Chemistry - energy

OCR - Unit 2, Module 3

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  • Chemical reaction are accompanied by enthalpy changes.
  • Delta H = negative - exothermic reaction.
  • Delta H = endothermic reaction.
  • Endothermic reactions.
  • Needs an imput of energy.
  • Eg: Thermal decomposition of limestone.
  • Photosynthesis.
  • Oxidation of fuels is the most exploited reaction.
  • This is a exothermic reaction.
  • Products have less enthalpy than the reactants and so the excess energy is released as heat.
  • Reaction of  carbohdrates is a important exothermic reation.
  • E.g - Respiration.
  • Respiration takes place over a series of steps.
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Enthalpy profile diagrams

  • Definition - Diagram for a reaction to compare the enthalpy of reactants with the enthalpy of the products.
  • Exothermic reactions - Enthalpy porducts < enthalpy of reactants.
  • Enthalpy change = negative.
  • Heat is released to the surroundings.
  • Reacting chemicals lose heat energy.
  • Heat lost by chemicals = gained by surroundings.
  • Endothermic reactions.
  • Enthalpy of products > Enthalpy of reactants.
  • Delta H = +/ve..
  • Heat taken in by the surroundings.
  • Reacting chemicals gain emergy.
  • Heat gained by chemicals - lost by the surroundings.
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Activation Energy

  • Definition - The minimum energy required to start a reaction by the breaking of bonds.
  • Ea- activation energy.
  • If there was ni Ea, then exothermic reaction would not exist because fuels would just combust spontanoeusly.
  • Exothermic reactions.
  • Still has to have an imput of energy to break the first bond and kick start the reaction.
  • Once a exothermic reaction begins then the Ea is regenerated and the reaction = self sufficient.
  • Ea - given by a spark or  by lighting reactants.
  • Products have lower energy than the reactants.
  • ENDOTHERMIC reactions - Reactants have lower energy compared to the products...
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Standard Conditions

  • Standard condition =
  • Temp - 25 degrees or 298 Kelvin.
  • Pressure - 100 KPa or 1 atmosphere pressure.
  • Concentration of 1 mol dm (cubed) for reactions within an aqueous solution.
  • Physical states of substances are under standard conditions... All reactants and products being in their natural state.
  • E.g. the natural state of Hydrogen = H2.
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Standard Enthalpy definitions

  • Standard enthalpy change of Combustion - Enthalpy change that takes place when 1 mole of substance reacts completely with oxygen under standard conditions. All reactants and products being in their standard state.
  • Standard enthalpy of Formation - Enthalpy change that takes place when one mole of a compound is formed from its constituent elements in their standard states . under standard conditions. 
  • Standard enthalpy of formation is problematic for an element.
  • Forming H2 for H2 requires no chemical change.
  • So all elements have a standard enthalpy of 0Kjmol -1.
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Enthalpy changes

  • Can't determine enthalpy of products, but can determine heat exchange with the surroundings.
  • Exothermic reaction - heat produced is trapped by a clorimetre increasing the temp of the solution.
  • Heat lost in a chemical is gained by the surroundings.
  • Heat gained by a chemical - lost by the surroundings.
  • Provided we know what happens in surroundings then we would know what happens in the chemical system.
  • Temp increase on the thermometre - then the reaction = exothermic.
  • Temp decreases on the thermometre - then the reaction = endothermic.
  • Endothermic - heat required is removed from the solution - decreasing the temp of the solution  .
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Endo / Exo

  • Exdothermic
  • Chemical system loses energy.
  • Temp rise in surroundings.
  • Delta H = -/ve.
  • Endothermic
  • Chemical system gains energy.
  • Surroundings lose heat energy.
  • Delta H = +/ve.
  • MC(delta)T ... (Remember MKAT).
  • C = specific heat capacity.
  • The energy required to heat 1g of a substance by 1 degree.
  • For H2O this is 4.18 kjmol-1 or 4.2.
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Bond Enthalpies

  • Definition - Average Bond Enthalpy : Average enthalpy that takes place when breaking by homolytics fission of 1 mol of a given type of a bond in a molecule of gaseous species.
  • Definition - Bond Enthalpy : Enthalpy change that takes place when breaking by homolytic fission of 1 mole of a given bond in a molecule of gaseous species.
  • Chemical bonds - Storehouse for chemical energies, strength of chemical bond determined by the bond enthalpy.
  •  A) Energy first needed to make the reactants.
  •  B) Bond breaking requires energy - endothermic.
  •  C) Energy released as new bonds are formed.
  •  D) Bond making - exothermic reaction.
  • Average bond enthalpy - chemical reaction - bond breaking followed by bond making.
  • Energy is needed to break bonds - = an endothermic process.
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How do you know if the overall bond enthalpy is po

  • What decides if the overall reaction is +/ve or -/ve ??.
  • Bond Enthalpy - relative strength of bonds being broken and bonds being made..
  • Exothermic - bonds formed = stronger than the bonds broken.
  • Endothermic - formed bonds = weaker and the bonds broken stronger.
  • How to use bond enthalpies?.
  • Sum of bond enthalpies of bonds being broken - sum of bonde enthalpies of bonds being formed.
  • ENDOTHERMIC change - Delta H = +/ve.
  • Temperature of the surroundings goes down.
  • EXOTHERMIC change - Delta H = -/ve.
  • Then temperature of the surrounding increase.
  • If the smae bond is formed as the bond broken ... then the enthalpy change will be of the same magnitude, but the signs go the opposite way.
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Standard Enthalpy change of combustion

  • Diagram shows alternative routes between reactants and products.
  • Shows indirect determination of enthalpy change from other known enthalpy cycle using hess's law.
  • Hess's law.
  • - If a reaction can take place more than 1 route and the intial and the final condition are the same then the total enthalpy = the same for each route.
  • The reason why you can't measure standard enthalpy of combustion directly is because the reaction may have a high activation energy, secondly there might be a slow reaction rate and lastly there may be more than one reaction taking place...
  • Using hess's law :
  • Route 1 ----A-----> Product
  • Route 2 ----B-----> Intermediate ----C-----> Product
  • So therefore, A = B +C
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