Chemistry - energy - Revision Cards in A Level and IB Chemistry

Enthalpy

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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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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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 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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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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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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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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