Standard enthalpy changes and Hess’ Law
d. recall the definition of standard enthalpy changes of reaction,
formation, combustion, neutralization and atomization and use
experimental data to calculate energy transferred in a reaction
and hence the enthalpy change of the reaction. This will be
limited to experiments where substances are mixed in an
insulated container, and combustion experiments
e. recall Hess’s Law and apply it to calculating enthalpy changes of
reaction from data provided, selected from a table of data or
obtained from experiments and understand why standard data is
necessary to carry out calculations of this type
g. demonstrate an understanding of the terms bond enthalpy and
mean bond enthalpy, and use bond enthalpies in Hess cycle
calculations and recognise their limitations. Understand that
bond enthalpy data gives some indication about which bond will
break first in a reaction, how easy or difficult it is and therefore
how rapidly a reaction will take place at room temperature.
Home Learning Task – Read pp43-45
and answer the questions therein.
Crowe2010
Connector - Calculations using the results from an
experimental calorimeter.
• 100 cm3 of water (100g) was
measured into the calorimeter.
• The spirit burner contained the fuel
ethanol C2H5OH and weighed 18.62g
at the start.
• After burning it weighed 17.14g and
the temperature of the water rose
from 18 to 89oC.
Method
1. Use q = mcΔT to calculate the energy released. ( c = 4.2 J g-1 K-1 )
2. q represents the heat released from the mass of ethanol burnt.
3. Calculate the energy released per g.
4. Calculate the energy released per mole
5. Now give the Enthalpy of Combustion of ethanol (ΔHc) in kJ mol-1.
• The temperature rise = 89 - 18 = 71oC (exothermic).
• Mass of fuel burned = 18.62-17.14 = 1.48g.
• Heat absorbed by the water
= mass of water x SHCwater x temperature
• = 100 x 4.2 x 71 = 29820 J (for 1.48g)
• heat energy released per g = energy supplied in J / mass of fuel
burned in g
• heat energy released on burning = 29820 / 1.48 = 20149 J/g of
C2H5OH
• this energy change can be also expressed on a molar basis.
• Relative atomic masses Ar: C = 12, H = 1, O = 16
• Mr(C2H5OH) = (2 x 12) + (1 x 5) + 16 + 16 = 46, so 1 mole = 46g
• Heat released (given out) by 1 mole of C2H5OH = 46 x 20149
= 926854 J/mole or 927 kJ/mol (3 sf)
• Enthalpy of combustion of ethanol = ΔHc = -927 kJmol-1
Standard enthalpy change of …..
Formation ΔHθf is the enthalpy change when 1 mole of a compound
is made from its elements in their standard states, under standard
conditions.
Combustion ΔHθc is the enthalpy change when 1 mole of a
substance burns completely in oxygen, under standard conditions.
Atomisation ΔHθat is the enthalpy change when 1 mole of gaseous
atoms is made from the element in its standard state, under standard
conditions.
Reaction ΔHθr is the enthalpy change when molar quantities of
reactants as stated in an equation react under standard conditions.
Neutralisation ΔHθn is the energy released when unit molar
quantities of acids and alkalis completely neutralise each other at
298K.
Note: Make sure that you know these definitions
Hess’s Law
The enthalpy change in turning any reactants into a set
of products is the same no matter what route we take.
C
b
A
c
a
d
D
e
B
a=b+c=d-e
Note: arrows in same direction as “a” are added,
Those in the opposite direction are subtracted.
Using enthalpy cycles
Some enthalpy changes cannot be measured directly:
e.g. The enthalpy of formation of ethanol
Standard Enthalpy of Formation, ΔHθf is the enthalpy change
when 1 mole of a compound is made from its elements in their
standard states, under standard conditions.
Write the equation for the formation of ethanol from its elements.
Why can’t ethanol be prepared in this way?
Using enthalpy cycles to calculate enthalpy changes
that cannot be measured directly
Although we cannot directly measure the enthalpy of
formation of ethanol, it can be calculated by using
other measurable enthalpy changes.
Where:
ΔH1 = 2x (enthalpy of combustion of graphite) + 3x (enthalpy of combustion of hydrogen)
ΔH3 = enthalpy of combustion of ethanol
These enthalpies of combustion can be measured, and so ΔH1 can be
calculated since, applying Hess’ Law:
ΔH1 = ΔH2 – ΔH3
Enthalpy of Formation of Ethanol
Use the information above to calculate
the enthalpy of formation of ethanol.
Calculating the heat of formation of ethanol
From the previous slide:
ΔH1 = ΔH2 – ΔH3
2C + 3H2 + ½ O2
C2H5OH
So:
ΔH2 = 2ΔHcθ[C, graphite(s)] + 3ΔHcθ[H2,(g)]
ΔH3 = ΔHcθ[C2H5OH(l)]
ΔH1 =
ΔH2
–
= [2(-393.5) + 3(-285.8)] –
= [-787 – 857.4] + 1367.3
= -1644.4 + 1367.3
= -277.4kJmol-1
ΔH3
[-1367.3]
Enthalpy Level Diagram
This is another way to show the enthalpy changes in an energy cycle:
Practical 1.6 Finding an enthalpy change that
cannot be measured directly
The standard enthalpy of formation of calcium carbonate, ΔHf [CaCO3] is
defined as the enthalpy change when one mole of calcium carbonate is
formed from its elements in their standard states, under standard
conditions.
It is given by the following equation:
Ca(s)
+ C(graphite)
+ 1½O2(g)
➔
CaCO3(s)
The reaction, as written, is unlikely to take place in the laboratory,
and so ΔHf [CaCO3] cannot be determined directly.
However, the enthalpy changes when calcium metal and calcium
carbonate react with hydrochloric acid, are measurable, and
applying the principles of Hess’s law, ΔHf [CaCO3] can then be
determined.
Practical 1.6 –
Finding an enthalpy change that cannot be
measured directly
Practical 1.6 Finding an enthalpy change that
cannot be measured directly
Useful information
 Usually one reagent is in excess to ensure a complete reaction
• So calculations should be based on the fully reacted reagent.
 Certain assumptions are made during the calculation
• The density of the solution and its specific heat capacity is
that of water.
• That no heat is lost to the surroundings.
 The main source of error in these experiments is
• heat loss to the surroundings – (atmosphere & equipment)
 Other sources of error include:
• incorrect measurements
• solution concentrations
• mass of reactants
Worksheets Exp 1.6 from teacher’s guide