ENDOTHERMIC AND
EXOTHERMIC REACTIONS
ENERGY
AND
REACTIONS
HEAT OF
REACTION
Heat of combustion
Heat of neutralization
BOND
ENERGY
HEAT OF
FORMATION
Hess’s law
OUR PLANET IS WARMING.
WHY?
EXOTHERMIC AND
ENDOTHERMIC REACTIONS
The study of the heat changes that accompany
chemical reactions is called thermochemistry
Burning fuels in air always produces heat, “burning” food
inside our bodies supplies us with the energy we need
Combustion of methane is
one example of a reaction
that produces heat
Any chemical reaction that produces heat is called an
exothermic reaction
EXOTHERMIC AND
ENDOTHERMIC REACTIONS
We note a drop in temperature if we hold a beaker of water at
room temperature and add some ammonium chloride (NH4Cl) to
it: in this case heat is taken in from the hand.
NH4Cl
Any chemical reaction that takes in heat from the surroundings
is called an endothermic reaction
HEAT OF REACTION
Hydrogen is a fuel which burns in oxygen to form water:
H2(g) + ½ O2(g)
H2O(g) D H = - 242 kJ
The symbol D H (delta H) indicates the variation of hentalpy: the heat change
taking place when a chemical reaction occurs at a constant pressure and a constant
temperature. This equation states that when one mole of hydrogen reacts with half
a mole of oxygen to form one mole of steam 242 kJ of heat are released.
A negative value of D H indicates an exothermic reaction
A positive value of D H indicates an endothermic reaction
The D H value above is referred to as the molar heat of reaction. If the number of
moles in the balanced equation are changed, then the heat of reaction also
changes:
2 H2(g) + O2(g)
2H2O(g) D H = - 484 kJ
Heat of reaction is the heat change when the moles of reactants,
indicated in the balanced equation, react completely.
HEAT OF COMBUSTION
A reaction in which a substance is burned is called a combustion reaction.
Consider the reaction in which carbon is burned in oxygen to form carbon
dioxide :
C(s) + O2(g)
CO2(g)
D H = - 393 kJ
The heat given out in the reaction is called the heat of
combustion.
Heat of combustion of a substance is the heat change
when one mole of the substance is completely burned
in excess oxygen.
HEAT OF COMBUSTION
• The phrase “is completely burned in excess oxygen” means that
the product of reaction must be carbon dioxide and not carbon
monoxide as it occurs in an incomplete combustion:
C(s) + ½ O2(g)
CO(g)
D H = - 111 kJ
The value of D H is not the heat of combustion of carbon.
•
Another important phrase is “one mole”.
Consider the thermochemical reaction in which butane is burned:
2 C4H10(g) + 13 O2(g)
kJ
8CO2(g) + 10H2O(g)
D H = - 5720
The heat of combustion of butane is not - 5720 kJ,
In fact the heat of combustion involves one mole of butane
being burned; therefore the correct value is
–5720/2 = -2860 kJ mol-1
HEAT OF COMBUSTION
Heat of combustion is accurately measured using a bomb calorimeter.
This instrument consists of a steel container (the bomb), with a screw-on cap. The
sample whose heat of combustion is to be measured is placed in a crucible inside
the bomb. The bomb is placed in a container of water (the calorimeter), oxygen is
pumped inside and the sample is ignited by electric wires. By measuring the
increase in temperature of the water it is possible to measure the heat of
combustion of the fuel.
HEAT OF COMBUSTION
In some cases, in industry, chemists measure the heat
given out by a fuel (or food) in terms of the mass of fuel
rather than one mole of substance, in order to compare the
efficiency of various fuels and to state if the fuel is of good
quality.
For example, if you read the side of a box of cornflakes, you will
find that 100g of cornflakes provides about 1550 kJ of energy.
Chemists use the term: kilogram calorific value.
The kilogram calorific value of a fuel is the heat, the energy,
produced when 1 kg of the fuel is completely burned in oxygen.
HEAT OF COMBUSTION
Some examples of the heat of combustion of various fuels
Fuel
D H (kJ mol-1) at 25 °C
Kilogram caloric
value (kJ kg-1)
Formula Heat of combustion
Methane
CH4(g)
-890
55.625
Ethane
Propane
Butane
Hexane
C2H6(g)
-1560
52.000
C3H8(g)
-2230
50.454
C4H10(g)
-2877
49.603
C6H14(l)
-4194
48.767
Ethanol
C2H5OH(l)
-1371
29.804
Carbon
Hydrogen
C (g)
-393
32.750
H2(g)
-286
143.000
HEAT OF NEUTRALISATION
Neutralisation: a reaction between an acid and a base to form a salt and water:
HCl +
NaOH
NaCl
+
H2O
Strong acids, strong bases and salts are fully dissociated in water, thus the essential
reaction taking place ingnores spectator ions Na+ and Cl-:
H+ + Cl- + Na+ + OHH+ + OH-
Na+ + Cl- + H2O
H2O
D H = - 57.1 kJ
Neutralisation reactions are nearly always exothermic.
The heat of neutralisation is the heat change when one mole of H+ ions
from an acid reacts with one mole of OH- ions from a base.
The heat of reaction in a reaction of neutralisation depends on the numbers of moles of
water formed, therefore the heat of reaction between sulphuric acid and sodium
hydroxide is double since two moles of water are formed:
H2SO4 + 2NaOH
Na2SO4 + 2H2O
D H = - 2 · 57 kJ = - 114 kJ
BOND ENERGY
What determines the value of D H for a particular reaction? Consider
the combustion of methane:
CH4(s) + 2O2(g)
CO2(g) + 2H2O(g)
As known, when a
chemical reaction occurs,
the bonds of reactants are
broken and the bonds of
products are formed.
Energy is required to break
bonds and energy is
released when bonds are
formed.
D H = - 890 kJ
Reactants
Products
Energy absorbed
to break bonds of
reactants.
Energy released
when bonds of
products are formed.
The bond energy is the energy required to break one mole of covalent
bonds and to separate the neutral atom completely from each other (the
same amount of energy is released when one mole of bonds is formed).
BOND ENERGY
Let us consider the energy required to break 4 C – H bonds and the energy
required to break the two double bonds in the oxygen molecules; these
values (table to the right) are considered positives:
+ 4 · 412 kJ = + 1648 kJ required
+ 2 · 496 kJ =+ 992 kJ required
Below are the values of the energy released
(negative values) when the bonds of products
are formed:
- 2 · 743 kJ = - 1486 kJ released
- 4 · 463 kJ = - 1852kJ released
Bond
Bond energy
( kJ mol-1)
C-H
412
C-C
348
C=C
612
O–H
463
O=O
496
C=O
743
BOND ENERGY
The D H for the combustion of methane is obtained making the sum of energy
required to break bonds of reactants and the energy released when bonds of
products are formed:
REACTANTS
+ 4 · 412 kJ = + 1648 kJ required
ENERGY
+
+ 2 · 496 kJ = + 992 kJ required
+
PRODUCTS
ENERGY
- 2 · 743 kJ = - 1486 kJ released
+
- 4 · 463 kJ = - 1852kJ released
D H = - 698kJ
HEAT OF FORMATION
The heat of formation of a compound is the heat change that takes
place when one mole of a compound, in its standard state, is
formed from its elements in their standard states.
The standard state of an element or compound is its normal form at 25 °C at one
atmosphere pressure (101 kPa)
The heat of formation of water is – 285.8 kJ / mol and may be represented as:
H2(g) + ½ O2(g)
H2O(l) D H = - 285.8 kJ / mol
The following equation does not represent the heat of formation of water:
H2(g) + ½ O2(g)
H2O(g) D H = - 241.8 kJ / mol
The reason is because the water formed in the above reaction is in the gaseous
state. At 25 °C the normal state of water is a liquid.
HEAT OF FORMATION
The heat of formation of any element in its standard state is zero. The heats of
formation of most compounds have negative value of D H: energy is given out
when the compound is formed from its elements. The more negative the value of
the heat of formation is, the more stable is the compound.
Substance
Formula
Normal state at
25 °C
CH4
Heat of formaton
D H (kJ mol-1)
- 74.9
Methane
Ethane
C2H6
- 84. 7
gas
Carbon monoxide
CO
- 111
gas
Carbon dioxide
CO2
- 393
gas
Ammonia
NH3
- 46.2
gas
Ethanol
C2H5OH
- 278
liquid
Iron
Fe
0
solid
Oxygen
O2
0
gas
gas
HESS’S LAW
Hess’s law states that if a chemical reaction takes place in a
number of stages, the sum of the heat changes in the separate
stages, is equal to the heat change when the reaction is carried
out in one stage.
In other words, the amount of energy liberated or absorbed in a
chemical reaction is the same whether the reaction takes place in
one o several steps, as shown below:
A
B
C
D
D HA  D = D HA  B + D HB  C + D HC  D
HESS’S LAW
Consider the two routes for converting carbon and
oxygen to carbon dioxide
C
+1/2 O2
D H2 =
?
CO
+ O2
D H 1 = - 394 kJ mol -1
C O2
HESS’S LAW
Hess’s law may be used to calculate the unknown heat of the reaction that
converts carbon to carbon monoxide: C + 1/2 O2
CO
C
+1/2 O2
+ O2
D H 1 = - 394 kJ mol -1
C O2
D H2 =
?
CO
D H1=D H2+D H3
C(s) + 1/2 O2(g)
- 394 = x - 283
CO(g)
x = - 394 + 283 = -111 kJ mol-1
D H 2 = - 111 kJ mol-1
HESS’S LAW
Hess’s Law is useful because:
• the
value of the heat of reaction can be
calculated indirectly, if it proves difficult to find it
directly in the laboratory.
(It is difficult to measure the heat of formation of CO
since some CO2 is also formed during the experiment).
• thermochemical
reactions can be added (+),
subtracted (-), multiplied (x) or divided (:) to
calculate unknown heat of reaction.
HESS’S LAW
Calculate the heat of formation of CH4 from the following heats of
combustion:
a) C(s) + O2(g)
b) H2(g) + 1/2 O2(g)
c) CH4(g) + 2 O2(g)
CO2 (g)
H2O (g)
CO2 (g) + 2 H2O (g)
D H = - 393 kJ mol -1
D H = - 286 kJ mol -1
D H = - 879 kJ mol -1
We need to write the balanced equation of formation of methane from
his elements:
C(s) + 2 H2(g)
CH4 (g)
D H=?
HESS’S LAW
We must build up the reaction of formation of methane using the
above mentioned thermochemical reactions. Introduce C
a) C(s) + O2(g)
D H = - 393 kJ mol -1
CO2 (g)
Introduce 2 H2. Multiply equation b) by 2. Do not forget to multiply
D H by 2 as well.
b) x 2 : 2 H2(g) + O2(g)
2 H2O (g)
D H = - 572 kJ mol -1
HESS’S LAW
Get the left-hand side of the required equation by adding the following
equations:
a)
C(s) + O2(g)
CO2 (g)
D H = - 393 kJ mol -
2 H2(g) + O2(g)
2 H2O (g)
D H = - 572 kJ mol -
1
b) x 2 :
1
d) C(s) + 2 H2(g) + 2 O2(g) CO2 (g)
+ 2 H2O (g) D H = - 965 kJ mol -1
HESS’S LAW
Introduce CH4 We have CH4 on the left-hand side of equation c). We need it
on the right-hand side of the equation for the formation of methane. Reverse
equation c), remembering to reverse the sign of D H too
c) reversed CO2 (g) + 2 H2O (g)
1
CH4(g) + 2 O2(g)
D H = + 879 kJ mol -
HESS’S LAW
Now add c) reversed to d) equation:
d)
1
C(s) + 2 H2(g) + 2 O2(g)
c) reversed CO2 (g) + 2 H2O (g)
CO2 (g) + 2 H2O (g) D H = - 965 kJ mol CH4(g) + 2 O2(g)
C(s) + 2 H2(g)+ 2 O2(g)+CO2 (g)+ 2 H2O (g)
O2(g)
D H = +879 kJ mol -1
CO2 (g)+ 2 H2O (g) + CH4(g) + 2
We obtain the required equation of formation of methane:
C(s) + 2 H2(g)
CH4 (g)
-1
-1
STUDENT EXPERIMENT
Calculate the heat liberated using the equation:
Heat liberated = (ms + me) · c · (t2 – t1)
ms= mass of solution in kg
me = mass equivalent in water of the calorimeter that absorbs heat (about 30 g)
c= specific heat capacity of solution in J kg-1 K-1 (assume 4060 J kg-1 K-1 for the
solution)
t2 – t1 = temperature rise
Note: Assume that the density of the solution is 1 g / cm 3
Calculate the heat of neutralisation
Knowing the heat liberated when this number of moles of acid was
neutralized, you can now calculate the amount of heat liberated when one
mole of the acid is neutralized.
STUDENT EXPERIMENT
Experiment
To determine the heat of reaction (heat of neutralisation) of
hydrochloric acid with sodium hydroxide.
Introduction
In this experiment , an equal number of moles of hydrocloric acid are
mixed with an equal number of moles of sodium hydroxide in a
calorimeter. The rise in temperature is then measured and from this
the heat of neutralization is calculated.
Apparatus required
Thermometer (reading to 0.1 °C), two beakers (100 ml), a calorimeter
(250 ml), two burettes (50 ml).
Chemicals required
Hydrochloric acid (1.0 M), sodium hydroxide (1.0 M).
STUDENT EXPERIMENT
Procedure
• Using a burette, place 50 ml of the 1 M hydrochloric acid solution into
the first beaker.
• Using the second burette, place 50 ml of the 1 M solution of sodium
hydroxide into the second beaker.
• Measure separately the temperature of the two solutions.
• When the two solutions are at the same temperature, place the acid into
the calorimeter and quickly add the base to the acid, stirring well.
• Place the lid on the calorimeter, to prevent heat loss, and record the
maximum temperature reached.
STUDENT EXPERIMENT
Summarize results as follows:
Temperature of HCl before mixing
=
°C
Temperature of NaOH before mixing
=
°C
Highest temperature reached after mixing
=
°C
Temperature rise
=
°C
Number of moles of acid used
=
mol
Number of moles of base used
=
mol
Mass of solution
=
Kg
CLIMATE CHANGE
Why is the planet warming?
• Huge amounts of energy are needed for heating,
generating electricity and for transport.
• Fuels give off numerous emissions that many say
are bringing about great climate changes, because of
the greenhouse effect of global warming.
• All economic activity that requires energy
consumption contributes to these emissions
producing greenhouse gases.
CLIMATE CHANGE
Who and what contributes to greenhouse gases?
Transport: CO2 emissions come from fuel combustion.
Agriculture: is the largest producer of CH4.
Industry: mainly contributes to greenhouse gas emissions through
energy use, including direct consumption of fossil fuels and use of
electricity (e.g. CO2 is a by-product of cement manufacture).
Industry also produces some CFC’s although they, along with
substances that were attacking the ozone layer, were banned by
the Montreal Protocol in 1989.
Energy: is the largest source of CO2 – mainly emissions
from fossil fuel combustion to produce electricity.
CLIMATE CHANGE
Why are greenhouse gases dangerous for the environment?
Greenhouse gas …
… is a gas in the atmosphere that freely allows
radiation from the sun to reach the earth’s
surface, but traps the heat radiated back from the
earth’s surface towards space.
The heating effect is analogous to the manner in
which the glass of a greenhouse traps the sun’s
radiation to warm the air inside the greenhouse.
CLIMATE CHANGE
Greenhouse gas
Sun’s
radiation
CO2 freely allows radiation
from the sun, but traps heat
radiated back from the
earth’s surface.
Heat
Vegetation subtracts CO2
from the atmosphere,
reducing damage by
greenhouse gases.
CLIMATE CHANGE
Solutions?
Hydrogen has the best potential of becoming the fuel of the future.
Hydrogen can be produced from sustainable, renewable sources
and may contribute to meeting the growth in world energy demand.
Hydrogen is a carbon-free energy carrier. When used in fuel cells,
there are no harmful emissions.
The use of energy may lead to climate changes. It is thus necessary
to make the transition to cleaner and environmentally favourable
energy carriers.