Chapter 6 thermochemistry
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CHEMISTRY. Chapter 6 Thermochemistry. Thermodynamics. - is the study of energy and the transformations it undergoes in chemical reactions. Units: Joules (J) calorie = 4.184 J calorie – the amount of E needed to raise the temperature of 1 gram of water 1 o C. Energy.

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Chapter 6 Thermochemistry

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Chapter 6 thermochemistry

CHEMISTRY

Chapter 6Thermochemistry

Chapter 5


Thermodynamics

Thermodynamics

- is the study of energy and the transformations it undergoes in chemical reactions.

Units: Joules (J)

calorie = 4.184 J

calorie – the amount of E needed to raise the temperature of 1 gram of water 1 oC

Chapter 5


Energy

Energy

  • Examples of energy:

  • Potential E

  • Kinetic E

  • Work – E needed to move an object against a force

  • Heat – E transferred from hot to cold objects - capacity to do work or transfer heat

Chapter 5


Today s topics

Today’s Topics

  • 4 Thermodynamic Functions

  • Definition of State Function

  • Internal Energy

  • Point of Views – System, Surroundings Universe

  • Sign Conventions

Chapter 5


4 thermodynamic functions

4 Thermodynamic Functions

  • E-Internal Energy

  • H-Enthalpy

  • S-Entropy

  • G-Gibb’s Free Energy

Chapter 5


What is a state function

What is a state function?

  • A state function is a property that depends on the present condition and not on how the change occurs.

  • Derived from calculating the change:D = [Final value – initial value]

Chapter 5


4 state functions

4 State Functions

  • Therefore:

  • DE = Ef -Ei

    DH = Hf - Hi

    DS = Sf - Si

    DG = Gf - Gi

Chapter 5


Internal energy e

Internal Energy (E)

  • Internal Energy - is the sum of the kinetic and potential energy of a system

  • Is the sum of heat (q) and work (w)

    DE = Efinal – Einitial

    DE = q + w

Chapter 5


Point of views

Point of Views

  • System - the reaction we are studying

  • Surroundings – anything else besides the reactionFor example: the container, you, etc….

  • Universe – system + surroundings

Chapter 5


Q heat and w work

q = heat and w = work

  • If q is (+), system gaining heat from surroundings (endo)

  • If q is (-), system giving up heat to the surroundings (exo)

  • If w is (+), system is the recipient of work from the surroundings. (in short, surroundings is doing work on the system.

Chapter 5


Thermodynamic functions

Thermodynamic Functions

  • Have value, unit and magnitude.

  • The sign of DE depends on the magnitude of q and w.

  • Knowing the value of DE does not tell us which variable is larger, q or w.

Chapter 5


Problem 1

Problem 1

  • Calculate the change in internal energy of the system for a process in which the system absorbs 140 J of heat from the surroundings and does 85 J of work on the surroundings.

Chapter 5


Problem 2

Problem 2

  • Consider the reaction of hydrogen and oxygen gases to produce water. As the reaction occurs, the system loses 1150 J of heat to the surroundings. The expanding gas does 480 J of work on the surroundings as it pushes against the atmosphere. Calculate the change in the internal energy of the system?

Chapter 5


Problem 3

Problem 3

  • Calculate DE and determine whether the process is endothermic or exothermic.

  • 1.)q = 1.62 kJ and w = -874 J

  • 2.)The system releases 113 kJ of heat to the surroundings and does 39 kJ of work.

  • 3.) The system absorbs 77.5 kJ of heat while doing 63.5 kJ of work on the surroundings.

Chapter 5


Thermodynamic functions1

Thermodynamic Functions

  • Have value, unit and magnitude.

  • The sign of DE depends on the magnitude of q and w.

  • Knowing the value of DE does not tell us which variable is larger, q or w.

Chapter 5


Enthalpy

Enthalpy

  • Enthalpy – accounts for heat flow in chemical reactions that occur at constant P when nothing other than P-V work are performed

    DH = Hfinal - Hinitial

  • If:H = E + PDV

  • Then:DH = DE + PDV

Chapter 5


Ways of measuring d h

Ways of Measuring DH

  • Calorimetry

  • Hess’s Law

  • Heats of Formation (DHof)

Chapter 5


Calorimetry

Calorimetry

  • - Is the measurement of heat flow

  • Specific Heat capacity (C) - the amount of heat needed to raise the temperature of 1 g of substance 1 oC.

  • The greater the heat capacity, the greater the heat required to produce a rise in temp.

Chapter 5


Calorimetry equation

Calorimetry Equation

  • q = mCDT

  • - qsystem = qsurroundings

  • - qsystem = qwater + qcalorimeter

  • - qsubstance = (mCDTsolution + mCDTcalorimeter)

  • Simplifies to:

  • - qsubstance = (mCDTsolution + CDTcalorimeter)

Chapter 5


Calorimetry problem

Calorimetry Problem

  • A 30.0 gram sample of water at 280 K is mixed with 50.0 grams of water at 330 K. Calculate the final temperature of the mixture assuming no heat loss to the surroundings. The specific heat capacity of the solution is 4.18 J/g-oC.

Chapter 5


Calorimetry problem1

Calorimetry Problem

  • A 46.2 gram sample of copper is heated to 95.4 oC and then placed in a calorimeter containing 75.0 gram of water at 19.6 oC. The final temperature of the metal and water is 21.8 oC. Calculate the specific heat capacity of copper, assuming that all the heat lost by the copper is gained by the water.

Chapter 5


Calorimetry problem2

Calorimetry Problem

  • A 15.0 gram sample of nickel metal is heated to 100.0 oC and dropped into 55.0 grams of water, initially at 23 oC. Assuming that no heat is lost to the calorimeter, calculate the final temperature of the nickel and water. The specific heat of nickel is 0.444 J/g-oC. The specific heat of water is 4.18 J/g-oC.

Chapter 5


Problem 4

Problem 4

  • The specific heat of water is 4.18 J/g-K.

  • How much heat is needed to warm 250 g of water from 22 oC to 98 oC?

  • What is the molar heat capacity of water?

  • Molar heat Capacity = C x molar mass

Chapter 5


Problem 5

Problem 5

  • Large beds of rocks are used in solar heated homes to store heat. Assume that the specific heat of rocks is 0.82 J/g-K.

  • A. Calculate the amount of heat absorbed by 50 kg. of rocks if their temperature rose by 12.0 oC.

  • B. What temperature change would these rocks undergo if they emitted 450 kJ of heat?

Chapter 5


Chapter 6 thermochemistry

  • Work – E needed to move an object against a force

  • When P is constant, P-V work is given byw = - PDV

Chapter 5


Enthalpy1

Enthalpy

  • Enthalpy of a reaction or heat of Reaction: DH = Hproducts - Hreactants

  • 1. sign of DH depends on the amount of reactant consumed

  • 2. DH sign is opposite for backwards reaction

  • 3. DHrxn depends on the physical state of the reactants and products.

Chapter 5


Problem 11

Problem 1

  • Given the reaction:

  • 2H2 (g) + O2 (g)  2 H2O (g) DH = -483 kJ

  • Calculate the DH value for:

  • 2 H2O (g)  2H2 (g) + O2 (g)

Chapter 5


Problem 21

Problem 2

  • Given the reaction:

  • 2H2 (g) + O2 (g)  2 H2O (g) DH = -483 kJ

  • How much heat is released when 10.5 grams of H2 is burned in a constant-pressure system?

Chapter 5


Ways of measuring d h1

Ways of Measuring DH

  • Calorimetry

  • Hess’s Law

  • Heats of Formation (DHof)

Chapter 5


Hess law

Hess’ Law

  • If a reaction is carried out in steps, DH for the reaction will equal the sum of the enthalpy changes for the individual steps.

Chapter 5


Problem 12

Problem 1

  • Given:

  • C (s) + O2 (g)  CO2 (g) DH = -393.5 kJ

  • CO(g) + ½ O2 (g)  CO2 (g) DH = -283.5 kJ

  • Calculate the enthalpy of combustion for:C(s) + ½ O2 (g)  CO (g)

Chapter 5


Problem 22

Problem 2

  • Given:

  • C(graphite) + O2 (g)  CO2 (g) DH = -393.5 kJ

  • C(diamond) + O2 (g)  CO2 (g) DH = -395.4 kJ

  • Calculate the enthalpy of combustion for:C(graphite)  C (diamond)

Chapter 5


Problem 31

Problem 3

  • Given:

  • C2H2 (g) + 5/2 O2 (g)  2CO2 (g) + H2O (l) DH = -1299.6 kJ

  • C (s) + O2 (g)  CO2 (g) DH = -395.4 kJ

  • H2 (g) + 1/2 O2 (g)  H2O (l) DH = -285.8 kJ

  • Calculate the enthalpy of combustion for:2C (s) + H2 (g)  C2H2 (g)

Chapter 5


Enthalpies of formation

Enthalpies of Formation

  • - known as heat of formation

  • - gives the energy needed for a compound to form

  • Standard enthalpy - is the enthalpy change (DH) when the reactants and products are in their standard state, usually 1 atm and 25 oC- denoted by DHo ( ex. DHof)

Chapter 5


Standard enthalpy of formation d h o f

Standard Enthalpy of Formation, (DHof)

  • By definition:

  • The standard enthalpy of formation ( ex. DHof) of the most stable form of any element is ZERO because there is no formation reaction needed when the element is in its standard state.

  • - important for diatomic molecules

  • - need knowledge of standard states of compounds

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

  • If there is more than one state for a substance under standard conditions, the more stable one is used.

  • Standard enthalpy of formation of the most stable form of an element is zero.

  • Using Enthalpies of Formation of Calculate Enthalpies of Reaction

  • We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

  • Using Enthalpies of Formation of Calculate Enthalpies of Reaction

  • For a reaction

Chapter 5


Next topic

Next Topic

  • DS = Entropy = disorder

Chapter 5


Thermodynamic question

Thermodynamic Question

  • Can a process occur?

  • Spontaneous or Non-spontaneous?

    • Forward vs. Reverse reactions

    • Example: Gas Expansion

  • Reversible or irreversible?

Chapter 5


Entropy

Entropy

  • Entropy, S, is a measure of the disorder of a system.

  • Spontaneous reactions proceed to lower energy or higher entropy.

  • In ice, the molecules are very well ordered because of the H-bonds.

  • Therefore, ice has a low entropy.

Chapter 5


Chapter 6 thermochemistry

  • As ice melts, the intermolecular forces are broken (requires energy), but the order is interrupted (so entropy increases).

  • Water is more random than ice, so ice spontaneously melts at room temperature.

  • Conclusion: The higher the entropy the more spontaneous the reaction.

Chapter 5


Chapter 6 thermochemistry

  • Generally, when an increase in entropy in one process is associated with a decrease in entropy in another, the increase in entropy dominates.

  • Entropy is a state function.

  • For a system, S = Sfinal - Sinitial.

  • If S > 0 the randomness increases, if S < 0 the order increases.

Chapter 5


Chapter 6 thermochemistry

Entropy and the Second Law of Thermodynamics

  • Entropy

  • Suppose a system changes reversibly between state 1 and state 2. Then, the change in entropy is given by

    • at constant T where qrev is the amount of heat added reversibly to the system. (Example: a phase change occurs at constant T with the reversible addition of heat.)

Chapter 5


Chapter 6 thermochemistry

Entropy and the Second Law of Thermodynamics

  • The Second Law of Thermodynamics

  • Spontaneous processes have a direction.

  • In any spontaneous process, the entropy of the universe increases.

  • Suniv = Ssys + Ssurr: the change in entropy of the universe is the sum of the change in entropy of the system and the change in entropy of the surroundings.

  • Entropy is not conserved: Suniv is increasing.

Chapter 5


Chapter 6 thermochemistry

Entropy and the Second Law of Thermodynamics

  • The Second Law of Thermodynamics

  • Reversible process: Suniv = 0.

  • Spontaneous process (i.e. irreversible): Suniv > 0.

  • Ssys for a spontaneous process can be less than 0as long asSsurr > 0.This would makeSuniv still (+).

  • For an isolated system, Ssys = 0 for a reversible process and Ssys > 0 for a spontaneous process.

Chapter 5


Third law of thermodynamics

Third Law of Thermodynamics

  • The entropy of a perfect crystal at 0 K is zero.

  • Entropy changes dramatically at a phase change.

  • As we heat a substance from absolute zero, the entropy must increase.

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Chapter 6 thermochemistry

Entropy Changes in Chemical Reactions

  • Absolute entropy can be determined from complicated measurements.

  • Standard molar entropy, S: entropy of a substance in its standard state. Similar in concept to H.

  • Units: J/mol-K. Note units of H: kJ/mol.

  • Standard molar entropies of elements (S ) are not 0.

  • For a chemical reaction which produces n moles of products from m moles of reactants:

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Chapter 6 thermochemistry

Gibbs Free Energy

  • For a spontaneous reaction the entropy of the universe must increase.

  • Reactions with large negative H values are spontaneous.

  • How do we correlate S and H to predict whether a reaction is spontaneous?

  • Gibbs free energy, G, of a state is

  • For a process occurring at constant temperature

Chapter 5


Gibb s free energy

Gibb’s Free Energy

  • Is the capacity to do maximum useful work.

  • Heat decreases the amount of useful work done.

Chapter 5


Chapter 6 thermochemistry

WORK

  • At constant P, wsys = - PDV

  • At constant temperature, wsys = - TDS

  • If DG = DH – TDS, for maximum useful work, DH must be = 0.

Chapter 5


Chapter 6 thermochemistry

Gibbs Free Energy

  • Three important conditions:

    • If G < 0 then the forward reaction is spontaneous.

    • If G = 0 then reaction is at equilibrium and no net reaction will occur.

    • If G > 0 then the forward reaction is not spontaneous. If G > 0, work must be supplied from the surroundings to drive the reaction.

  • For a reaction the free energy of the reactants decreases to a minimum (equilibrium) and then increases to the free energy of the products.

Chapter 5


Problem

Problem

  • Correlate S and H to predict whether a reaction is non-spontaneous or spontaneous. If spontaneous, determine whether the reaction will be spontaneous at all temperature, at high temperature, or at low temperature.

  • A. If S is (-) and H is (+).

  • B. If S is (-) and H is (-).

  • C. If S is (+) and H is (+).

  • D. If S is (+) and H is (-).

Chapter 5


Chapter 6 thermochemistry

Gibbs Free Energy

  • Consider the formation of ammonia from N2 and H2.

  • Initially ammonia will be produced spontaneously (Q < Keq).

  • After some time, the ammonia will spontaneously react to form N2 and H2 (Q > Keq).

  • At equilibrium, ∆G = 0 and Q = Keq.

Chapter 5


Chapter 6 thermochemistry

Gibbs Free Energy

  • Standard Free-Energy Changes

  • Gf

  • Standard states are: pure solid, pure liquid, 1 atm (gas), 1 M concentration (solution), and G = 0 for elements.

  • G for a process is given by

  • The quantity G for a reaction tells us whether a mixture of substances will spontaneously react to produce more reactants (G > 0) or products (G < 0).

Chapter 5


Chapter 6 thermochemistry

Free Energy and Temperature

Chapter 5


Chapter 6 thermochemistry

Free Energy and The Equilibrium Constant

  • Recall that G and K (equilibrium constant) apply to standard conditions.

  • Recall that G and Q (equilibrium quotient) apply to any conditions.

  • It is useful to determine whether substances under any conditions will react:

Chapter 5


Chapter 6 thermochemistry

Free Energy and The Equilibrium Constant

  • At equilibrium, Q = K and G = 0, so

  • From the above we can conclude:

    • If G < 0, then K > 1.

    • If G = 0, then K = 1.

    • If G > 0, then K < 1.

Chapter 5


Calculating d h d s d g

Calculating DH, DS, DG

  • Use Hess’ Law

  • Standard Heats of Formation (DHo, DSo, DGo )

  • Equations

Chapter 5


Chapter 6 thermochemistry

  • Next Chapter: ELECTROCHEMISTRY

Chapter 5


Chapter 6 thermochemistry

  • Galvanic Cell – spontaneous

  • Electrochemical Cell – non-spontaneous

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Hess law1

Hess’ Law

  • If a reaction is carried out in a series of steps, DH for the reaction will equal the sum of the enthalpy changes for the individual steps

Chapter 5


Chapter 6 thermochemistry

Chapter 5


Chapter 6 thermochemistry

The Nature of Energy

  • Kinetic Energy and Potential Energy

  • Kinetic energy is the energy of motion:

  • Potential energy is the energy an object possesses by virtue of its position.

  • Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.

Chapter 5


Chapter 6 thermochemistry

The Nature of Energy

  • Systems and Surroundings

  • System: part of the universe we are interested in.

  • Surroundings: the rest of the universe.

  • We sometimes use the calorie instead of the joule:

  • 1 cal = 4.184 J (exactly)

  • A nutritional Calorie:

  • 1 Cal = 1000 cal = 1 kcal

Chapter 5


Chapter 6 thermochemistry

The Nature of Energy

  • Transferring Energy: Work and Heat

  • Force is a push or pull on an object.

  • Work is the product of force applied to an object over a distance:

  • Energy is the work done to move an object against a force.

  • Heat is the transfer of energy between two objects.

  • Energy is the capacity to do work or transfer heat.

Chapter 5


Chapter 6 thermochemistry

The First Law of Thermodynamics

  • Internal Energy

  • Internal Energy: total energy of a system.

  • Cannot measure absolute internal energy.

  • Change in internal energy,

Chapter 5


Chapter 6 thermochemistry

The First Law of Thermodynamics

  • Relating DE to Heat and Work

  • Energy cannot be created or destroyed.

  • Energy of (system + surroundings) is constant.

  • Any energy transferred from a system must be transferred to the surroundings (and vice versa).

  • From the first law of thermodynamics:

    • when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:

Chapter 5


Chapter 6 thermochemistry

The First Law of Thermodynamics


Chapter 6 thermochemistry

The First Law of Thermodynamics

  • Exothermic and Endothermic Processes

  • Endothermic: absorbs heat from the surroundings.

  • Exothermic: transfers heat to the surroundings.

  • An endothermic reaction feels cold.

  • An exothermic reaction feels hot.

  • State function: depends only on the initial and final states of system, not on how the internal energy is used.

Chapter 5


Chapter 6 thermochemistry

Enthalpy

  • Chemical reactions can absorb or release heat.

  • However, they also have the ability to do work.

  • For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work.

  • Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)

  • The work performed by the above reaction is called pressure-volume work.

  • When the pressure is constant,

Chapter 5


Chapter 6 thermochemistry

Enthalpy

  • Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.

  • Enthalpy is a state function.

  • If the process occurs at constant pressure,

Chapter 5


Chapter 6 thermochemistry

Enthalpy

  • Since we know that

  • We can write

  • When DH, is positive, the system gains heat from the surroundings.

  • When DH, is negative, the surroundings gain heat from the system.

Chapter 5


Chapter 6 thermochemistry

Enthalpy

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Reaction

  • For a reaction:

  • Enthalpy is an extensive property (magnitude DH is directly proportional to amount):

  • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)DH = -802 kJ

  • 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) DH = -1604 kJ

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Reaction

  • When we reverse a reaction, we change the sign of DH:

  • CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)DH = +802 kJ

  • Change in enthalpy depends on state:

  • H2O(g)  H2O(l)DH = -88 kJ

Chapter 5


Chapter 6 thermochemistry

Calorimetry

  • Heat Capacity and Specific Heat

  • Calorimetry = measurement of heat flow.

  • Calorimeter = apparatus that measures heat flow.

  • Heat capacity = the amount of energy required to raise the temperature of an object (by one degree).

  • Molar heat capacity = heat capacity of 1 mol of a substance.

  • Specific heat = specific heat capacity = heat capacity of 1 g of a substance.

Chapter 5


Chapter 6 thermochemistry

Calorimetry

  • Constant Pressure Calorimetry

  • Atmospheric pressure is constant!

Chapter 5


Chapter 6 thermochemistry

Hess’s Law

  • Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.

  • For example:

  • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)H = -802 kJ

  • 2H2O(g)  2H2O(l) H = -88 kJ

  • CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)H = -890 kJ

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

  • If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof .

  • Standard conditions (standard state): 1 atm and 25 oC (298 K).

  • Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state.

  • Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

  • If there is more than one state for a substance under standard conditions, the more stable one is used.

  • Standard enthalpy of formation of the most stable form of an element is zero.

  • Using Enthalpies of Formation of Calculate Enthalpies of Reaction

  • We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

Chapter 5


Chapter 6 thermochemistry

Enthalpies of Formation

  • Using Enthalpies of Formation of Calculate Enthalpies of Reaction

  • For a reaction

Chapter 5


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