9 - Enthalpy Flashcards
enthalpy H
the measure of heat energy in a chemical system
-> cannot be measured but enthalpy changes can
enthalpy change ΔH
ΔH= H(products) - H(reactants)
+ = products have more energy thanreactnts
conservation of energy
energy cannot be created or destroyed
where is heat transferred between in a reaction
the system and surroundings
the system- reactants and products
surroundings- apparatus, the lab
universe= system + surrounding
exothermic
ΔH is negative
chemical system loses energy
surroundings gain energy
temp of surroundings increase
endothermic
ΔH is positive
chemical system gains energy
surroundings loses energy
temp of surroundings decreases
activation energy
minimum energy required for a reaction to take place
-> the energy required to break bonds acts as an energy barrier to the reaction
standard conditions
100 kPa
298 K (25’c)
1 moldm-3
standard state = the physical state of a substance under standard conditions
standard enthalpy change of reaction ΔrH°
enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions and states
ΔrH° of Mg and O2
Mg (s) + 1/2O2(g) -> MgO (s) ΔrH°= -602
or
2Mg (s) + O2(g) -> 2MgO (s) ΔrH°= -602 x 2= -1204
standard enthalpy change of formation ΔfH°
enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions and states
NOTE; writing this formula must show that one mole of the substance is produced
ΔrH° of MgO(s)
Mg (s) + 1/2O2(g) -> MgO (s) ΔfH°= -602
standard enthalpy change of combustion ΔcH°
the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions and states
ΔcH° of butane
C4H10 + 13/2 O2 -> 4CO2 + 5H20
standard enthalpy change of neutralisation ΔneutH°
enthalpy change when an acid and a base react to form one mole of H2O(l) under standard conditions and states
- this value is the same for all neutralisation reactions
Calculating an energy change - the mass of the surroundings
m = of the mass that changes temperature in grams
NOTE - not the mass of the reactants
Calculating an energy change - specific heat capacity
the energy required to raise the temperature of 1g of a substance by 1K
-> in most experiments we will be measuring c for water as this is usually the substance that will have a temp change
c= 4.18 Jg-1K-1
Calculating an energy change - temperature change
ΔT=T(final) - T(initial)
formula for heat energy
q = mcΔT
m in grams
ΔT - doesn’t matter as long final and initial are the same units
unit of q usually calculated is J as C= Jg-1K-1
J-> KJ divide by 1000
method to determine ΔcH° of methanol
CH3OH(l) + 3/2O2(g) -> CO2(g) + H2O(l)
Materials and Apparatus:
Methanol (fuel)
Spirit burner- wick and methanol
Thermometer
Measuring cylinder
clamp
Water
Calorimeter (insulated container)
Experimental Setup:
Set up the spirit burner on a stable surface.
Measure the initial mass of the spirit burner (with methanol) and record it.
Measure a known volume of water using the measuring cylinder and pour it into the calorimeter.
Measure the initial temperature of the water in the calorimeter and record it.
Combustion of Methanol:
Ignite the spirit burner containing methanol and let it burn
Stir the methanol with the thermometer
Temperature Measurement:
After the combustion, measure the final temperature of the water in the calorimeter and record it.
Measure the weight of the spirit burner
use q=mcΔT to calculate energy change of the water
then calculate the mole of CH3OH burnt
calculate ΔcH
->ΔcH is the enthalpy change of the complete combustion of 1 mole of methanol
-> in the experiment x mol of methanol transfers y KJ of energy to the water
-> the water gained y KJ of energy from the combustion of x of methanol
->therefore 1 mole of methanol has lost y/x kj of energy
-> so ΔcH of methanol is -y/x
NOTE= ΔcH is negative as system has lost energy due to the combustion of methanol (exothermic0
reasons why the experimental ΔcH is not accurate
- heat loss to the surroundings other than water. such as the beaker and the air surrounding the flame
- incomplete combustion of methanol- black layer of soot
- evaporation of methanol from the wick - the burner must be weight soon as possible after extinguishing the flame otherwise methanol may have evaporated
- non-standard conditions
all but the last reasons would lead to a value of ΔcH that is less exothermic
average bond enthalpy
the average energy required to break a particular type of bond in a gaseous molecule, averaged over a range of different compounds
bond enthalpies are always exo or end?
endothermic
-> bond enthalpies is breaking bonds which requires energy
-> always a positive value
to break bonds energy is __________
required
to form bonds energy is __________
released
limitations of average bond enthalpies
actual bond enthalpies can vary depending on its chemical environment
calculate ΔrH°
Σ(bond enthalpies in reactants) - Σ(bond enthalpies in products)
calculations using average ______ _______ need all species to be _________ molecules. If you produce H2O(g) rather than H2O(l), this means that the calculate _____ is not a ______ enthalpy change.
what do you do to overcome this?
bond
enthalpies
gaseous
ΔrH°
standard
work out the standard enthalpy change but consider the enthalpy change for H2O(g) condensing into H2O(l)
Hess’ Law
is a reaction can take place by two routes, and the starting and finishing conditions are the same, the total enthalpy change is the same for each route
why is Hess’ law helpful
enthalpy changes of reactions are often hard to determine directly
