词条 | Exothermic reaction |
释义 |
An exothermic reaction is a chemical reaction that releases energy through light or heat. It is the opposite of an endothermic reaction.[1] Expressed in a chemical equation: reactants → products + energy. Exothermic Reaction means "exo" (derived from the greek word: "έξω", literally translated to "out") meaning releases and "thermic" means heat. So the reaction in which there is release of heat with or without light is called exothermic reaction. OverviewAn exothermic reaction is a chemical reaction that releases heat. It gives net energy to its surroundings. That is, the energy needed to initiate the reaction is less than the energy released.[2] When the medium in which the reaction is taking place collects heat, the reaction is exothermic. When using a calorimeter, the total amount of heat that flows into (or through) the calorimeter is the negative of the net change in energy of the system. The absolute amount of energy in a chemical system is difficult to measure or calculate. The enthalpy change, Δ{{var|H}}, of a chemical reaction is much easier to work with. The enthalpy change equals the change in internal energy of the system plus the work needed to change the volume of the system against constant ambient pressure. A bomb calorimeter is very suitable for measuring the energy change, Δ{{var|H}}, of a combustion reaction. Measured and calculated ΔH values are related to bond energies by: Δ{{var|H}} = (energy used in forming product bonds) − (energy released in breaking reactant bonds) In an exothermic reaction, by definition, the enthalpy change has a negative value: Δ{{var|H}} < 0 since a larger value (the energy released in the reaction) is subtracted from a smaller value (the energy used for the reaction). For example, when hydrogen burns: 2H2 (g) + O2 (g) → 2H2O (g) Δ{{var|H}} = −483.6 kJ/mol of O2 [3] In an adiabatic system, the temperature raise due to enthalpy change can be expressed as {{math|1= −Δ{{var|H}}298.15 K = {{intmath|int|{{var|T}}0|{{var|T}}1}}{{var|C}}p, pd{{var|T}} + {{intmath|int|298 K|{{var|T}}0}}({{var|C}}p, p−{{var|C}}p, r)d{{var|T}}}}[4] where Δ{{var|H}}298.15 K is the standard enthalpy of reaction at 298 K, {{var|T}}0 and {{var|T}}1 are the initial and final temperature of the system, respectively, and {{var|C}}p,p and {{var|C}}p,r are the heat capacities of the product and reactant, respectively. Assuming the heat capacity of the system remains as a constant value {{var|C}}p,p={{var|C}}p,r={{var|C}}p, the change of temperature Δ{{var|T}}={{var|T}}1−{{var|T}}0 can be expressed as {{math|1= −Δ{{var|H}}298.15 K = {{intmath|int|{{var|T}}0|{{var|T}}0+Δ{{var|T}}}}{{var|C}}p, pd{{var|T}} = Δ{{var|TC}}p, p}}[4] The most commonly available hand warmers make use of the oxidation of iron to achieve an exothermic reaction: 4Fe (s) + 3O2 (g) → 2Fe2O3 (s). Examples of exothermic reactions
Other points to think about
MeasurementHeat production or absorption in either a physical process or chemical reaction is measured using calorimetry. One common laboratory instrument is the reaction calorimeter, where the heat flow into or from the reaction vessel is monitored. The technique can be used to follow chemical reactions as well as physical processes such as crystallization and dissolution. Energy released is measured in Joule per mole. The reaction has a negative ΔH(heat change) value due to heat loss. e.g.: -123 J/mol See also
References1. ^Article written by Anne Marie Helmenstine, Ph.D on exothermic and endothermic reactions {{cite web |url=http://chemistry.about.com/cs/generalchemistry/a/aa051903a.htm |title=Archived copy |accessdate=2016-04-05 |deadurl=no |archiveurl=https://web.archive.org/web/20160318203441/http://chemistry.about.com/cs/generalchemistry/a/aa051903a.htm |archivedate=2016-03-18 |df= }} 2. ^{{cite web|url=http://chemistry.about.com/cs/generalchemistry/a/aa051903a.htm|title=Endothermic and Exothermic Reactions|access-date=5 April 2016|website=About Chemistry|date=3 February 2013|author=|deadurl=no|archiveurl=https://web.archive.org/web/20160318203441/http://chemistry.about.com/cs/generalchemistry/a/aa051903a.htm|archivedate=18 March 2016|df=}} 3. ^{{cite web |url=http://chemistry.osu.edu/~woodward/ch121/ch5_enthalpy.htm |title=Archived copy |accessdate=2013-07-20 |deadurl=no |archiveurl=https://web.archive.org/web/20130708030319/http://chemistry.osu.edu/~woodward/ch121/ch5_enthalpy.htm |archivedate=2013-07-08 |df= }} 4. ^1 {{cite journal|last1=Yin|first1=Xi|last2=Wu|first2=Jianbo|last3=Li|first3=Panpan|last4=Shi|first4=Miao|last5=Yang|first5=Hong|title=Self-Heating Approach to the Fast Production of Uniform Metal Nanostructures|journal=ChemNanoMat|date=January 2016|volume=2|issue=1|pages=37–41|doi=10.1002/cnma.201500123}} External links{{Portal bar|Chemistry}} 1 : Thermochemistry |
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