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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

21 22 23 24 25 26 27 28 29 30

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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

21 22 23 24 25 26 27 28 29 30

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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

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

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

21 22 23 24 25 26 27 28 29 30

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)