thermodynamics. study of energy changes that accompany physical and chemical processes....
TRANSCRIPT
Thermodynamics
Thermodynamics
• study of energy changes that accompany physical and chemical processes.
• Thermochemistry is one component of thermodynamics which focuses on how energy changes are measured and predicted.
Energy
• ability to do work or transfer heat.1. kinetic energy-energy of motion (example: electrical) KE = 1/2mv2
2. potential energy-stored energy due to composition (example: chemical)
Types of Changes
• Endothermic- a reaction in which energy is absorbed from the surroundings.
• Exothermic-a reaction in which energy is released.
First Law of Thermodynamics
• the total amount of energy in the universe is constant.
• Also known as the Law of Conservation of Energy which states that energy cannot be created or destroyed, but can only change form.
Thermodynamic Terms
• System-the substances being studied.• Surroundings- everything in the
system’s environment• Universe-the system plus the
surroundings.• State Function-property whose value
depends only on the state of the system-not on the pathway it took to get there.
Examples of State Functions• Pressure, volume, and temperature are
all examples of state functions.• Example: Ti= 30oC and Tf= 22oC
∆ T = -8oCHow the temperature change occurred
does not matter• ∆ X = Xf – Xi If there is an increase in
X, ∆X > 0. If there is a decrease in X, ∆X < 0.
Enthalpy
• Enthalpy change (∆H) is the quantity of heat (q) transferred in or out of a system.
• ∆H = Hfinal – Hinitial
• ∆H = Hproducts - Hreactants
Exothermic Reactions
• ∆H = energy of products – energy of reactants• If the reaction is exothermic, ∆H < 0.
Endothermic Reactions
• ∆H = energy of products – energy of reactants• If the reaction is endothermic, ∆H > 0.
Measuring Energy Changes• Calorimetry-process of measuring
energy changes• Calorimeter-device used to
measure heat• Heat released by reaction = heat
gained by calorimeter + heat gained by the solution
• q = m(∆T)C
Internal Energy
• Internal Energy, E, is all of the energy contained within a substance.
1) kinetic energy of the molecules2) energy of attraction and
repulsion between subatomic particles, molecules, ions, etc.
3) other forms of energy• Internal energy is a state function
• ∆E = Efinal – Einitial
= Eproducts – Ereactants
= q + wq represents heat and w represents
work• ∆E = heat absorbed by system +
work done on the system
The following conventions apply to the signs of q and w • q is positive: heat is absorbed by the
system from the surroundings (endothermic)
• q is negative: heat is released by the system to the surroundings (exothermic)
• w is positive: work is done on the system by the surroundings
• w is negative: work is done by the system on the surroundings.
Writing Equations
• When ∆E < 0, energy is released by the system and can be written as a product.
• When ∆E > 0, energy is absorbed by the system and can be written as a reactant.
Effect of Pressure on Work
• Work done on a system = -P∆V• If volume decreases (could be due to a
decrease in the number of moles of gas), work is done on the system so the sign of w is positive.
• If volume increases (could be due to an increase in the number of moles of gas), work is done by the system so the sign of w is negative.
• In constant volume reactions, no work is done so E = q
Relationship between ∆H and ∆E.
• ∆H = ∆E + P∆V (constant T and P) *use with physical changes
• ∆H = qp (constant T and P)
• ∆H= ∆E + (∆n)RT or
∆E = ∆H – (∆n)RT (constant T and P) *use with chemical changes
Hess’s Law
• Since enthalpy is a state function, the change in enthalpy in going from some initial state to some final state is independent of the pathway.
• The change in enthalpy in going from a particular set of reactants to a particular set of products is the same whether the reaction takes place in one step or a series of steps.
Characteristics of Enthalpy Changes
• If the reaction is reversed, the sign of ∆H is also reversed.
• The magnitude of ∆H is directly proportional to the quantities of reactants and products. If the coefficients are multiplied by an integer, the value of ∆H is multiplied by the same number.
Standard Enthalpies of Formation
• For a reaction under standard conditions of constant pressure, enthalpy changes can be measured using a calorimeter.
• Because some reactions proceed too slowly, a process is needed that allows the enthalpy change to be calculated.
Standard Enthalpies of Formation• The standard enthalpy of formation
(∆Hfo) is defined as the change in
enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states.
• The standard state is a precisely defined reference state.
• See page 246 for definitions of standard states.
Bond Energy and Enthalpy
• Bond energy values can be used to calculate approximate energies for reactions also.
• ∆H = sum of energies required to break old bonds (positive sign/endothermic) plus the sum of energies released in formation of new bonds (negative sign/exothermic)
Spontaneous Processes
• A process is spontaneous if it occurs without outside intervention.
• Spontaneous processes may be either fast or slow.
• Thermodynamics lets us predict whether a process will occur but gives no information about the amount of time required for the process.
Entropy
• Entropy (S) is a measure of the molecular randomness or disorder.
• The driving force of spontaneous processes is an increase in entropy of the universe.
• The natural progression of things is from order to disorder, from lower entropy to higher entropy.
Entropy (continued)
• Entropy describes the number of arrangements (positions/energy levels) that are available to a system existing in a given state.
• Nature spontaneously proceeds toward the states that have the highest probabilities of existing.
• The states with the highest probabilities of existing is that which has the greatest disorder. (Sgas> Sliquid > Ssolid ; in the gaseous state, molecules have many more positions available to them and are therefore more disordered).
Predicting Entropy Changes• ∆S = Sfinal - Sinitial
• If the entropy increases, ∆S is > 0• If the entropy decreases, ∆S is < 0
Which of the following has the greatest entropy?
0%
0%
0%
10
1. 1 mole of solid CO2
2. 1 mole of gaseous CO2
3. Both are equal
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Which has the greatest entropy?
0%
0%
0% 1. 1 mole of N2 gas at 1 atm
2. 1 mole of N2 gas at 0.01 atm
3. Both have the same entropy
10
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21 22 23 24 25 26 27 28 29 30
What is the sign for the entropy change when solid sugar is added to water to form a solution?
0%
0% 1. Positive
2. Negative
10
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21 22 23 24 25 26 27 28 29 30
What is the sign for the entropy change when iodine vapor condenses on a cold surface to form crystals?
0%
0% 1. Positive
2. Negative
10
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Second Law of Thermodynamics• The second law of thermodynamics states that
in any spontaneous process, there is always an increase in the entropy of the universe.
• In other words, the entropy of the universe is increasing and is not conserved.
• ∆Suniv = ∆Ssys + ∆Ssurr • If ∆Suniv > 0, the process is spontaneous as
written.• If ∆Suniv < 0, the process is spontaneous in the
opposite direction.
The Effect of Temperature on Spontaneity
• An exothermic process in the system causes heat to flow to the surroundings, increasing the random motions and entropy of the surroundings. ∆Ssurr >0
• The opposite is true for endothermic processes.
• As a result, nature tends to seek the lowest possible energy.
Effect of temp (continued)
• The magnitude of ∆Ssurr depends on the temperature.
• ∆Ssurr depends directly on the quantity of heat transferred and inversely on the temperature.
• ∆Ssurr = -∆H
T
Predict the sign of ∆Ssurr for the following process: H2O(l) H2O (g)
0%
0% 1. Positive
2. Negative
10
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21 22 23 24 25 26 27 28 29 30
Predict the sign of ∆Ssurr for the following process: CO2(g) CO2 (s)
0%
0% 1. Positive
2. Negative
10
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Third Law of Thermodynamics• The Third Law of Thermodynamics states that
the entropy of a perfect crystal at 0 K is zero.• The standard entropy values So of many
common substances at 298 K and 1 atm are listed in Appendix 4.
• Because entropy is a state function,∆So
rxn = ΣnpSoproducts – ΣnrSo
reactants
• Generally, the more complex the molecule, the higher the standard entropy value.
Free Energy
• Free energy (G) is the energy that is available to do work.
• ∆G = ∆ H – T ∆ S where H is enthalpy, T is Kelvin temp, and S is entropy.
• A process is spontaneous (at constant T and P) in the direction in which the free energy decreases (- ∆G means + ∆Suniv)
• See the table on page 761• At the melting point and boiling point,
∆G = 0.
Free Energy and Chemical Reactions
• For chemical reactions, we are often interested in the standard free energy change (∆Go), the change in free energy that will occur if the reactants in their standard states are converted to the products in their standard states.
Calculating Free Energy
• ∆Go cannot be measured directly, but can be calculated from other quantities.
(∆Go = ∆Ho - T ∆So)• Free energy is a state function and can be determined
using similar procedures as those for finding ∆H using Hess’s law.
• Free energy can also be calculated using standard free energies of formation (the free energy that accompanies the formation of 1 mole of that substance from its constituent elements with all the reactants and products in their standard states.(∆Go
rxn = ΣnpGfoproducts – ΣnrGf
oreactants)