energetics Flashcards

1
Q

exothermic reactions

A
  • heat energy is released to surroundings

- products will have less energy than the reactants

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2
Q

endothermic reactions

A
  • heat energy is absorbed from the surroundings

- products will have more energy than the reactants

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3
Q

enthalpy change

A
  • the amount of heat given out/taken in during a reaction carried out at constant pressure
  • symbol: ∆H
  • units: kJmol-1
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4
Q

enthalpy of combustion

A
  • the heat released during the complete combustion of one mole of that substance
  • always a -ve value
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5
Q

bond energies

A
  • a measure of the average bond strength for a particular covalent bond
  • measured in compounds in the gaseous state so the values only apply in calculations where all reactants and all products are in the gaseous state
  • ΔH = [sum of energy breaking bonds] + [sum of energy forming bonds]
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6
Q

bond breaking

A
  • to break a covalent bond between 2 atoms, energy must be supplied
  • bond breaking is an endothermic process (+ve value)
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7
Q

bond forming

A
  • when a covalent bond between 2 atoms is made, energy is released
  • bond forming is an exothermic process (-ve value)
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8
Q

limitations of bond energies

A
  • bond energies are average values so values calculated using them are not specific to the molecules used
  • thus, the actual value is likely to have a difference from the calculated value of over 10%.
  • calculating the enthalpy change can only be done for reactions which take place entirely in the gaseous state
  • to obtain more accurate values for the ΔH, calorimetry experiments can be used
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9
Q

energy formula

A

Eh = cmΔT

  • Eh = energy gained by water (in kJ)
  • c = specific heat capacity of water (4.2)
  • m = mass of water (in kg) 1l of water = 1kg of water
  • ΔT = change in temp of water

ΔH =Eh/n (n=no. of moles), must have a +/- sign

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10
Q

calorimetry in aqueous solutions

A
  • a measured volume of a known conc of solution is placed in a polystyrene cup (an insulator so less heat is lost/gained to/from the surroundings)
  • the temp of the solution is taken
  • a measured volume of a known conc of another solution is added
  • the mixture is stirred and the highest temp rise is found. -ΔT, can then be calculated and used in the formula
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11
Q

assumptions made in energy calculations

A
  • at room temp and pressure, the density of water is 1.00g cm−3 and the density of a solution is assumed to be the same even though it will not be quite the same
  • the specific heat capacity is particular to the substance being heated. As most of an aqueous mixture is water, it can be assumed that the specific heat capacity of a solution is the same as that for pure water.
  • as water is the major component, it can be assumed that only the water has absorbed any heat and that the calorimeter doesn’t absorb any energy
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12
Q

sources of error in calorimetry calculations

A
  • incomplete combustion: especially true when larger molecules are used as the fuel
  • heat loss to surroundings is more likely in an open setting: heat can be lost before it reaches the calorimeter, and if the calorimeter not insulated
  • heat is lost to the calorimeter: since it is a conductor, some heat will be absorbed by the calorimeter, (this error can be removed by calculating the energy change of the calorimeter and the water and adding these values.
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