- 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.
- 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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- 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 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.
- STANARD STATES.
- 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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- 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
- Chemical system loses energy.
- Temp rise in surroundings.
- Delta H = -/ve.
- 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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- 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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