Thermochemistry - Horton High School

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Thermochemistry
-a study of energy

Look at the world around us. The world we have
built depends on our use of energy. The cars we
drive, the electricity that powers our lights and
machines, the dynamite we use in construction,
and the way we heat our homes, all depend on
our efficient use of energy. The way we use this
energy is therefore very important to us.


Two aspects of energy use should be considered:
a) Sources of energy we use. b) The
technologies we have created to use the sources
of energy efficiently.

Nitroglycerin C3H5(NO3)3 (l) gives off 1500 KJ of energy per
mole when it decomposes.

Canadian Minute

It is the speed of the decomposition reaction which makes
nitroglycerin such a violent explosive. Unlike burning,
which can only travel as fast as the flame front can move
through the material, high explosives are decomposed
almost instantaneously by a supersonic shock wave
passing through the material. This instantaneous
destruction of all the molecules in the sample is called a
detonation, and the rapid expansion of hot gases that
results is what causes the destructive blast. In fact, 4
moles of nitroglycerin produces 35 moles of hot gases.
Nitroglycerin gives off 1500 KJ of energy
per mole when it decomposes.
 Alfred Nobel found a way to stabilize
nitroglycerin when he invented dynamite.


Energy changes accompany every
chemical and physical process from the
decomposition of nitroglycerin to the
vaporization of water. We will explore the
energy changes that accompany various
processes.



Energy is defined as the ability to do work (we
are really talking about the movement that
occurs against a restraining force and is equal to
the force x the distance over which the motion
occurs) or produce heat.
Energy comes in many forms- solar, nuclear, and
electrical are just a few examples. In chemistry,
heat energy is often what we are interested in.
In previous science classes you have most likely
learned that the individual atoms and molecules
that make up all substances are in motion. This
motion is called kinetic energy. The energy
associated with this motion is heat, or thermal
energy.

Heat can also be defined as the flow of
energy. This energy transfer is always from
the hotter substance to the colder object. If
you hold an ice cube in your hand, your
hand feels colder. Your hand is giving heat
(or losing heat) to the ice cube. The ice
cube does not transfer “coldness” to your
hand.

In addition to kinetic energy there is
another basic form of energy- potential
energy. Potential energy is energy of
position (stored energy).


We will be concerned with one type of
potential energy in this unit – chemical
bond energy.
Chemical bond energy comes from the
attractive forces between molecules, and
between atoms. This form of energy plays an
important role in chemical reactions.
The standard unit for energy is the joule (J). Another standard unit
for energy is the calorie.
Calories (with a capital C) are used to measure the energy in food;
1 Calorie (capital C) = 1000 calories (lowercase c).
. One joule is the amount of energy used by a 100 W light bulb in
0.1 s.
. One joule is approximately equal to the amount of energy you
expend when you bring a cheeseburger to your mouth.
. An ordinary paper match releases a little over 1,000 J when
completely burned.
Temperature describes the amount of motion that the molecules or
atoms in a material have. Fast movement represents high
temperature and slow movement represents low temperature.
Temperature then is a measure of the average kinetic energy of
the particles that make up the substance.
Kelvin = oC + 273.15
oC
= K – 273.15

Thermodynamics and Thermochemistry

Thermodynamics is the study of heat and
how it can be converted into other energy
forms. In thermodynamics, the system is
a specific part, often the system is a
chemical reaction. The surrounding are
typically everything outside the test tube
where the chemical reaction is taking
place.



An open system can exchange both energy
and matter with its surroundings. An active
volcano is an example of an open system.
A closed system can exchange energy but
not matter. A sealed greenhouse is an
example of a closed system.
An isolated system can exchange neither
energy nor matter. A sealed thermos bottle
is an isolated system.

The First Law of Thermodynamics
also known as Law of Conservation of Energy
The first law states that energy cannot be created
or destroyed. In other words, when a system
gains or loses energy from the surroundings, the
total energy will be constant.
. When octane (C3H8, the main component of
gasoline) is burned, chemical bond energy
(potential energy) is converted into mechanical
energy (pistons moving in the car); kinetic
energy) and heat.
. When we eat, our bodies convert the chemical
energy of the food into movement of our
muscles; heat is also produced as a product of
this conversion.

This concept can be expressed as:
ΔE = q + w
The E represents the internal energy. This includes but is not limited to all kinetic
energy and potential energy possessed by all components of a system. ΔE is
the change in internal energy of a system as a result of heat flow and work.
Heat flow is indicated by q, and the work done by the system is w.
. q is a positive value if heat is absorbed by the system.
. q is a negative value if heat is lost by the system.
. w is a positive value if work is done on the system.
. w is a negative value if work is done by the system.
"In this house we obey the laws of thermodynamics!" (after Lisa constructs a
perpetual motion machine whose energy increases with time) — Homer
Simpson


The Second Law of Thermodynamics states
that heat energy always travels
spontaneously from a warmer body (the body
with the higher temperature) to a colder
body.
You have known since you were a child that if
you touch a hot stove you will get burned.
The stove is warmer than your body and
therefore transfer heat to your hand!

If you go outside in the very cold
without gloves you can get frostbite.
The heat is moving away from your
body into the surroundings. When your
body can no longer replace the heat as
quickly as it is being lost, your skin
freezes and you get frostbite.

Endothermic and Exothermic Reactions
All chemical reactions (and physical
changes and nuclear changes as well)
involve energy changes. Exothermic
reactions are reactions in which there is a
net release of energy. When energy is
released, an energy term will appear on
the product side of the equation.
Fe2O3(s) + 2 Al
(s)
 Al2O3
(s)
+ 2Fe(l) + 847.6 kJ

Endothermic reactions are reactions that
require a net input of energy. This is
indicated by writing the energy term on
the reactant side of the equation.
2S03
(g)
+ 198 kJ  2 SO2
(g)
+ O2
(g)

Measuring Heat: Calorimetry
An instrument called a calorimeter
measures the heat which is exchanged
during a reaction.

A calorimeter is a thermally insulated container
where a reaction system is contained and the
energy exchange between the system and its
surroundings can be measured.

If the reaction is exothermic, energy is given off
causing the temperature of the water in the
calorimeter to rise.
If the reaction is endothermic, energy will be
required. This causes the temperature of the
water to drop.
By noting the temperature change of the water
in the calorimeter, we can calculate how much
energy was exchanged.


The specific heat constant (c) of a substance is
defined as the energy required to raise the
temperature of one gram of a substance by one
degree Celsius.
 Every substance has a characteristic specific heat
constant. The specific heat constant of water is 4.19
J/goC.
 The First Law of Thermodynamics says that the heat
energy of the system is equal to the negative heat
energy of the surroundings.



q surroundings = -q system
The heat, q entering or exiting can be determined by
using the heat equation for the surroundings.
q(heat)=(mass) (specific heat)(final temperature – initial temperature)
 q = mcΔt


Water has a relatively high specific heat which allows it to absorb and
release large quantities of heat. This is why temperature changes are
often more moderate in areas near large lakes and other large bodies of
water. There is a reason why people buy cottages other than boating!
Specific Heats of various
substances at 25oC
Substance
Water (liquid)
Water (solid)
Water (steam)
Ethanol (l)
Aluminum (s)
Lead (s)
Iron (s)
Silver (s)
Specific Heat
J/(goC)
4.184
2.03
2.01
2.44
0.897
0.129
0.449
0.235

Q= mcΔT
Where
Q = amount of heat transferred, in kJ
m = mass of the sample in grams
ΔT = change in temperature, in oC
Note:
We indicate an increase in temperature as a positive
change and a decrease in temperature as a negative
change.
If the temperature started at 50 oC and ended at 40 oC,
ΔT would have a value of – 10oC.
The Calculations

A sample of 100 grams of water is placed in a calorimeter
at a temperature of 20.0 oC. Ammonium nitrate is
dissolved in the water lowering the temperature to 10.0 oC.
Calculate the heat change for the water. Specific heat of
water = 4.19 J/goC.

Q = mcΔT
= 100 g  4.19 J/goC  -10oC
= - 4190 J
The qwater is negative indicating that the heat flows from the
surroundings toward the solution. The solution is an
endothermic solution ΔHsolution = +4190 J or + 4.190 kJ


A 1000.0 g mass of water had a starting
temperature of 50oC. It lost 33,600 J of heat
over a 5 minute period. What was the
temperature of the water at the end of this
period?
Q = mcΔT so
ΔT = Q
mc
=
33,600 J
1000.0 g  4.184 J/goC
= 8 oC
The temperature decreased 8oC so the fianl
temperature was 50 – 8 = 42oC.

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There are energy changes during a physical change. They are changes of state or phase.
During the heating of a solid, liquid or a gas, there is a temperature change. During the
state change itself, the temperature remains constant. These heats are hidden (latent).
The latent heat of fusion is the energy change that occurs during the phase change
between solid and liquid. The latent heat of vaporization is the energy change that
occurs during the phase change between a liquid and a gas.
Molar heat of fusion is the amount of energy required to melt one mole of substance.
Molar heat of vaporization is the amount of energy required to vaporize one mole of
substance.


Fusion (melting) the substance absorbs heat, ΔHfus is positive
Freezing the substance releases energy, ΔHsolid is negative

ΔHfus = -ΔHsolid


Vaporization the substance absorbs energy, ΔHvap is positive
Condensation the substance releases energy, ΔH cond is negative

ΔHvap = -ΔHcond

To calculate the energy for the phase
change you need multiply the number of
moles (n) by the molar heat for the phase
change:

q = nΔHphase
Potential Energy Diagram of a
Physical Change

Examples of Energy Diagrams
1. Potential Energy Profile for a Phase Change
NH4NO3(s)  NH4NO3 (aq) ΔH = +25 kJ
NaOH(S)  NaOH
(aq)
ΔH = -44.2 kJ

An equation that indicates that the heat change is a
thermochemical equation.

NH4NO3

NaOH

Another way to show heat change is to use the ΔH
notation, the heat change for the system is shown outside
the equation. A positive ΔH represents an endothermic
process (system increases in energy). A negative ΔH
represents an exothermic process (system decreases in
energy).
Enthalpy or ΔH is written in terms of the heat change of
the system.

(s)
(s)
+ 25 kJ  NH4NO3
 NaOH
(aq)
(aq)
+ 44.2 kJ






Enthalpy, H is a state function used to describe the heat
changes that occur in a reaction under constant pressure.
When a reaction is allowed to take place in an open
container, a quantity of heat proportional to the quantity of
matter present, will be released or absorbed.
This flow of heat is the enthalpy change, ΔH. The units
for ΔH are kJ or kJ/mol.
Reactions that release heat are termed exothermic. They
have negative values of ΔH.
Reactions that absorb heat are termed endothermic. They
have positive values of ΔH.
Generally processes that feel cold (like an ice pack) are
endothermic and processes that feel hot (like a fire) are
exothermic.
Enthalpy

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Enthalpy Changes (Energy changes during phase, chemical
and nuclear change)
When a phase, chemical or nuclear change takes place not
only does the kinetic energy of the system change but so
does the potential energies of the molecules.
Chemists therefore consider the total energy of the system
when calculating these energy changes. Chemists define
the total internal energy of a system at constant pressure
as Enthalpy .
Enthalpy is given the symbol H.
Chemists have no method of measuring the total energy of
a system easily, but they can measure the changes that
occur in this energy. They do this by examining what
happens to the surroundings of the system. These changes
are termed Δ H or enthalpy change.
SATP
The ΔH of a reaction changes with varying conditions of
temperature and pressure, so chemists define a set of
conditions called standard atmospheric temperature and
pressure (SATP). Under these conditions they measure the
value of H. The standard atmospheric temperature is 25
oC and the standard pressure is 100 KPa.



Comparison of Enthalpy change
The enthalpy change for phase changes are typically smaller than those for
chemical changes. Enthalpy changes for nuclear changes are typically very
large compared to chemical changes. See below.
Phase change ( dissolving )
NH4NO3 (s)NH4+ (aq)+NO3 - (aq) H=+ 27 kJ
Chemical change ( decomposition )

2 Fe2O3 (s)4 Fe (s)+3 O2 (g)
H=1625 kJ

Representing (communicating) Enthalpy Changes in Reactions
Chemists can represent the change in enthalpy using several methods:
1. Using a graph of Enthalpy change vs the course of the reaction.
2. Using the balanced equation and enthalpy as a term in the equation.
3. Using the balanced equation and enthalpy written separately as a
"H".
1. Graphical
Graph of Endothermic Change for the
decomposition of water
H
H2(g)- ½ O2 (g)
ΔH = +285.8 KJ
H2O (l)
Course of Reaction

2) Using a balanced equation and energy as a term.
There is a direct relationship between the amount of substances that reacts and forms in an
equation. The factor that determines the exact relationship is the equation factor or mole ratio
(balance in equation). For example in the reaction below:
2H2 (g) + 1 O2 (g) ------> 2 H2O (l)
If one mole of hydrogen reacts with excess oxygen, one mole of water forms. If two moles of
oxygen reacts with excess hydrogen, 4 moles of water forms.
What relationship exists between the amount of reactant or product substance and the change in
enthalpy. (ΔH )? Complete the dry lab assignment to determine the answer to this question.
The heat absorbed and produced in a chemical reaction also varies directly as the amount of
substance that reacts and the exact amount is determined by the heat change for the reaction
(ΔH) .
If you double the amount of substance reacted then you will double the heat change.
Examine the following reaction between Hydrogen and oxygen to form water
2H2 (g) + 1 O2 (g) -----> 2H2O (l)
If one mole of oxygen when reacted releases 572 kJ of energy , then 2 moles will release
1046 kJ
Because the relationship is a direct one, energy can be written as a term in the equation.
If the reaction is exothermic the energy is written the products.
If the reaction is endothermic the energy is written in the reactants.

3) " ΔH" notation
The third method used by chemists to represent enthalpy change is
called " H" notation. In this case the enthalpy is written separately
from the reaction.
If the reaction is exothermic the ΔH is negative.
If the reaction is endothermic the ΔH is positive.
Using the equations above:
Exothermic: (-H value)
2Al (s) + 3Cl2 (g) ------> 2 AlCl3 (s)
Δ H = - 1408 kJb)
Endothermic: (+H value)
2 H2O ------> 2 H2 + 1 O2 (g)
Δ H = + 572 kJ
What would happen if we cut the # of moles of reactant and product in
half in each of the above reactions?
Enthalpy change would be cut in half.
Complete the assignment “Communicating Enthalpies”

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The heat absorbed and produced in a chemical
reaction also varies directly as the amount of
substance that reacts and the exact amount is
determined by the heat change for the reaction
(H) . If you double the amount of substance
reacted then you will double the heat change.
Examine the following reaction between
Hydrogen and oxygen to form water
2H2 (g) + 1 O2 (g) -----> 2H2O (l)
If one mole of oxygen when reacted releases 572
kJ of energy , then 2 moles will release 1046 kJ
Because the relationship is a direct one, energy
can be written as a term in the equation.
If the reaction is exothermic the energy is
written the products. If the reaction is
endothermic the energy is written in the
reactants.
Example:
a) Exothermic: written in the products.
2Al (s) + 3Cl2 (g) ----> 2 AlCl3 (s) + 1408 kJ
 If we double the amount of the reactants and
products, the heat doubles
4Al (s) + 6Cl2 (g) ----> 4 AlCl3 (s) + 2816 kJ
b) Endothermic: written in the reactants.
H2O + 286 kJ ------> H2 + 1/2 O2 (g)
If we double the amount of the reactants and
products we double the heat.
2 H2O + 572 kJ ------> 2 H2 + 1 O2 (g)

Energy Profile for a Chemical Change
To keep comparisons fair the chemist will measure heat
changes under standard conditions (SATP 25 ºC and
101kPa). The reactants and products are at SATP. To show
that the ΔH was measured at SATP we use the symbol
ΔHº.
CaO(s) + H2O(l) → Ca(OH)2(s) +
65.2kJ (exothermic)
Another way to show heat change is to use
the ΔH notation, the heat change for the
system is shown outside the equation. A
negative ΔH represents an exothermic
process (system decreases in energy).
CaO(s) + H2O(l) → Ca(OH)2(s) ΔHº = -65.2kJ/mol

can be represented by the energy
profile:
This

2NaHCO3 + 129kJ → Na2CO3(s) +H2O(g) + CO2(g) (endothermic)
2NaHCO3 → Na2CO3(s) + H2O(g) + CO2 (g) ΔHº = +129kJ
 A positive ΔH represents an endothermic process (system
increases in energy). This reaction can be represented by the
energy profile:


The endothermic reaction shown below indicates that 92.2 kJ are
absorbed when 2 moles of NH3 decompose to form 1 mole of N2
and 3 moles of H2.
2NH3 (g)  N2 (g) + 3H2 (g)
92.2 kJ + 2NH3 (g)

ΔH = + 92.2 kJ
or
N2 (g) + 3H2 (g)
Question :
In the reaction above, 100 g of NH3 are allowed to react to produce N2 and H2, how many kJ of
heat will be absorbed?
Answer:
The reaction tells us that for every 2 moles of NH3 consumed, 92.9 kJ will be absorbed. This
information will be used to calculate the conversion from moles of NH3 to kJ.
100 g NH3
x
1 mol NH3
17 g NH3
x
92.2 kJ
2 mol NH3
=
271 kJ
When a reaction is reversed, the sign of ΔH is changed:
N2 (g)
+
3H2 (g)

2NH3 (g)
ΔH = -92.2 kJ
Enthalpy Calculations

Hess’s Law states that the enthalpies of
reactions may be added when these
reactions are added. Substances
appearing on the same side are added,
while those on opposite sides are
subtracted. Typically, some reactions will
need to be reversed and multiplied
through by a number, so that when they
are combined the desired equation will
result.
Hess’s Law

In math we can add equations because
we are adding the same to each side of
the equation. Common terms on the
opposite sides of the equations can be
cancelled out before we add the equations
(b and c terms can be cancelled)
a+b=c+d
a+c=b+d
2a =
2d
The heat of reaction #3 could be found
using a colorimeter or it can be found
indirectly by calculating the sum of Reaction
#1 and Reaction #2. The additive property
of chemical reactions is called Hess’s Law.


Find the enthalpy of the equation:
C2H2 (g) + 5

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