Standard enthalpy of formation

Last updated March 19, 2024

In chemistry and thermodynamics, 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 in their reference state, 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.^[1] There is no standard temperature. Its symbol is Δ_fH^⦵. The superscript Plimsoll on this symbol indicates that the process has occurred under standard conditions at the specified temperature (usually 25 °C or 298.15 K).

Standard states are defined for various types of substances. For a gas, it is the hypothetical state the gas would assume if it obeyed the ideal gas equation at a pressure of 1 bar. For a gaseous or solid solute present in a diluted ideal solution, the standard state is the hypothetical state of concentration of the solute of exactly one mole per liter (1 M) at a pressure of 1 bar extrapolated from infinite dilution. 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.

For elements that have multiple allotropes, the reference state usually is chosen to be 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.^[2]

For example, the standard enthalpy of formation of carbon dioxide is the enthalpy of the following reaction under the above conditions:

{\ce {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 reference 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 $Δ f H ⦵ 298 K$ .

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

For 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,

{\ce {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):

$H sub$ , the standard enthalpy of atomization (or sublimation) of solid lithium.
$IE Li$ , the first ionization energy of gaseous lithium.
$B(F-F)$ , the standard enthalpy of atomization (or bond energy) of fluorine gas.
$EA F$ , the electron affinity of a fluorine atom.
$U L$ , the lattice energy of lithium fluoride.

The sum of these enthalpies give the standard enthalpy of formation ( $Δ f H$ ) 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 to evaluate the lattice energy:^[3]

-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 measurable using bomb calorimetry. The standard enthalpy of formation is then determined using Hess's law. The combustion of methane:

{\ce {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 (CO₂) and water (H₂O):

{\ce {CH4 -> C + 2H2}}

{\ce {C + O2 -> CO2}}

{\ce {2H2 + O2 -> 2H2O}}

Applying Hess's law,

\Delta _{\text{comb}}H^{\ominus }({\text{CH}}_{4})=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{\text{O}})]-\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4}).

Solving for the standard of enthalpy of formation,

\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{\text{O}})]-\Delta _{\text{comb}}H^{\ominus }({\text{CH}}_{4}).

The value of $\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})$ 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:^[4]

\Delta _{\text{r}}H^{\ominus }=\sum \nu \Delta _{\text{f}}H^{\ominus }({\text{products}})-\sum \nu \Delta _{\text{f}}H^{\ominus }({\text{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, ${\ce {CH4 + 2O2 -> CO2 + 2H2O}}$ :

\Delta _{\text{r}}H^{\ominus }=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{}{\text{O}})]-\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4})+2\Delta _{\text{f}}H^{\ominus }({\text{O}}_{2})].

However ${\ce {O2}}$ is an element in its standard state, so that $\Delta _{\text{f}}H^{\ominus }({\text{O}}_{2})=0$ , and the heat of reaction is simplified to

\Delta _{\text{r}}H^{\ominus }=[\Delta _{\text{f}}H^{\ominus }({\text{CO}}_{2})+2\Delta _{\text{f}}H^{\ominus }({\text{H}}_{2}{}{\text{O}})]-\Delta _{\text{f}}H^{\ominus }({\text{CH}}_{4}),

which is the equation in the previous section for the enthalpy of combustion $\Delta _{\text{comb}}H^{\ominus }$ .

Key concepts for enthalpy calculations

When a reaction is reversed, the magnitude of ΔH stays the same, but the sign changes.
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.
The change in enthalpy for a reaction can be calculated from the enthalpies of formation of the reactants and the products
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 °C

Thermochemical properties of selected substances at 298.15 K and 1 atm

Inorganic substances

Species	Phase	Chemical formula	Δ_fH^⦵ /(kJ/mol)
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
Barium chloride	Solid	BaCl₂	−858.6
Barium carbonate	Solid	BaCO₃	−1216
Barium hydroxide	Solid	Ba(OH)₂	−944.7
Barium oxide	Solid	BaO	−548.1
Barium sulfate	Solid	BaSO₄	−1473.3
Beryllium	Solid	Be	0
Beryllium hydroxide	Solid	Be(OH)₂	−903
Beryllium oxide	Solid	BeO	−609.4
Boron trichloride	Solid	BCl₃	−402.96
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	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
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
Calcium	Solid	Ca	0
Calcium	Gas	Ca	178.2
Calcium(II) ion	Gas	Ca²⁺	1925.90
Calcium(II) ion	Aqueous	Ca²⁺	−542.7
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
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	CO₃^2–	−675.23
Monatomic chlorine	Gas	Cl	121.7
Chloride ion	Aqueous	Cl⁻	−167.2
Chlorine	Gas	Cl₂	0
Chromium	Solid	Cr	0
Copper	Solid	Cu	0
Copper(II) oxide	Solid	CuO	−155.2
Copper(II) sulfate	Aqueous	CuSO₄	−769.98
Fluorine	Gas	F₂	0
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	Solid	I₂	0
Iodine	Gas	I₂	62.438
Iodine	Aqueous	I₂	23
Iodide ion	Aqueous	I⁻	−55
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	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
Lithium fluoride	Solid	LiF	−616.93
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	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⁻ ₄	−543
Mercury(II) oxide (red)	Solid	HgO	−90.83
Mercury sulfide (red, cinnabar)	Solid	HgS	−58.2
Nitrogen	Gas	N₂	0
Ammonia (ammonium hydroxide)	Aqueous	NH₃ (NH₄OH)	−80.8
Ammonia	Gas	NH₃	−46.1
Ammonium nitrate	Solid	NH₄NO₃	−365.6
Ammonium chloride	Solid	NH₄Cl	−314.55
Nitrogen dioxide	Gas	NO₂	33.2
Hydrazine	Gas	N₂H₄	95.4
Hydrazine	Liquid	N₂H₄	50.6
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
Nitric acid	Aqueous	HNO₃	−207
Monatomic oxygen	Gas	O	249
Oxygen	Gas	O₂	0
Ozone	Gas	O₃	143
White phosphorus	Solid	P₄	0
Red phosphorus	Solid	P	−17.4^[5]
Black phosphorus	Solid	P	−39.3^[5]
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	−1505.5^[6]
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 nitrate	Solid	KNO₃	−494.5
Potassium perchlorate	Solid	KClO₄	−430.12
Silicon	Gas	Si	368.2
Silicon carbide	Solid	SiC	−74.4,^[7] −71.5^[8]
Silicon tetrachloride	Liquid	SiCl₄	−640.1
Silica (Quartz)	Solid	SiO₂	−910.86
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	Solid	Na	0
Sodium	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 (monoclinic)	Solid	S₈	0.3
Sulfur (rhombic)	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
Titanium	Gas	Ti	468
Titanium tetrachloride	Gas	TiCl₄	−763.2
Titanium tetrachloride	Liquid	TiCl₄	−804.2
Titanium dioxide	Solid	TiO₂	−944.7
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

Formula	Name	Δ_fH^⦵ /(kcal/mol)	Δ_fH^⦵ /(kJ/mol)
Straight-chain
CH₄	Methane	−17.9	−74.9
C₂H₆	Ethane	−20.0	−83.7
C₂H₄	Ethylene	12.5	52.5
C₂H₂	Acetylene	54.2	226.8
C₃H₈	Propane	−25.0	−104.6
C₄H₁₀	n-Butane	−30.0	−125.5
C₅H₁₂	n-Pentane	−35.1	−146.9
C₆H₁₄	n-Hexane	−40.0	−167.4
C₇H₁₆	n-Heptane	−44.9	−187.9
C₈H₁₈	n-Octane	−49.8	−208.4
C₉H₂₀	n-Nonane	−54.8	−229.3
C₁₀H₂₂	n-Decane	−59.6	−249.4
C₄ Alkane branched isomers
C₄H₁₀	Isobutane (methylpropane)	−32.1	−134.3
C₅ Alkane branched isomers
C₅H₁₂	Neopentane (dimethylpropane)	−40.1	−167.8
C₅H₁₂	Isopentane (methylbutane)	−36.9	−154.4
C₆ Alkane branched isomers
C₆H₁₄	2,2-Dimethylbutane	−44.5	−186.2
C₆H₁₄	2,3-Dimethylbutane	−42.5	−177.8
C₆H₁₄	2-Methylpentane (isohexane)	−41.8	−174.9
C₆H₁₄	3-Methylpentane	−41.1	−172.0
C₇ Alkane branched isomers
C₇H₁₆	2,2-Dimethylpentane	−49.2	−205.9
C₇H₁₆	2,2,3-Trimethylbutane	−49.0	−205.0
C₇H₁₆	3,3-Dimethylpentane	−48.1	−201.3
C₇H₁₆	2,3-Dimethylpentane	−47.3	−197.9
C₇H₁₆	2,4-Dimethylpentane	−48.2	−201.7
C₇H₁₆	2-Methylhexane	−46.5	−194.6
C₇H₁₆	3-Methylhexane	−45.7	−191.2
C₇H₁₆	3-Ethylpentane	−45.3	−189.5
C₈ Alkane branched isomers
C₈H₁₈	2,3-Dimethylhexane	−55.1	−230.5
C₈H₁₈	2,2,3,3-Tetramethylbutane	−53.9	−225.5
C₈H₁₈	2,2-Dimethylhexane	−53.7	−224.7
C₈H₁₈	2,2,4-Trimethylpentane (isooctane)	−53.5	−223.8
C₈H₁₈	2,5-Dimethylhexane	−53.2	−222.6
C₈H₁₈	2,2,3-Trimethylpentane	−52.6	−220.1
C₈H₁₈	3,3-Dimethylhexane	−52.6	−220.1
C₈H₁₈	2,4-Dimethylhexane	−52.4	−219.2
C₈H₁₈	2,3,4-Trimethylpentane	−51.9	−217.1
C₈H₁₈	2,3,3-Trimethylpentane	−51.7	−216.3
C₈H₁₈	2-Methylheptane	−51.5	−215.5
C₈H₁₈	3-Ethyl-3-Methylpentane	−51.4	−215.1
C₈H₁₈	3,4-Dimethylhexane	−50.9	−213.0
C₈H₁₈	3-Ethyl-2-Methylpentane	−50.4	−210.9
C₈H₁₈	3-Methylheptane	−60.3	−252.5
C₈H₁₈	4-Methylheptane	?	?
C₈H₁₈	3-Ethylhexane	?	?
C₉ Alkane branched isomers (selected)
C₉H₂₀	2,2,4,4-Tetramethylpentane	−57.8	−241.8
C₉H₂₀	2,2,3,3-Tetramethylpentane	−56.7	−237.2
C₉H₂₀	2,2,3,4-Tetramethylpentane	−56.6	−236.8
C₉H₂₀	2,3,3,4-Tetramethylpentane	−56.4	−236.0
C₉H₂₀	3,3-Diethylpentane	−55.7	−233.0

Other organic compounds

Species	Phase	Chemical formula	Δ_fH^⦵ /(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

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In chemistry, the standard molar entropy is the entropy content of one mole of pure substance at a standard state of pressure and any temperature of interest. These are often chosen to be the standard temperature and pressure.

A chemical equation is the symbolic representation of a chemical reaction in the form of symbols and chemical formulas. The reactant entities are given on the left-hand side and the product entities are on the right-hand side with a plus sign between the entities in both the reactants and the products, and an arrow that points towards the products to show the direction of the reaction. The chemical formulas may be symbolic, structural, or intermixed. The coefficients next to the symbols and formulas of entities are the absolute values of the stoichiometric numbers. The first chemical equation was diagrammed by Jean Beguin in 1615.

In chemistry and thermodynamics, the enthalpy of neutralization is the change in enthalpy that occurs when one equivalent of an acid and a base undergo a neutralization reaction to form water and a salt. It is a special case of the enthalpy of reaction. It is defined as the energy released with the formation of 1 mole of water. When a reaction is carried out under standard conditions at the temperature of 298 K and 1 atm of pressure and one mole of water is formed, the heat released by the reaction is called the standard enthalpy of neutralization.

In thermodynamics, the Gibbs free energy is a thermodynamic potential that can be used to calculate the maximum amount of work, other than pressure-volume work, that may be performed by a thermodynamically closed system at constant temperature and pressure. It also provides a necessary condition for processes such as chemical reactions that may occur under these conditions. The Gibbs free energy is expressed as

Hess’ law of constant heat summation, also known simply as Hess' law, is a relationship in physical chemistry named after Germain Hess, a Swiss-born Russian chemist and physician who published it in 1840. The law states that the total enthalpy change during the complete course of a chemical reaction is independent of the sequence of steps taken.

The standard enthalpy of reaction for a chemical reaction is the difference between total product and total reactant molar enthalpies, calculated for substances in their standard states. The value can be approximately interpreted in terms of the total of the chemical bond energies for bonds broken and bonds formed.

Chemical energy is the energy of chemical substances that is released when the substances undergo a chemical reaction and transform into other substances. Some examples of storage media of chemical energy include batteries, food, and gasoline. Breaking and re-making chemical bonds involves energy, which may be either absorbed by or evolved from a chemical system. If reactants with relatively weak electron-pair bonds convert to more strongly bonded products, energy is released. Therefore, relatively weakly bonded and unstable molecules store chemical energy.

The Born–Haber cycle is an approach to analyze reaction energies. It was named after two German scientists, Max Born and Fritz Haber, who developed it in 1919. It was also independently formulated by Kasimir Fajans and published concurrently in the same journal. The cycle is concerned with the formation of an ionic compound from the reaction of a metal with a halogen or other non-metallic element such as oxygen.

The equilibrium constant of a chemical reaction is the value of its reaction quotient at chemical equilibrium, a state approached by a dynamic chemical system after sufficient time has elapsed at which its composition has no measurable tendency towards further change. For a given set of reaction conditions, the equilibrium constant is independent of the initial analytical concentrations of the reactant and product species in the mixture. Thus, given the initial composition of a system, known equilibrium constant values can be used to determine the composition of the system at equilibrium. However, reaction parameters like temperature, solvent, and ionic strength may all influence the value of the equilibrium constant.

In chemical kinetics, a reaction rate constant or reaction rate coefficient is a proportionality constant which quantifies the rate and direction of a chemical reaction by relating it with the concentration of reactants.

In chemical thermodynamics, the reaction quotient (Q_r or just Q) is a dimensionless quantity that provides a measurement of the relative amounts of products and reactants present in a reaction mixture for a reaction with well-defined overall stoichiometry, at a particular point in time. Mathematically, it is defined as the ratio of the activities (or molar concentrations) of the product species over those of the reactant species involved in the chemical reaction, taking stoichiometric coefficients of the reaction into account as exponents of the concentrations. In equilibrium, the reaction quotient is constant over time and is equal to the equilibrium constant.

The Van 't Hoff equation relates the change in the equilibrium constant, $K eq$ , of a chemical reaction to the change in temperature, T, given the standard enthalpy change, $Δ r H ⊖$ , for the process. The subscript $means "reaction" and the superscript means "standard". It was proposed by Dutch chemist Jacobus Henricus van 't Hoff in 1884 in his book Études de Dynamique chimique .$

Thermodynamic databases contain information about thermodynamic properties for substances, the most important being enthalpy, entropy, and Gibbs free energy. Numerical values of these thermodynamic properties are collected as tables or are calculated from thermodynamic datafiles. Data is expressed as temperature-dependent values for one mole of substance at the standard pressure of 101.325 kPa, or 100 kPa. Both of these definitions for the standard condition for pressure are in use.

In chemistry, transition state theory (TST) explains the reaction rates of elementary chemical reactions. The theory assumes a special type of chemical equilibrium (quasi-equilibrium) between reactants and activated transition state complexes.

Equilibrium chemistry is concerned with systems in chemical equilibrium. The unifying principle is that the free energy of a system at equilibrium is the minimum possible, so that the slope of the free energy with respect to the reaction coordinate is zero. This principle, applied to mixtures at equilibrium provides a definition of an equilibrium constant. Applications include acid–base, host–guest, metal–complex, solubility, partition, chromatography and redox equilibria.

References

↑ IUPAC , Compendium of Chemical Terminology , 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) " standard pressure ". doi : 10.1351/goldbook.S05921
↑ Oxtoby, David W; Pat Gillis, H; Campion, Alan (2011). Principles of Modern Chemistry. p. 547. ISBN 978-0-8400-4931-5.
↑ Moore, Stanitski, and Jurs. Chemistry: The Molecular Science. 3rd edition. 2008. ISBN 0-495-10521-X. pages 320–321.
↑ "Enthalpies of Reaction". www.science.uwaterloo.ca. Archived from the original on 25 October 2017. Retrieved 2 May 2018.
1 2 Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 392. ISBN 978-0-13-039913-7.
↑ Green, D.W., ed. (2007). Perry's Chemical Engineers' Handbook (8th ed.). Mcgraw-Hill. pp. 2–191. ISBN 9780071422949.
↑ Kleykamp, H. (1998). "Gibbs Energy of Formation of SiC: A contribution to the Thermodynamic Stability of the Modifications". Berichte der Bunsengesellschaft für physikalische Chemie. 102 (9): 1231–1234. doi:10.1002/bbpc.19981020928.
↑ "Silicon Carbide, Alpha (SiC)". March 1967. Retrieved 5 February 2019.

Zumdahl, Steven (2009). Chemical Principles (6th ed.). Boston. New York: Houghton Mifflin. pp. 384–387. ISBN 978-0-547-19626-8.

External links

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[1] IUPAC , Compendium of Chemical Terminology , 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) " standard pressure ". doi : 10.1351/goldbook.S05921

[2] Oxtoby, David W; Pat Gillis, H; Campion, Alan (2011). Principles of Modern Chemistry. p. 547. ISBN 978-0-8400-4931-5.

[3] Moore, Stanitski, and Jurs. Chemistry: The Molecular Science. 3rd edition. 2008. ISBN 0-495-10521-X. pages 320–321.

[4] "Enthalpies of Reaction". www.science.uwaterloo.ca. Archived from the original on 25 October 2017. Retrieved 2 May 2018.

[HS-5] 1 2 Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 392. ISBN 978-0-13-039913-7.

[6] Green, D.W., ed. (2007). Perry's Chemical Engineers' Handbook (8th ed.). Mcgraw-Hill. pp. 2–191. ISBN 9780071422949.

[7] Kleykamp, H. (1998). "Gibbs Energy of Formation of SiC: A contribution to the Thermodynamic Stability of the Modifications". Berichte der Bunsengesellschaft für physikalische Chemie. 102 (9): 1231–1234. doi:10.1002/bbpc.19981020928.

[8] "Silicon Carbide, Alpha (SiC)". March 1967. Retrieved 5 February 2019.

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