Chemistry_Energetics_Lecture_1.Chemistry_Energetics_Lecture

FC304
Chemistry
Theme: 3 Hour: 1 Energetics 1

The lecture notes included in this presentation have
been adapted from the resources accompanying the
module textbooks:
• Lister, T and Renshaw, J (2015) AQA Chemistry AS
Student’s book, London: Nelson Thornes
• Lister, T and Renshaw, J (2015) AQA Chemistry A2
Student’s book, London: Nelson Thornes

Introduction to Theme 3
In the this theme we will cover:
 Heat, specific heat and enthalpy (Session 1)
 Standard enthalpy changes and Hess’s law (Session 2)
 Bond enthalpies (Session 3)
 Born-Haber cycles (Session 5)
 Enthalpy of solution (Session 6)
 Entropy and Free Energy (Session 7)
 Exam style questions (Sessions 4 and 8)

Introduction to Theme 3
In this theme we will cover the following academic literacies:
• Find, evaluate and synthesise data from multiple sources
• Find and explain patterns in, interpret, evaluate and draw
conclusions from data
• Use subject specific vocabulary effectively
• Demonstrate numeracy skills of benefit to study and future
employability

At the end of this lesson you
should be able to:
• Relate heat and specific heat.
• Calculate ∆H from calorimetric data.

Thermochemistry
• Thermodynamics is the science of the relationship
between heat and other forms of energy.
• Thermochemistry is the area of Thermodynamics
related to the heat changes involved in chemical
reactions.

All chemical reactions obey the First Law of
Thermodynamics (also known as The Law of
Conservation of Energy):
“Energy may be converted from one form to
another, but the total quantities of energy remain
constant”.
Energy and Chemical Reactions

When a chemical reaction happens:
• energy is required to break bonds
• energy is released when bonds are formed.

• In chemical reactions, energy is often transferred from
the “system” to its “surroundings,” or vice versa.
• The substance or mixture of substances under study
in which a change occurs is called the
thermodynamic system (or simply system.)
• The surroundings are everything outside the
thermodynamic system.

Heat Energy and Chemical
Reactions
• Heat can be defined as the energy that flows into or out
of a system because of a difference in temperature
between the system and its surroundings.
• The Heat energy is transferred from a region of higher
temperature to one of lower temperature; once the
temperatures become equal, heat flow stops.

Heat Energy and Chemical
Reactions
Exothermic
Heat “out of” a system
Endothermic
Heat “into” a system
Energy
System
Surroundings
Energy
System
Surroundings

Enthalpy, H
The heat absorbed or evolved by a reaction
depends on the conditions under which it occurs,
such as pressure.
Instead of the term “heat”, scientists prefer to refer
to a related absolute property:
Enthalpy.
The symbol used for enthalpy is H

Enthalpy and Enthalpy Change
• When a reaction happens, reactants change into
products.
• The reactants have a particular enthalpy (energy)
and the products have a different enthalpy
(energy).
• The difference between these two energies is the
“heat of reaction” or “enthalpy of the reaction” ΔH
ΔH = Hproducts - Hreactants

• There is no way of measuring the enthalpy of
any single substance directly.
• For this reason we can only discuss enthalpy
changes in a reaction and this can be measured
as an amount of heat given out (exothermic
reaction) or taken in (endothermic reaction).

We can represent the level of enthalpy possessed
by reactants and products on a “Enthalpy
Diagram”.
Enthalpy, H
Progress of Reaction
Reactants
Products
Δ H

Exothermic reactions
If the potential energy diagram shows that the
energy of the reactants is higher than that of the
products, the reaction will release energy.
ΔH is negative. It will be an Exothermic reaction.
Enthalpy, H
Reactants
Products
Δ H is -ve

Enthalpy Change - Exothermic
• E.g. When concrete is mixed, the main reaction is
between calcium oxide and water.
• This reaction gives out a large amount of heat –
so much that when building a large concrete
structure such as a dam, cooling pipes must be
included to carry away the heat.
CaO(s) + H2O(l) → Ca(OH)2(s) ΔH = −65.2 kJ mol-1

Endothermic reactions
If the potential energy diagram shows that the
energy of the reactants is lower than that of the
products, the reaction will take in energy.
ΔH is positive. It will be an Endothermic reaction.
Enthalpy, H
Reactants
Products
Δ H is +ve

Enthalpy Change - Endothermic
One example of an endothermic reaction is:
2NH4NO3(s)+ Ba(OH)2.8H2O(s) 2NH3(g)+ 10H2O(l) + Ba(NO3)2(aq)
ΔH = +17.44 kJ mol-1
Watch this at:
http://www.youtube.com/watch?v=GmiZ0huvZzs

Obviously, these amounts of heat energy released or
absorbed must be dependent on the quantity of substances
reacting so we must define the enthalpy changes in terms
of;
• “energy changes (kilojoules) per amount of substance
(the mole)”.
Therefore, the units of enthalpy and enthalpy change are
kJ mol-1

Thermochemical equations
It is important to give the exact reaction equation
when quoting the associated energy change.
2Mg(s)+ O2(g)  2MgO(s) ΔH⦵ = - 1204 kJ mol-1
Mg(s)+
1
2
O2(g)  2MgO(s) ΔH⦵ = - 602 kJ mol-1

Thermochemical equations
• Note that you must also always give the states of the
reactants and products when quoting H for a
reaction.
Pb(s)+ Cl2(g)  PbCl2(s) H ⦵ = - 359 kJ mol-1
• There would be a large extra energy change needed
to change Pb(s) into Pb(g)!

Standard State Enthalpy Changes
• In order that chemists worldwide can compare notes on
thermochemical experiments, a series of standard conditions
have been agreed.
• The term “standard state enthalpy” refers to an enthalpy
change for a reaction in which reactants and products are
considered to be in their standard states
• Standard states are (most stable state of the substance) at a
specified temperature 298 K (25°C) and pressure (one
atmosphere )

Standard State Enthalpies
• The enthalpy change for a reaction in which
reactants are in their standard states is denoted
as;
∆H⦵
• You will see enthalpy changes without the circle
but you can assume that all enthalpy changes in
this module are standard state unless you are
specifically told otherwise.

Measuring Enthalpy changes
• Enthalpy changes are calculated by measuring the
heat required to change the temperature of a
surrounding substance (usually water).
• The heat required to raise the temperature of a
substance is its heat capacity.

Specific heat capacity
The specific heat capacity, c, of a substance is:
The amount of heat needed to raise the temperature of one
gram of a substance by one degree
Substance Specific heat capacity (J g-1 K-1)
Aluminium 0.901
Iron 0.449
Water 4.18

Measuring Enthalpy Changes:
Calorimetry
• A bomb calorimeter is a device used to measure
the heat absorbed or released during a physical or
chemical change.
• Everything inside the “bomb” is the system. The
water is the surroundings.

Allow the reaction to heat (or cool) a known mass of water,
measure the temperature change of the water then calculate
the energy required using the formula;
Q = c m ∆T
Q= Heat /energy change (J)
∆T = the temperature change (K) (final temperature – initial temperature)
m = mass of water (g)
c = specific heat capacity of water (4.18 J g-1 K-1 OR kJ K-1 kg-1)
Measuring the Enthalpy Change

Step to find enthalpy by Calorimetric Method
1. Use q = c m ∆T, to calculate the given quantities' energy change.
2. Divide the answer of step 1 by a thousand to convert J to kJ
3. Work out the moles of the reactants used
4. Divide q by the number of moles of the reactant, to give ΔH
5. Add a sign and unit (kJmol-1)
Measuring the Enthalpy Change

Example
0.253g of ethanol is burned in a calorimeter. The temperature of the
surrounding 150g of water rises 10K.
Calculate the enthalpy of combustion for 1 mole of ethanol:
Step1:
Heat is given out by the ethanol Q = c m ΔT
= 4.18 x 150 x 10
= 6270 J
Step 2: 6270 J (6.27 kJ)
• Heat has been given out so the reaction is exothermic = - 6.27 kJ

Example
Step 3:
The molar mass of ethanol = 46g
Moles of ethanol used = 0.253/46 = 0.0055 moles
Step 4:
Assuming all the heat produced by the combustion was absorbed
by the water: 0.0055 moles of ethanol burns to give -6.27kJ of heat.
Therefore 1 mole of ethanol should burn to give
(-6.27 x 1/0.0055) kJ = -1140 kJ heat
Step 5:
Estimated enthalpy of combustion (∆Hc) = - 1140 kJ mol-1

Question
• Enthalpy changes are not just for combustion (burning)
reactions – the heat evolved when a substance dissolves
can be measured this way too.
• Calculate the enthalpy of solution for dissolving 0.80 g of
Sodium Hydroxide in 25cm3 of water if the temperature
rise observed is 5.2°C. (c = 4.18 J g-1 K-1)
• Note that the density of water = 1g cm-3

Summary
• Reactions can be exothermic or endothermic.
• Laboratory methods can be used for determining enthalpy
changes using the equation: energy transferred = c m ∆T.

Self-study and homework
 Complete Heat and Enthalpy worksheet

Chemistry_Energetics_Lecture_1.Chemistry_Energetics_Lecture

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