Enthalpy Changes Flashcards
the energy of a system
is the sum of the kinetic energy and potential energy of the particles in the system
what is enthalpy
measure of the potential energy of particles
what do reactions usually involve
enthalpy changes and so heat energy is taken in or given out to balance the enthalpy change
exothermic reactions
- give out energy to surroundings
- system loses energy so is deltaH negative
- products more stable than reactants
- often spontaneous reaction or small energy input needed
examples of exothermic reactions
combustion or respiration
endothermic reactions
- take in energy from surroundings
- system gains energy so deltaH positive
- reaction vessel feels cool
examples of endothermic reactions
photosynthesis or thermal decomposition
what does deltaH =
Hproducts - Hreactants
what is the activation energy
the minimum energy required for a reaction to take place
definition of standard enthalpy change of a reaction ΔrH θ
enthalpy change associated with a stated equation reacting under standard conditions
standard enthalpy change of combustion ΔcH θ
enthalpy change when 1 mike of a substance is burnt completely in o2 under standard conditions
standard enthalpy change of formation ΔfH θ
enthalpy change when 1 mole of a substance is formed when its constituent elements in their standard states under standard conditions
ΔfH θ for elements
zero
standard enthalpy change of atomisation ΔatH θ
enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard states under standard conditions
standard enthalpy of neutralisation ΔneutH θ
enthalpy change when 1 mole of water is formed according to the following equation under standard conditions
ΔneutH θ = H+ (aq) + OH- (aq) -> H2O (l)
standard state
usual physical state under standard conditions eg magnesium mg(s)
standard conditions
100kPa (1atm)
298k (25°c)
1 moldm-3 solution
thermochemical equation
a chemical equation with a ΔH value after it
for weak acids that are partially dissociated
ΔneutH θ is smaller as some energy is used up in breaking the O-H bond in order for some of the H+ ions to be available for neutralisation
calorimetry
measures heat produced or absorbed by a chemical equation
calorimetry equation
q=mC ΔT
where: q= the energy given out or taken in in joules
m= mass of whatever is being heated in g
c= specific heat capacity and is 4.18Jg-1k-1
hess’ law
the overall enthalpy change for converting given reactants into given products is the same whatever path is taken, as long as the initial conditions are the same and the final conditions are the same
average bond enthalpy
enthalpy change on breaking 1 mole of bonds in 1 mole of gaseous molecules forming uncharged products
always endothermic
why are bond enthalpies never completely accurate
- bond enthalpies quoted in data booklets are averaged for many different compounds so will not be precise for any particular compound
- bond energy is defined for gaseous compounds so won’t be accurate for solids or liquids
factors that affect bond enthalpy
- bond length - shorter bonds are stronger because of the shared e-’s closer to the positive nuclei hence cl-cl > br-br > i-i
