSO4 −769.98
Fluorine
| Fluorine| Gas| F2| 0
Hydrogen
| Monatomic hydrogen| Gas| H| 218
| Hydrogen| Gas| H2| 0
| Water| Gas| H2O| −241.818
| Water| Liquid| H2O| −285.8
| Hydrogen ion| Aqueous| H+| 0
| Hydroxide ion| Aqueous| OH−| −230
| Hydrogen peroxide| Liquid| H2O2| −187.8
| Phosphoric acid| Liquid| H3PO4| −1288
| Hydrogen cyanide| Gas| HCN| 130.5
| Hydrogen bromide| Liquid| HBr| −36.3
| Hydrogen chloride| Gas| HCl| −92.30
| Hydrogen chloride| Aqueous| HCl| −167.2
| Hydrogen fluoride| Gas| HF| −273.3
| Hydrogen iodide| Gas| HI| 26.5
Iodine
| Iodine| Solid| I2| 0
| Iodine| Gas| I2| 62.438
| Iodine| Aqueous| I2| 23
| Iodide ion| Aqueous| I−| −55
Iron
| Iron| Solid| Fe| 0
| Iron carbide (Cementite)| Solid| Fe3C| 5.4
| Iron(II) carbonate (Siderite)| Solid| FeCO3| −750.6
| Iron(III) chloride| Solid| FeCl3| −399.4
| Iron(II) oxide (Wüstite)| Solid| FeO| −272
| Iron(II,III) oxide (Magnetite)| Solid| Fe3O4| −1118.4
| Iron(III) oxide (Hematite)| Solid| Fe2O3| −824.2
| Iron(II) sulfate| Solid| FeSO4| −929
| Iron(III) sulfate| Solid| Fe2(SO4)3| −2583
| Iron(II) sulfide| Solid| FeS| −102
| Pyrite| Solid| FeS2| −178
Lead
| Lead| Solid| Pb| 0
| Lead dioxide| Solid| PbO2| −277
| Lead sulfide| Solid| PbS| −100
| Lead sulfate| Solid| PbSO4| −920
| Lead(II) nitrate| Solid| Pb(NO3)2| −452
| Lead(II) sulfate| Solid| PbSO4| −920
Magnesium
| Magnesium| Solid| Mg| 0
| Magnesium ion| Aqueous| Mg2+| −466.85
| Magnesium carbonate| Solid| MgCO3| −1095.797
| Magnesium chloride| Solid| MgCl2| −641.8
| Magnesium hydroxide| Solid| Mg(OH)2| −924.54
| Magnesium hydroxide| Aqueous| Mg(OH)2| −926.8
| Magnesium oxide| Solid| MgO| −601.6
| Magnesium sulfate| Solid| MgSO4| −1278.2
Manganese
| Manganese| Solid| Mn| 0
| Manganese(II) oxide| Solid| MnO| −384.9
| Manganese(IV) oxide| Solid| MnO2| −519.7
| Manganese(III) oxide| Solid| Mn2O3| −971
| Manganese(II,III) oxide| Solid| Mn3O4| −1387
| Permanganate| Aqueous
MnO−}}| −543
Mercury
| Mercury(II) oxide (red)| Solid| HgO| −90.83
| Mercury sulfide (red, cinnabar)| Solid| HgS| −58.2
Nitrogen
| Nitrogen| Gas| N2| 0
| Ammonia| Aqueous| NH3| −80.8
| Ammonia| Gas| NH3| −45.90
| Ammonium chloride| Solid| NH4Cl| −314.55
| Nitrogen dioxide| Gas| NO2| 33.2
| Nitrous oxide| Gas| N2O| 82.05
| Nitric oxide| Gas| NO| 90.29
| Dinitrogen tetroxide| Gas| N2O4| 9.16
| Dinitrogen pentoxide| Solid| N2O5| −43.1
| Dinitrogen pentoxide| Gas| N2O5| 11.3
Oxygen
|Monatomic oxygen|Gas|O|249
|Oxygen|Gas|O2|0
|Ozone|Gas|O3|143
Phosphorus
| White phosphorus| Solid| P4| 0
| Red phosphorus| Solid| P
page=392}}
| Black phosphorus| Solid| P| –39.3
| Phosphorus trichloride| Liquid| PCl3| −319.7
| Phosphorus trichloride| Gas| PCl3| −278
| Phosphorus pentachloride| Solid| PCl5| −440
| Phosphorus pentachloride| Gas| PCl5| −321
| Phosphorus pentoxide| Solid| P2O5
TITLE=PERRY'S CHEMICAL ENGINEERS' HANDBOOK PUBLISHER=MCGRAW-HILL PAGE=2–191 REF=P2O5,
Potassium
| Potassium bromide| Solid| KBr| −392.2
| Potassium carbonate| Solid| K2CO3| −1150
| Potassium chlorate| Solid| KClO3| −391.4
| Potassium chloride| Solid| KCl| −436.68
| Potassium fluoride| Solid| KF| −562.6
| Potassium oxide| Solid| K2O| −363
| Potassium perchlorate| Solid| KClO4| −430.12
Silicon
| Silicon| Gas| Si| 368.2
| Silicon carbide| Solid| SiC
TITLE=GIBBS ENERGY OF FORMATION OF SIC: A CONTRIBUTION TO THE THERMODYNAMIC STABILITY OF THE MODIFICATIONS DATE=1998 REF=SIC_AAUTHOR= ACCESS-DATE=5 FEBRUARY 2019, SiC_b,
| Silicon tetrachloride| Liquid| SiCl4| −640.1
| Silica (Quartz)| Solid| SiO2| −910.86
Silver
| Silver bromide| Solid| AgBr| −99.5
| Silver chloride| Solid| AgCl| −127.01
| Silver iodide| Solid| AgI| −62.4
| Silver oxide| Solid| Ag2O| −31.1
| Silver sulfide| Solid| Ag2S| −31.8
Sodium
Sodium >| Solid Na 0
| Gas Na +107.5
Sodium bicarbonate >| Solid NaHCO3 −950.8
| Sodium carbonate| Solid| Na2CO3| −1130.77
| Sodium chloride| Aqueous| NaCl| −407.27
| Sodium chloride| Solid| NaCl| −411.12
| Sodium chloride| Liquid| NaCl| −385.92
| Sodium chloride| Gas| NaCl| −181.42
|Sodium chlorate|Solid|NaClO3| -365.4
| Sodium fluoride| Solid| NaF| −569.0
| Sodium hydroxide| Aqueous| NaOH| −469.15
| Sodium hydroxide| Solid| NaOH| −425.93
|Sodium hypochlorite|Solid|NaOCl
347.1
| Sodium nitrate| Aqueous| NaNO3| −446.2
| Sodium nitrate| Solid| NaNO3| −424.8
| Sodium oxide| Solid| Na2O| −414.2
Sulfur
| Solid S8 0.3
| Solid S8 0
| Hydrogen sulfide| Gas| H2S| −20.63
| Sulfur dioxide| Gas| SO2| −296.84
| Sulfur trioxide| Gas| SO3| −395.7
| Sulfuric acid| Liquid| H2SO4| −814
Tin
Titanium
| Titanium| Gas| Ti| 468
| Titanium tetrachloride| Gas| TiCl4| −763.2
| Titanium tetrachloride| Liquid| TiCl4| −804.2
| Titanium dioxide| Solid| TiO2| −944.7
Zinc
| Zinc| Gas| Zn| 130.7
| Zinc chloride| Solid| ZnCl2| −415.1
| Zinc oxide| Solid| ZnO| −348.0
Zinc sulfate>| Solid ZnSO4 −980.14

Aliphatic hydrocarbons{| class"wikitable sortable"

! Formula !! Name !! ΔfH⦵ /(kcal/mol) !! ΔfH⦵ /(kJ/mol)
Straight-chain
Methane >| −74.9
Ethane >| −83.7
Ethylene >| 52.5
Acetylene >| 226.8
Propane >| −104.6
n-Butane >| −125.5
n-Pentane >| −146.9
n-Hexane >| −167.4
n-Heptane >| −187.9
n-Octane >| −208.4
n-Nonane >| −229.3
n-Decane >| −249.4
C4 Alkane branched isomers
Isobutane (methylpropane) >| −134.3
C5 Alkane branched isomers
Neopentane (dimethylpropane) >| −167.8
Isopentane (methylbutane) >| −154.4
C6 Alkane branched isomers
2,2-Dimethylbutane >| −186.2
2,3-Dimethylbutane >| −177.8
2-Methylpentane (isohexane) >| −174.9
3-Methylpentane >| −172.0
C7 Alkane branched isomers
2,2-Dimethylpentane >| −205.9
2,2,3-Trimethylbutane >| −205.0
3,3-Dimethylpentane >| −201.3
2,3-Dimethylpentane >| −197.9
2,4-Dimethylpentane >| −201.7
2-Methylhexane >| −194.6
3-Methylhexane >| −191.2
3-Ethylpentane >| −189.5
C8 Alkane branched isomers
2,3-Dimethylhexane >| −230.5
2,2,3,3-Tetramethylbutane >| −225.5
2,2-Dimethylhexane >| −224.7
2,2,4-Trimethylpentane (isooctane) >| −223.8
2,5-Dimethylhexane >| −222.6
2,2,3-Trimethylpentane >| −220.1
3,3-Dimethylhexane >| −220.1
2,4-Dimethylhexane >| −219.2
2,3,4-Trimethylpentane >| −217.1
2,3,3-Trimethylpentane >| −216.3
2-Methylheptane >| −215.5
3-Ethyl-3-Methylpentane >| −215.1
3,4-Dimethylhexane >| −213.0
3-Ethyl-2-Methylpentane >| −210.9
3-Methylheptane >| −252.5
4-Methylheptane >| ?
3-Ethylhexane >| ?
C9 Alkane branched isomers (selected)
2,2,4,4-Tetramethylpentane >| −241.8
2,2,3,3-Tetramethylpentane >| −237.2
2,2,3,4-Tetramethylpentane >| −236.8
2,3,3,4-Tetramethylpentane >| −236.0
3,3-Diethylpentane >| −233.0

Other organic compounds

{| class="wikitable sortable"! Species! Phase! Chemical formula! ΔfH⦵ /(kJ/mol)
Acetone >| Liquid C3H6O −248.4
| Benzene| Liquid| C6H6| 48.95
| Benzoic acid| Solid| C7H6O2| −385.2
| Carbon tetrachloride| Liquid| CCl4| −135.4
| Carbon tetrachloride| Gas| CCl4| −95.98
| Ethanol| Liquid| C2H5OH| −277.0
| Ethanol| Gas| C2H5OH| −235.3
| Glucose| Solid| C6H12O6| −1271
Isopropanol >| Gas C3H7OH −318.1
| Methanol (methyl alcohol)| Liquid| CH3OH| −238.4
| Methanol (methyl alcohol)| Gas| CH3OH| −201.0
| Methyl linoleate (Biodiesel)| Gas| C19H34O2| −356.3
Sucrose >| Solid C12H22O11 −2226.1
| Trichloromethane (Chloroform)| Liquid| CHCl3| −134.47
| Trichloromethane (Chloroform)| Gas| CHCl3| −103.18
| Vinyl chloride| Solid| C2H3Cl| −94.12

See also

References

{{reflist}}
  • BOOK, Zumdahl, Steven, 2009, Chemical Principles, 6th, 384–387, Houghton Mifflin, Boston. New York, 978-0-547-19626-8,

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