Chapter: 5 Energetics and Thermochemistry
Title: Lesson 1 Exothermic and Endothermic
Learning Objectives:
– Reflect on prior knowledge of the energy changes in chemical reactions
– Understand the difference between endothermic and exothermic reactions
– Draw enthalpy level diagrams showing the relative stabilities of products and
reactants
– Complete an experiment investigating endothermic and exothermic
reactions
Energy and Heat Transfer Energy

Energy is a measure of the ability to do work.

We will be focusing on reactions involving heat changes.

Heat is a mode of energy transfer which occurs as a result of a temperature
difference.

Heat increases the average kinetic energy of the molecules in a disordered
fashion.

However, work can be contrasted as being a more ordered process of
transferring energy.
Work on lifting a
beaker of water lifts
all the water
molecules in the
same way!
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System and Surroundings

When chemical changes happen it is useful to
distinguish between the system (area of interest)
and the surroundings (everything else!)

Most reactions take place in an open system.
Energy and matter can be exchanged with the
surroundings.

A closed system can exchange energy but matter
cannot be exchanged with the surroundings.

Total energy cannot change during the process.
Any energy lost by the system is gained by the
surroundings and vice versa.
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Enthalpy, H

This is a measure of the energy locked up inside
chemicals

We can only measure changes in enthalpy, ∆H, not
enthalpy itself

Substances with lower enthalpy are more stable
than those with higher enthalpy

Enthalpy of reactants and products are not the
same (energy is either taken in or given out during
the reaction)

Enthalpy level diagrams show the changes in
enthalpy over the course of a reaction.

∆H can be observed as change in temperature
(measured or calculated)
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
∆H is positive when heat is added to the system

∆H is negative when heat is released from the system into the
surroundings
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ENDOTHERMIC & EXOTHERMIC
Chemical reactions can either release energy to
their surroundings, EXOTHERMIC
__________, or energy can be
transferred to them from the surroundings,
http://www.youtube.com/watch?v=pENDOTHERMIC
____________.
27I_osoaw
KEY IDEA
•Bonds contain enthalpy.
•Energy is absorbed to break bonds apart in the REACTANTS in a
chemical reaction.
•Energy is released when new bonds form in the PRODUCTS in a
chemical reaction.
•When bonds break, energy is absorbed (endothermic – get cooler) .
Self sustaining until reactants run out.
•When bonds form, energy is released (exothermic – get hotter).
Needs energy for the reaction to keep going or runs out when
reactants run out.
ENDOTHERMIC & EXOTHERMIC
EXAMPLES:
Exothermic:
A reaction endothermic in one
direction is always exothermic
in the other
• Neutralising an acid with an alkali
• Burning magnesium
• Adding water to anhydrous copper sulphate
Endothermic:
• Thermal decomposition of limestone
• Photosynthesis
• Heating hydrated copper sulphate
BBC Bitesize Hydrated Copper Sulphate Explanation
Good examples of exothermic and endothermic reactions used in everyday
life are the HOT PACKS and COLD PACKS used for treatment of sprains
and other muscular disorders.
HOT PACKS
COLD PACKS
A hot pack is a plastic bag filled with
saturated sodium ethanoate
(CH3COONa).
A typical cold pack contains the ionic
compound ammonium nitrate
(NH4NO3) and water.
It also includes a small concave metal
disc.
When allowed to mix, the nitrate goes
into solution and the electrostatic
forces of attraction between the
ammonium ions (NH4+) and nitrate
ions (NO3-) are broken.
Twisting the disc causes “nucleation “
(first phase of crystallisation) and the
sodium acetate begins to crystallize
very rapidly.
Since bonds are being formed, energy
is released and the pack becomes hot.
The hot pack is 'recharged' by boiling
it in water for several minutes.
Energy is taken in from the
surroundings to overcome these
forces, thus causing the cooling effect.
ENTHALPY CHANGES
ENTHALPY (H)
= the energy content of a
substance measured at
constant pressure.
H = enthalpy change
= enthalpy of PRODUCTS –
enthalpy of REACTANTS
Enthalpy (H)
Products (P)
A
Reactants (R)
H +ve
H -ve
B
For reaction A
Reaction Path (ie time)
For reaction B
 H is POSITIVE
 H is NEGATIVE
 ENDOTHERMIC reaction
 EXOTHERMIC reaction
 heat energy ABSORBED
 heat energy RELEASED
 net bond BREAKING
 net bond FORMATION
MgCO3(s)  MgO(s) + CO2(g)
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
The units are kilojoules per mole (kJmol-1)
ENTHALPY
Enthalpy Level Diagrams
Endothermic
Exothermic
Heat
taken in
Energy level of products
is higher than reactants
so heat taken in from
surroundings.
Heat
given out
Energy level of
products is lower than
reactants so heat given
out to surroundings.
EXOTHERMIC
An exothermic
enthalpy change is
always given a
negative value, as
energy is lost to
the surroundings.
ΔH = -kJmol-1
ΔH = ΔH products - ΔH reactants
Small number
Bigger number
ENDOTHERMIC
An endothermic
enthalpy change is
always given a
positive value, as the
energy is gained by
the system from the
surroundings.
ΔH = + kJmol-1.
ΔH = ΔH products - ΔH reactants
Big number
Silly Video
Small number
Enthalpy level diagrams

These show the changes in enthalpy over the course of a
reaction
H
H

∆H is negative (H goes down)

∆H is positive (H goes up)

Activation energy small

Activation energy large

Products more stable

Products less stable

Break weaker bonds, make stronger bonds

Break stronger bonds, make weaker bonds

Chemical energy turned into heat energy

Heat energy turned into chemical energy
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Simplified diagrams

These are the level needed by the IB

Less useful as they tell you nothing about activation energy and so on
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Intermediates

Some reactions form intermediate products


A product that quickly turns into something else
This can lead to more complicated enthalpy level diagrams
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STANDARD ENTHALPY CHANGES
Standard conditions:
•Pressure 100kPa
•Concentration of 1 mol dm-3 for all solutions
•All substances in their standard states
•Temp 298K
•ΔHθ  ‘delta H standard’
ENDOTHERMIC & EXOTHERMIC
QUANTITIES:
• Energy is measured in kilojoules per mole
• kJmol-1
CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l)
• Combustion of methane  one mole of methane reacts
with 2 moles of oxygen
• It gives out 890kJ of energy
ENTHALPY
Question
ENTHALPY
Answer
ENTHALPY
PHYSICAL STATES
Whether products and reactants are solids, liquids or
gases affects the enthalpy change of a reaction.
Heat is put in to change a liquid to a gas (endo)
Heat is given out when a gas turns to a liquid (exo)
H2 (g) + ½O2 (g)  H2O (l)
ΔH = -285.8kJmol-1
H2 (g) + ½O2 (g)  H2O (g)
ΔH = -241.8kJmol-1
Difference in the ΔH’s represents the amount of heat
needed to turn one mole of water into 1 mole of steam
ENTHALPY
1.Balance each of the equations below
2.Name the compounds or elements present in each one
3.Identify each one as exothermic or endothermic.
4.Draw the energy level diagram for each
CH4(s) + O2(g) → CO2(g) + H2O(l)
ΔH = -890 kJmol-1
HCl(g) → H2(g) + Cl2(g)
ΔH = 185 kJmol-1
NH3(g) + O2(g) → NO(g) + H2O(l)
ΔH = -1169 kJmol-1
ENTHALPY
ANSWERS
CH4(s) + 2O2(g) → CO2(g) + 2H2O(l)
Methane
Oxygen
Carbon
dioxide
Water
ENDOTHERMIC
2HCl(g) → H2(g) + Cl2(g)
Hydrogen
chloride
Hydrogen
Chlorine
4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(l)
Ammonia
Oxygen
EXOTHERMIC
Nitrogen
monoxide
EXOTHERMIC
Water
For each of the reactions above, construct a simple
enthalpy level diagram showing the enthalpy change.
CH4(s) + 2O2(g)
ΔH =-890 kJmol-1
CO2(g) + 2H2O(l)
H2(g) + Cl2(g)
ΔH = +185 kJmol-1
2HCl(g)
4NH3(g) + 5O2(g)
ΔH =-1169 kJmol-1
4NO(g) + 6H2O(l)
ENTHALPY
PLENARY
Cut out the boxes and stick them in the right place to
summarise enthalpy changes for exo and endothermic
reactions
Key Points
Exothermic reactions
Endothermic reactions
Give out energy
Absorb energy
Decrease in enthalpy
Increase in enthalpy
Products more stable
Reactants more stable
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