Unit 11 Thermodynamics

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Unit 11 Thermodynamics

• Chapter 16

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Thermodynamics

• Definition
• A study of heat transfer that accompanies
  chemical changes
• Concerned with overall chemical changes
• Chemical Change involves
• A change in energy
• A degree of disorder

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In Thermodynamics

• System refers to the reaction itself
• Surroundings refers to everything else
• Standard Conditions
• 25C (298 K)
• 1 atmosphere
• 1 molar solution

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Enthalpy

• Symbol H
• Measure of heat content (energy) of a system at
  constant pressure
• Cant be measure directly, can only measure the
  change in enthalpy.
• We call this the Heat of Reaction, ?H
• Measure of the heat released or absorbed in a
  chemical reaction!
• ?Hrxn ?Hproducts - ?Hreactants

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Hesss Law

• This summer your parents decide they are going to
  take you on a road trip to CA. You drive 400 mi
  day 1, 350 mi day 2, 275 mi day 3, and 100 mi day
  4. What was the total mileage of the trip?
• You add your daily totals,
• 400 mi 350 mi 275 mi 100 mi 1125 mi total
• In the same way, if the reaction we want requires
  two or more equations we can sum the enthalpy
  changes to get the total energy change!

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Hesss Law

• Say we want to find the ?H of the reaction A
  D ? E
• We know the following
• A B ? C ?H 27 kJ
• C D ? B E ?H -15 kJ
• Using Hesss Law we can calculate the ?H we want
  by taking the sum of the reaction (27 kJ) (-15
  kJ) 12 kJ

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• Hess's law states that the change in enthalpy of
  the reaction equals the sum of the enthalpy
  change for the intermediate steps of the
  reaction. Hess's law could also be stated " as
  the heat evolved or absorbed in a chemical
  process is the same whether the process takes
  place in one or several steps. Hess's law is also
  noted as the law of constant heat summation.

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Standard Heat of Formation

• Change in enthalpy from the formation of 1 mol of
  a compound, in its standard state, from its
  elements.
• Symbol
• ? refers to standard conditions
• Units kJ/mol
• Example
• S(s) O2(g) ? SO2(g) -297 kJ/mol
• Table 16-7 on page 510 lists Standard Hf

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Chemical Reactions

• Compare and contrast the following graphs
• What did you notice?
• How would you write these reactions in standard
  format?

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Standard Heat of Formation

• What can we tell from values?
• Positive value means?
• Negative value means?
• Thermochemical equations
• 4Fe(s) 3O2(g) ? 2Fe2O3(s) 1625 kJ
• NH4NO3(s) 27 kJ ? NH4(aq) NO3-(aq)

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Where do Standard Heats of formation come from?

• When you state the height of a mountain, it is
  relative to another point (usually sea level).
• In the same way enthalpies of formation are
  stated based on the following arbitrary standard
• Every free element in its standard state has a
  value of exactly 0.0 kJ.
• That way when the heat of formation is negative
  the system has lost heat, when positive the
  system has gained heat!

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Check for understanding!

• Do elements in their standard states possess zero
  energy?
• Why are elements in their standard states
  assigned enthalpies of zero?
• What does the 33.2 on the graph tell you?
• What does the -396 on the graph tell you?

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Enthalpy change from Standard Heat of Formation

• Use standard heat of formation to calculate
  ?Hrxn? for the combustion of methane
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• We can summarize Hesss Law into the following
  equation
• ?Hrxn? S?Hf?(products) - S?Hf?(reactants)
• The symbol S means to take the sum of the terms.

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Enthalpy change from Standard Heat of Formation

• Use standard heat of formation to calculate
  ?Hrxn? for the combustion of methane
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• First look up ?Hf? values
• Second Use the formula and multiply each term by
  the coefficient of the substance in the balanced
  chemical equation
• Third Do the math

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Enthalpy change from Standard Heat of Formation

• Use standard heat of formation to calculate
  ?Hrxn? for the combustion of methane
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• ?Hf?(CO2) -394 kJ ?Hf?(H2O) -286 kJ
  ?Hf?(CH4) -75 kJ ?Hf?(O2) 0.0 kJ
• ?Hrxn? (-394 kJ) (2)(-286 kJ) (-75 kJ)
  (2)(0.0 kJ)
• products - reactants
• ?Hrxn? -966 kJ -75 kJ -891 kJ
• Is this reaction endothermic or exothermic?

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Exothermic Reactions
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Exothermic Reactions

• Potential Energy converts to Kinetic Energy as
  you release the energy to the surroundings.
• Where is the Potential energy stored?
• How do we know there is a shift to Kinetic
  energy?
• Most reactions are exothermic and spontaneous

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Endothermic Reactions
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Endothermic Reactions

• Energy content of products is greater than
  reactants
• If the products have more energy why are the
  surroundings cold? (Hint PE?)
• The surroundings feel cold because the bonds
  absorb the heat energy from the surroundings, so
• Kinetic energy converts to Potential energy

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Reaction Spontaneity

• What happens when you leave a iron nail outside
  for a few months?
• What happens when you light a gas stove?
• Do these reactions take place spontaneously?
  (without outside intervention)
• Will the reverse of these reactions take place
  spontaneously?
• What about ice melting at room temperature?
• There is something more than ?H determining
  spontaneity!

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Entropy

• Symbol S
• Measure of the disorder or randomness of the
  particles that make up the system
• Molecules are more likely to exist in a high
  state of disorder than in a low state of
  disorder.
• Change in entropy is similar to change in
  enthalpy
• ?Ssystem Sproducts - Sreactants

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Entropy

• Do you expect the ?S of the phase change shown
  below to be positive or negative?
• Sproducts gt Sreactants ?Ssystem positive
• Sproducts lt Sreactants ?Ssystem negative

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Entropy

• What happens to molecules when you increase their
  temperature?
• Do you think this will increase or decrease their
  entropy?
• ?S positive more entropy, ie more disorder
• ?S negative less entropy, ie less disorder
• Reactions tend to go spontaneous towards
  increased entropy.

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Entropy Practice

• Predict the ?S for the following changes
• H2O(l) ? H2O(g)
• CO2(g) ? CO2(aq)
• 2SO3(g) ? 2SO2(g) O2(g)
• NaCl(s) ? Na(aq) Cl-(aq)
• CH3OH(l) ? CH3OH(aq)

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Gibbs Free Energy

• By calculating free energy (energy that is
  available to do work) we can determine if a
  reaction is spontaneous.
• Just like Enthalpy and Entropy we can only
  measure the free energy as a change.
• ?Gsystem ?Hsystem T?Ssystem
• The sign of ?Gsystem tells you if the reaction is
  spontaneous
• Negative spontaneous, will occur
• Postive nonspontaneous, will NOT occur

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Gibbs Free Energy

• How do the enthalpies and entropies affect
  reaction spontaneity?
• What is happening if ?Gsystem 0?

-?Hsystem ?Hsystem
?Ssystem Always spontaneous, -?Gsystem Spontaneous only at high temperatures, or - ?Gsystem
-?Ssystem Spontaneous only at low temperatures, or - ?Gsystem Never spontaneous ?Gsystem
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Review of Symbols

• ?H Tells us if a reaction is endothermic or
  exothermic (measure of change in energy)
• Postive endothermic
• Negative exothermic
• ?S Tells us if the reaction is more or less
  ordered (randomness of particles)
• Positive more disordered
• Negative more ordered
• ?G Tells us if the reaction is spontaneous by
  determining the amount of energy available to do
  work.
• Positive nonspontaneous
• Negative spontaneous