Physical Science 20 Heat Unit
Thermodynamics
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Thermodynamics is the study of energy transformation. There are 2 laws that govern the movement of heat in
nature between objects. These are termed the 2 laws of thermodynamics
o First law of thermodynamics: law of conservation of energy states that energy is never lost or gained, it
only changes form
 Energy can have many forms (heat, light, sound), and it can change between forms and is never
lost in a closed system
 When an exothermic reaction occurs, some of the energy of the system is converted into heat
or light. The energy is released by the system to its surroundings. The total energy of the system
decreases, but the energy of the surroundings increases by the same amount:
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Energysystem  Energysurroundings
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When an endothermic occurs some of the energy of the surroundings is absorbed and
converted to molecular enthalpy
Second law of thermodynamics: heat energy always travels spontaneously from a warmer body (body
with a higher temperature) to a colder body
Heat and Temperature
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In studying energy changes in systems we need to make clear distinctions between the terms heat and
temperature
o Heat: it the thermal energy (kinetic energy) that is transferred from one body to another. It is measured
in metric unit termed Joules (J). As mentioned above heat is transferred spontaneously from objects of
higher temperature to ones of lower temperature (warmer to colder bodies)
o Temperature: it is a measurement of the average kinetic energy of the particles that make up the
substance.
Substances that have the same temperature but very different amounts of heat.
o Energy in the lab is most easily measured as heat. The amount of heat transferred can be readily
determined by changes in temperature of the surroundings.
Form of energy are heat, light, electricity, and energy of motion (kinetic energy)
The energy required to raise the temperature of one gram of water one Celsius degree is defined as one Calorie.
It is not an SI metric unit and therefore has been replaced by Joule.
Calculate Energy Changes
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When a substance undergoes temperature without a phase change, three factors determine the amount of heat
(Q) a substance will absorb or release.
o Mass of a specific substance (m) measured in grams (g)
o
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Specific heat capacity (c) of a substance (on a table), units (
J
)
g  oC
o Temperature change (ΔT) that the substance undergoes (oC)
Heat Calculation formula
o The factors above all have a direct relationship with the amount of heat transferred
o Q=mc ΔT
Enthalpy Changes
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When a phase or chemical change takes place not only does kinetic energy of a system change by so does the
potential energies of the molecule. Chemists therefore consider the total energy of the system when calculating
the energy changes. Chemists define the total internal energy of a system at constant pressure as Enthalpy (H).
You can’t measure this directly, but you can measure it by examining what happens to the surrounding s of the
system, these changes are termed enthalpy change (ΔH)
If energy is absorbed (endothermic reaction) the enthalpy change is positive (+ΔH). If the energy is released
(exothermic) the enthalpy decreases and the enthalpy change is negative (-ΔH)
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), it is 25oC and standard
pressure is 100KPa.
Comparison of Enthalpy change
o The enthalpy change for phase changes are typically smaller than those for a chemical change
 Physical Change (dissolving)
 NH4NO3 (s)  NH4+(aq) + NO3-(aq)
ΔH= +27KJ
 Chemical Change (decomposition)
 2Fe2O3 (s)  4 Fe (s) + 3O2(g)
ΔH = +1625KJ
Representing Enthalpy changes in reactions
o Graph – enthalpy change vs. course of the reaction
o Enthalpy change as a term in the equation
 The heat absorbed and produced in a chemical reaction also varies directly with the amount of a
substance that reacts and the exact amount is determined by the heat change for that reaction
(ΔH). If you double the amount of substance that is reacted, then you will double the enthalpy
change
 2H2(g) + O2 (g)  2H2O(l)
ΔH = 572 KJ
 4H2 (g) + 2O2 (g)  4H2O(l)
ΔH = 1144 KJ
 If the reaction is exothermic the energy will appear on products side
 If the reaction is endothermic the energy will appear on the products side
 H2(g) + O2 (g) + 572 KJ  2H2O(l)
 2H2O(l)  H2(g) + O2 (g) + 572 KJ
o
Enthalpy change written separately as ΔH
 2Al(s) + 3Cl2(g)  2AlCl3(s)
ΔH = -1408KJ
Molar Enthalpy
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Molar enthalpy for many substance in certain reactions have been determined by chemists in tables
Q = n ΔH
o Q – heat
o n – number of moles
o ΔH – molar enthalpy (enthalpy change per mole)
The table tells you 2 things
o The amount of energy required to raise the temperature of one kilogram of a substance one Kelvin
varies from one substance to another
o It is also noted that the heat required to raise the temperature of a substance varies with the state of
the substance (solid, liquid or gas)
Calorimetry
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Is the experimental calculation of molar enthalpy
Theory – chemists determine the enthalpy change in a system by examining the energy changes in the
surroundings. Using the first law of thermodynamics chemists assume that the energy change in the system will
be equal and opposite to the energy change in the surroundings.
Energysystem  Energysurroundings
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To determine the energy of the surroundings chemists use substance that will change temperature. Therefore
they calculate the energy change in the surroundings (water) using:
o Q=mc ΔT
 Q – energy change
 M - mass of a water
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 ΔT - Temperature change of the surroundings (water)
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Energy of the system is determined by
o Q = n ΔH
So, when we put them both together we can calculate the energy of the substance
o
H 
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C - Specific heat capacity of a its surroundings (water) (on a table), units
mcT
n
M – mass of the surroundings
C – specific heat capacity of the surroundings (water)
ΔT – change in temperature of the surroundings
N – moles of the substance in the system (reaction that occurs)
Experiment – Calorimetry
o Takes place in a calorimeter (an insulated cup and water)
o Insulate so that the energy stays within the system
J
g  oC
Hess’s Law: Theoretical Calculation of Enthalpies of Reactions
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Theory: Calorimetry cannot be used to measure the enthalpies of the following:
o Reactions that can’t be isolated from others and occur at the same time.
 For example, the formation of carbon monoxide from carbon and oxygen. When carbon burns in
air it forms both carbon dioxide and carbon monoxide. The formation of carbon monoxide can’t
be separated from carbon monoxide, then how can we measure the heat of formation of carbon
dioxide?
o Reactions that proceed too slowly to effectively measure with a calorimeter.
 For example, rusting of iron is a very slow process, and the temperature changes in the
surroundings are too small to measure.
In order to measure these heats of these reactions (enthalpies), chemists have developed a theoretical method
based on a simple idea.
o Some aspects of a system are independent of the way the system changes from start to finish
Example to prove my point: Personal Wealth
o 3 scenarios
 A man works diligently his whole life and retires at 55, with a personal wealth of 2 million dollars
 A man makes and loses a fortune of 2 million dollars several times through his life and
eventually ends up at 55, with a personal wealth of 2 million dollars
 A man lives in poverty until the day he retires at 55, and wins a lottery worth 2 million dollars.
o Each man lived a lifetime, each retired at 55 and in each case lived a very different life, but they all
ended up with 2 million dollars.
o Their wealth is therefore independent of how their lives were lived
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Calculating Hess’s Law
o In order to calculate a ΔH overall (net) using Hess’s Law we need to be able to manipulate the steps of a
reaction so that they can all add up to the overall reaction.
 When using heats of formation we can ONLY manipulate 2 things
 Reversing the reaction to change the location of products and reactants, and therefore
changing the sign on the ΔH
 Multiply the equation by a common factor. When we do this we must also multiply the
ΔH by the same factor.
o Using Heats of Formation
 Steps
 Write the overall reaction for the reaction if not given
 Manipulate the given equations (table) for the steps of the reaction so they add up to
get the overall reaction
 Add up the equations cancelling common substance in reactants and products
 Add up the heats of formation of the steps
 Heats of formations of a compound is the amount of heat required to form a compound from its
elements
 Heats of formation of elements are zero as they cannot be made and therefore do not appear
on a table
o Using Formula
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H overall   H products   H reac tan ts
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Steps
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Balance the overall reaction
Find the heats of formation from the table
Put ΔH’s into formula and solve
Calculating Enthalpy Changes
Ratio Method
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Calculate the number of moles of a substance reacted for produced
Create a proportion using the balances and heat in the chemical reaction
Solve the missing quantity
E.g.
o 2C4H10(l) + 13O2(g)  8CO2 (g) + 10H2O(g)
ΔH= -5315KJ
 Calculate the amount of heat released when 25g of C4H10 (l) is burned in oxygen using the
equation above
Equation Method (Q = n ΔH)
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Determine the information given
Calculate the number of moles of a substance and the molar enthalpy from the equation, or take it from the
table of values if needed
Calculate the missing quantity
E.g.
o 2C4H10(l) + 13O2(g)  8CO2 (g) + 10H2O(g)
ΔH= -5315KJ
o How much heat will be released if 65 grams of butane is burned in a lighter according to the above
equations